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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid .
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how can you determine if the bacteria have picked up your vector ( plasmid ) dna ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid .
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and how can you determine if you have inserted any dna into the plasmid ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria .
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how can you determine which clone contains the dna of interest to you ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene .
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what are the possible reasons for why a gene is unsuccessfully cloned ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today .
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should n't that be exxon less genes rather than intron less genes ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria .
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who , when invented the way of dna cloning ?
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key points : dna cloning is a molecular biology technique that makes many identical copies of a piece of dna , such as a gene . in a typical cloning experiment , a target gene is inserted into a circular piece of dna called a plasmid . the plasmid is introduced into bacteria via process called transformation , and bacteria carrying the plasmid are selected using antibiotics . bacteria with the correct plasmid are used to make more plasmid dna or , in some cases , induced to express the gene and make protein . introduction when you hear the word “ cloning , ” you may think of the cloning of whole organisms , such as dolly the sheep . however , all it means to clone something is to make a genetically exact copy of it . in a molecular biology lab , what ’ s most often cloned is a gene or other small piece of dna . if your friend the molecular biologist say that her “ cloning ” isn ’ t working , she 's almost certainly talking about copying bits of dna , not making the next dolly ! overview of dna cloning dna cloning is the process of making multiple , identical copies of a particular piece of dna . in a typical dna cloning procedure , the gene or other dna fragment of interest ( perhaps a gene for a medically important human protein ) is first inserted into a circular piece of dna called a plasmid . the insertion is done using enzymes that “ cut and paste ” dna , and it produces a molecule of recombinant dna , or dna assembled out of fragments from multiple sources . next , the recombinant plasmid is introduced into bacteria . bacteria carrying the plasmid are selected and grown up . as they reproduce , they replicate the plasmid and pass it on to their offspring , making copies of the dna it contains . what is the point of making many copies of a dna sequence in a plasmid ? in some cases , we need lots of dna copies to conduct experiments or build new plasmids . in other cases , the piece of dna encodes a useful protein , and the bacteria are used as “ factories ” to make the protein . for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria . the basic steps are : cut open the plasmid and `` paste '' in the gene . this process relies on restriction enzymes ( which cut dna ) and dna ligase ( which joins dna ) . transform the plasmid into bacteria . use antibiotic selection to identify the bacteria that took up the plasmid . grow up lots of plasmid-carrying bacteria and use them as `` factories '' to make the protein . harvest the protein from the bacteria and purify it . let 's take a closer look at each step . 1 . cutting and pasting dna how can pieces of dna from different sources be joined together ? a common method uses two types of enzymes : restriction enzymes and dna ligase . a restriction enzyme is a dna-cutting enzyme that recognizes a specific target sequence and cuts dna into two pieces at or near that site . many restriction enzymes produce cut ends with short , single-stranded overhangs . if two molecules have matching overhangs , they can base-pair and stick together . however , they wo n't combine to form an unbroken dna molecule until they are joined by dna ligase , which seals gaps in the dna backbone . our goal in cloning is to insert a target gene ( e.g. , for human insulin ) into a plasmid . using a carefully chosen restriction enzyme , we digest : the plasmid , which has a single cut site the target gene fragment , which has a cut site near each end then , we combine the fragments with dna ligase , which links them to make a recombinant plasmid containing the gene . 2 . bacterial transformation and selection plasmids and other dna can be introduced into bacteria , such as the harmless e. coli used in labs , in a process called transformation . during transformation , specially prepared bacterial cells are given a shock ( such as high temperature ) that encourages them to take up foreign dna . a plasmid typically contains an antibiotic resistance gene , which allows bacteria to survive in the presence of a specific antibiotic . thus , bacteria that took up the plasmid can be selected on nutrient plates containing the antibiotic . bacteria without a plasmid will die , while bacteria carrying a plasmid can live and reproduce . each surviving bacterium will give rise to a small , dot-like group , or colony , of identical bacteria that all carry the same plasmid . not all colonies will necessarily contain the right plasmid . that ’ s because , during a ligation , dna fragments don ’ t always get “ pasted ” in exactly the way we intend . instead , we must collect dna from several colonies and see whether each one contain the right plasmid . methods like restriction enzyme digestion and pcr are commonly used to check the plasmids . 3 . protein production once we have found a bacterial colony with the right plasmid , we can grow a large culture of plasmid-bearing bacteria . then , we give the bacteria a chemical signal that instructs them to make the target protein . the bacteria serve as miniature “ factories , '' churning out large amounts of protein . for instance , if our plasmid contained the human insulin gene , the bacteria would start transcribing the gene and translating the mrna to produce many molecules of human insulin protein . once the protein has been produced , the bacterial cells can be split open to release it . there are many other proteins and macromolecules floating around in bacteria besides the target protein ( e.g. , insulin ) . because of this , the target protein must be purified , or separated from the other contents of the cells by biochemical techniques . the purified protein can be used for experiments or , in the case of insulin , administered to patients . uses of dna cloning dna molecules built through cloning techniques are used for many purposes in molecular biology . a short list of examples includes : biopharmaceuticals . dna cloning can be used to make human proteins with biomedical applications , such as the insulin mentioned above . other examples of recombinant proteins include human growth hormone , which is given to patients who are unable to synthesize the hormone , and tissue plasminogen activator ( tpa ) , which is used to treat strokes and prevent blood clots . recombinant proteins like these are often made in bacteria . gene therapy . in some genetic disorders , patients lack the functional form of a particular gene . gene therapy attempts to provide a normal copy of the gene to the cells of a patient ’ s body . for example , dna cloning was used to build plasmids containing a normal version of the gene that 's nonfunctional in cystic fibrosis . when the plasmids were delivered to the lungs of cystic fibrosis patients , lung function deteriorated less quickly $ ^2 $ . gene analysis . in basic research labs , biologists often use dna cloning to build artificial , recombinant versions of genes that help them understand how normal genes in an organism function . these are just a few examples of how dna cloning is used in biology today . dna cloning is a very common technique that is used in a huge variety of molecular biology applications .
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for instance , the human insulin gene is expressed in e. coli bacteria to make insulin used by diabetics . steps of dna cloning dna cloning is used for many purposes . as an example , let 's see how dna cloning can be used to synthesize a protein ( such as human insulin ) in bacteria .
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could you passably create another organism with the process of dna cloning ?
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what are these ritual objects ? the vajra ( tibetan : dorjie ) and bell ( sanskrit : ghanta ; tibetan : drilbu ) are the most important ritual objects of tibetan buddhism . most every lama has a pair and knows how to use them . they represent “ method ” ( vajra ) and “ wisdom ” ( bell ) . combined together they symbolize enlightenment as they embody the union of all dualities : bliss and emptiness , compassion and wisdom , appearance and reality , conventional truth and ultimate truth , and male and female , etc . what is meant by method and wisdom ? method indicates the compassionate activities of the bodhisattva that relieve living beings of their miseries . it is the skillful means that brings about the elimination of ignorance , greed , cruelty , etc . in living beings and causes them to follow the path to enlightenment . wisdom is the direct insight into ultimate reality ; it is the wisdom that realizes emptiness . by combining method and wisdom , the bodhisattva accumulates merit and insight and eventually attains buddhahood . what is the symbolism of the vajra and bell ? most vajras have five prongs that symbolize the five wisdoms that are attained through the transcendance of five kleshas ( greed , anger , delusion , pride and envy ) . the hub between them signifies emptiness . this one has eight prongs plus the central hub . vajra is a sanskrit word , in tibetan it is called a dorje . it is related to the word for diamond , and appears to be similar to the thunderbolt weapon carried by the vedic god indra , and the olympian zeus . as a thunderbolt weapon it destroys both internal and external enemies . as a diamond it symbolizes the indestructible and all-penetrating mind of enlightenment . the sound of the bell calls to mind the empty nature of all things . that is , according to the buddha , nothing whatsoever can exist independently , all phenomena are empty of true or inherent existence . by being profoundly aware of the empty nature of all things , we become free of attachment and aversion , and are liberated from the painful cycle of birth and death ( samsara ) . the bell is also a musical instrument its sound , together with other sacred instruments such as the hand-drum ( damaru ) , are played in rituals as musical offerings to the buddhas and other gods . how are they used ? the vajra and bell are often seen represented in the hands of deities in art , and in practice are held in the hands of the monks during rituals , the vajra in the right hand , the bell in the left . they are moved in prescribed movements . when the arms are crossed this symbolizes that the two are united—representing enlightenment . the sound of the bell is considered by tibetan buddhists as the most beautiful music . this music is presented as one of eight offerings to the deity that is invoked during the ritual . what are the eight offerings presented in rituals ? when tibetans buddhist begin meditation , they will invoke the presence of the deity , bow , and make offerings . for peaceful deities , the offerings are as follows : pure water for the deity to drink water for the deity to wash with scented oil for the deity to be anointed with flowers incense butter lamps food music , played on the ghanta ( bell ) and the damaru , a small two-faced drum with clappers attached by string , played by twisting back and forth in the hand this thunderbolt and bell were cast for the chinese emperor yongle ( 1403–1424 ) as a gift for a distinguished lama of tibet . the emperor possibly wished to gain merit for the commission . this and other gifts like it show the relationship between the tibetan lamas and the emperors of china . known as the priest-patron relationship , this was one way that ideas and artistic styles spread between china and tibet . artists working in china in imperial workshops were ordered to make tibetan style objects for either the personal use of the emperor or to send to important lamas in tibet , who were often considered to be their spiritual teachers .
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the bell is also a musical instrument its sound , together with other sacred instruments such as the hand-drum ( damaru ) , are played in rituals as musical offerings to the buddhas and other gods . how are they used ? the vajra and bell are often seen represented in the hands of deities in art , and in practice are held in the hands of the monks during rituals , the vajra in the right hand , the bell in the left .
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if vajras and bells are to be used for offerings and enlightenment , then how come great teachers like milarepa were enlightened without using these tools ?
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what are these ritual objects ? the vajra ( tibetan : dorjie ) and bell ( sanskrit : ghanta ; tibetan : drilbu ) are the most important ritual objects of tibetan buddhism . most every lama has a pair and knows how to use them . they represent “ method ” ( vajra ) and “ wisdom ” ( bell ) . combined together they symbolize enlightenment as they embody the union of all dualities : bliss and emptiness , compassion and wisdom , appearance and reality , conventional truth and ultimate truth , and male and female , etc . what is meant by method and wisdom ? method indicates the compassionate activities of the bodhisattva that relieve living beings of their miseries . it is the skillful means that brings about the elimination of ignorance , greed , cruelty , etc . in living beings and causes them to follow the path to enlightenment . wisdom is the direct insight into ultimate reality ; it is the wisdom that realizes emptiness . by combining method and wisdom , the bodhisattva accumulates merit and insight and eventually attains buddhahood . what is the symbolism of the vajra and bell ? most vajras have five prongs that symbolize the five wisdoms that are attained through the transcendance of five kleshas ( greed , anger , delusion , pride and envy ) . the hub between them signifies emptiness . this one has eight prongs plus the central hub . vajra is a sanskrit word , in tibetan it is called a dorje . it is related to the word for diamond , and appears to be similar to the thunderbolt weapon carried by the vedic god indra , and the olympian zeus . as a thunderbolt weapon it destroys both internal and external enemies . as a diamond it symbolizes the indestructible and all-penetrating mind of enlightenment . the sound of the bell calls to mind the empty nature of all things . that is , according to the buddha , nothing whatsoever can exist independently , all phenomena are empty of true or inherent existence . by being profoundly aware of the empty nature of all things , we become free of attachment and aversion , and are liberated from the painful cycle of birth and death ( samsara ) . the bell is also a musical instrument its sound , together with other sacred instruments such as the hand-drum ( damaru ) , are played in rituals as musical offerings to the buddhas and other gods . how are they used ? the vajra and bell are often seen represented in the hands of deities in art , and in practice are held in the hands of the monks during rituals , the vajra in the right hand , the bell in the left . they are moved in prescribed movements . when the arms are crossed this symbolizes that the two are united—representing enlightenment . the sound of the bell is considered by tibetan buddhists as the most beautiful music . this music is presented as one of eight offerings to the deity that is invoked during the ritual . what are the eight offerings presented in rituals ? when tibetans buddhist begin meditation , they will invoke the presence of the deity , bow , and make offerings . for peaceful deities , the offerings are as follows : pure water for the deity to drink water for the deity to wash with scented oil for the deity to be anointed with flowers incense butter lamps food music , played on the ghanta ( bell ) and the damaru , a small two-faced drum with clappers attached by string , played by twisting back and forth in the hand this thunderbolt and bell were cast for the chinese emperor yongle ( 1403–1424 ) as a gift for a distinguished lama of tibet . the emperor possibly wished to gain merit for the commission . this and other gifts like it show the relationship between the tibetan lamas and the emperors of china . known as the priest-patron relationship , this was one way that ideas and artistic styles spread between china and tibet . artists working in china in imperial workshops were ordered to make tibetan style objects for either the personal use of the emperor or to send to important lamas in tibet , who were often considered to be their spiritual teachers .
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what are these ritual objects ? the vajra ( tibetan : dorjie ) and bell ( sanskrit : ghanta ; tibetan : drilbu ) are the most important ritual objects of tibetan buddhism .
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when did these tools first appear ?
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overview during the classical period , states in mesoamerica and the andes were composed of a variety of kingdoms that traded and often came into conflict with one another . some states , such as teotihuacán near modern-day mexico city , held more power than others . in what is today peru and bolivia , andean states used the mountains , rivers , and coastline to their advantage when farming and creating a food supply for their societies . city of the gods northeast of mexico city , surrounded by lakes , the ruins of a once-massive city still inspire awe . a four-mile-long avenue runs through the remains of a complex grid pattern of apartments , colorful murals , a pyramid that ’ s over 700 feet tall and 700 feet wide , and another pyramid that once housed the remains of 200 people , possibly as tributes to accompany a wealthy leader into the afterlife . the city was named teotihuacán by the aztecs who discovered it after its collapse . the name translates roughly to “ city of the gods. ” unfortunately , no written records or art depicting specific rulers survive from teotihuacán itself . we do know , through other archeological methods , that the city center wielded enormous power between 300 and 600 ce . teotihuacán likely housed 200,000 people in the city itself , governed the surrounding 10,000 square miles directly , and used its armies to colonize other regions as far as 600 miles away . what made teotihuacán so powerful ? despite many unanswered questions about this city , we know that its growth and feats of urban planning wouldn ’ t have been possible without a powerful centralized government . the intricate urban grid and the massive temples must have required a coordinated effort and significant funds . we also know that there was trade between teotihuacán and other societies . for example , tools made of obsidian—a black , shiny material made of volcanic glass—were found in mayan territory . similarly , teotihuacán-style pottery has been found throughout mesoamerica . this serves as evidence of teotihuacán ’ s reach and expansion in the region . the city ’ s reach is also evident in documents from the mayan city of tikal recording the arrival of the teotihuacán military in 378 ce . these writings also suggest that teotihuacán became involved in the local government . this pattern echoes other examples in the early americas where leaders sought to build states in order to control resources and create cohesive societies . question : what do the urban planning and pyramids at teotihuacán suggest about its state power ? the maya : independent city-states the maya , a group of people who inhabited mesoamerica after the olmec , lived in what is today southern mexico , guatemala , honduras , belize , and el salvador . complex maya societies—including city-states—arose throughout these different areas , and local lords struggled with one another for power and access to trade routes and goods . surviving artwork from the time suggests that these rulers held considerable power in their own kingdoms and were possibly seen as divine . starting in the third century bce , mayan people settled in the fertile highlands of current day guatemala . one major city , kaminaljuyú , was located near what is now guatemala city ; it boasted impressive temples and access to trade routes into central mexico . in the fourth century ce , however , teotihuacán colonized the kaminaljuyú . the lack of a cohesive empire across mesoamerica was probably due in part to the large number of rulers jostling for power and difficult geography . mayan cities were located in diverse environments ranging from rainforest to highlands , which made governing over multiple cities difficult . what might teotihuacán ’ s colonization of kaminaljuyú signal to leaders of nearby city-states ? what inferences can we make about the decentralized maya city-states versus the city of teotihuacán ? why do you think the maya formed independent city-states rather than an empire ? the andes : innovations in agriculture early societies in the andes , in what is today peru and bolivia , faced geographical hurdles just as the maya did . the towering andes mountains and coastal deserts made unification difficult , but snow runoff from the mountains trickled into a network of rivers , facilitating agriculture . the moche civilization on the northern coast of peru developed a successful irrigation system and harvested maize , cotton , beans , and squash . farther east , the wari in the northern highlands used the mountains themselves as a means of tiered agriculture ; they irrigated the hillsides using the melted snow that flowed down from the mountains . evidence also shows the wari kingdom developed city planning and roadways that connected its major city to trade routes . andean kingdoms sought to control trade and imports such as seafood from the coast and potatoes and quinoa from the high plains . kingdoms situated their large cities at lower elevations in order to gain access to trade routes and imports more easily . question : what do the wari systems of agriculture and infrastructure tell us about the kingdom ? although they were not successful in creating expansive empires , early states in mesoamerica and the andes did build states with centralized governments ; areas with irrigation for crops ; roadways for travel and trade to bring wealth back to the capitals ; and armies to defend and enlarge their territories . we don ’ t know the full reach of power of cities like teotihuacán , but we can conclude that their success must have been the result of a centralized government and technology .
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overview during the classical period , states in mesoamerica and the andes were composed of a variety of kingdoms that traded and often came into conflict with one another . some states , such as teotihuacán near modern-day mexico city , held more power than others . in what is today peru and bolivia , andean states used the mountains , rivers , and coastline to their advantage when farming and creating a food supply for their societies .
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did any of the civilizations listed above reach some part of modern day argentina ?
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overview during the classical period , states in mesoamerica and the andes were composed of a variety of kingdoms that traded and often came into conflict with one another . some states , such as teotihuacán near modern-day mexico city , held more power than others . in what is today peru and bolivia , andean states used the mountains , rivers , and coastline to their advantage when farming and creating a food supply for their societies . city of the gods northeast of mexico city , surrounded by lakes , the ruins of a once-massive city still inspire awe . a four-mile-long avenue runs through the remains of a complex grid pattern of apartments , colorful murals , a pyramid that ’ s over 700 feet tall and 700 feet wide , and another pyramid that once housed the remains of 200 people , possibly as tributes to accompany a wealthy leader into the afterlife . the city was named teotihuacán by the aztecs who discovered it after its collapse . the name translates roughly to “ city of the gods. ” unfortunately , no written records or art depicting specific rulers survive from teotihuacán itself . we do know , through other archeological methods , that the city center wielded enormous power between 300 and 600 ce . teotihuacán likely housed 200,000 people in the city itself , governed the surrounding 10,000 square miles directly , and used its armies to colonize other regions as far as 600 miles away . what made teotihuacán so powerful ? despite many unanswered questions about this city , we know that its growth and feats of urban planning wouldn ’ t have been possible without a powerful centralized government . the intricate urban grid and the massive temples must have required a coordinated effort and significant funds . we also know that there was trade between teotihuacán and other societies . for example , tools made of obsidian—a black , shiny material made of volcanic glass—were found in mayan territory . similarly , teotihuacán-style pottery has been found throughout mesoamerica . this serves as evidence of teotihuacán ’ s reach and expansion in the region . the city ’ s reach is also evident in documents from the mayan city of tikal recording the arrival of the teotihuacán military in 378 ce . these writings also suggest that teotihuacán became involved in the local government . this pattern echoes other examples in the early americas where leaders sought to build states in order to control resources and create cohesive societies . question : what do the urban planning and pyramids at teotihuacán suggest about its state power ? the maya : independent city-states the maya , a group of people who inhabited mesoamerica after the olmec , lived in what is today southern mexico , guatemala , honduras , belize , and el salvador . complex maya societies—including city-states—arose throughout these different areas , and local lords struggled with one another for power and access to trade routes and goods . surviving artwork from the time suggests that these rulers held considerable power in their own kingdoms and were possibly seen as divine . starting in the third century bce , mayan people settled in the fertile highlands of current day guatemala . one major city , kaminaljuyú , was located near what is now guatemala city ; it boasted impressive temples and access to trade routes into central mexico . in the fourth century ce , however , teotihuacán colonized the kaminaljuyú . the lack of a cohesive empire across mesoamerica was probably due in part to the large number of rulers jostling for power and difficult geography . mayan cities were located in diverse environments ranging from rainforest to highlands , which made governing over multiple cities difficult . what might teotihuacán ’ s colonization of kaminaljuyú signal to leaders of nearby city-states ? what inferences can we make about the decentralized maya city-states versus the city of teotihuacán ? why do you think the maya formed independent city-states rather than an empire ? the andes : innovations in agriculture early societies in the andes , in what is today peru and bolivia , faced geographical hurdles just as the maya did . the towering andes mountains and coastal deserts made unification difficult , but snow runoff from the mountains trickled into a network of rivers , facilitating agriculture . the moche civilization on the northern coast of peru developed a successful irrigation system and harvested maize , cotton , beans , and squash . farther east , the wari in the northern highlands used the mountains themselves as a means of tiered agriculture ; they irrigated the hillsides using the melted snow that flowed down from the mountains . evidence also shows the wari kingdom developed city planning and roadways that connected its major city to trade routes . andean kingdoms sought to control trade and imports such as seafood from the coast and potatoes and quinoa from the high plains . kingdoms situated their large cities at lower elevations in order to gain access to trade routes and imports more easily . question : what do the wari systems of agriculture and infrastructure tell us about the kingdom ? although they were not successful in creating expansive empires , early states in mesoamerica and the andes did build states with centralized governments ; areas with irrigation for crops ; roadways for travel and trade to bring wealth back to the capitals ; and armies to defend and enlarge their territories . we don ’ t know the full reach of power of cities like teotihuacán , but we can conclude that their success must have been the result of a centralized government and technology .
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the lack of a cohesive empire across mesoamerica was probably due in part to the large number of rulers jostling for power and difficult geography . mayan cities were located in diverse environments ranging from rainforest to highlands , which made governing over multiple cities difficult . what might teotihuacán ’ s colonization of kaminaljuyú signal to leaders of nearby city-states ?
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so teotihuacan exercised influence over the mayan cities , so what social , economical , or religious affects did it have ?
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a serious departure when it appeared at the royal academy annual exhibition of 1850 christ in the house of his parents must have seemed a serious departure from standard religious imagery . painted by the young john everett millais , a member of the pre-raphaelite brotherhood ( p.r.b . ) , christ in the house of his parents focuses on the ideal of truth to nature that was to become the hallmark of the brotherhood . the picture centers on the young christ whose hand has been injured , being cared for by the virgin , his mother . christ ’ s wound , a perforation in his palm , foreshadows his ultimate end on the cross . a young st. john the baptist carefully brings a bowl of water to clean the wound , symbolic of christ washing the feet of his disciples . joseph , st anne ( the virgin ’ s mother ) and a carpenter ’ s assistant also react to christ ’ s accident . at a time when most religious paintings of the holy family were calm and tranquil groupings , this active event in the young life of the savior must have seemed extremely radical . the same can be said for millais ’ handling of the figures and the setting in the painting . mary ’ s wrinkled brow and the less than clean feet of some of the figures are certainly not idealized . according to the principles of the p.r.b. , the attention to detail is incredible . each individual wood shaving on the floor is exquisitely painted , and the rough-hewn table is a more functional than beautiful . the tools of the carpenters trade are evident hanging on the wall behind , while stacks of wood line the walls . the setting is a place of work , not a sacred spot . painted in a carpenter 's shop william michael rossetti recorded in the p.r.b . journal that millais started to work on the subject in november 1849 and began the actual painting at the end of december . we know from rossetti and the reminiscences of fellow brotherhood member william holman hunt that millais worked on location in a carpenter ’ s shop on oxford street , catching cold while working there in january . millais ’ son tells us that his father purchased sheep heads from a butcher to use as models for the sheep in the upper left of the canvas . he did not show the finished canvas to his friends until april of 1850 . scathing reviews although millais ’ exhibit at the royal academy in 1849 , isabella , had been well received , the critics blasted christ in the house of his parents . the most infamous review , however , was the one by charles dickens that appeared in his magazine household words in june 1850 . in it he described christ as a hideous , wry-necked , blubbering , red-haired boy in a nightgown , who appears to have received a poke playing in an adjacent gutter , and to be holding it up for the contemplation of a kneeling woman , so horrible in her ugliness that ( supposing it were possible for any human creature to exist for a moment with that dislocated throat ) she would stand out from the rest of the company as a monster in the vilest cabaret in france or in the lowest gin-shop in england . the commentary in the times was equally unfavorable , stating that millais ’ “ attempt to associate the holy family with the meanest details of a carpenter ’ s shop , with no conceivable omission of misery , of dirt , of even disease , all finished with loathsome minuteness , is disgusting. ” the painting proved to be so controversial that queen victoria asked that it be removed from the exhibition and brought to her so she could examine it . at the royal academy the attacks on millais ’ painting were undoubtedly unsettling for the young artist . millais had been born in 1829 on the island of jersey , but his parents eventually moved to london to benefit their son ’ s artistic education . when millais began at the royal academy school in 1840 he had the distinction of being the youngest person ever to have been admitted . at the royal academy , millais became friendly with the young william holman hunt , who is turn introduced millais to dante gabriel rossetti , and the idea for the pre-raphaelite brotherhood was born . the young artists exhibited their first set of paintings in 1849 , all of which were well received , but the paintings shown in 1850 were universally criticized , although none with as much fervor as christ in the house of his parents . millais ’ christ in the house of his parents is a remarkable religious painting for its time . it presents the holy family in a realistic manner , emphasizing the small details that bring the tableau to life . it is a scene we can easily imagine happening , but it is still laced with the symbolism expected of a christian subject . it is millais ’ marriage of these two ideas that makes christ in the house of his parents such a compelling image , and at the same time , made it so reprehensible to millais ’ contemporaries . essay by dr. rebecca jeffrey easby additional resources this painting at tate britain millais on the victorian web sir john everett millais in the google art project albert boime , `` sources for sit john everett millais 's christ in the house of his parents , '' gazette des beaux-arts , pp . 71-83 .
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scathing reviews although millais ’ exhibit at the royal academy in 1849 , isabella , had been well received , the critics blasted christ in the house of his parents . the most infamous review , however , was the one by charles dickens that appeared in his magazine household words in june 1850 . in it he described christ as a hideous , wry-necked , blubbering , red-haired boy in a nightgown , who appears to have received a poke playing in an adjacent gutter , and to be holding it up for the contemplation of a kneeling woman , so horrible in her ugliness that ( supposing it were possible for any human creature to exist for a moment with that dislocated throat ) she would stand out from the rest of the company as a monster in the vilest cabaret in france or in the lowest gin-shop in england .
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why would charles dickens have been so critical of this painting ?
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a serious departure when it appeared at the royal academy annual exhibition of 1850 christ in the house of his parents must have seemed a serious departure from standard religious imagery . painted by the young john everett millais , a member of the pre-raphaelite brotherhood ( p.r.b . ) , christ in the house of his parents focuses on the ideal of truth to nature that was to become the hallmark of the brotherhood . the picture centers on the young christ whose hand has been injured , being cared for by the virgin , his mother . christ ’ s wound , a perforation in his palm , foreshadows his ultimate end on the cross . a young st. john the baptist carefully brings a bowl of water to clean the wound , symbolic of christ washing the feet of his disciples . joseph , st anne ( the virgin ’ s mother ) and a carpenter ’ s assistant also react to christ ’ s accident . at a time when most religious paintings of the holy family were calm and tranquil groupings , this active event in the young life of the savior must have seemed extremely radical . the same can be said for millais ’ handling of the figures and the setting in the painting . mary ’ s wrinkled brow and the less than clean feet of some of the figures are certainly not idealized . according to the principles of the p.r.b. , the attention to detail is incredible . each individual wood shaving on the floor is exquisitely painted , and the rough-hewn table is a more functional than beautiful . the tools of the carpenters trade are evident hanging on the wall behind , while stacks of wood line the walls . the setting is a place of work , not a sacred spot . painted in a carpenter 's shop william michael rossetti recorded in the p.r.b . journal that millais started to work on the subject in november 1849 and began the actual painting at the end of december . we know from rossetti and the reminiscences of fellow brotherhood member william holman hunt that millais worked on location in a carpenter ’ s shop on oxford street , catching cold while working there in january . millais ’ son tells us that his father purchased sheep heads from a butcher to use as models for the sheep in the upper left of the canvas . he did not show the finished canvas to his friends until april of 1850 . scathing reviews although millais ’ exhibit at the royal academy in 1849 , isabella , had been well received , the critics blasted christ in the house of his parents . the most infamous review , however , was the one by charles dickens that appeared in his magazine household words in june 1850 . in it he described christ as a hideous , wry-necked , blubbering , red-haired boy in a nightgown , who appears to have received a poke playing in an adjacent gutter , and to be holding it up for the contemplation of a kneeling woman , so horrible in her ugliness that ( supposing it were possible for any human creature to exist for a moment with that dislocated throat ) she would stand out from the rest of the company as a monster in the vilest cabaret in france or in the lowest gin-shop in england . the commentary in the times was equally unfavorable , stating that millais ’ “ attempt to associate the holy family with the meanest details of a carpenter ’ s shop , with no conceivable omission of misery , of dirt , of even disease , all finished with loathsome minuteness , is disgusting. ” the painting proved to be so controversial that queen victoria asked that it be removed from the exhibition and brought to her so she could examine it . at the royal academy the attacks on millais ’ painting were undoubtedly unsettling for the young artist . millais had been born in 1829 on the island of jersey , but his parents eventually moved to london to benefit their son ’ s artistic education . when millais began at the royal academy school in 1840 he had the distinction of being the youngest person ever to have been admitted . at the royal academy , millais became friendly with the young william holman hunt , who is turn introduced millais to dante gabriel rossetti , and the idea for the pre-raphaelite brotherhood was born . the young artists exhibited their first set of paintings in 1849 , all of which were well received , but the paintings shown in 1850 were universally criticized , although none with as much fervor as christ in the house of his parents . millais ’ christ in the house of his parents is a remarkable religious painting for its time . it presents the holy family in a realistic manner , emphasizing the small details that bring the tableau to life . it is a scene we can easily imagine happening , but it is still laced with the symbolism expected of a christian subject . it is millais ’ marriage of these two ideas that makes christ in the house of his parents such a compelling image , and at the same time , made it so reprehensible to millais ’ contemporaries . essay by dr. rebecca jeffrey easby additional resources this painting at tate britain millais on the victorian web sir john everett millais in the google art project albert boime , `` sources for sit john everett millais 's christ in the house of his parents , '' gazette des beaux-arts , pp . 71-83 .
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in it he described christ as a hideous , wry-necked , blubbering , red-haired boy in a nightgown , who appears to have received a poke playing in an adjacent gutter , and to be holding it up for the contemplation of a kneeling woman , so horrible in her ugliness that ( supposing it were possible for any human creature to exist for a moment with that dislocated throat ) she would stand out from the rest of the company as a monster in the vilest cabaret in france or in the lowest gin-shop in england . the commentary in the times was equally unfavorable , stating that millais ’ “ attempt to associate the holy family with the meanest details of a carpenter ’ s shop , with no conceivable omission of misery , of dirt , of even disease , all finished with loathsome minuteness , is disgusting. ” the painting proved to be so controversial that queen victoria asked that it be removed from the exhibition and brought to her so she could examine it . at the royal academy the attacks on millais ’ painting were undoubtedly unsettling for the young artist .
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what did queen victoria think of the painting ?
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“ has anyone seen my wallet ? ” “ what time do i have to go to the dentist ? ” “ what ’ s the name of that movie we went to last week ? ” we all ask these sorts of questions from time-to-time no matter our age ; and not having a perfect memory is perfectly normal . it 's also normal that our longer term memory , as well as our ability to remember things that only occur occasionally , fade a little as we age . nevertheless , as we get into our sixties , many of us worry that forgetfulness or difficulty paying attention are signs of worse things to come . one of the most widespread fears as we age is dementia due to alzheimer ’ s disease . both dementia ( including alzheimer ’ s ) and delirium are both common causes of memory loss , impaired thinking and understanding ( cognition ) and impaired behaviour . they are distinct disorders but may be difficult to tell apart : dementia is a group of symptoms that mainly affects memory , cognition and social interactions , and the ability to do everyday tasks . symptoms start gradually often with no clear beginning , and are usually permanent . delirium , on the other hand , typically begins suddenly with a noticeable start point . it mainly affects attention , and often resolves after a few days or weeks , although it can last longer . structure , function and brain damage in dementia your brain is your body ’ s control centre . it is full of millions of interconnected brain cells , called neurons that receive , process , and send messages that control and coordinate almost everything that you do . your brain is divided into three main parts , the brain stem , cerebellum and cerebrum . these different regions are responsible for different functions , such as movement , speech , sense perception , emotions , heart rate , cognition and memory . information travels back and forth from the brain to the body along neurons , in the form of electrical impulses . when a nerve impulse reaches the end of a neuron a chemical neurotransmitter is released , which stimulates the electrical nerve impulse in the next neuron , allowing it to “ jump ” from one neuron to the next . neurons in different parts of the brain use different neurotransmitters to transfer information . some neurotransmitters stimulate brain activity , others calm the brain , and some do both . when your brain is functioning well , these neurotransmitters are all in balance . dementia may occur anytime neurons get damaged . the damage may be anywhere within the brain , and in more than one area at the same time . most dementias are caused by neurodegenerative diseases , most commonly alzheimer ’ s disease , lewy body dementia and frontotemporal dementia . these diseases cause clumps of abnormal proteins to build up inside neurons , damaging them , and causing them to slowly degenerate and die . not surprisingly , this disrupts the production of neurotransmitters and interrupts brain signals . in addition to these , vascular dementia is another common cause of progressive dementia . in this case , brain damage occurs when the blood supply to the neurons is reduced or blocked , again causing them to malfunction or die . blood vessel damage and blockage may result during a stroke , or may be caused by high blood pressure or other blood vessel disorders . there are also several other rare neurodegenerative diseases , and many other less common causes of dementia , such as infections , or dietary deficiencies . delirium is an acute , transient , and usually reversible brain malfunction . what exactly happens inside the brain of someone with delirium is not well understood ; although , it is thought to be brought on by multiple neurotransmitter imbalances , which alter your normal brain activity . symptoms of dementia and delirium dementia symptoms affect people differently depending on the area or areas of the brain that are affected , and may be progressive or reversible depending upon the cause of the damage . sometimes the symptoms are very similar to those of delirium , making diagnosis somewhat tricky , and it is fairly common for a person with dementia to develop delirium . dementia : symptoms usually start gradually , are fairly constant on a day-to-day basis , and slowly and steadily become worse over the course of about a decade | delirium : symptoms usually appear over a few hours to a few days , and may fluctuate on and off during the day , and often occurs at night . | - | - | - cognitive symptoms include : | cognitive symptoms include : | memory loss | memory loss | difficulty speaking and communicating | difficulty speaking and communicating | difficulty with complex tasks | rambling or nonsense speech | difficulty planning and organizing | difficulty reading and writing | disorientation | disorientation | loss of coordination | wandering attention | | becoming easily distracted | | becoming withdrawn | | | psychological symptoms include : | psychological symptoms include : | personality changes | inability to focus | inability to reason | inability to reason | inappropriate behaviour | reduced awareness of the environment | paranoia | agitation | agitation | hallucinations | hallucinations | disturbed sleep | fear , anxiety , anger or depression | | delirium is usually transient rather than permanent , although how well you recover often depends on how well you were beforehand — if you were in good health before it happened , you are more likely to have a full recovery . unfortunately , for some people , particularly those who are critically ill , delirium may lead to significant memory loss and decrease in thinking skills , as well as a general decline in health , poor recovery and increased risk of death . what causes dementia and delirium ? dementia four diseases account for most cases of dementia : alzheimer ’ s disease ( about 50-60 % of cases ) , vascular dementia ( about 15-20 % of cases ) , lewy body dementia , and frontotemporal dementia. $ ^1 $ each of these have different characteristics : alzheimer ’ s symptoms usually start in your mid to late 60s , although a small proportion of people get early onset alzheimer 's , with symptoms starting in their 40s or 50s . this form of alzheimer ’ s has a strong genetic link and typically runs in families . vascular dementia is rare if you are under 65 , but usually starts more suddenly than alzheimer ’ s . diagnosis may be complicated by the fact that you have alzheimer ’ s or another dementia at the same time . people at risk for vascular dementia often have a history of smoking , cardiac dysrhythmia , hypertension , diabetes , and coronary artery disease . lewy body dementia : lewy bodies are abnormal clumps of protein that build up in neurons causing brain signalling malfunctions ; they are often found in the brains of people with alzheimer ’ s and parkinson ’ s disease ( primarily a movement disorder ) . lewy body dementia is similar to alzheimer ’ s , but can be distinguished by unique symptoms such as rapid eye movement sleep behaviour disorder , and fluctuations between confusion and clarity . frontotemporal dementia often has an earlier onset than alzheimer ’ s , generally in your 50s or early 60s . the brain damage occurs at the front of your brain in regions that control personality , behaviour , and language . exactly why we get these dementias remains a puzzle , although there are likely genetic links in many cases . other rare conditions have also been linked to dementia including huntington ’ s disease , creutzfeldt-jakob disease , and parkinson ’ s disease , as well as traumatic brain injury . additionally , there are a variety of other reasons you may get dementia when the symptoms are often reversible . these include infections ( e.g. , meningitis , or syphilis ) , nutritional deficiencies , reactions to medications , brain bleeds , substance abuse , and poisoning . it is also possible that malfunction of other vital organs including the liver or kidneys can disrupt brain function causing dementia . as an older adult , any condition that ends up in a visit to the hospital increases your risk for delirium , and up to 80 % of people who are critically ill will be delirious at some time during their hospital stay. $ ^2 $ there are many different conditions that can trigger delirium with most common including dehydration , acute infections , such as urinary tract infection or pneumonia , prescribed drugs including antidepressants , sleeping pills and narcotics , and alcohol or substance abuse or withdrawal . how many people have dementia and delirium ? dementia : worldwide , there are around 50 million people living with dementia , and with 8 million new cases every year ; as the average lifespan is steadily increasing , dementia is increasing fast in many countries. $ ^3 $ overall , almost 1 in 10 people over the age of 60 have dementia , and the prevalence increases rapidly with age. $ ^1 $ women are at slightly more risk for it than men , and your chances increase if you are obese , have diabetes , high blood pressure , or other cardiovascular risk factors. $ ^1 $ delirium : delirium is very common in the hospital setting , especially when you are older and unwell . in fact , almost half of older patients are delirious when they are admitted , or develop delirium while they are there.2 like dementia , delirium increases significantly with age.4 in contrast to dementia , delirium is a little more common among men than women . are there ways to prevent dementia and delirium ? dementia : we don ’ t know for sure how to prevent dementia , but a healthy active lifestyle that includes keeping your mind and body active , and that keeps you socially connected may be of benefit . other steps you can take that may also reduce your risk are to quit smoking , lower your blood pressure , and eat a healthy diet that maintains a good balance of nutrients and vitamins , all of which may indirectly reduce your risk . delirium : the best way to prevent delirium is to remove the triggers . in a hospital setting this could include systems that maintain as calm and quiet an environment as possible , while providing the individualized support and medical care that is necessary , and when possible avoiding medications known to trigger delirium . what are the treatments for dementia and delirium ? dementia : memory loss and other symptoms of dementia may have many causes , so after reviewing your medical history , and symptoms , your doctor will likely want you to undergo other tests to evaluate your thinking and movement skills , and possibly your mental health . in addition to this , you may need a brain scan to check to see if you have had a stroke or brain bleed , or whether or not you have a brain tumour , as well as various laboratory tests to check for any metabolic and nutritional problems . most types of dementia are incurable , but there are medications that regulate brain chemicals involved in brain functions including memory , judgement , and learning that may be helpful with controlling or reducing your symptoms . two of the more commonly prescribed types of medicine are cholinesterase inhibitors and memantine , which is the first in a new class of drugs for the treatment of alzheimer ’ s . other medications are also available to treat any accompanying symptoms , and occupational therapists are experts at helping with strategies that improve coping with day-to-day challenges and quality of life . delirium : your doctor will likely follow a similar process to that outlined above for dementia , to decide whether or not you are delirious . if you are , your treatment will be aimed at eliminating the underlying cause or causes , which for example could include treating an infection , or stopping or changing a particular medicine that you are taking to treat something else . your doctor may also recommend additional supportive care , to ensure you stay hydrated , treat any pain you may have , keep you oriented with your surroundings , and help you get up and about again . consider the following : memantine is the first drug in a new class of drugs ( called n-methyl-d-aspartate receptor , or nmda receptor antagonists ) that has been approved to treat moderate to severe alzheimer ’ s disease . while cholinesterase inhibitors used to treat alzheimer ’ s help raise the levels of the neurotransmitter known as acetylcholine , memantine works by regulating a chemical called glutamate . glutamate is the salt of glutamic acid . it is an essential amino acid that is involved in many chemical processes that occur in living organisms , that lead to growth , production of energy , and elimination of waste . in the brain however , glutamate is an important neurotransmitter that ’ s involved in regulating almost everything your brain does including memory , cognition , and learning . for your brain to work well , glutamate has to be in the right place , at the right time , and in the right amount — too much or too little glutamate is harmful . in alzheimer ’ s disease , excess glutamate accumulates when neurons get damaged , which can be devastating for functioning neurons . memantine works by blocking the effects of too much glutamate , preventing cell death , and delaying some of the symptoms of moderate to severe alzheimer ’ s disease . because memantine has a different mechanism of action , it may be used in combination with cholinesterase inhibitors , which may improve outcomes for longer periods of time .
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nevertheless , as we get into our sixties , many of us worry that forgetfulness or difficulty paying attention are signs of worse things to come . one of the most widespread fears as we age is dementia due to alzheimer ’ s disease . both dementia ( including alzheimer ’ s ) and delirium are both common causes of memory loss , impaired thinking and understanding ( cognition ) and impaired behaviour .
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what causes alzheimer 's disease ?
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“ has anyone seen my wallet ? ” “ what time do i have to go to the dentist ? ” “ what ’ s the name of that movie we went to last week ? ” we all ask these sorts of questions from time-to-time no matter our age ; and not having a perfect memory is perfectly normal . it 's also normal that our longer term memory , as well as our ability to remember things that only occur occasionally , fade a little as we age . nevertheless , as we get into our sixties , many of us worry that forgetfulness or difficulty paying attention are signs of worse things to come . one of the most widespread fears as we age is dementia due to alzheimer ’ s disease . both dementia ( including alzheimer ’ s ) and delirium are both common causes of memory loss , impaired thinking and understanding ( cognition ) and impaired behaviour . they are distinct disorders but may be difficult to tell apart : dementia is a group of symptoms that mainly affects memory , cognition and social interactions , and the ability to do everyday tasks . symptoms start gradually often with no clear beginning , and are usually permanent . delirium , on the other hand , typically begins suddenly with a noticeable start point . it mainly affects attention , and often resolves after a few days or weeks , although it can last longer . structure , function and brain damage in dementia your brain is your body ’ s control centre . it is full of millions of interconnected brain cells , called neurons that receive , process , and send messages that control and coordinate almost everything that you do . your brain is divided into three main parts , the brain stem , cerebellum and cerebrum . these different regions are responsible for different functions , such as movement , speech , sense perception , emotions , heart rate , cognition and memory . information travels back and forth from the brain to the body along neurons , in the form of electrical impulses . when a nerve impulse reaches the end of a neuron a chemical neurotransmitter is released , which stimulates the electrical nerve impulse in the next neuron , allowing it to “ jump ” from one neuron to the next . neurons in different parts of the brain use different neurotransmitters to transfer information . some neurotransmitters stimulate brain activity , others calm the brain , and some do both . when your brain is functioning well , these neurotransmitters are all in balance . dementia may occur anytime neurons get damaged . the damage may be anywhere within the brain , and in more than one area at the same time . most dementias are caused by neurodegenerative diseases , most commonly alzheimer ’ s disease , lewy body dementia and frontotemporal dementia . these diseases cause clumps of abnormal proteins to build up inside neurons , damaging them , and causing them to slowly degenerate and die . not surprisingly , this disrupts the production of neurotransmitters and interrupts brain signals . in addition to these , vascular dementia is another common cause of progressive dementia . in this case , brain damage occurs when the blood supply to the neurons is reduced or blocked , again causing them to malfunction or die . blood vessel damage and blockage may result during a stroke , or may be caused by high blood pressure or other blood vessel disorders . there are also several other rare neurodegenerative diseases , and many other less common causes of dementia , such as infections , or dietary deficiencies . delirium is an acute , transient , and usually reversible brain malfunction . what exactly happens inside the brain of someone with delirium is not well understood ; although , it is thought to be brought on by multiple neurotransmitter imbalances , which alter your normal brain activity . symptoms of dementia and delirium dementia symptoms affect people differently depending on the area or areas of the brain that are affected , and may be progressive or reversible depending upon the cause of the damage . sometimes the symptoms are very similar to those of delirium , making diagnosis somewhat tricky , and it is fairly common for a person with dementia to develop delirium . dementia : symptoms usually start gradually , are fairly constant on a day-to-day basis , and slowly and steadily become worse over the course of about a decade | delirium : symptoms usually appear over a few hours to a few days , and may fluctuate on and off during the day , and often occurs at night . | - | - | - cognitive symptoms include : | cognitive symptoms include : | memory loss | memory loss | difficulty speaking and communicating | difficulty speaking and communicating | difficulty with complex tasks | rambling or nonsense speech | difficulty planning and organizing | difficulty reading and writing | disorientation | disorientation | loss of coordination | wandering attention | | becoming easily distracted | | becoming withdrawn | | | psychological symptoms include : | psychological symptoms include : | personality changes | inability to focus | inability to reason | inability to reason | inappropriate behaviour | reduced awareness of the environment | paranoia | agitation | agitation | hallucinations | hallucinations | disturbed sleep | fear , anxiety , anger or depression | | delirium is usually transient rather than permanent , although how well you recover often depends on how well you were beforehand — if you were in good health before it happened , you are more likely to have a full recovery . unfortunately , for some people , particularly those who are critically ill , delirium may lead to significant memory loss and decrease in thinking skills , as well as a general decline in health , poor recovery and increased risk of death . what causes dementia and delirium ? dementia four diseases account for most cases of dementia : alzheimer ’ s disease ( about 50-60 % of cases ) , vascular dementia ( about 15-20 % of cases ) , lewy body dementia , and frontotemporal dementia. $ ^1 $ each of these have different characteristics : alzheimer ’ s symptoms usually start in your mid to late 60s , although a small proportion of people get early onset alzheimer 's , with symptoms starting in their 40s or 50s . this form of alzheimer ’ s has a strong genetic link and typically runs in families . vascular dementia is rare if you are under 65 , but usually starts more suddenly than alzheimer ’ s . diagnosis may be complicated by the fact that you have alzheimer ’ s or another dementia at the same time . people at risk for vascular dementia often have a history of smoking , cardiac dysrhythmia , hypertension , diabetes , and coronary artery disease . lewy body dementia : lewy bodies are abnormal clumps of protein that build up in neurons causing brain signalling malfunctions ; they are often found in the brains of people with alzheimer ’ s and parkinson ’ s disease ( primarily a movement disorder ) . lewy body dementia is similar to alzheimer ’ s , but can be distinguished by unique symptoms such as rapid eye movement sleep behaviour disorder , and fluctuations between confusion and clarity . frontotemporal dementia often has an earlier onset than alzheimer ’ s , generally in your 50s or early 60s . the brain damage occurs at the front of your brain in regions that control personality , behaviour , and language . exactly why we get these dementias remains a puzzle , although there are likely genetic links in many cases . other rare conditions have also been linked to dementia including huntington ’ s disease , creutzfeldt-jakob disease , and parkinson ’ s disease , as well as traumatic brain injury . additionally , there are a variety of other reasons you may get dementia when the symptoms are often reversible . these include infections ( e.g. , meningitis , or syphilis ) , nutritional deficiencies , reactions to medications , brain bleeds , substance abuse , and poisoning . it is also possible that malfunction of other vital organs including the liver or kidneys can disrupt brain function causing dementia . as an older adult , any condition that ends up in a visit to the hospital increases your risk for delirium , and up to 80 % of people who are critically ill will be delirious at some time during their hospital stay. $ ^2 $ there are many different conditions that can trigger delirium with most common including dehydration , acute infections , such as urinary tract infection or pneumonia , prescribed drugs including antidepressants , sleeping pills and narcotics , and alcohol or substance abuse or withdrawal . how many people have dementia and delirium ? dementia : worldwide , there are around 50 million people living with dementia , and with 8 million new cases every year ; as the average lifespan is steadily increasing , dementia is increasing fast in many countries. $ ^3 $ overall , almost 1 in 10 people over the age of 60 have dementia , and the prevalence increases rapidly with age. $ ^1 $ women are at slightly more risk for it than men , and your chances increase if you are obese , have diabetes , high blood pressure , or other cardiovascular risk factors. $ ^1 $ delirium : delirium is very common in the hospital setting , especially when you are older and unwell . in fact , almost half of older patients are delirious when they are admitted , or develop delirium while they are there.2 like dementia , delirium increases significantly with age.4 in contrast to dementia , delirium is a little more common among men than women . are there ways to prevent dementia and delirium ? dementia : we don ’ t know for sure how to prevent dementia , but a healthy active lifestyle that includes keeping your mind and body active , and that keeps you socially connected may be of benefit . other steps you can take that may also reduce your risk are to quit smoking , lower your blood pressure , and eat a healthy diet that maintains a good balance of nutrients and vitamins , all of which may indirectly reduce your risk . delirium : the best way to prevent delirium is to remove the triggers . in a hospital setting this could include systems that maintain as calm and quiet an environment as possible , while providing the individualized support and medical care that is necessary , and when possible avoiding medications known to trigger delirium . what are the treatments for dementia and delirium ? dementia : memory loss and other symptoms of dementia may have many causes , so after reviewing your medical history , and symptoms , your doctor will likely want you to undergo other tests to evaluate your thinking and movement skills , and possibly your mental health . in addition to this , you may need a brain scan to check to see if you have had a stroke or brain bleed , or whether or not you have a brain tumour , as well as various laboratory tests to check for any metabolic and nutritional problems . most types of dementia are incurable , but there are medications that regulate brain chemicals involved in brain functions including memory , judgement , and learning that may be helpful with controlling or reducing your symptoms . two of the more commonly prescribed types of medicine are cholinesterase inhibitors and memantine , which is the first in a new class of drugs for the treatment of alzheimer ’ s . other medications are also available to treat any accompanying symptoms , and occupational therapists are experts at helping with strategies that improve coping with day-to-day challenges and quality of life . delirium : your doctor will likely follow a similar process to that outlined above for dementia , to decide whether or not you are delirious . if you are , your treatment will be aimed at eliminating the underlying cause or causes , which for example could include treating an infection , or stopping or changing a particular medicine that you are taking to treat something else . your doctor may also recommend additional supportive care , to ensure you stay hydrated , treat any pain you may have , keep you oriented with your surroundings , and help you get up and about again . consider the following : memantine is the first drug in a new class of drugs ( called n-methyl-d-aspartate receptor , or nmda receptor antagonists ) that has been approved to treat moderate to severe alzheimer ’ s disease . while cholinesterase inhibitors used to treat alzheimer ’ s help raise the levels of the neurotransmitter known as acetylcholine , memantine works by regulating a chemical called glutamate . glutamate is the salt of glutamic acid . it is an essential amino acid that is involved in many chemical processes that occur in living organisms , that lead to growth , production of energy , and elimination of waste . in the brain however , glutamate is an important neurotransmitter that ’ s involved in regulating almost everything your brain does including memory , cognition , and learning . for your brain to work well , glutamate has to be in the right place , at the right time , and in the right amount — too much or too little glutamate is harmful . in alzheimer ’ s disease , excess glutamate accumulates when neurons get damaged , which can be devastating for functioning neurons . memantine works by blocking the effects of too much glutamate , preventing cell death , and delaying some of the symptoms of moderate to severe alzheimer ’ s disease . because memantine has a different mechanism of action , it may be used in combination with cholinesterase inhibitors , which may improve outcomes for longer periods of time .
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the damage may be anywhere within the brain , and in more than one area at the same time . most dementias are caused by neurodegenerative diseases , most commonly alzheimer ’ s disease , lewy body dementia and frontotemporal dementia . these diseases cause clumps of abnormal proteins to build up inside neurons , damaging them , and causing them to slowly degenerate and die .
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what do you do when your grandparents has dementia or alzheimer 's disease but you or your family do n't know that they have dementia or alzheimer 's disease ?
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conceived by performance and conceptual artist vito acconci , following piece was an activity that took place everyday on the streets of new york , between october 3rd and 25th , 1969 . it was part of other performance and conceptual events sponsored by the architectural league of new york that occurred during those three weeks . the terms of the exhibition “ street works iv ” were to do a piece , sometime during the month , that used a street in new york city . so acconci decided to follow people around the streets and document his following of them . but why would he do this ? why would acconci follow random people around new york ? acconci ’ s work is typical of performance and conceptual art made during this period in the way that he uses his body as the object of his art in order to explore some specific idea . in essence , following piece was concerned with the language of our bodies , not so much in a private manner , but in a deeply public manner . by selecting a passer-by at random until they entered a private space , acconci submitted his own movements to the movements of others , showing how our bodies are themselves always subject to external forces that we may or may not be able to control . in his notes that the artist kept during the performance , acconci wrote : following piece , potentially , could use all the time allotted and all the space available : i might be following people , all day long , everyday , through all the streets in new york city . in actuality , following episodes ranged from two or three minutes when someone got into a car and i couldn ’ t grab a taxi , i couldn ’ t follow – to seven or eight hours – when a person went to a restaurant , a movie . in terms of the art work , rather than being just another object that we look at in the gallery , following piece was part of the revolution that took place in the art world in the late 1960s that tried to bring art out of the gallery and into the street in order to explore real issues such as space , time , and the human body . many artists , such as acconci , used their bodies as their chosen medium . look at some of acconci ’ s notes of the period which he wrote before , during and after the event : • i need a scheme ( follow the scheme , follow a person ) • i add myself to another person ( i give up control/i don ’ t have to control myself ) • subjective relationship ; subjunctive relationship • a way to get around . ( a way to get myself out of the house . ) get into the middle of things . • out of space . out of time . ( my time and space are taken up , out of myself , into a larger system ) . all of these ideas were influenced by acconci ’ s readings . as many other artists of the period , acconci wanted to get away from specific art problems and engage with social problems . acconci read books such as edward hall ’ s the hidden dimension ( 1969 ) , erving goffmann ’ s the presentation of the self in every day life ( 1959 ) , and kurt lewin ’ s in principles of topological psychological ( 1936/1966 ) . all of these books explored the ways in which the individual and the social are interlinked in terms of complex codes that structure the way we act and live everyday . with regard to the influence on these texts on following piece , acconci ’ s use of diagrams specifically refers to lewin ’ s notion of “ field theory ” : that is , a model that sought to explain human behaviour in terms of relations and in relation to its environment and surroundings . lewin placed behaviour in a “ field ” in order to examine it in a theoretical manner . the diagrams drawn by acconci are an imaginative engagement with this idea of human relations as engaging in a specific field or space . so by following someone around new york , acconci could perhaps experience what it was like to relinquish self control to others and also explore the intersecting systems that grouped different people together in one field . as we can see from the diagrams , acconci ’ s intentions were not subjective but much more systematic—they constituted an exploration of the private and public fields that occur in every social space . ironically , for all the effort to get out of the gallery , much of acconci ’ s documentation of following piece , for example , the texts , photographs ( which were taken after the event ! ) , and diagrams , now constitutes a work of art in its own right . moma owns several of the photographs of following piece and other “ versions ” of this work are also in existence . so , even though acconci ’ s following piece was a performance that occurred in a very specific period ( 3rd to 25th october , 1969 ) , the reproduction and circulation of the work continues . this fact not only teaches us important things about the nature of performance art and its relationship to the art world , but also how the context of the art work is also never exactly fixed and each time it is presented something new occurs with the work itself . essay by jp mcmahon additional resources : following piece photographs at the metropolitan museum of art willoughby sharp video interview with vito acconci ( 1973 ) includes discussion of following piece following piece documentation
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but why would he do this ? why would acconci follow random people around new york ? acconci ’ s work is typical of performance and conceptual art made during this period in the way that he uses his body as the object of his art in order to explore some specific idea . in essence , following piece was concerned with the language of our bodies , not so much in a private manner , but in a deeply public manner .
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perhaps acconci 's work is a good example of the discipline of sociology poeticly entering the art scene in the 1960s ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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i am confused , which force is stronger , hydrogen bonds or dipole-dipole attraction ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules .
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why ca n't we say that h2s also has hydrogen bond along with london dispersion bond and dipole-dipole attraction ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules .
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the article said dipole-dipole interactions and hydrogen bonding are equally strong and hydrogen bonding is a type of dipole-dipole interaction , so how come covalent compounds containing hydrogen bonds have higher boiling and melting points than polar covalent compounds ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine .
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difference between inter and intramolecular bonds ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ .
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is there hydrogen bonding in hcl ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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how would you identify the intermolecular and intramolecular forces in h2s ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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would the inter-molecular force be dipole-dipole , because the dipoles do n't cancel out ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds .
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and would the intramolecular force be polar covalent because the difference in electronegativity is greater than 0 but less then 1.7 ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules .
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then what are dipole-induced dipole forces , ion-dipole forces , and ion-induced dipole forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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can someone please tell me how london dispersion force exists within h2s bond ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion .
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how can we have partially charged ions if electric charge is quantized ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable .
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why is ionic bond stronger than covalent bond ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules .
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do molecules that experience hydrogen bonding also have dipole-dipole forces and dispersion forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction .
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why loose bonds are stronger than simple dipole dipole bonds ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine .
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can substances have more than one type of intermolecular interaction ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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between which atoms exactly does london dispersion force exist ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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why it takes more energy to break an intramuscular force than it does an intermolecular force ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so which bond , intermolecular or intramolecular , is more important when determining the melting and boiling points for covalent compounds ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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can covalent , ionic , and hydrogen bonds be intermolecular or intramolecular ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons .
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in the example of ch3oh , why was c positively charged ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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how do i know if a compound is an ion that can make it an ion ion imf or an ion dipole imf ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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what is the major distinguishing factor between ionic/covalent bonding and intermolecular forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules .
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are carbon-hydrodgen bonds intramolecular forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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do metals have intermolecular forces or are they different ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable .
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break down of covalent bond of h2 require ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine .
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why is the example mislabeled ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization .
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so if the question ask for the strength of bond of a molecule , do we answer between intra or inter ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization .
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if the question ask to compare the strength of bond between inter and inter what is the best answer ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later .
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i know the greek letter with the plus or minus sign represents the partial charge of each individual atom in a molecule , but i can not remember the terminology for each charged `` end '' of a polar molecule- is that the dipole or is it called something else ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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how would you draw dispersion forces between methanol molecules ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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is induced dipole the same as london dispersion ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions .
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is intermolecular bonding only for non-metals ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 .
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why is n't h2s considered to have a polar covalent bond ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ .
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the properties we observe in substances at the bulk scale , such as melting and boiling points , or surface tension , are a direct result of what ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons .
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how will molecules composed of different atoms will be bonded ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion .
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what type of bond does ethylene glycol form ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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why do all molecules exert london forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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does the intramolecular force have any effect against intermolecular force ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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the strongest interactions between molecules ionic bonds / covalent bonds / hydrogen bonds / dipole-dipole interactions / dispersion forces ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures .
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in the london dispersion forces paragraph , does bromine has a higher melting point than chlorine ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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why does intermolecular forces are identical to bonding between atoms in a single molecule ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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how does intermolecular forces of attraction applied in construction nails ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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which one is stronger from ion dipole force and dipole dipole force ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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at what point do the london dispersion forces become stronger than the dipole-dipole intermolecular bonds ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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why do molecules with weak intermolecular forces have low freezing points ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules .
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would n't that mean that the molecule is liquid at relatively low temperatures , because its freezing point is very low ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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so if it only freezes at very low temperatures , does n't that mean that its intermolecular forces are strong ; strong so that the forces do n't break and cause the molecule to freeze ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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in the first example/practice question , what is the difference between a dipole dipole and a london dispersion ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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i understand ionic forces in intramolecular bonds where there is the donating of an electron from one element to another forming an anion and a cation but how would ionic intermolecular forces work ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures .
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or , does it mean if you put carbonate ( co3 2- ) with say ammonium ( nh4+ ) ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization .
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is there any specific type of intramolecular bond between halogens , h , n , and o or is it just non-polar covalent ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons .
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what is the difference between ionic substances , charged particles and polar molecules ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms .
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how to intermolecular forces affect earth and how would life be without them ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions .
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what is the theory of intermolecular bonding ?
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there are two kinds of forces , or attractions , that operate in a molecule—intramolecular and intermolecular . let 's try to understand this difference through the following example . we have six towels—three are purple in color , labeled hydrogen and three are pink in color , labeled chlorine . we are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel . after sewing , we now have three pairs of towels : hydrogen sewed to chlorine . the next step is to attach these three pairs of towels to each other . for this we use velcro as shown above . so , the result of this exercise is that we have six towels attached to each other through thread and velcro . now if i ask you to pull this assembly from both ends , what do you think will happen ? the velcro junctions will fall apart while the sewed junctions will stay as is . the attachment created by velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together . a slight force applied to either end of the towels can easily bring apart the velcro junctions without tearing apart the sewed junctions . exactly the same situation exists in molecules . just imagine the towels to be real atoms , such as hydrogen and chlorine . these two atoms are bound to each other through a polar covalent bond—analogous to the thread . each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to velcro . we ’ ll talk about dipole-dipole interactions in detail a bit later . the polar covalent bond is much stronger in strength than the dipole-dipole interaction . the former is termed an intramolecular attraction while the latter is termed an intermolecular attraction . so now we can define the two forces : intramolecular forces are the forces that hold atoms together within a molecule . intermolecular forces are forces that exist between molecules . types of intramolecular forces of attraction ionic bond : this bond is formed by the complete transfer of valence electron ( s ) between atoms . it is a type of chemical bond that generates two oppositely charged ions . in ionic bonds , the metal loses electrons to become a positively charged cation , whereas the nonmetal accepts those electrons to become a negatively charged anion . covalent bond : this bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons . because both atoms have similar affinity for electrons and neither has a tendency to donate them , they share electrons in order to achieve octet configuration and become more stable . a nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5 . a polar covalent bond is formed when atoms of slightly different electronegativities share electrons . the difference in electronegativity between bonded atoms is between 0.5 and 1.9 . hydrogen chloride , $ \text { hcl } $ ; the $ \text { o } - { h } $ bonds in water , $ \text { h } _ { 2 } \text { o } $ ; and hydrogen fluoride , $ \text { hf } $ , are all examples of polar covalent bonds . metallic bonding : this type of covalent bonding specifically occurs between atoms of metals , in which the valence electrons are free to move through the lattice . this bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions . metallic bonds are present in samples of pure elemental metals , such as gold or aluminum , or alloys , like brass or bronze . the freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity . relative strength of the intramolecular forces intramolecular force | basis of formation | relative strength : - : | : - : | : - : metallic bond | metal cations to delocalized electrons | 1 , strongest ionic bond | cations to anions | 2 polar covalent bond | partially charged cation to partially charged anion | 3 nonpolar covalent bond | nuclei to shared electrons | 4 , weakest intermolecular forces of attraction now let ’ s talk about the intermolecular forces that exist between molecules . intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point , melting point , density , and enthalpies of fusion and vaporization . types of intermolecular forces that exist between molecules dipole-dipole interactions : these forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule . the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom . the partially positive end of hydrogen is attracted to the partially negative end of the oxygen , nitrogen , or fluorine of another molecule . hydrogen bonding is a relatively strong force of attraction between molecules , and considerable energy is required to break hydrogen bonds . this explains the exceptionally high boiling points and melting points of compounds like water , $ \text { h } _ { 2 } \text { o } $ , and hydrogen fluoride , $ \text { hf } $ . hydrogen bonding plays an important role in biology ; for example , hydrogen bonds are responsible for holding nucleotide bases together in $ \text { dna } $ and $ \text { rna } $ . london dispersion forces , under the category of van der waal forces : these are the weakest of the intermolecular forces and exist between all types of molecules , whether ionic or covalent—polar or nonpolar . the more electrons a molecule has , the stronger the london dispersion forces are . for example , bromine , $ \text { br } { 2 } $ , has more electrons than chlorine , $ \text { cl } { 2 } $ , so bromine will have stronger london dispersion forces than chlorine , resulting in a higher boiling point for bromine , 59 $ ^\text { o } $ c , compared to chlorine , –35 $ ^\text { o } $ c . also , the breaking of london dispersion forces doesn ’ t require that much energy , which explains why nonpolar covalent compounds like methane— $ \text { ch } _ { 4 } $ —oxygen , and nitrogen—which only have london dispersion forces of attraction between the molecules—freeze at very low temperatures . relative strength of intermolecular forces of attraction intermolecular force | occurs between … | relative strength : - : | : - : | : - : dipole-dipole attraction | partially oppositely charged ions | strongest hydrogen bonding | $ \text { h } $ atom and $ \text { o } $ , $ \text { n } $ / or $ \text { f } $ atom | as strong as dipole-dipole attraction london dispersion attraction | temporary or induced dipoles | weakest how forces of attraction affect properties of compounds polar covalent compounds—like hydrogen chloride , $ \text { hcl } $ , and hydrogen iodide , $ \text { hi } $ —have dipole-dipole interactions between partially charged ions and london dispersion forces between molecules . nonpolar covalent compounds—like methane $ \text { ch } { 4 } $ and nitrogen gas , $ \text { n } { 2 } $ ) —only have london dispersion forces between molecules . the rule of thumb is that the stronger the intermolecular forces of attraction , the more energy is required to break those forces . this translates into ionic and polar covalent compounds having higher boiling and melting points , higher enthalpy of fusion , and higher vaporization than covalent compounds . boiling and melting points of compounds depend on the type and strength of the intermolecular forces present , as tabulated below : type of compound | intermolecular forces present | relative order of boiling and melting points -|-|- ionic compounds | ion to ion attraction between ions , london dispersion forces | 1 , highest ) covalent compounds containing hydrogen bonds | hydrogen bonds , london dispersion forces | 2 polar covalent compounds | dipole-dipole attraction between dipoles created by partially charged ions , london dispersion forces | 3 nonpolar covalent compounds | london dispersion forces | 4 , lowest let ’ s try to identify the different kinds of intermolecular forces present in some molecules . $ \text { h } _ { 2 } \text { s } $ —london dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction $ \text { ch } _ { 3 } \text { oh } $ —london dispersion force , dipole-dipole attraction , and hydrogen bonding $ \text { c } { 2 } \text { h } { 6 } $ —london dispersion forces—it ’ s a nonpolar covalent compound— and no other intermolecular attractions
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the prerequisite for this type of attraction to exist is partially charged ions—for example , the case of polar covalent bonds such as hydrogen chloride , $ \text { hcl } $ . dipole-dipole interactions are the strongest intermolecular force of attraction . hydrogen bonding : this is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen , nitrogen , or fluorine atom .
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does ionic bonding count as an intermolecular force ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude .
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why are n't these paintings new when they are called modern ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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does the term avant-garde only apply to modern art or can it also be applied to current art of this decade ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists .
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when did the avant-garde artists begin to appear ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes .
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were woman aloud to paint in the third centery ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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is modern art from the realist movement generally symbolic or is it pretty straight-forward ?
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people use the term “ modern ” in a variety of ways , often very loosely , with a lot of implied associations of new , contemporary , up-to-date , and technological . we know the difference between a modern country and a third world country and it usually has less to do with art and more to do with technology and industrial progress , things like indoor plumbing , easy access to consumer goods , freedom of expression , and voting rights . in the 19th century , however , modernity and its connection with art had certain specific associations that people began recognizing and using as barometers to distinguish themselves and their culture from earlier nineteenth century ways and attitudes . chronologically , modernism refers to the period from 1850 to 1960 . it begins with the realist movement and ends with abstract expressionism . that ’ s just a little over one hundred years . during that period the western world experienced some significant changes that transformed europe and the united states from traditional societies that were agriculturally based into modern ones with cities and factories and mass transportation . here are some important features that all modern societies share . capitalism capitalism replaced landed fortunes and became the economic system of modernity in which people exchanged labor for a fixed wage and used their wages to buy ever more consumer items rather than produce such items themselves . this economic change dramatically affected class relations because it offered opportunities for great wealth through individual initiative , industrialization and technology—somewhat like the technological and dot.com explosion of the late 20th and early 21st century . the industrial revolution which began in england in the late 18th century and rapidly swept across europe ( hit the u.s. immediately following the civil war ) transformed economic and social relationships , offered an ever increasing number of cheaper consumer goods , and changed notions of education . who needed the classics when a commercial/technically oriented education was the key to financial success ? the industrial revolution also fostered a sense of competition and progress that continues to influence us today . urban culture urban culture replaced agrarian culture as industrialization and cities grew . cities were the sites of new wealth and opportunity with their factories and manufacturing potential . people moving from small farms , towns to large cities helped to breakdown traditional culture and values . there were also new complications such as growing urban crime , prostitution , alienation , and depersonalization . in a small town you probably knew the cobbler who made your shoes and such a personal relationship often expanded into everyday economics—you might be able to barter food or labor for a new pair of shoes or delay payments . these kinds of accommodations that formed a substructure to agrarian life were swept away with urbanization . city dwellers bought shoes that were manufactured , transported by railroads , displayed in shop windows , and purchased only for cash . assembly lines , anonymous labor , and advertising created more consumer items but also a growing sense of depersonalization . the gap between the “ haves ” and the “ have nots ” increased and were more visible in the city . technology technological advances such as industrialization , railroads , gas lighting , streetcars , factory systems , indoor plumbing , appliances , and scientific advances were rapidly made and these changes dramatically affected the way people lived and thought about themselves . one consequence was that people in industrialized areas thought of themselves as progressive and modern and considered undeveloped cultures in undeveloped countries as primitive and backward . secularism modernity is characterized by increasing secularismand diminished religious authority . people did not abandon religion but they paid less attention to it . organized religions were increasingly less able to dictate standards , values , and subject matter . fine art moved from representing human experience and its relationship to god 's creation , to a focus on personal emotions and individual spiritual experiences that were not based in any organized and institutionalized religion . optimism the modern world was extremely optimistic—people saw these changes as positive . they welcomed innovation and championed progress . change became a signifier of modernity . anything that was traditional and static signaled outmoded , old-fashioned , conservative and was to be avoided by the new modern public . modern europe and the u.s. internalized these positions and used modernity as a way of determining and validating their superiority . the nineteenth century was also a period of tremendous colonial growth and expansion , in the name of progress and social benefit and all of these activities were spearheaded by newly industrialized western countries . many artists closely identified with modernity and embraced the new techniques and innovations , the spirit of progress , invention , discovery , creativity and change . they wanted to participate in creating the modern world and they were anxious to try out new ideas rather than following the more conservative guidelines of academic art . this is not to say that these mid-nineteenth century artists were the first to challenge an older generation or set of ideas . many academic artists had argued over formal issues , styles and subject matter but this was much like a good natured agreement within a club ; everyone in the group agreed to disagree . a middle-class audience by the mid-1850 ’ s polite academic disagreements were being taken out of the academy and onto the street . artists were looking increasingly to the private sector for patronage , tapping into that growing group of bourgeois or middle class collectors with money to spend and houses to fill with paintings . this new middle class audience that made its money through industrialization and manufacturing had lots of “ disposable income ” , and they wanted pictures that they could understand , that were easy to look at , fit into their homes , addressed subjects they liked . not for them the historical cycles of gods , saints and heroes with their complex intellectual associations and references ; instead , they wanted landscapes , genre scenes , and still life . they were not less educated than earlier buyers , but educated with a different focus and set of priorities . reality was here and now , progress was inevitable , and the new hero of modern life was the modern man . modernity is then a composite of contexts : a time , a space , and an attitude . what makes a place or an object “ modern ” depends on these conditions . the avant-garde throughout the 19th century there were artists who produced pictures that we do not label “ modern art ” generally because the techniques or subjects were associated with the conservative academic styles , techniques and approaches . on the other hand , modern artists were often called the “ avant garde. ” this was originally a military term that described the point man ( the first soldier out ) —the one to take the most risk . the french socialist henri de saint-simon first used the term in the early 1820 ’ s to describe an artist whose work would serve the needs of the people , of a socialist society rather than the ruling classes . the avant garde is also used to identify artists whose painting subjects and techniques were radical , marking them off from the more traditional or academic styles , but not with any particular political ideology in mind . avant garde became a kind of generic term for a number of art movements centered on the idea of artistic autonomy and independence . in some cases the avant garde was closely associated with political activism , especially socialist or communist movements ; in other cases , the avant garde was pointedly removed from politics and focused primarily on aesthetics . the avant garde was never a cohesive group of artists and what was avant garde in one nation was not necessarily the same in others . finally , although modern artists were working throughout many countries in europe and the united states , most 19th art and much 20th century modern art is centered in france and produced by french artists . unlike england which was politically stable in the 19th century , france went through a variety of governments and insurrections all of which provided a unique political and cultural environment that fostered what we know as modern art . essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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essay by dr. parme giuntini additional resources : modernism from christopher l.c.e . witcombe impressionism : art and modernity on the metropolitan museum of art 's heilbrunn timeline of art history modern art from moma learning
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did religion get in the way of art at this time ?
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introduction 21st-century art is a burgeoning field of practice , research , and publication , making it an incredibly dynamic field of study . many important topics have been resonating in the new century and inspiring new thinking and scholarly debate , such as the surge of bio art in response to scientific research in the life sciences , and the critical theory known as relational aesthetics that developed in response to an increase in art that invites viewers ’ participation and interaction . other topics that were much-discussed in the late 20th century remain vital for the analysis of 21st-century art and visual culture , including semiotics , post-modernism , and feminism . art of the 21st century emerges from a vast variety of materials and means . these include the latest electronic technologies , such as digital imaging and the internet ; familiar genres with a long history that continue to be practiced with great vigor , such as painting ( see , for example , the work of julie mehretu and shahzia sikander ) ; and materials and processes once associated primarily with handicrafts , re-envisioned to express new concepts . many artists regularly and freely mix media and forms , making the choices that best serve their concepts and purposes . activities vary from spectacular projects accomplished with huge budgets and extraordinary production values to modest endeavors that emphasize process , ephemeral experiences , and a do-it-yourself approach . the notion of influences has also shifted with changes in communications and technology ; every location around the world has artists who respond to local geographies and histories as well as the sway of global visual culture . globalization a key feature of the art scene in the 21st century ( and of many sectors of 21st-century life ) is the impact of globalization – the accelerating interconnectivity of human activity and information across time and space . aided by the internet and mass media , awareness of the vitality of contemporary art in localities around the globe has grown exponentially . anyone with access to the internet can follow developments in shanghai , sydney , são paulo , or nairobi . simultaneously the increased movement of artists across borders and oceans has added to the intermixing of influences and artistic vocabularies . for example , wangechi mutu , originally from kenya , pursued further education in south wales and then in the united states . her collaged images of women are informed by african tribal arts , 20th-century european and american collage artists , and the latest illustrations from fashion , pornography , and medical sources . the meaning and consequences of globalization are much debated by scholars . economically and politically , is globalization a force for growth and freedom in societies everywhere , or does it contribute to further exploitation of developing regions by the wealthy ? does globalization work in different ways in different localities ? regarding globalization and art , do practices in asia , africa , the middle east , and elsewhere challenge the traditional assumptions and value judgments that are the basis of the western canon ? are western institutions rethinking that canon or simply adding art from other places to their rosters in a token and uncritical gesture of inclusivity ? how do curated exhibitions that address themes of globalization represent artists from various parts of the world ? the expanding art market and the proliferation of biennials and art fairs helped a select group of artists from every continent to gain an international presence ; but have the underlying structure and values of the art market changed otherwise ? visual culture in the 21st century visual culture has grown as a recognized interdisciplinary field of study , taking a multi-faceted approach to understanding how images of all types communicate and participate in the construction of identity , gender , class , power relationships , and other social and political meanings and values . medicine , science , politics , consumer culture , and religion and spirituality are some of the arenas that visual culture studies examine along with art . visual culture scholars analyze film , television , graphic novels , fashion design , and other forms of popular culture in addition to established fine art media such as painting , and they draw upon many methodologies and theories , including semiotics , sociology , psychoanalysis , reception theory , feminism , and the concept of the gaze , to name a few . just as visual culture scholars are examining images and media of all types so , too , are 21st-century artists drawing inspiration , imagery , materials , and concepts from diverse areas of culture , moving well beyond influences from the history of fine art and design . the world of professional sports and fanatic fans has been a topic for paul pfeiffer , while the commercial television industry has informed various video installations by christian jankowski . most contemporary artists do not draw rigid distinctions between high art and popular culture . for instance , a number of contemporary artists embrace traditional techniques of fiber art but use them to create unorthodox forms or address current social and political issues . along these lines , ghada amer has used thread to embroider on canvas repeated motifs of nude women engaged in sexual acts , then partially obscured the embroidered images with gestural painted brushstrokes . her themes include the expression and repression of female sexuality and eroticism in both western and islamic societies . another example of intermixing visual cultures is the complex array of interactions between science and contemporary art , with many artists engaging with scientific imagery and ideas in their practice . for example , wim delvoye ’ s ongoing series called cloaca imagines humans as cyborgs , representing the human digestive system as a kind of biomechanical contraption . finally , many 21st-century artists are deeply affected by their immersion in global visual culture , which is now made vividly present through online networks . many artists maintain a personal website , and some create art expressly for dissemination through social media . as always , new technologies provide new opportunities and challenges . public and participatory art public art was a well-established genre by the late 20th century , attracting both traditional and experimental practitioners . public art in the 21st century has expanded even more as a field of activity in which creative investigation can take place . in addition to continuing familiar forms such as site-specific monuments , murals , graffiti , and collaborations between artists , engineers , and architects , public art encompasses new purposes , forms , and locations , including pop-up art shops , street parades , and online projects . public artists in the 21st century might use established approaches such as installation and performance but introduce new variations . for instance , it is now common for artists to hire other people , sometimes with special skills , to undertake performances on their behalf . in this vein , vanessa beecroft hired fashion models for performances , and the collaborative artists allora & amp ; caldazilla directed professional athletes as performers in some of their installations . a pronounced tendency in the 21st century has been art that is participatory , in which the social interactions prompted by a work become its content . often called relational art , the work literally engages the public in some way . for instance , carsten höller has installed giant slides in museums for visitors to slide down , and rirkrit tiravanija has prepared thai food and served it to gallery goers . artists attracted by the immediacy and connectivity of globally networked media often create online projects that invite social interaction . relational aesthetics has developed ( and been contested ) as a critical theory for analyzing and evaluating such undertakings . key questions in these debates include : does it matter if the social interactions prompted by such works promote a better world or are conviviality and entertainment sufficient goals ? to what extent should the physical products of relational art ( such as höller ’ s slides ) be evaluated aesthetically as well as for their social effects ? the 21st century is just beginning—issues and ideas are evolving rapidly and new artists are constantly gaining attention and influence . essay by jean robertson , chancellor ’ s professor of art history , herron school of art and design , indiana university-purdue university indianapolis this content was first developed for oxford art online and appears courtesy of oxford university press . visit to learn more about contemporary art and see a list of significant 21st-century artists . additional resources : wangechi mutu 's website wim delvoye 's website about contemporary art ( from the getty museum ) postmodernism : recent developments in art in india on the metropolitan museum of art 's heilbrunn timeline of art history postmodernism : recent developments in art in pakistan and bangladesh on the metropolitan museum of art 's heilbrunn timeline of art history art and photography : 1990s-present on the metropolitan museum of art 's heilbrunn timeline of art history
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many important topics have been resonating in the new century and inspiring new thinking and scholarly debate , such as the surge of bio art in response to scientific research in the life sciences , and the critical theory known as relational aesthetics that developed in response to an increase in art that invites viewers ’ participation and interaction . other topics that were much-discussed in the late 20th century remain vital for the analysis of 21st-century art and visual culture , including semiotics , post-modernism , and feminism . art of the 21st century emerges from a vast variety of materials and means .
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what are the characteristics of 21st century portraits ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic .
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in example 2 , i did n't understand why [ h ] = 6.3x10 ... + x , why was x added to the concentration of h what does it have to do with that ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ .
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why have the units been dropped ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic .
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in try 2 , why is the contribution of [ h+ } from autoionization not 1 x 10^7 ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq .
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if the equilibrium constant kw tells us the extent of reaction or 'equilibrium position ' , how can it remain constant if the equilibrium position is changing ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ?
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if ammonia is the general term for nh3 , is there a general term for `` ph '' ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew !
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where did the 6.3x10^-8 m come from ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ?
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how does temperature affect ph ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base .
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when there is acid rain what does the ph of the water is ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic .
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when the bloke subs in the kw constant as 10^-14 is that the same as 1.0 x 10^-14 ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ .
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for example 1 why is [ h3o } equal to 10^-10 m ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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is the new equilibrium still at 10^-14 ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ?
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in the le chatelier 's principle part , could someone tell me what exactly is the reverse reaction being favored to reduce the concentration of hydroxide ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium .
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this is silly , but why is it that the contribution of autoionization of water becomes significant when the solution is slightly basic/acidic ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water .
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if it were stated that the water 's temperature is 25 degrees , could we just use 1x10^-7 instead of x ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent .
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and since in this case the answer would be 6.78 ( instead of 6.87 ) , is it safe to assume that water in actual problem was colder than 25 degrees ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article .
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for practice 2 , how do we know that pure water is neutral at 0 celsius ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ .
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suppose i need to find the ph of 10^-8m aqueous soln of hcl.how do i approach this problem ?
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key points water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions . water is amphoteric water is one of the most common solvents for acid-base reactions . as we discussed in a previous article on brønsted-lowry acids and bases , water is also amphoteric , capable of acting as either a brønsted-lowry acid or base . practice $ 1 $ : identifying the role of water in a reaction in the following reactions , identify if water is playing the role of an acid , a base , or neither . autoionization of water since acids and bases react with each other , this implies that water can react with itself ! while that might sound strange , it does happen $ - $ water molecules exchange protons with one another to a very small extent . we call this process the autoionization , or self-ionization , of water . the proton exchange can be written as the following balanced equation : $ \qquad\qquad\text { h } _2\text { o } ( l ) +\text { h } _2\text { o } ( l ) \rightleftharpoons\text { h } _3\text { o } ^+ ( aq ) +\text { oh } ^- ( aq ) $ one water molecule is donating a proton and acting as a bronsted-lowry acid , while another water molecule accepts the proton , acting as a bronsted-lowry base . this results in the formation of hydronium and hydroxide ions in a $ 1:1 $ molar ratio . for any sample of pure water , the molar concentrations of hydronium , $ \text { h } _3\text { o } ^+ $ , and hydroxide , $ \text { oh } ^- $ , must be equal : $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] ~~\text { in pure water } $ note that this is process is readily reversible . because water is a weak acid and a weak base , the hydronium and hydroxide ions exist in very , very small concentrations relative to that of non-ionized water . just how small are these concentrations ? let 's find out by examining the equilibrium constant for this reaction ( also called the autoionization constant ) , which has the special symbol $ k_\text { w } $ . the autoionization constant , $ k_\text { w } $ the expression for the autoionization constant is $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] \quad\quad\text { ( eq . 1 ) } $ remember that when writing equilibrium expressions , the concentrations of solids and liquids are not included . therefore , our expression for $ k_\text { w } $ does not include the concentration of water , which is a pure liquid . we can calculate the value of $ k_\text { w } $ at $ 25\ , ^\circ\text { c } $ using $ [ \text { h } _3\text { o } ^+ ] $ , which is related to the $ \text { ph } $ of water . at $ 25\ , ^\circ\text { c } $ , the $ \text { ph } $ of pure water is $ 7 $ . therefore , we can calculate the concentration of hydronium ions in pure water : $ [ \text { h } _3\text { o } ^+ ] =10^ { -\text { ph } } =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ in the last section , we saw that hydronium and hydroxide form in a $ 1:1 $ molar ratio during the autoionization of pure water . we can use that relationship to calculate the concentration of hydroxide in pure water at $ 25^\circ\text { c } $ : $ [ \text { oh } ^- ] = [ \text { h } _3\text { o } ^+ ] =10^ { -7 } \text { m } ~~\text { at } 25\ , ^\circ\text { c } $ this is a little tough to visualize , but $ 10^ { -7 } $ is an extremely small number ! within a sample of water , only a small fraction of the water molecules will be in the ionized form . now that we know $ [ \text { oh } ^- ] $ and $ [ \text { h } 3\text { o } ^+ ] $ , we can use these values in our equilibrium expression to calculate $ k\text { w } $ at $ 25^\circ\text { c } $ : $ k_\text { w } = ( 10^ { -7 } ) \times ( 10^ { -7 } ) =10^ { -14 } ~~\text { at } 25\ , ^\circ\text { c } $ concept check : how many hydroxide and hydronium ions are in one liter of water at $ 25^\circ\text { c } $ ? relationship between the autoionization constant , $ \text { ph } $ , and $ \text { poh } $ the fact that $ k_\text { w } $ is equal to $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ leads to an interesting and useful new equation . if we take the negative logarithm of both sides of $ \text { eq . 1 } $ in the previous section , we get the following : $ \begin { align } -\log { k_\text { w } } & amp ; =-\log ( { [ \text { h } _3\text { o } ^+ } ] [ \text { oh } ^- ] ) \ \ & amp ; =-\big ( \log [ \text { h } _3\text { o } ^+ ] +\log [ \text { oh } ^- ] \big ) \ \ & amp ; =-\log [ \text { h } _3\text { o } ^+ ] + ( -\log [ \text { oh } ^- ] ) \ \ & amp ; =\text { ph } +\text { poh } \end { align } $ we can abbreviate $ -\log { k_\text { w } } $ as $ \text { p } k_\text w $ , which is equal to $ 14 $ at $ 25\ , ^\circ\text { c } $ : $ \text { p } k_\text { w } =\text { ph } +\text { poh } =14~~\text { at } 25\ , ^\circ \text c\quad\quad\text { ( eq . 2 } ) $ therefore , the sum of $ \text { ph } $ and $ \text { poh } $ will always be $ 14 $ for any aqueous solution at $ 25\ , ^\circ\text { c } $ . keep in mind that this relationship will not hold true at other temperatures , because $ k_\text { w } $ is temperature dependent ! example $ 1 $ : calculating $ [ \text { oh } ^- ] $ from $ \text { ph } $ an aqueous solution has a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . what is the concentration of hydroxide ions in the solution ? method $ 1 $ : using eq . $ 1 $ one way to solve this problem is to first find $ [ \text { h } ^+ ] $ from the $ \text { ph } $ : $ \begin { align } [ \text { h } _3\text { o } ^+ ] & amp ; =10^ { -\text { ph } } \ \ & amp ; =10^ { -10 } \ , \text m\\end { align } $ we can then calculate $ [ \text { oh } ^- ] $ using eq . 1 : $ \begin { align } k_\text { w } & amp ; = [ \text { h } 3\text { o } ^+ ] [ \text { oh } ^- ] ~~~\quad\quad\text { rearrange to solve for } [ \text { oh } ^- ] \ \ [ \text { oh } ^- ] & amp ; =\dfrac { k\text { w } } { [ \text { h } 3\text { o } ^+ ] } \qquad\quad\qquad\text { plug in values for } k\text w \ , \text { and [ h } _3 \text o^+ ] \ \ & amp ; =\dfrac { 10^ { -14 } } { 10^ { -10 } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ method $ 2 $ : using eq . $ 2 $ another way to calculate $ [ \text { oh } ^- ] $ is to calculate it from the $ \text { poh } $ of the solution . we can use eq . 2 to calculate the $ \text { poh } $ of our solution from the $ \text { ph } $ . rearranging eq . 2 and solving for the $ \text { poh } $ , we get : $ \begin { align } \text { poh } & amp ; =14-\text { ph } \ \ & amp ; =14-10\ \ & amp ; =4\end { align } $ we can now use the equation for $ \text { poh } $ to solve for $ [ \text { oh } ^- ] $ . $ \begin { align } [ \text { oh } ^- ] & amp ; =10^ { -\text { poh } } \ \ & amp ; =10^ { -4 } \text { m } \end { align } $ using either method of solving the problem , the hydroxide concentration is $ 10^ { -4 } \text { m } $ for an aqueous solution with a $ \text { ph } $ of $ 10 $ at $ 25\ , ^\circ\text { c } $ . definitions of acidic , basic , and neutral solutions we have seen that the concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ are equal in pure water , and both have a value of $ 10^ { -7 } \text { m } $ at $ 25\ , ^\circ\text { c } $ . when the concentrations of hydronium and hydroxide are equal , we say that the solution is neutral . aqueous solutions can also be acidic or basic depending on the relative concentrations of $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ autoionization and le chatelier 's principle we also know that in pure water , the concentrations of hydroxide and hydronium are equal . most of the time , however , we are interested in studying aqueous solutions containing other acids and bases . in that case , what happens to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ ? the moment we dissolve other acids or bases in water , we change $ [ \text { h } 3\text { o } ^+ ] $ and/or $ [ \text { oh } ^- ] $ such that the product of the concentrations is no longer is equal to $ k\text { w } $ . that means the reaction is no longer at equilibrium . in response , le chatelier 's principle tells us that the reaction will shift to counteract the change in concentration and establish a new equilibrium . for example , what if we add an acid to pure water ? while pure water at $ 25\ , ^\circ \text c $ has a hydronium ion concentration of $ 10^ { -7 } \ , \text m $ , the added acid increases the concentration of $ \text { h } _3\text { o } ^+ $ . in order to get back to equilibrium , the reaction will favor the reverse reaction to use up some of the extra $ \text { h } _3\text { o } ^+ $ . this causes the concentration of $ \text { oh } ^- $ to decrease until the product of $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ is once again equal to $ 10^ { -14 } $ . once the reaction reaches its new equilibrium state , we know that : $ [ \text h^+ ] & gt ; [ \text { oh } ^- ] $ because the added acid increased $ [ \text h^+ ] $ . thus , our solution is acidic ! $ [ \text { oh } ^- ] & lt ; 10^ { -7 } \ , \text m $ because favoring the reverse reaction decreased $ [ \text { oh } ^- ] $ to get back to equilibrium . the important thing to remember is that any aqueous acid-base reaction can be described as shifting the equilibrium concentrations for the autoionization of water . this is really useful , because that means we can apply eq . 1 and eq . 2 to all aqueous acid-base reactions , not just pure water ! autoionization matters for very dilute acid and base solutions the autoionization of water is usually introduced when first learning about acids and bases , and it is used to derive some extremely useful equations that we 've discussed in this article . however , we will often calculate $ [ \text h^+ ] $ and $ \text { ph } $ for aqueous solutions without including the contribution from the autoionization of water . the reason we can do this is because autoionization usually contributes relatively few ions to the overall $ [ \text h^+ ] $ or $ [ \text { oh } ^- ] $ compared to the ions from additional acid or base . the only situation when we need to remember the autoionization of water is when the concentration of our acid or base is extremely dilute . in practice , this means that we need to include the contribution from autoionization when the concentration of $ \text h^+ $ or $ \text { oh } ^- $ is within ~ $ 2 $ orders of magnitude ( or less than ) of $ \text { 10 } ^ { -7 } \ , \text m $ . we will now go through an example of how to calculate the $ \text { ph } $ of a very dilute acid solution . example $ 2 $ : calculating the $ \text { ph } $ of a very dilute acid solution let 's calculate the $ \text { ph } $ of a $ \text { hcl } $ solution with a hydronium ion concentration of $ 6.3 \times 10^ { -8 } \ , \text m $ . try 1 : ignoring the autoionization of water if we ignore the autoionization of water and simply use the formula for $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } ] \ \ & amp ; =7.20\end { align } $ easy ! we have an aqueous acid solution with a $ \text { ph } $ that is greater than $ 7 $ . but , wait , would n't that make it a basic solution ? that ca n't be right ! try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic . whew ! summary water can undergo autoionization to form $ \text { h } _3\text { o } ^+ $ and $ \text { oh } ^- $ ions . the equilibrium constant for the autoionization of water , $ k_\text { w } $ , is $ 10^ { -14 } $ at $ 25\ , ^\circ\text { c } $ . in a neutral solution , $ [ \text { h } _3\text { o } ^+ ] = [ \text { oh } ^- ] $ in an acidic solution , $ [ \text { h } _3\text { o } ^+ ] & gt ; [ \text { oh } ^- ] $ in a basic solution , $ [ \text { oh } ^- ] & gt ; [ \text { h } _3\text { o } ^+ ] $ for aqueous solutions at $ 25\ , ^\circ\text { c } $ , the following relationships are always true : $ k_\text { w } = [ \text { h } _3\text { o } ^+ ] [ \text { oh } ^- ] =10^ { -14 } $ $ \text { ph } +\text { poh } =14 $ the contribution of the autoionization of water to $ [ \text { h } _3\text { o } ^+ ] $ and $ [ \text { oh } ^- ] $ becomes significant for extremely dilute acid and base solutions .
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try 2 : including the contribution from autoionization to $ [ \text { h } ^+ ] $ since the concentration of this solution is extremely dilute , the concentration of the hydronium from the hydrochloric acid is close to the $ [ \text { h } ^+ ] $ contribution from the autoionization of water . that means : we have to include the contribution from autoionization to $ [ \text { h } ^+ ] $ since the autoionization of water is an equilibrium reaction , we must solve for the overall $ [ \text { h } ^+ ] $ using the expression for $ k_\text { w } $ : $ k_\text { w } = [ \text h^+ ] [ \text { oh } ^- ] =1.0\times10^ { -14 } $ if we say that $ x $ is the contribution of autoionization to the equilibrium concentration of $ \text h^+ $ and $ \text { oh } ^- $ , the concentrations at equilibrium will be as follows : $ [ \text h^+ ] =6.3 \times 10^ { -8 } \ , \text m+x $ $ [ \text { oh } ^- ] =x $ plugging these concentrations into our equilibrium expression , we get : $ \begin { align } k_\text { w } & amp ; = ( 6.3 \times 10^ { -8 } \ , \text m+x ) x=1.0\times10^ { -14 } \ \ & amp ; =x^2+6.3 \times 10^ { -8 } x\end { align } $ rearranging this expression so that everything is equal to $ 0 $ gives the following quadratic equation : $ 0=x^2+6.3 \times 10^ { -8 } x-1.0\times10^ { -14 } $ we can solve for $ x $ using the quadratic formula , which gives the following solutions : $ x=7.3 \times 10^ { -8 } \ , \text m , -1.4 \times 10^ { -7 } \ , \text m $ since the concentration of $ \text { oh } ^- $ ca n't be negative , we can eliminate the second solution . if we plug in the first value of $ x $ to get the equilibrium concentration of $ \text h^+ $ and calculate $ \text { ph } $ , we get : $ \begin { align } \text { ph } & amp ; =-\text { log } [ \text h^+ ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +x ] \ \ & amp ; =-\text { log } [ 6.3 \times 10^ { -8 } +7.3 \times 10^ { -8 } ] \ \ & amp ; =-\text { log } [ 1.36 \times 10^ { -7 } ] \ \ & amp ; =6.87\end { align } $ thus we can see that once we include the autoionization of water , our very dilute $ \text { hcl } $ solution has a $ \text { ph } $ that is weakly acidic .
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how did you solve for x in the quadratic equation ?
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up to around the year 1200 , members of religious houses—monks and nuns—were the primary consumers of books . they produced the objects themselves and in high numbers , because religious houses could not function without them . the 13th century saw a sharp rise in the production and consumption of books outside the monasteries . books were now also made for profit in urban shops , both for use by citizens , students and even monks . most medieval bookstores were empty because books were too expensive to have in stock . instead , each customer would have a long talk with the shopkeeper , who would ask how much he wanted to spend , what materials he preferred , what kind of writing style should be used for the text—and so on . the medieval book is therefore always one of a kind . users often modified the manuscript post-production , bringing it even more in tune with their needs . bookmarks could be added for quick access to favorite chapters , while nota signs and maniculae placed in the margin marked important passages . moreover , glosses and slips with notes were inserted where the text needed clarification . fossilized taste ( bookmarks ) the bookmark guided the reader to the beginning of a favorite chapter or a significant section of the book . flowers or leaves , which were sometimes drawn into the manuscript , marked the page in the most elegant and practical manner . more permanent and secure , however , are bookmarks like the one seen here : a piece of parchment that was pasted onto the page . what a great thought , that medieval fingers inserted these stubs to get to their literary fix ! some bookmarks are fragments of redundant manuscripts , from which they were cut as a form of recycling . the stubs are interesting to book historians because they show what texts of passages were enjoyed by individuals that lived hundreds of years ago . they are the fossilized remains of medieval literary taste . note to self ( nota ) while the bookmark guided the reader to an important chapter or text , the nota-sign marked a significant passage or sentence on the page . from time to time readers noticed something in the text worth highlighting . in such cases they wrote the latin `` nota '' in the margin , which means to `` examine '' or `` inspect '' . while some of these nota-signs may have served as a reminder to check something , others appear to express a more generic `` attention ! `` —like the manicula did ( see below ) . as this example shows , the sign is not written like a normal word . rather , its four letters are reshuffled and stretched so as to form a unique symbol . this was likely done to distinguish a reader 's passages from those marked by other users of the book , such as his fellow brethren in the monastery . helping hands ( maniculae ) with a nota-sign the reader expressed that a passage was noteworthy or deserved a closer read . the manicula ( latin for `` little hand '' ) was another means to do so . as with nota-signs , the actual form of the pointing finger varies considerably . readers may have had their own unique design to distinguish their hands from those of other readers . the hands are sometimes accompanied by short notes , which the reader may have written in response to the text . as with bookmarks and nota-signs , maniculae show us what information were deemed important or relevant to an individual long ago . in that sense they lend a helping hand to the book historian as much as they did to the medieval reader . add-on ( glosses ) from the moment we learn to read as a child we are told not to write in our books . still , we often do , especially when we use the book for school . if it is not our own copy , we may write with pencil , so that our personal thoughts may be removed after use . medieval readers had no problem writing in the margins , and apparently had no qualms about using permanent ink . in fact , it is hard to point out a manuscript that does not contain any such add-ons . some of these notes ( or glosses ) were extensive , and the scribe had to extend the margins to accommodate them . the glosses seen here provide an alternative meaning ( mixta vel bibita ) , a clarification ( posita ) and an encouragement to check something out in another book ( the letter `` r '' for `` require '' ) . sticky note more extensive notes were sometimes written on tiny paper or parchment slips like the one seen here . students are known to have used them to take down notes in the classroom or when they were studying a text at home . few of them survive today . not only were they easy to lose , but many of them were actually thrown out , similar to the fate of our modern day `` sticky notes . '' in some manuscripts they survive because they were tucked in between the pages , as seen here . essay by dr. erik kwakkel
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as with nota-signs , the actual form of the pointing finger varies considerably . readers may have had their own unique design to distinguish their hands from those of other readers . the hands are sometimes accompanied by short notes , which the reader may have written in response to the text .
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did the body of the dragon in the maniculae change also , or just the hand with different readers ?
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the search for identity is an unending one . as we peer across big lenses of time , such as those that separate us from the ancient mediterranean world , one of the questions that occurs again and again is “ who were these people ? ” in the case of palmyra , a prosperous caravan city located in the syrian desert , a remarkable assemblage of funerary portraiture grants us a glimpse at the self-styled identity of a number of the city ’ s former occupants . among the tomb types from roman syria are the curious “ tower tombs ” of palmyra , which find no comparison in roman architecture from the western empire ( above and left ) . they were the main tomb typology from the final quarter of the first century until the middle of the second century c.e. , at which point the underground hypogeum became the preferred tomb type . a hypogeum is a subterranean tomb , most often excavated directly from the bedrock , thus creating an underground chamber or chambers for burials . the tower tombs tend to occupy high ground and were likely built for kinship groupings . these tall , slender structures enclose tiers of niches or loculi in which human remains would be deposited . each loculus would then be sealed with a stone slab , often carved with a relief portrait of the descendent . an exceptionally fine ( and early ) example is the tower tomb of iamblichus ( dated to 83 c.e . on the basis of epigraphic evidence—evidence from inscriptions ) . another is the well preserved tomb of elahbel ( above ) with its elaborately coffered ceilings ( below ) . the tomb also includes a small balcony ( below ) . individual portraiture the individual loculus relief sculptures present a rich range of iconographic information about the people of palmyra . these individualized reliefs are formatted as portrait reliefs and depict their subjects intimately , often with symbols of their status and social position . the portrait of a priest dating c. 50-150 c.e . ( above left ) provides a good example of this practice . the priest holds ritual vessels ( a bowl and a jug ) and wears the traditional polos hat—a high , cylindrical hat worn by both men and women and derived from the divine crowns of the goddesses of the ancient near east and anatolia . a fragmentary female figure stands behind the priest ’ s right shoulder . the funerary bust of tamma , c. 50-150 c.e . ( above right ) demonstrates similar traits . tamma is richly dressed , perhaps indicating worldly wealth , and holds a spindle and distaff , perhaps indicating that she produced fabric in her household . the inscription identifies her as “ tamma , daughter of shamshi geram , son of malku , son of nashum. ” the bust portrait of a couple ( above ) shows a pair of decedents . this portrait carries a greek inscription , which differs from the typical aramaic inscriptions . the text identifies the two individuals as viria phoebe and gaius virius alcimus . this pair have the same clan name , a possible indication they are the former slaves of a brother and sister . alcimus holds a book-roll , while the woman holds the spindle and distaff ( both implements associated with cloth production ) . these objects may be meant to evoke their respective roles . group reliefs a third century c.e . funerary relief from palmyra now in the british museum ( below ) depicts a funeral banquet . elite tombs of this period demonstrate a mixture of roman and near eastern motifs . in this particular relief that depicts a funeral banquet , the reclining male is attended by a seated female ; perhaps the pair are meant to be husband and wife . the idea of the funeral banquet is a roman motif adopted by local craftsmen . the male—presumably the deceased—reclines on a couch while holding an open vessel . he is depicted at a slightly larger scale than the attendant female . his costume is of parthian origin , a sort of pant-suit . the parthian empire , c. 247 b.c.e.-224 c.e . was a major political power of ancient iran located on the eastern margin of the roman empire . reliefs such as this one would be arranged in groups of three in communal tombs , thereby giving the tomb chamber the resemblance of a roman-style dining room ( triclinium ) in which a real banquet would have taken place . the funerary reliefs from palmyra form a profoundly evocative body of evidence . the individualized treatment of the sculptures themselves still serves to convey important elements about the identities of these individuals . we can glean information about wealth , social status , role in the community , familial relationships—all of which help to enrich our reconstruction of ancient palmyrene society . the reliefs also demonstrate the degree to which palmyra existed in a multicultural and multilingual landscape , one in which the traits , trends , styles , and languages of the graeco-roman world and the near eastern world not only overlapped but intertwined , producing new , unique cultural objects . this is an important realization , one that helps remind us of the degree to which the ancient world was diverse and varied and that a great deal of the material culture of the ancient world resulted from shared cultural influence and hybridization . the tombs of palmyra embody and evoke this climate of cultural diversity as they still stand as monuments to palmyrene identity . essay by dr. jeffrey becker additional resources w. ball , rome in the east : the transformation of an empire ( london : routledge , 2001 ) . m.a.r . colledge , the art of palmyra ( london : westview press , 1976 ) . michael danti , “ palmyrene funerary sculptures at penn , ” expedition 43.3 ( november 2001 ) . palmyrene funerary sculptures at penn m. k. heyn , “ gesture and identity in the funerary art of palmyra ” american journal of archaeology 114.4 ( october 2010 ) pp . 631-61 . doi : 10.3764/aja.114.4.631 andreas j. kropp , images and monuments of near eastern dynasts , 100 bc - ad 100 ( oxford : oxford university press , 2013 ) . heilbrunn timeline of art history , the metropolitan museum of art the lady marti , funerary portrait of a woman , palmyra , c. ad 170-190 , ny carlsberg glyptotek , copenhagen funerary portrait of a priest from palmyra , walters art museum the wisconsin palmyrene aramaic inscription project palmyra portrait project , aarhus university portrait bust , penn museum portrait of a deceased woman , liebieghaus , frankfurt
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the tomb also includes a small balcony ( below ) . individual portraiture the individual loculus relief sculptures present a rich range of iconographic information about the people of palmyra . these individualized reliefs are formatted as portrait reliefs and depict their subjects intimately , often with symbols of their status and social position . the portrait of a priest dating c. 50-150 c.e .
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given the current status of isis ( alternatively called isil ) and all the chaos within syria ... what is the status of palmyra ?
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the search for identity is an unending one . as we peer across big lenses of time , such as those that separate us from the ancient mediterranean world , one of the questions that occurs again and again is “ who were these people ? ” in the case of palmyra , a prosperous caravan city located in the syrian desert , a remarkable assemblage of funerary portraiture grants us a glimpse at the self-styled identity of a number of the city ’ s former occupants . among the tomb types from roman syria are the curious “ tower tombs ” of palmyra , which find no comparison in roman architecture from the western empire ( above and left ) . they were the main tomb typology from the final quarter of the first century until the middle of the second century c.e. , at which point the underground hypogeum became the preferred tomb type . a hypogeum is a subterranean tomb , most often excavated directly from the bedrock , thus creating an underground chamber or chambers for burials . the tower tombs tend to occupy high ground and were likely built for kinship groupings . these tall , slender structures enclose tiers of niches or loculi in which human remains would be deposited . each loculus would then be sealed with a stone slab , often carved with a relief portrait of the descendent . an exceptionally fine ( and early ) example is the tower tomb of iamblichus ( dated to 83 c.e . on the basis of epigraphic evidence—evidence from inscriptions ) . another is the well preserved tomb of elahbel ( above ) with its elaborately coffered ceilings ( below ) . the tomb also includes a small balcony ( below ) . individual portraiture the individual loculus relief sculptures present a rich range of iconographic information about the people of palmyra . these individualized reliefs are formatted as portrait reliefs and depict their subjects intimately , often with symbols of their status and social position . the portrait of a priest dating c. 50-150 c.e . ( above left ) provides a good example of this practice . the priest holds ritual vessels ( a bowl and a jug ) and wears the traditional polos hat—a high , cylindrical hat worn by both men and women and derived from the divine crowns of the goddesses of the ancient near east and anatolia . a fragmentary female figure stands behind the priest ’ s right shoulder . the funerary bust of tamma , c. 50-150 c.e . ( above right ) demonstrates similar traits . tamma is richly dressed , perhaps indicating worldly wealth , and holds a spindle and distaff , perhaps indicating that she produced fabric in her household . the inscription identifies her as “ tamma , daughter of shamshi geram , son of malku , son of nashum. ” the bust portrait of a couple ( above ) shows a pair of decedents . this portrait carries a greek inscription , which differs from the typical aramaic inscriptions . the text identifies the two individuals as viria phoebe and gaius virius alcimus . this pair have the same clan name , a possible indication they are the former slaves of a brother and sister . alcimus holds a book-roll , while the woman holds the spindle and distaff ( both implements associated with cloth production ) . these objects may be meant to evoke their respective roles . group reliefs a third century c.e . funerary relief from palmyra now in the british museum ( below ) depicts a funeral banquet . elite tombs of this period demonstrate a mixture of roman and near eastern motifs . in this particular relief that depicts a funeral banquet , the reclining male is attended by a seated female ; perhaps the pair are meant to be husband and wife . the idea of the funeral banquet is a roman motif adopted by local craftsmen . the male—presumably the deceased—reclines on a couch while holding an open vessel . he is depicted at a slightly larger scale than the attendant female . his costume is of parthian origin , a sort of pant-suit . the parthian empire , c. 247 b.c.e.-224 c.e . was a major political power of ancient iran located on the eastern margin of the roman empire . reliefs such as this one would be arranged in groups of three in communal tombs , thereby giving the tomb chamber the resemblance of a roman-style dining room ( triclinium ) in which a real banquet would have taken place . the funerary reliefs from palmyra form a profoundly evocative body of evidence . the individualized treatment of the sculptures themselves still serves to convey important elements about the identities of these individuals . we can glean information about wealth , social status , role in the community , familial relationships—all of which help to enrich our reconstruction of ancient palmyrene society . the reliefs also demonstrate the degree to which palmyra existed in a multicultural and multilingual landscape , one in which the traits , trends , styles , and languages of the graeco-roman world and the near eastern world not only overlapped but intertwined , producing new , unique cultural objects . this is an important realization , one that helps remind us of the degree to which the ancient world was diverse and varied and that a great deal of the material culture of the ancient world resulted from shared cultural influence and hybridization . the tombs of palmyra embody and evoke this climate of cultural diversity as they still stand as monuments to palmyrene identity . essay by dr. jeffrey becker additional resources w. ball , rome in the east : the transformation of an empire ( london : routledge , 2001 ) . m.a.r . colledge , the art of palmyra ( london : westview press , 1976 ) . michael danti , “ palmyrene funerary sculptures at penn , ” expedition 43.3 ( november 2001 ) . palmyrene funerary sculptures at penn m. k. heyn , “ gesture and identity in the funerary art of palmyra ” american journal of archaeology 114.4 ( october 2010 ) pp . 631-61 . doi : 10.3764/aja.114.4.631 andreas j. kropp , images and monuments of near eastern dynasts , 100 bc - ad 100 ( oxford : oxford university press , 2013 ) . heilbrunn timeline of art history , the metropolitan museum of art the lady marti , funerary portrait of a woman , palmyra , c. ad 170-190 , ny carlsberg glyptotek , copenhagen funerary portrait of a priest from palmyra , walters art museum the wisconsin palmyrene aramaic inscription project palmyra portrait project , aarhus university portrait bust , penn museum portrait of a deceased woman , liebieghaus , frankfurt
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the search for identity is an unending one . as we peer across big lenses of time , such as those that separate us from the ancient mediterranean world , one of the questions that occurs again and again is “ who were these people ? ” in the case of palmyra , a prosperous caravan city located in the syrian desert , a remarkable assemblage of funerary portraiture grants us a glimpse at the self-styled identity of a number of the city ’ s former occupants .
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were the faces purposefully erased ?
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the search for identity is an unending one . as we peer across big lenses of time , such as those that separate us from the ancient mediterranean world , one of the questions that occurs again and again is “ who were these people ? ” in the case of palmyra , a prosperous caravan city located in the syrian desert , a remarkable assemblage of funerary portraiture grants us a glimpse at the self-styled identity of a number of the city ’ s former occupants . among the tomb types from roman syria are the curious “ tower tombs ” of palmyra , which find no comparison in roman architecture from the western empire ( above and left ) . they were the main tomb typology from the final quarter of the first century until the middle of the second century c.e. , at which point the underground hypogeum became the preferred tomb type . a hypogeum is a subterranean tomb , most often excavated directly from the bedrock , thus creating an underground chamber or chambers for burials . the tower tombs tend to occupy high ground and were likely built for kinship groupings . these tall , slender structures enclose tiers of niches or loculi in which human remains would be deposited . each loculus would then be sealed with a stone slab , often carved with a relief portrait of the descendent . an exceptionally fine ( and early ) example is the tower tomb of iamblichus ( dated to 83 c.e . on the basis of epigraphic evidence—evidence from inscriptions ) . another is the well preserved tomb of elahbel ( above ) with its elaborately coffered ceilings ( below ) . the tomb also includes a small balcony ( below ) . individual portraiture the individual loculus relief sculptures present a rich range of iconographic information about the people of palmyra . these individualized reliefs are formatted as portrait reliefs and depict their subjects intimately , often with symbols of their status and social position . the portrait of a priest dating c. 50-150 c.e . ( above left ) provides a good example of this practice . the priest holds ritual vessels ( a bowl and a jug ) and wears the traditional polos hat—a high , cylindrical hat worn by both men and women and derived from the divine crowns of the goddesses of the ancient near east and anatolia . a fragmentary female figure stands behind the priest ’ s right shoulder . the funerary bust of tamma , c. 50-150 c.e . ( above right ) demonstrates similar traits . tamma is richly dressed , perhaps indicating worldly wealth , and holds a spindle and distaff , perhaps indicating that she produced fabric in her household . the inscription identifies her as “ tamma , daughter of shamshi geram , son of malku , son of nashum. ” the bust portrait of a couple ( above ) shows a pair of decedents . this portrait carries a greek inscription , which differs from the typical aramaic inscriptions . the text identifies the two individuals as viria phoebe and gaius virius alcimus . this pair have the same clan name , a possible indication they are the former slaves of a brother and sister . alcimus holds a book-roll , while the woman holds the spindle and distaff ( both implements associated with cloth production ) . these objects may be meant to evoke their respective roles . group reliefs a third century c.e . funerary relief from palmyra now in the british museum ( below ) depicts a funeral banquet . elite tombs of this period demonstrate a mixture of roman and near eastern motifs . in this particular relief that depicts a funeral banquet , the reclining male is attended by a seated female ; perhaps the pair are meant to be husband and wife . the idea of the funeral banquet is a roman motif adopted by local craftsmen . the male—presumably the deceased—reclines on a couch while holding an open vessel . he is depicted at a slightly larger scale than the attendant female . his costume is of parthian origin , a sort of pant-suit . the parthian empire , c. 247 b.c.e.-224 c.e . was a major political power of ancient iran located on the eastern margin of the roman empire . reliefs such as this one would be arranged in groups of three in communal tombs , thereby giving the tomb chamber the resemblance of a roman-style dining room ( triclinium ) in which a real banquet would have taken place . the funerary reliefs from palmyra form a profoundly evocative body of evidence . the individualized treatment of the sculptures themselves still serves to convey important elements about the identities of these individuals . we can glean information about wealth , social status , role in the community , familial relationships—all of which help to enrich our reconstruction of ancient palmyrene society . the reliefs also demonstrate the degree to which palmyra existed in a multicultural and multilingual landscape , one in which the traits , trends , styles , and languages of the graeco-roman world and the near eastern world not only overlapped but intertwined , producing new , unique cultural objects . this is an important realization , one that helps remind us of the degree to which the ancient world was diverse and varied and that a great deal of the material culture of the ancient world resulted from shared cultural influence and hybridization . the tombs of palmyra embody and evoke this climate of cultural diversity as they still stand as monuments to palmyrene identity . essay by dr. jeffrey becker additional resources w. ball , rome in the east : the transformation of an empire ( london : routledge , 2001 ) . m.a.r . colledge , the art of palmyra ( london : westview press , 1976 ) . michael danti , “ palmyrene funerary sculptures at penn , ” expedition 43.3 ( november 2001 ) . palmyrene funerary sculptures at penn m. k. heyn , “ gesture and identity in the funerary art of palmyra ” american journal of archaeology 114.4 ( october 2010 ) pp . 631-61 . doi : 10.3764/aja.114.4.631 andreas j. kropp , images and monuments of near eastern dynasts , 100 bc - ad 100 ( oxford : oxford university press , 2013 ) . heilbrunn timeline of art history , the metropolitan museum of art the lady marti , funerary portrait of a woman , palmyra , c. ad 170-190 , ny carlsberg glyptotek , copenhagen funerary portrait of a priest from palmyra , walters art museum the wisconsin palmyrene aramaic inscription project palmyra portrait project , aarhus university portrait bust , penn museum portrait of a deceased woman , liebieghaus , frankfurt
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( above left ) provides a good example of this practice . the priest holds ritual vessels ( a bowl and a jug ) and wears the traditional polos hat—a high , cylindrical hat worn by both men and women and derived from the divine crowns of the goddesses of the ancient near east and anatolia . a fragmentary female figure stands behind the priest ’ s right shoulder .
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'the priest holds ritual vessels ( a bowl and a jug ) ... ' what meaning do these vessels have ?
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