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Hematoxylin and eosin stain ( or haematoxylin and eosin stain or hematoxylin-eosin stain ; often abbreviated as H&E stain or HE stain ) is one of the principal tissue stains used in histology . [ 1 ] [ 2 ] [ 3 ] It is the most widely used stain in medical diagnosis [ 1 ] and is often the gold standard . [ 4 ] For example, when a pathologist looks at a biopsy of a suspected cancer , the histological section is likely to be stained with H&E.
H&E is the combination of two histological stains: hematoxylin and eosin . The hematoxylin stains cell nuclei a purplish blue, and eosin stains the extracellular matrix and cytoplasm pink, with other structures taking on different shades, hues, and combinations of these colors. [ 5 ] [ 6 ] Hence a pathologist can easily differentiate between the nuclear and cytoplasmic parts of a cell, and additionally, the overall patterns of coloration from the stain show the general layout and distribution of cells and provides a general overview of a tissue sample's structure. [ 7 ] Thus, pattern recognition, both by expert humans themselves and by software that aids those experts (in digital pathology ), provides histologic information.
This stain combination was introduced in 1877 by chemist Nicolaus Wissozky at the Kazan Imperial University in Russia. [ 7 ] [ 8 ]
The H&E staining procedure is the principal stain in histology [ 3 ] [ 7 ] [ 2 ] [ 5 ] in part because it can be done quickly, [ 7 ] is not expensive, and stains tissues in such a way that a considerable amount of microscopic anatomy [ 9 ] [ 10 ] is revealed, [ 7 ] [ 5 ] [ 4 ] and can be used to diagnose a wide range of histopathologic conditions. [ 8 ] The results from H&E staining are not overly dependent on the chemical used to fix the tissue or slight inconsistencies in laboratory protocol, [ 11 ] and these factors contribute to its routine use in histology. [ 7 ]
H&E staining does not always provide enough contrast to differentiate all tissues, cellular structures, or the distribution of chemical substances, [ 9 ] and in these cases more specific stains and methods are used. [ 10 ] [ 7 ]
There are many ways to prepare the hematoxylin solutions (formulation) used in the H&E procedure, [ 11 ] [ 12 ] [ 6 ] in addition, there are many laboratory protocols for producing H&E stained slides, [ 9 ] some of which may be specific to a certain laboratory. [ 7 ] Although there is no standard procedure, [ 11 ] [ 9 ] the results by convention are reasonably consistent in that cell nuclei are stained blue and the cytoplasm and extracellular matrix are stained pink. [ 7 ] Histology laboratories may also adjust the amount or type of staining for a particular pathologist. [ 7 ]
After tissues have been collected (often as biopsies ) and fixed, they are typically dehydrated and embedded in melted paraffin wax , the resulting block is mounted on a microtome and cut into thin slices. [ 6 ] The slices are affixed to microscope slides at which point the wax is removed with a solvent and the tissue slices attached to the slides are rehydrated and are ready for staining. [ 6 ] Alternatively, H&E stain is the most used stain in Mohs surgery in which tissues are typically frozen, cut on a cryostat (a microtome that cuts frozen tissue), fixed in alcohol, and then stained. [ 9 ]
The H&E staining method involves application of haematoxylin mixed with a metallic salt, or mordant , often followed by a rinse in a weak acid solution to remove excess staining ( differentiation ), followed by bluing in mildly alkaline water. [ 13 ] [ 8 ] [ 14 ] After the application of haematoxylin, the tissue is counterstained with eosin (most commonly eosin Y ). [ 6 ] [ 8 ] [ 7 ]
Hematoxylin principally colors the nuclei of cells blue or dark-purple, [ 6 ] [ 15 ] [ 14 ] along with a few other tissues, such as keratohyalin granules and calcified material. Eosin stains the cytoplasm and some other structures including extracellular matrix such as collagen [ 5 ] [ 7 ] [ 14 ] in up to five shades of pink. [ 8 ] The eosinophilic (substances that are stained by eosin) [ 5 ] structures are generally composed of intracellular or extracellular proteins . The Lewy bodies and Mallory bodies are examples of eosinophilic structures. Most of the cytoplasm is eosinophilic and is rendered pink. [ 10 ] [ 15 ] Red blood cells are stained intensely red. [ citation needed ]
Although hematein , an oxidized form of hematoxylin, [ 5 ] [ 16 ] [ 14 ] is the active colorant (when combined with a mordant), the stain is still referred to as hematoxylin . [ 8 ] [ 13 ] Hematoxylin, when combined with a mordant (most commonly aluminum alum ) is often considered to "resemble" [ 10 ] a basic, positively charged, or cationic stain. [ 5 ] Eosin is an anionic (negatively charged) and acidic stain. [ 5 ] [ 10 ] The staining of nuclei by hemalum (a combination of aluminum ions and hematein) [ 14 ] is ordinarily due to binding of the dye-metal complex to DNA, but nuclear staining can be obtained after extraction of DNA [ 14 ] from tissue sections. The mechanism is different from that of nuclear staining by basic (cationic) dyes such as thionine or toluidine blue . [ 10 ] Staining by basic dyes occurs only from solutions that are less acidic than hemalum, and it is prevented by prior chemical or enzymatic extraction of nucleic acids. There is evidence to indicate that co-ordinate bonds, similar to those that hold aluminium and hematein together, bind the hemalum complex to DNA and to carboxy groups of proteins in the nuclear chromatin . [ citation needed ]
The structures do not have to be acidic or basic to be called basophilic and eosinophilic; the terminology is based on the affinity of cellular components for the dyes. Other colors, e.g. yellow and brown, can be present in the sample; they are caused by intrinsic pigments such as melanin . Basal laminae need to be stained by PAS stain or some silver stains , if they have to be well visible. Reticular fibers also require silver stain. Hydrophobic structures also tend to remain clear; these are usually rich in fats, e.g. adipocytes , myelin around neuron axons , and Golgi apparatus membranes. [ citation needed ] | https://en.wikipedia.org/wiki/H&E_stain |
Carborane acids H(CXB 11 Y 5 Z 6 ) (X, Y, Z = H, Alk, F, Cl, Br, CF 3 ) are a class of superacids , [ 1 ] some of which are estimated to be at least one million times stronger than 100% pure sulfuric acid in terms of their Hammett acidity function values ( H 0 ≤ −18) and possess computed p K a values well below −20, establishing them as some of the strongest known Brønsted acids. [ 2 ] [ 3 ] [ 4 ] The best-studied example is the highly chlorinated derivative H(CHB 11 Cl 11 ) . The acidity of H(CHB 11 Cl 11 ) was found to vastly exceed that of triflic acid , CF 3 SO 3 H , and bistriflimide , (CF 3 SO 2 ) 2 NH , compounds previously regarded as the strongest isolable acids.
Their high acidities stem from the extensive delocalization of their conjugate bases, carboranate anions (CXB 11 Y 5 Z 6 − ), which are usually further stabilized by electronegative groups like Cl, F, and CF 3 . Due to the lack of oxidizing properties and the exceptionally low nucleophilicity and high stability of their conjugate bases, they are the only superacids known to protonate C 60 fullerene without decomposing it. [ 5 ] [ 6 ] Additionally, they form stable, isolable salts with protonated benzene , C 6 H 7 + , the parent compound of the Wheland intermediates encountered in electrophilic aromatic substitution reactions.
The fluorinated carborane acid, H(CHB 11 F 11 ) , is even stronger than chlorinated carborane acid. It is able to protonate butane to form tert -butyl cation at room temperature and is the only known acid to protonate carbon dioxide to give the bridged cation, [H(CO 2 ) 2 ] + , making it possibly the strongest known acid. In particular, CO 2 does not undergo observable protonation when treated with the mixed superacids HF-SbF 5 or HSO 3 F-SbF 5 . [ 7 ] [ 8 ] [ 9 ]
As a class, the carborane acids form the most acidic group of well-defined, isolable substances known, far more acidic than previously known single-component strong acids like triflic acid or perchloric acid . In certain cases, like the nearly perhalogenated derivatives mentioned above, their acidities rival (and possibly exceed) those of the traditional mixed Lewis-Brønsted superacids like magic acid and fluoroantimonic acid . (However, a head-to-head comparison has not been possible thus far, due to the lack of a measure of acidity that is suitable for both classes of acids: p K a values are ill-defined for the chemically complex mixed acids while H 0 values cannot be measured for the very high melting carborane acids).
A Brønsted–Lowry acid's strength corresponds with its ability to release a hydrogen ion. One common measure of acid strength for concentrated, superacidic liquid media is the Hammett acidity function, H 0 . Based on its ability to quantitatively protonate benzene, the chlorinated carborane acid H(CHB 11 Cl 11 ) was conservatively estimated to have an H 0 value at or below −18, leading to the common assertion that carborane acids are at least a million times stronger than 100% sulfuric acid ( H 0 = −12). [ 11 ] [ 12 ] However, since the H 0 value measures the protonating ability of a liquid medium, the crystalline and high-melting nature of these acids precludes direct measurement of this parameter. In terms of p K a , a slightly different measure of acidity defined as the ability of a given solute to undergo ionization in a solvent, carborane acids are estimated to have p K a values below −20, even without electron-withdrawing substituents on the boron atoms (e.g., H(CHB 11 H 11 ) is estimated to have a p K a of −24), [ 13 ] with the (yet unknown) fully fluorinated analog H(CB 11 F 12 ) having a calculated p K a of −46. [ 4 ] The known acid H(CHB 11 F 11 ) with one fewer fluorine is expected to be only slightly weaker (p K a < −40).
In the gas phase, H(CHB 11 F 11 ) has a computed acidity of 216 kcal/mol, compared to an experimentally determined acidity of 241 kcal/mol (in reasonable agreement with the computed value of 230 kcal/mol) for H(CHB 11 Cl 11 ) . In contrast, HSbF 6 (a simplified model for the proton donating species in fluoroantimonic acid ) has a computed gas phase acidity of 255 kcal/mol, while the previous experimentally determined record holder was (C 4 F 9 SO 2 ) 2 NH, a congener of bistriflimide , at 291 kcal/mol. Thus, H(CHB 11 F 11 ) is likely the most acidic substance so far synthesized in bulk, in terms of its gas phase acidity. In view of its unique reactivity, it is also a strong contender for being the most acidic substance in the condensed phase (see above). Some even more strongly acidic derivatives have been predicted, with gas phase acidities < 200 kcal/mol. [ 14 ] [ 15 ]
Carborane acids differ from classical superacids in being well-defined one component substances. In contrast, classical superacids are often mixtures of a Brønsted acid and Lewis acid (e.g. HF/SbF 5 ). [ 6 ] Despite being the strongest acid, the boron-based carborane acids are described as being "gentle", cleanly protonating weakly basic substances without further side reactions. [ 11 ] Whereas conventional superacids decompose fullerenes due to their strongly oxidizing Lewis acidic component, carborane acid has the ability to protonate fullerenes at room temperature to yield an isolable salt. [ 16 ] [ 6 ] Furthermore, the anion that forms as a result of proton transfer is nearly completely inert. This property is what makes the carborane acids the only substances that are comparable in acidity to the mixed superacids that can also be stored in a glass bottle, as various fluoride-donating species (which attack glass) are not present or generated. [ 17 ] [ 16 ]
Carborane acid was first discovered and synthesized by Professor Christopher Reed and his colleagues in 2004 at the University of California, Riverside. [ 6 ] The parent molecule from which carborane acid is derived, an icosahedral carboranate anion, HCB 11 H − 11 , was first synthesized at DuPont in 1967 by Walter Knoth. Research into this molecule's properties was put on hiatus until the mid 1980s when the Czech group of boron scientists, Plešek, Štíbr, and Heřmánek improved the process for halogenation of carborane molecules. These findings were instrumental in developing the current procedure for carborane acid synthesis. [ 16 ] [ 18 ] The process consists of treating Cs + [HCB 11 H 11 ] − with SO 2 Cl 2 , refluxing under dry argon to fully chlorinate the molecule yielding carborane acid, but this has been shown to fully chlorinate only under select conditions. [ 19 ] [ 16 ] [ 20 ]
In 2010, Reed published a guide giving detailed procedures for the synthesis of carborane acids and their derivatives. [ 21 ] Nevertheless, the synthesis of carborane acids remains lengthy and difficult and requires a well-maintained glovebox and some specialized equipment. The starting material is commercially available decaborane(14) , a highly toxic substance. The most well-studied carborane acid H(CHB 11 Cl 11 ) is prepared in 13 steps. The last few steps are especially sensitive and require a glovebox at < 1 ppm H 2 O without any weakly basic solvent vapors, since bases as weak as benzene or dichloromethane will react with carborane-based electrophiles and Brønsted acids. The final step of the synthesis is the metathesis of the μ-hydridodisilylium carboranate salt with excess liquid, anhydrous hydrogen chloride, presumably driven by the formation of strong Si–Cl and H–H bonds in the volatile byproducts:
The product was isolated by evaporation of the byproducts and was characterized by its infrared (ν CH = 3023 cm −1 ) and nuclear magnetic resonance (δ 4.55 (s, 1H, CH), 20.4 (s, 1H, H + ) in liquid SO 2 ) spectra (note the extremely downfield chemical shift of the acidic proton). [ 21 ] Although the reactions used in the synthesis are analogous, obtaining a pure sample of the more acidic H(CHB 11 F 11 ) turned out to be even more difficult, requiring extremely rigorous procedures to exclude traces of weakly basic impurities. [ 7 ]
Carborane acid consists of 11 boron atoms; each boron atom is bound to a chlorine atom. The chlorine atoms serve to enhance acidity and act as shields against attacks from the outside due to the steric hindrance they form around the cluster. The cluster, consisting of the 11 borons, 11 chlorines, and a single carbon atom, is paired with a hydrogen atom, bound to the carbon atom. The boron and carbon atoms are allowed to form six bonds due to boron's ability to form three-center, two-electron bonds. [ 18 ]
Although the structure of the carborane acid differs greatly from conventional acids, both distribute charge and stability in a similar fashion. The carboranate anion distributes its charge by delocalizing the electrons throughout the 12 cage atoms. [ 22 ] This was shown in a single crystal X-ray diffraction study revealing shortened bond lengths in the heterocyclic portion of the ring suggesting electronic delocalization. [ 23 ]
The chlorinated carba- closo -dodecaborate anion HCB 11 Cl − 11 is an outstandingly stable anion with what has previously been described as "substitutionally inert" B–Cl vertices.
The descriptor closo indicates that the molecule is formally derived (by B-to-C + replacement) from a borane of stoichiometry and charge [B n H n ] 2− ( n = 12 for known carborane acids). [ 24 ] The cagelike structure formed by the 11 boron atoms and 1 carbon atom allows the electrons to be highly delocalized through the 3D cage (the special stabilization of the carborane system has been termed "σ-aromaticity"), and the high energy required to disrupt the boron cluster portion of the molecule is what gives the anion its remarkable stability. [ 24 ] Because the anion is extremely stable, it will not behave as a nucleophile toward the protonated substrate, while the acid itself is completely non-oxidizing, unlike the Lewis acidic components of many superacids like antimony pentafluoride. Hence, sensitive molecules like C 60 can be protonated without decomposition. [ 25 ] [ 26 ]
There are many proposed applications for the boron-based carborane acids. For instance, they have been proposed as catalysts for hydrocarbon cracking and isomerization of n -alkanes to form branched isoalkanes ("isooctane", for example). Carborane acids may also be used as strong, selective Brønsted acids for fine chemical synthesis, where the low nucleophilicity of the counteranion may be advantageous. In mechanistic organic chemistry, they may be used in the study of reactive cationic intermediates. [ 27 ] In inorganic synthesis, their unparalleled acidity may allow for the isolation of exotic species like salts of protonated xenon. [ 17 ] [ 18 ] [ 28 ] | https://en.wikipedia.org/wiki/H(CHB11Cl11) |
Carborane acids H(CXB 11 Y 5 Z 6 ) (X, Y, Z = H, Alk, F, Cl, Br, CF 3 ) are a class of superacids , [ 1 ] some of which are estimated to be at least one million times stronger than 100% pure sulfuric acid in terms of their Hammett acidity function values ( H 0 ≤ −18) and possess computed p K a values well below −20, establishing them as some of the strongest known Brønsted acids. [ 2 ] [ 3 ] [ 4 ] The best-studied example is the highly chlorinated derivative H(CHB 11 Cl 11 ) . The acidity of H(CHB 11 Cl 11 ) was found to vastly exceed that of triflic acid , CF 3 SO 3 H , and bistriflimide , (CF 3 SO 2 ) 2 NH , compounds previously regarded as the strongest isolable acids.
Their high acidities stem from the extensive delocalization of their conjugate bases, carboranate anions (CXB 11 Y 5 Z 6 − ), which are usually further stabilized by electronegative groups like Cl, F, and CF 3 . Due to the lack of oxidizing properties and the exceptionally low nucleophilicity and high stability of their conjugate bases, they are the only superacids known to protonate C 60 fullerene without decomposing it. [ 5 ] [ 6 ] Additionally, they form stable, isolable salts with protonated benzene , C 6 H 7 + , the parent compound of the Wheland intermediates encountered in electrophilic aromatic substitution reactions.
The fluorinated carborane acid, H(CHB 11 F 11 ) , is even stronger than chlorinated carborane acid. It is able to protonate butane to form tert -butyl cation at room temperature and is the only known acid to protonate carbon dioxide to give the bridged cation, [H(CO 2 ) 2 ] + , making it possibly the strongest known acid. In particular, CO 2 does not undergo observable protonation when treated with the mixed superacids HF-SbF 5 or HSO 3 F-SbF 5 . [ 7 ] [ 8 ] [ 9 ]
As a class, the carborane acids form the most acidic group of well-defined, isolable substances known, far more acidic than previously known single-component strong acids like triflic acid or perchloric acid . In certain cases, like the nearly perhalogenated derivatives mentioned above, their acidities rival (and possibly exceed) those of the traditional mixed Lewis-Brønsted superacids like magic acid and fluoroantimonic acid . (However, a head-to-head comparison has not been possible thus far, due to the lack of a measure of acidity that is suitable for both classes of acids: p K a values are ill-defined for the chemically complex mixed acids while H 0 values cannot be measured for the very high melting carborane acids).
A Brønsted–Lowry acid's strength corresponds with its ability to release a hydrogen ion. One common measure of acid strength for concentrated, superacidic liquid media is the Hammett acidity function, H 0 . Based on its ability to quantitatively protonate benzene, the chlorinated carborane acid H(CHB 11 Cl 11 ) was conservatively estimated to have an H 0 value at or below −18, leading to the common assertion that carborane acids are at least a million times stronger than 100% sulfuric acid ( H 0 = −12). [ 11 ] [ 12 ] However, since the H 0 value measures the protonating ability of a liquid medium, the crystalline and high-melting nature of these acids precludes direct measurement of this parameter. In terms of p K a , a slightly different measure of acidity defined as the ability of a given solute to undergo ionization in a solvent, carborane acids are estimated to have p K a values below −20, even without electron-withdrawing substituents on the boron atoms (e.g., H(CHB 11 H 11 ) is estimated to have a p K a of −24), [ 13 ] with the (yet unknown) fully fluorinated analog H(CB 11 F 12 ) having a calculated p K a of −46. [ 4 ] The known acid H(CHB 11 F 11 ) with one fewer fluorine is expected to be only slightly weaker (p K a < −40).
In the gas phase, H(CHB 11 F 11 ) has a computed acidity of 216 kcal/mol, compared to an experimentally determined acidity of 241 kcal/mol (in reasonable agreement with the computed value of 230 kcal/mol) for H(CHB 11 Cl 11 ) . In contrast, HSbF 6 (a simplified model for the proton donating species in fluoroantimonic acid ) has a computed gas phase acidity of 255 kcal/mol, while the previous experimentally determined record holder was (C 4 F 9 SO 2 ) 2 NH, a congener of bistriflimide , at 291 kcal/mol. Thus, H(CHB 11 F 11 ) is likely the most acidic substance so far synthesized in bulk, in terms of its gas phase acidity. In view of its unique reactivity, it is also a strong contender for being the most acidic substance in the condensed phase (see above). Some even more strongly acidic derivatives have been predicted, with gas phase acidities < 200 kcal/mol. [ 14 ] [ 15 ]
Carborane acids differ from classical superacids in being well-defined one component substances. In contrast, classical superacids are often mixtures of a Brønsted acid and Lewis acid (e.g. HF/SbF 5 ). [ 6 ] Despite being the strongest acid, the boron-based carborane acids are described as being "gentle", cleanly protonating weakly basic substances without further side reactions. [ 11 ] Whereas conventional superacids decompose fullerenes due to their strongly oxidizing Lewis acidic component, carborane acid has the ability to protonate fullerenes at room temperature to yield an isolable salt. [ 16 ] [ 6 ] Furthermore, the anion that forms as a result of proton transfer is nearly completely inert. This property is what makes the carborane acids the only substances that are comparable in acidity to the mixed superacids that can also be stored in a glass bottle, as various fluoride-donating species (which attack glass) are not present or generated. [ 17 ] [ 16 ]
Carborane acid was first discovered and synthesized by Professor Christopher Reed and his colleagues in 2004 at the University of California, Riverside. [ 6 ] The parent molecule from which carborane acid is derived, an icosahedral carboranate anion, HCB 11 H − 11 , was first synthesized at DuPont in 1967 by Walter Knoth. Research into this molecule's properties was put on hiatus until the mid 1980s when the Czech group of boron scientists, Plešek, Štíbr, and Heřmánek improved the process for halogenation of carborane molecules. These findings were instrumental in developing the current procedure for carborane acid synthesis. [ 16 ] [ 18 ] The process consists of treating Cs + [HCB 11 H 11 ] − with SO 2 Cl 2 , refluxing under dry argon to fully chlorinate the molecule yielding carborane acid, but this has been shown to fully chlorinate only under select conditions. [ 19 ] [ 16 ] [ 20 ]
In 2010, Reed published a guide giving detailed procedures for the synthesis of carborane acids and their derivatives. [ 21 ] Nevertheless, the synthesis of carborane acids remains lengthy and difficult and requires a well-maintained glovebox and some specialized equipment. The starting material is commercially available decaborane(14) , a highly toxic substance. The most well-studied carborane acid H(CHB 11 Cl 11 ) is prepared in 13 steps. The last few steps are especially sensitive and require a glovebox at < 1 ppm H 2 O without any weakly basic solvent vapors, since bases as weak as benzene or dichloromethane will react with carborane-based electrophiles and Brønsted acids. The final step of the synthesis is the metathesis of the μ-hydridodisilylium carboranate salt with excess liquid, anhydrous hydrogen chloride, presumably driven by the formation of strong Si–Cl and H–H bonds in the volatile byproducts:
The product was isolated by evaporation of the byproducts and was characterized by its infrared (ν CH = 3023 cm −1 ) and nuclear magnetic resonance (δ 4.55 (s, 1H, CH), 20.4 (s, 1H, H + ) in liquid SO 2 ) spectra (note the extremely downfield chemical shift of the acidic proton). [ 21 ] Although the reactions used in the synthesis are analogous, obtaining a pure sample of the more acidic H(CHB 11 F 11 ) turned out to be even more difficult, requiring extremely rigorous procedures to exclude traces of weakly basic impurities. [ 7 ]
Carborane acid consists of 11 boron atoms; each boron atom is bound to a chlorine atom. The chlorine atoms serve to enhance acidity and act as shields against attacks from the outside due to the steric hindrance they form around the cluster. The cluster, consisting of the 11 borons, 11 chlorines, and a single carbon atom, is paired with a hydrogen atom, bound to the carbon atom. The boron and carbon atoms are allowed to form six bonds due to boron's ability to form three-center, two-electron bonds. [ 18 ]
Although the structure of the carborane acid differs greatly from conventional acids, both distribute charge and stability in a similar fashion. The carboranate anion distributes its charge by delocalizing the electrons throughout the 12 cage atoms. [ 22 ] This was shown in a single crystal X-ray diffraction study revealing shortened bond lengths in the heterocyclic portion of the ring suggesting electronic delocalization. [ 23 ]
The chlorinated carba- closo -dodecaborate anion HCB 11 Cl − 11 is an outstandingly stable anion with what has previously been described as "substitutionally inert" B–Cl vertices.
The descriptor closo indicates that the molecule is formally derived (by B-to-C + replacement) from a borane of stoichiometry and charge [B n H n ] 2− ( n = 12 for known carborane acids). [ 24 ] The cagelike structure formed by the 11 boron atoms and 1 carbon atom allows the electrons to be highly delocalized through the 3D cage (the special stabilization of the carborane system has been termed "σ-aromaticity"), and the high energy required to disrupt the boron cluster portion of the molecule is what gives the anion its remarkable stability. [ 24 ] Because the anion is extremely stable, it will not behave as a nucleophile toward the protonated substrate, while the acid itself is completely non-oxidizing, unlike the Lewis acidic components of many superacids like antimony pentafluoride. Hence, sensitive molecules like C 60 can be protonated without decomposition. [ 25 ] [ 26 ]
There are many proposed applications for the boron-based carborane acids. For instance, they have been proposed as catalysts for hydrocarbon cracking and isomerization of n -alkanes to form branched isoalkanes ("isooctane", for example). Carborane acids may also be used as strong, selective Brønsted acids for fine chemical synthesis, where the low nucleophilicity of the counteranion may be advantageous. In mechanistic organic chemistry, they may be used in the study of reactive cationic intermediates. [ 27 ] In inorganic synthesis, their unparalleled acidity may allow for the isolation of exotic species like salts of protonated xenon. [ 17 ] [ 18 ] [ 28 ] | https://en.wikipedia.org/wiki/H(CHB11F11) |
The helium hydride ion , hydridohelium(1+) ion , or helonium is a cation ( positively charged ion ) with chemical formula HeH + . It consists of a helium atom bonded to a hydrogen atom, with one electron removed. It can also be viewed as protonated helium. It is the lightest heteronuclear ion, and is believed to be the first compound formed in the Universe after the Big Bang . [ 3 ]
The ion was first produced in a laboratory in 1925. It is stable in isolation, but extremely reactive, and cannot be prepared in bulk, because it would react with any other molecule with which it came into contact. Noted as the strongest known acid —stronger than even fluoroantimonic acid —its occurrence in the interstellar medium had been conjectured since the 1970s, [ 4 ] and it was finally detected in April 2019 using the airborne SOFIA telescope . [ 5 ] [ 6 ]
The helium hydrogen ion is isoelectronic with molecular hydrogen ( H 2 ). [ 7 ]
Unlike the dihydrogen ion H + 2 , the helium hydride ion has a permanent dipole moment , which makes its spectroscopic characterization easier. [ 8 ] The calculated dipole moment of HeH + is 2.26 or 2.84 D . [ 9 ] The electron density in the ion is higher around the helium nucleus than the hydrogen. 80% of the electron charge is closer to the helium nucleus than to the hydrogen nucleus. [ 10 ]
Spectroscopic detection is hampered, because one of its most prominent spectral lines, at 149.14 μm , coincides with a doublet of spectral lines belonging to the methylidyne radical ⫶ CH. [ 3 ]
The length of the covalent bond in the ion is 0.772 Å [ 11 ] or 77.2 pm .
The helium hydride ion has six relatively stable isotopologues , that differ in the isotopes of the two elements, and hence in the total atomic mass number ( A ) and the total number of neutrons ( N ) in the two nuclei:
They all have three protons and two electrons. The first three are generated by radioactive decay of tritium in the molecules HT = 1 H 3 H , DT = 2 H 3 H , and T 2 = 3 H 2 , respectively. The last three can be generated by ionizing the appropriate isotopologue of H 2 in the presence of helium-4. [ 7 ]
The following isotopologues of the helium hydride ion, of the dihydrogen ion H + 2 , and of the trihydrogen ion H + 3 have the same total atomic mass number A :
The masses in each row above are not equal, though, because the binding energies in the nuclei are different. [ 16 ]
Unlike the helium hydride ion, the neutral helium hydride molecule HeH is not stable in the ground state. However, it does exist in an excited state as an excimer (HeH*), and its spectrum was first observed in the mid-1980s. [ 19 ] [ 20 ] [ 21 ]
The neutral molecule is the first entry in the Gmelin database . [ 4 ]
Since HeH + reacts with every substance, it cannot be stored in any container. As a result, its chemistry must be studied by creating it in situ .
Reactions with organic substances can be studied by substituting hydrogen in the desired organic compound with tritium . The decay of tritium to 3 He + followed by its extraction of a hydrogen atom from the compound yields 3 HeH + , which is then surrounded by the organic material and will in turn react. [ 22 ] [ 23 ]
HeH + cannot be prepared in a condensed phase , as it would donate a proton to any anion , molecule or atom that it came in contact with. It has been shown to protonate O 2 , NH 3 , SO 2 , H 2 O , and CO 2 , giving HO + 2 , NH + 4 , HSO + 2 , H 3 O + , and HCO + 2 respectively. [ 22 ] Other molecules such as nitric oxide , nitrogen dioxide , nitrous oxide , hydrogen sulfide , methane , acetylene , ethylene , ethane , methanol and acetonitrile react but break up due to the large amount of energy produced. [ 22 ]
In fact, HeH + is the strongest known acid , with a proton affinity of 177.8 kJ/mol, or a p K a of −63. [ 24 ]
Additional helium atoms can attach to HeH + to form larger clusters such as He 2 H + , He 3 H + , He 4 H + , He 5 H + and He 6 H + . [ 22 ]
The dihelium hydride cation, He 2 H + , is formed by the reaction of dihelium cation with molecular hydrogen:
It is a linear ion with hydrogen in the centre. [ 22 ]
The hexahelium hydride ion, He 6 H + , is particularly stable. [ 22 ]
Other helium hydride ions are known or have been studied theoretically. Helium dihydride ion, or dihydridohelium(1+) , HeH + 2 , has been observed using microwave spectroscopy. [ 25 ] It has a calculated binding energy of 25.1 kJ/mol, while trihydridohelium(1+) , HeH + 3 , has a calculated binding energy of 0.42 kJ/mol. [ 26 ]
Hydridohelium(1+), specifically [ 4 He 1 H] + , was first detected indirectly in 1925 by T. R. Hogness and E. G. Lunn. They were injecting protons of known energy into a rarefied mixture of hydrogen and helium, in order to study the formation of hydrogen ions like H + , H + 2 and H + 3 . They observed that H + 3 appeared at the same beam energy (16 eV ) as H + 2 , and its concentration increased with pressure much more than that of the other two ions. From these data, they concluded that the H + 2 ions were transferring a proton to molecules that they collided with, including helium. [ 7 ]
In 1933, K. Bainbridge used mass spectrometry to compare the masses of the ions [ 4 He 1 H] + (helium hydride ion) and [ 2 H 2 1 H] + (twice-deuterated trihydrogen ion) in order to obtain an accurate measurement of the atomic mass of deuterium relative to that of helium. Both ions have 3 protons, 2 neutrons, and 2 electrons. He also compared [ 4 He 2 H] + (helium deuteride ion) with [ 2 H 3 ] + ( trideuterium ion), both with 3 protons and 3 neutrons. [ 16 ]
The first attempt to compute the structure of the HeH + ion (specifically, [ 4 He 1 H] + ) by quantum mechanical theory was made by J. Beach in 1936. [ 27 ] Improved computations were sporadically published over the next decades. [ 28 ] [ 29 ]
H. Schwartz observed in 1955 that the decay of the tritium molecule T 2 = 3 H 2 should generate the helium hydride ion [ 3 HeT] + with high probability.
In 1963, F. Cacace at the Sapienza University of Rome conceived the decay technique for preparing and studying organic radicals and carbenium ions. [ 30 ] In a variant of that technique, exotic species like methanium are produced by reacting organic compounds with the [ 3 HeT] + that is produced by the decay of T 2 that is mixed with the desired reagents. Much of what we know about the chemistry of [HeH] + came through this technique. [ 31 ]
In 1980, V. Lubimov (Lyubimov) at the ITEP laboratory in Moscow claimed to have detected a mildly significant rest mass (30 ± 16) eV for the neutrino , by analyzing the energy spectrum of the β decay of tritium. [ 32 ] The claim was disputed, and several other groups set out to check it by studying the decay of molecular tritium T 2 . It was known that some of the energy released by that decay would be diverted to the excitation of the decay products, including [ 3 HeT] + ; and this phenomenon could be a significant source of error in that experiment. This observation motivated numerous efforts to precisely compute the expected energy states of that ion in order to reduce the uncertainty of those measurements. [ citation needed ] Many have improved the computations since then, and now there is quite good agreement between computed and experimental properties; including for the isotopologues [ 4 He 2 H] + , [ 3 He 1 H] + , and [ 3 He 2 H] + . [ 18 ] [ 13 ]
In 1956, M. Cantwell predicted theoretically that the spectrum of vibrations of that ion should be observable in the infrared; and the spectra of the deuterium and common hydrogen isotopologues ( [ 3 HeD] + and [ 3 He 1 H] + ) should lie closer to visible light and hence easier to observe. [ 12 ] The first detection of the spectrum of [ 4 He 1 H] + was made by D. Tolliver and others in 1979, at wavenumbers between 1,700 and 1,900 cm −1 . [ 33 ] In 1982, P. Bernath and T. Amano detected nine infrared lines between 2,164 and 3,158 waves per cm. [ 17 ]
HeH + has been conjectured since the 1970s to exist in the interstellar medium . [ 34 ] Its first detection, in the nebula NGC 7027 , was reported in an article published in the journal Nature in April 2019. [ 5 ]
The helium hydride ion is formed during the decay of tritium in the molecule HT or tritium molecule T 2 . Although excited by the recoil from the beta decay, the molecule remains bound together. [ 35 ]
It is believed to be the first compound to have formed in the universe, [ 3 ] and is of fundamental importance in understanding the chemistry of the early universe. [ 36 ] This is because hydrogen and helium were almost the only types of atoms formed in Big Bang nucleosynthesis . Stars formed from the primordial material should contain HeH + , which could influence their formation and subsequent evolution. In particular, its strong dipole moment makes it relevant to the opacity of zero-metallicity stars . [ 3 ] HeH + is also thought to be an important constituent of the atmospheres of helium-rich white dwarfs, where it increases the opacity of the gas and causes the star to cool more slowly. [ 37 ]
HeH + could be formed in the cooling gas behind dissociative shocks in dense interstellar clouds, such as the shocks caused by stellar winds , supernovae and outflowing material from young stars. If the speed of the shock is greater than about 90 kilometres per second (56 mi/s), quantities large enough to detect might be formed. If detected, the emissions from HeH + would then be useful tracers of the shock. [ 38 ]
Several locations had been suggested as possible places HeH + might be detected. These included cool helium stars , [ 3 ] H II regions , [ 39 ] and dense planetary nebulae , [ 39 ] like NGC 7027 , [ 36 ] where, in April 2019, HeH + was reported to have been detected. [ 5 ] | https://en.wikipedia.org/wiki/H-He+ |
Diimide , also called diazene or diimine , is a compound having the formula HN=NH. It exists as two geometric isomers , E ( trans ) and Z ( cis ). The term diazene is more common for organic derivatives of diimide. Thus, azobenzene is an example of an organic diazene.
A traditional route to diimide involves oxidation of hydrazine with hydrogen peroxide or air. [ 1 ]
Alternatively the hydrolysis of diethyl azodicarboxylate or azodicarbonamide affords diimide: [ 2 ]
Nowadays, diimide is generated by thermal decomposition of 2,4,6‐triisopropylbenzenesulfonylhydrazide. [ 3 ]
Because of its instability, diimide is generated and used in-situ . A mixture of both the cis ( Z- ) and trans ( E- ) isomers is produced. Both isomers are unstable, and they undergo a slow interconversion. The trans isomer is more stable, but the cis isomer is the one that reacts with unsaturated substrates, therefore the equilibrium between them shifts towards the cis isomer due to Le Chatelier's principle . Some procedures call for the addition of carboxylic acids, which catalyse the cis–trans isomerization. [ 4 ] Diimide decomposes readily. Even at low temperatures, the more stable trans isomer rapidly undergoes various disproportionation reactions, primarily forming hydrazine and nitrogen gas : [ 5 ]
Because of this competing decomposition reaction, reductions with diimide typically require a large excess of the precursor reagent.
Diimide is occasionally useful as a reagent in organic synthesis . [ 4 ] It hydrogenates alkenes and alkynes with selective delivery of hydrogen from one face of the substrate resulting in the same stereoselectivity as metal-catalysed syn addition of H 2 . The only coproduct released is nitrogen gas. Although the method is cumbersome, the use of diimide avoids the need for high pressures or hydrogen gas and metal catalysts, which can be expensive. [ 6 ] The hydrogenation mechanism involves a six-membered C 2 H 2 N 2 transition state:
Diimide is advantageous because it selectively reduces alkenes and alkynes and is unreactive toward many functional groups that would interfere with normal catalytic hydrogenation . Thus, peroxides , alkyl halides , and thiols are tolerated by diimide, but these same groups would typically be degraded by metal catalysts. The reagent preferentially reduces alkynes and unhindered or strained alkenes [ 1 ] to the corresponding alkenes and alkanes. [ 4 ]
The dicationic form, H−N + ≡N + −H (diazynediium, diprotonated dinitrogen), is calculated to have the strongest known chemical bond. This ion can be thought of as a doubly protonated nitrogen molecule. The relative bond strength order (RBSO) is 3.38. [ 7 ] F−N + ≡N + −H (fluorodiazynediium ion) and F−N + ≡N + −F (difluorodiazynediium ion) have slightly lower strength bonds. [ 7 ]
In the presence of strong bases, diimide deprotonates to form the pernitride anion, N − =N − . | https://en.wikipedia.org/wiki/H-N=N-H |
Carbonic acid is a chemical compound with the chemical formula H 2 C O 3 . The molecule rapidly converts to water and carbon dioxide in the presence of water. However, in the absence of water, it is quite stable at room temperature . [ 5 ] [ 6 ] The interconversion of carbon dioxide and carbonic acid is related to the breathing cycle of animals and the acidification of natural waters . [ 4 ]
In biochemistry and physiology, the name "carbonic acid" is sometimes applied to aqueous solutions of carbon dioxide . These chemical species play an important role in the bicarbonate buffer system , used to maintain acid–base homeostasis . [ 7 ]
In chemistry , the term "carbonic acid" strictly refers to the chemical compound with the formula H 2 CO 3 . Some biochemistry literature effaces the distinction between carbonic acid and carbon dioxide dissolved in extracellular fluid.
In physiology , carbon dioxide excreted by the lungs may be called volatile acid or respiratory acid .
At ambient temperatures, pure carbonic acid is a stable gas. [ 6 ] There are two main methods to produce anhydrous carbonic acid: reaction of hydrogen chloride and potassium bicarbonate at 100 K in methanol and proton irradiation of pure solid carbon dioxide . [ 3 ] Chemically, it behaves as a diprotic Brønsted acid . [ 8 ] [ 9 ]
Carbonic acid monomers exhibit three conformational isomers : cis–cis, cis–trans, and trans–trans. [ 10 ]
At low temperatures and atmospheric pressure , solid carbonic acid is amorphous and lacks Bragg peaks in X-ray diffraction . [ 11 ] But at high pressure, carbonic acid crystallizes, and modern analytical spectroscopy can measure its geometry.
According to neutron diffraction of dideuterated carbonic acid ( D 2 CO 3 ) in a hybrid clamped cell ( Russian alloy / copper-beryllium ) at 1.85 GPa, the molecules are planar and form dimers joined by pairs of hydrogen bonds . All three C-O bonds are nearly equidistant at 1.34 Å , intermediate between typical C-O and C=O distances (respectively 1.43 and 1.23 Å). The unusual C-O bond lengths are attributed to delocalized π bonding in the molecule's center and extraordinarily strong hydrogen bonds. The same effects also induce a very short O—O separation (2.13 Å), through the 136° O-H-O angle imposed by the doubly hydrogen-bonded 8-membered rings. [ 4 ] Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid , where the distances exceed 2.4 Å. [ 11 ]
In even a slight presence of water, carbonic acid dehydrates to carbon dioxide and water , which then catalyzes further decomposition. [ 6 ] For this reason, carbon dioxide can be considered the carbonic acid anhydride .
The hydration equilibrium constant at 25 °C is [ H 2 CO 3 ]/[CO 2 ] ≈ 1.7×10 −3 in pure water [ 12 ] and ≈ 1.2×10 −3 in seawater . [ 13 ] Hence the majority of carbon dioxide at geophysical or biological air-water interfaces does not convert to carbonic acid, remaining dissolved CO 2 gas. However, the uncatalyzed equilibrium is reached quite slowly: the rate constants are 0.039 s −1 for hydration and 23 s −1 for dehydration.
In the presence of the enzyme carbonic anhydrase , equilibrium is instead reached rapidly, and the following reaction takes precedence: [ 14 ] HCO 3 − + H + ↽ − − ⇀ CO 2 + H 2 O {\displaystyle {\ce {HCO3^- {+}H^+ <=> CO2 {+}H2O}}}
When the created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium CO 2 ( soln ) ↽ − − ⇀ CO 2 ( g ) {\displaystyle {\ce {CO_2 (soln) <=> CO_2 (g)}}} must also be taken into consideration. The equilibrium constant for this reaction is defined by Henry's law .
The two reactions can be combined for the equilibrium in solution: HCO 3 − + H + ↽ − − ⇀ CO 2 ( soln ) + H 2 O K 3 = [ H + ] [ HCO 3 − ] [ CO 2 ( soln ) ] {\displaystyle {\begin{aligned}{\ce {HCO3^{-}{}+ H+{}<=> CO2(soln){}+ H2O}}&&K_{3}={\frac {[{\ce {H+}}][{\ce {HCO3^-}}]}{[{\ce {CO2(soln)}}]}}\end{aligned}}} When Henry's law is used to calculate the denominator care is needed with regard to units since Henry's law constant can be commonly expressed with 8 different dimensionalities. [ 15 ]
In wastewater treatment and agriculture irrigation, carbonic acid is used to acidify the water similar to sulfuric acid and sulfurous acid produced by sulfur burners. [ 16 ]
In the beverage industry , sparkling or "fizzy water" is usually referred to as carbonated water . It is made by dissolving carbon dioxide under a small positive pressure in water. Many soft drinks treated the same way effervesce .
Significant amounts of molecular H 2 CO 3 exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres) in planetary interiors. [ 17 ] [ 18 ] Pressures of 0.6–1.6 GPa at 100 K , and 0.75–1.75 GPa at 300 K are attained in the cores of large icy satellites such as Ganymede , Callisto , and Titan , where water and carbon dioxide are present. Pure carbonic acid, being denser, is expected to have sunk under the ice layers and separate them from the rocky cores of these moons. [ 19 ]
Carbonic acid is the formal Brønsted–Lowry conjugate acid of the bicarbonate anion, stable in alkaline solution . The protonation constants have been measured to great precision, but depend on overall ionic strength I . The two equilibria most easily measured are as follows: CO 3 2 − + H + ↽ − − ⇀ HCO 3 − β 1 = [ HCO 3 − ] [ H + ] [ CO 3 2 − ] CO 3 2 − + 2 H + ↽ − − ⇀ H 2 CO 3 β 2 = [ H 2 CO 3 ] [ H + ] 2 [ CO 3 2 − ] {\displaystyle {\begin{aligned}{\ce {CO3^{2-}{}+ H+{}<=> HCO3^-}}&&\beta _{1}={\frac {[{\ce {HCO3^-}}]}{[{\ce {H+}}][{\ce {CO3^{2-}}}]}}\\{\ce {CO3^{2-}{}+ 2H+{}<=> H2CO3}}&&\beta _{2}={\frac {[{\ce {H2CO3}}]}{[{\ce {H+}}]^{2}[{\ce {CO3^{2-}}}]}}\end{aligned}}} where brackets indicate the concentration of species . At 25 °C, these equilibria empirically satisfy [ 20 ] log ( β 1 ) = 0 .54 I 2 − 0 .96 I + 9 .93 log ( β 2 ) = − 2 .5 I 2 − 0 .043 I + 16 .07 {\displaystyle {\begin{alignedat}{6}\log(\beta _{1})=&&0&.54&I^{2}-0&.96&I+&&9&.93\\\log(\beta _{2})=&&-2&.5&I^{2}-0&.043&I+&&16&.07\end{alignedat}}} log( β 1 ) decreases with increasing I , as does log( β 2 ) . In a solution absent other ions (e.g. I = 0 ), these curves imply the following stepwise dissociation constants : p K 1 = log ( β 2 ) − log ( β 1 ) = 6.77 p K 2 = log ( β 1 ) = 9.93 {\displaystyle {\begin{alignedat}{3}p{\text{K}}_{1}&=\log(\beta _{2})-\log(\beta _{1})&=6.77\\p{\text{K}}_{2}&=\log(\beta _{1})&=9.93\end{alignedat}}} Direct values for these constants in the literature include p K 1 = 6.35 and p K 2 - p K 1 = 3.49 . [ 21 ]
To interpret these numbers, note that two chemical species in an acid equilibrium are equiconcentrated when p K = p H . In particular, the extracellular fluid ( cytosol ) in biological systems exhibits p H ≈ 7.2 , so that carbonic acid will be almost 50%-dissociated at equilibrium.
The Bjerrum plot shows typical equilibrium concentrations, in solution, in seawater , of carbon dioxide and the various species derived from it, as a function of pH . [ 8 ] [ 9 ] As human industrialization has increased the proportion of carbon dioxide in Earth's atmosphere , the proportion of carbon dioxide dissolved in sea- and freshwater as carbonic acid is also expected to increase. This rise in dissolved acid is also expected to acidify those waters, generating a decrease in pH. [ 22 ] [ 23 ] It has been estimated that the increase in dissolved carbon dioxide has already caused the ocean's average surface pH to decrease by about 0.1 from pre-industrial levels. | https://en.wikipedia.org/wiki/H-O-COOH |
In statistical mechanics of continuous systems, a potential for a many-body system is called H-stable (or simply stable ) if the potential energy per particle is bounded below by a constant that is independent of the total number of particles. In many circumstances, if a potential is not H-stable, it is not possible to define a grand canonical partition function in finite volume, because of catastrophic configurations with infinite particles located in a finite space.
Consider a system of particles in positions x 1 , x 2 , … ∈ R ν {\displaystyle x_{1},x_{2},\ldots \in R^{\nu }} ; the interaction or potential between a particle in position x i {\displaystyle x_{i}} and a particle in position x j {\displaystyle x_{j}} is
where ϕ ( x ) {\displaystyle \phi (x)} is a real, even (possibly unbounded) function. Then ϕ ( x ) {\displaystyle \phi (x)} is H-stable if there exists B > 0 {\displaystyle B>0} such that, for any n ≥ 1 {\displaystyle n\geq 1} and any x 1 , x 2 , … , x n ∈ R ν {\displaystyle x_{1},x_{2},\ldots ,x_{n}\in R^{\nu }} ,
The notion of H-stability in quantum mechanics is more subtle.
While in the classical case the kinetic part of the Hamiltonian is not important as it can be zero independently of the position of the particles, in the quantum case the kinetic term plays an important role in the lower bound for the total energy because of the uncertainty principle . (In fact, stability of matter was the historical reason for introducing such a principle in mechanics.)
The definition of stability is :
where E 0 is the ground state energy.
Classical H-stability implies quantum H-stability, but the converse is false.
The criterion is especially useful in statistical mechanics , where H-stability is necessary to the existence of thermodynamics , i.e. if a system is not H-stable, the thermodynamic limit does not exist. | https://en.wikipedia.org/wiki/H-stable_potential |
In classical statistical mechanics , the H -theorem , introduced by Ludwig Boltzmann in 1872, describes the tendency of the quantity H (defined below) to decrease in a nearly- ideal gas of molecules. [ 1 ] As this quantity H was meant to represent the entropy of thermodynamics, the H -theorem was an early demonstration of the power of statistical mechanics as it claimed to derive the second law of thermodynamics —a statement about fundamentally irreversible processes —from reversible microscopic mechanics. It is thought to prove the second law of thermodynamics , [ 2 ] [ 3 ] [ 4 ] albeit under the assumption of low-entropy initial conditions. [ 5 ]
The H -theorem is a natural consequence of the kinetic equation derived by Boltzmann that has come to be known as Boltzmann's equation . The H -theorem has led to considerable discussion about its actual implications, [ 6 ] with major themes being:
Boltzmann in his original publication writes the symbol E (as in entropy ) for its statistical function . [ 1 ] Years later, Samuel Hawksley Burbury , one of the critics of the theorem, [ 7 ] wrote the function with the symbol H, [ 8 ] a notation that was subsequently adopted by Boltzmann when referring to his "H- theorem". [ 9 ] The notation has led to some confusion regarding the name of the theorem. Even though the statement is usually referred to as the " Aitch theorem " , sometimes it is instead called the " Eta theorem", as the capital Greek letter Eta ( Η ) is indistinguishable from the capital version of Latin letter h ( H ) . [ 10 ] Discussions have been raised on how the symbol should be understood, but it remains unclear due to the lack of written sources from the time of the theorem. [ 10 ] [ 11 ] Studies of the typography and the work of J.W. Gibbs [ 12 ] seem to favour the interpretation of H as Eta . [ 13 ]
The H value is determined from the function f ( E , t ) dE , which is the energy distribution function of molecules at time t . The value f ( E , t ) dE is the number of molecules that have kinetic energy between E and E + dE . H itself is defined as
For an isolated ideal gas (with fixed total energy and fixed total number of particles), the function H is at a minimum when the particles have a Maxwell–Boltzmann distribution ; if the molecules of the ideal gas are distributed in some other way (say, all having the same kinetic energy), then the value of H will be higher. Boltzmann's H -theorem, described in the next section, shows that when collisions between molecules are allowed, such distributions are unstable and tend to irreversibly seek towards the minimum value of H (towards the Maxwell–Boltzmann distribution).
(Note on notation: Boltzmann originally used the letter E for quantity H ; most of the literature after Boltzmann uses the letter H as here. Boltzmann also used the symbol x to refer to the kinetic energy of a particle.)
Boltzmann considered what happens during the collision between two particles. It is a basic fact of mechanics that in the elastic collision between two particles (such as hard spheres), the energy transferred between the particles varies depending on initial conditions (angle of collision, etc.).
Boltzmann made a key assumption known as the Stosszahlansatz ( molecular chaos assumption), that during any collision event in the gas, the two particles participating in the collision have 1) independently chosen kinetic energies from the distribution, 2) independent velocity directions, 3) independent starting points. Under these assumptions, and given the mechanics of energy transfer, the energies of the particles after the collision will obey a certain new random distribution that can be computed.
Considering repeated uncorrelated collisions, between any and all of the molecules in the gas, Boltzmann constructed his kinetic equation ( Boltzmann's equation ). From this kinetic equation, a natural outcome is that the continual process of collision causes the quantity H to decrease until it has reached a minimum.
Although Boltzmann's H -theorem turned out not to be the absolute proof of the second law of thermodynamics as originally claimed (see Criticisms below), the H -theorem led Boltzmann in the last years of the 19th century to more and more probabilistic arguments about the nature of thermodynamics. The probabilistic view of thermodynamics culminated in 1902 with Josiah Willard Gibbs 's statistical mechanics for fully general systems (not just gases), and the introduction of generalized statistical ensembles .
The kinetic equation and in particular Boltzmann's molecular chaos assumption inspired a whole family of Boltzmann equations that are still used today to model the motions of particles, such as the electrons in a semiconductor. In many cases the molecular chaos assumption is highly accurate, and the ability to discard complex correlations between particles makes calculations much simpler.
The process of thermalisation can be described using the H-theorem or the relaxation theorem . [ 14 ]
There are several notable reasons described below why the H -theorem, at least in its original 1871 form, is not completely rigorous. As Boltzmann would eventually go on to admit, the arrow of time in the H -theorem is not in fact purely mechanical, but really a consequence of assumptions about initial conditions. [ 15 ]
Soon after Boltzmann published his H theorem, Johann Josef Loschmidt objected that it should not be possible to deduce an irreversible process from time-symmetric dynamics and a time-symmetric formalism. If the H decreases over time in one state, then there must be a matching reversed state where H increases over time ( Loschmidt's paradox ). The explanation is that Boltzmann's equation is based on the assumption of " molecular chaos ", i.e., that it follows from, or at least is consistent with, the underlying kinetic model that the particles be considered independent and uncorrelated. It turns out that this assumption breaks time reversal symmetry in a subtle sense, and therefore begs the question . Once the particles are allowed to collide, their velocity directions and positions in fact do become correlated (however, these correlations are encoded in an extremely complex manner). This shows that an (ongoing) assumption of independence is not consistent with the underlying particle model.
Boltzmann's reply to Loschmidt was to concede the possibility of these states, but noting that these sorts of states were so rare and unusual as to be impossible in practice. Boltzmann would go on to sharpen this notion of the "rarity" of states, resulting in his entropy formula of 1877.
As a demonstration of Loschmidt's paradox, a modern counterexample (not to Boltzmann's original gas-related H -theorem, but to a closely related analogue) is the phenomenon of spin echo . [ 16 ] In the spin echo effect, it is physically possible to induce time reversal in an interacting system of spins.
An analogue to Boltzmann's H for the spin system can be defined in terms of the distribution of spin states in the system. In the experiment, the spin system is initially perturbed into a non-equilibrium state (high H ), and, as predicted by the H theorem the quantity H soon decreases to the equilibrium value. At some point, a carefully constructed electromagnetic pulse is applied that reverses the motions of all the spins. The spins then undo the time evolution from before the pulse, and after some time the H actually increases away from equilibrium (once the evolution has completely unwound, the H decreases once again to the minimum value). In some sense, the time reversed states noted by Loschmidt turned out to be not completely impractical.
In 1896, Ernst Zermelo noted a further problem with the H theorem, which was that if the system's H is at any time not a minimum, then by Poincaré recurrence , the non-minimal H must recur (though after some extremely long time). Boltzmann admitted that these recurring rises in H technically would occur, but pointed out that, over long times, the system spends only a tiny fraction of its time in one of these recurring states.
The second law of thermodynamics states that the entropy of an isolated system always increases to a maximum equilibrium value. This is strictly true only in the thermodynamic limit of an infinite number of particles. For a finite number of particles, there will always be entropy fluctuations. For example, in the fixed volume of the isolated system, the maximum entropy is obtained when half the particles are in one half of the volume, half in the other, but sometimes there will be temporarily a few more particles on one side than the other, and this will constitute a very small reduction in entropy. These entropy fluctuations are such that the longer one waits, the larger an entropy fluctuation one will probably see during that time, and the time one must wait for a given entropy fluctuation is always finite, even for a fluctuation to its minimum possible value. For example, one might have an extremely low entropy condition of all particles being in one half of the container. The gas will quickly attain its equilibrium value of entropy, but given enough time, this same situation will happen again. For practical systems, e.g. a gas in a 1-liter container at room temperature and atmospheric pressure, this time is truly enormous, many multiples of the age of the universe, and, practically speaking, one can ignore the possibility.
Since H is a mechanically defined variable that is not conserved, then like any other such variable (pressure, etc.) it will show thermal fluctuations . This means that H regularly shows spontaneous increases from the minimum value. Technically this is not an exception to the H theorem, since the H theorem was only intended to apply for a gas with a very large number of particles. These fluctuations are only perceptible when the system is small and the time interval over which it is observed is not enormously large.
If H is interpreted as entropy as Boltzmann intended, then this can be seen as a manifestation of the fluctuation theorem . [ citation needed ]
H is a forerunner of Shannon's information entropy . Claude Shannon denoted his measure of information entropy H after the H-theorem. [ 17 ] The article on Shannon's information entropy contains an explanation of the discrete counterpart of the quantity H , known as the information entropy or information uncertainty (with a minus sign). By extending the discrete information entropy to the continuous information entropy , also called differential entropy , one obtains the expression in the equation from the section above, Definition and Meaning of Boltzmann's H , and thus a better feel for the meaning of H .
The H -theorem's connection between information and entropy plays a central role in a recent controversy called the Black hole information paradox .
Richard C. Tolman 's 1938 book The Principles of Statistical Mechanics dedicates a whole chapter to the study of Boltzmann's H theorem, and its extension in the generalized classical statistical mechanics of Gibbs . A further chapter is devoted to the quantum mechanical version of the H -theorem.
We let q i and p i be our generalized canonical coordinates for a set of r {\displaystyle r} particles. Then we consider a function f {\displaystyle f} that returns the probability density of particles, over the states in phase space . Note how this can be multiplied by a small region in phase space, denoted by δ q 1 . . . δ p r {\displaystyle \delta q_{1}...\delta p_{r}} , to yield the (average) expected number of particles in that region.
Tolman offers the following equations for the definition of the quantity H in Boltzmann's original H theorem.
Here we sum over the regions into which phase space is divided, indexed by i {\displaystyle i} . And in the limit for an infinitesimal phase space volume δ q i → 0 , δ p i → 0 ∀ i {\displaystyle \delta q_{i}\rightarrow 0,\delta p_{i}\rightarrow 0\;\forall \,i} , we can write the sum as an integral.
H can also be written in terms of the number of molecules present in each of the cells.
An additional way to calculate the quantity H is:
where P is the probability of finding a system chosen at random from the specified microcanonical ensemble . It can finally be written as:
where G is the number of classical states. [ clarification needed ]
The quantity H can also be defined as the integral over velocity space [ citation needed ] :
H = d e f ∫ P ( ln P ) d 3 v = ⟨ ln P ⟩ {\displaystyle \displaystyle H\ {\stackrel {\mathrm {def} }{=}}\ \int {P({\ln P})\,d^{3}v}=\left\langle \ln P\right\rangle }
where P ( v ) is the probability distribution.
Using the Boltzmann equation one can prove that H can only decrease.
For a system of N statistically independent particles, H is related to the thermodynamic entropy S through: [ 23 ]
So, according to the H -theorem, S can only increase.
In quantum statistical mechanics (which is the quantum version of classical statistical mechanics), the H-function is the function: [ 24 ]
where summation runs over all possible distinct states of the system, and p i is the probability that the system could be found in the i -th state.
This is closely related to the entropy formula of Gibbs ,
and we shall (following e.g., Waldram (1985), p. 39) proceed using S rather than H .
First, differentiating with respect to time gives
(using the fact that Σ dp i / dt = 0, since Σ p i = 1, so the second term vanishes. We will see later that it will be useful to break this into two sums.)
Now Fermi's golden rule gives a master equation for the average rate of quantum jumps from state α to β; and from state β to α. (Of course, Fermi's golden rule itself makes certain approximations, and the introduction of this rule is what introduces irreversibility. It is essentially the quantum version of Boltzmann's Stosszahlansatz .) For an isolated system the jumps will make contributions
where the reversibility of the dynamics ensures that the same transition constant ν αβ appears in both expressions.
So
The two differences terms in the summation always have the same sign. For example:
then
so overall the two negative signs will cancel.
Therefore,
for an isolated system.
The same mathematics is sometimes used to show that relative entropy is a Lyapunov function of a Markov process in detailed balance , and other chemistry contexts.
Josiah Willard Gibbs described another way in which the entropy of a microscopic system would tend to increase over time. [ 25 ] Later writers have called this "Gibbs' H -theorem" as its conclusion resembles that of Boltzmann's. [ 26 ] Gibbs himself never called it an H -theorem, and in fact his definition of entropy—and mechanism of increase—are very different from Boltzmann's. This section is included for historical completeness.
The setting of Gibbs' entropy production theorem is in ensemble statistical mechanics, and the entropy quantity is the Gibbs entropy (information entropy) defined in terms of the probability distribution for the entire state of the system. This is in contrast to Boltzmann's H defined in terms of the distribution of states of individual molecules, within a specific state of the system.
Gibbs considered the motion of an ensemble which initially starts out confined to a small region of phase space, meaning that the state of the system is known with fair precision though not quite exactly (low Gibbs entropy). The evolution of this ensemble over time proceeds according to Liouville's equation . For almost any kind of realistic system, the Liouville evolution tends to "stir" the ensemble over phase space, a process analogous to the mixing of a dye in an incompressible fluid. [ 25 ] After some time, the ensemble appears to be spread out over phase space, although it is actually a finely striped pattern, with the total volume of the ensemble (and its Gibbs entropy) conserved. Liouville's equation is guaranteed to conserve Gibbs entropy since there is no random process acting on the system; in principle, the original ensemble can be recovered at any time by reversing the motion.
The critical point of the theorem is thus: If the fine structure in the stirred-up ensemble is very slightly blurred, for any reason, then the Gibbs entropy increases, and the ensemble becomes an equilibrium ensemble. As to why this blurring should occur in reality, there are a variety of suggested mechanisms. For example, one suggested mechanism is that the phase space is coarse-grained for some reason (analogous to the pixelization in the simulation of phase space shown in the figure). For any required finite degree of fineness the ensemble becomes "sensibly uniform" after a finite time. Or, if the system experiences a tiny uncontrolled interaction with its environment, the sharp coherence of the ensemble will be lost. Edwin Thompson Jaynes argued that the blurring is subjective in nature, simply corresponding to a loss of knowledge about the state of the system. [ 27 ] In any case, however it occurs, the Gibbs entropy increase is irreversible provided the blurring cannot be reversed.
The exactly evolving entropy, which does not increase, is known as fine-grained entropy . The blurred entropy is known as coarse-grained entropy . Leonard Susskind analogizes this distinction to the notion of the volume of a fibrous ball of cotton: [ 28 ] On one hand the volume of the fibers themselves is constant, but in another sense there is a larger coarse-grained volume, corresponding to the outline of the ball.
Gibbs' entropy increase mechanism solves some of the technical difficulties found in Boltzmann's H -theorem: The Gibbs entropy does not fluctuate nor does it exhibit Poincare recurrence, and so the increase in Gibbs entropy, when it occurs, is therefore irreversible as expected from thermodynamics. The Gibbs mechanism also applies equally well to systems with very few degrees of freedom, such as the single-particle system shown in the figure. To the extent that one accepts that the ensemble becomes blurred, then, Gibbs' approach is a cleaner proof of the second law of thermodynamics . [ 27 ]
Unfortunately, as pointed out early on in the development of quantum statistical mechanics by John von Neumann and others, this kind of argument does not carry over to quantum mechanics. [ 29 ] In quantum mechanics, the ensemble cannot support an ever-finer mixing process, because of the finite dimensionality of the relevant portion of Hilbert space. Instead of converging closer and closer to the equilibrium ensemble (time-averaged ensemble) as in the classical case, the density matrix of the quantum system will constantly show evolution, even showing recurrences. Developing a quantum version of the H -theorem without appeal to the Stosszahlansatz is thus significantly more complicated. [ 29 ] | https://en.wikipedia.org/wiki/H-theorem |
H. A. Rey (born Hans Augusto Reyersbach ; September 16, 1898 – August 26, 1977) was a German-born American illustrator and author, known best for the series of children's picture books that he and his wife Margret Rey created about Curious George . [ 1 ] [ 2 ]
Hans Augusto Reyersbach was born in Hamburg , German Empire on September 16, 1898. He and his wife, Margret, were both German Jews . They first met in Hamburg at Margret's sister's 16th birthday party. They met again in Brazil , where Rey was working as a salesman of bathtubs and Margret had gone to escape the rise of Nazism in Germany . They got married in 1935 and moved to Paris, France in August of that year. [ 3 ] They lived in Montmartre and fled Paris in June 1940 on bicycles, carrying the Curious George manuscript with them. [ 4 ] [ 5 ] They received visas in Bayonne under instructions from the Portuguese consul in Bordeaux , Aristides de Sousa Mendes , which enabled them to leave Europe via Portugal. [ 6 ]
While in Paris, Rey's animal drawings came to the attention of a French publisher, who commissioned him to write a children's book. The characters in Cecily G. and the Nine Monkeys included an impish monkey named Curious George , and the couple then decided to write a book focused entirely on him. The outbreak of World War II interrupted their work. Being Jews , the Reys decided to flee Paris before the Nazis invaded the city. Hans assembled two bicycles, and they left the city just a few hours before it fell. Among the meager possessions they brought with them was the illustrated manuscript of Curious George . [ 4 ] [ 7 ]
The Reys' odyssey took them to Bayonne, France , where they were issued life-saving visas signed by Portuguese Vice-Consul Manuel Vieira Braga (following instructions from Aristides de Sousa Mendes ) on June 20, 1940. [ 8 ] They crossed the Spanish border, where they bought train tickets to Lisbon . From there, they returned to Brazil, where they had met five years earlier, but this time they continued on to New York. The Reys escaped Europe carrying the manuscript to the first Curious George book, which was published in New York by Houghton Mifflin in 1941. They originally planned to use watercolor illustrations, but since they were responsible for the color separation, Rey changed these to the cartoon-like images that continue to be featured in each of the books. A collector's edition with the original watercolors has since been released. [ 9 ]
Curious George was an instant success, and the Reys were commissioned to write more adventures of the mischievous monkey and his friend, the Man with the Yellow Hat. They wrote seven stories in all, with Hans mainly doing the illustrations and Margret working mostly on the stories, though they both admitted to sharing the work and cooperating fully in every stage of development. At first, however, covers omitted Margret's name. In later editions, this was changed, and Margret now receives full credit for her role in developing the stories. [ 3 ]
Curious George Takes a Job was named to the Lewis Carroll Shelf Award list in 1960.
In 1963, the Reys relocated to Cambridge, Massachusetts , [ 3 ] in a house near Harvard Square , and lived there until Rey died on August 26, 1977.
In the 1990s, friends of the Reys founded a children's bookstore named Curious George & Friends (formerly Curious George Goes to Wordsworth), which operated in Harvard Square until 2011. [ 10 ] A new Curious George themed store opened in 2012, The World's Only Curious George Store, which moved to Central Square in 2019. [ citation needed ]
Rey's interest in astronomy began during World War I and led to his desire to redraw constellation diagrams, which Rey found difficult to remember, so that they were more intuitive. This led to the 1952 publication of The Stars: A New Way to See Them ( ISBN 0-395-24830-2 ). His constellation diagrams were adopted widely and now appear in many astronomy guides, such as Donald H. Menzel 's A Field Guide to the Stars and Planets . As of 2008 The Stars: A New Way to See Them and a simplified presentation for children called Find the Constellations are still in print. A new edition of Find the Constellations was released in 2008, updated with modern fonts, the new status of Pluto , and some more current measurements of planetary sizes and orbital radii. [ 11 ]
The University of Oregon holds H. A. Rey papers dated 1940 to 1961, dominated by correspondence, primarily between Rey and his American and British publishers. [ 3 ]
The de Grummond Children's Literature Collection in Hattiesburg, Mississippi, holds more than 300 boxes of Rey papers dated 1973 to 2002. [ 13 ]
Dr. Lena Y. de Grummond, a professor in the field of library science at the University of Southern Mississippi , contacted the Reys in 1966 about USM's new children's literature collection. H. A. and Margret donated a pair of sketches at the time. When Margret Rey died in 1996, her will designated that the entire literary estate of the Reys be donated to the de Grummond Collection. | https://en.wikipedia.org/wiki/H._A._Rey |
The H. Milton Stewart School of Industrial and Systems Engineering is a department in the Georgia Institute of Technology 's College of Engineering dedicated to education and research in industrial engineering . The school is named after H. Milton Stewart , a local philanthropist and successful businessman who formerly graduated from the BSIE undergraduate program.
Unlike similar programs at other schools, the School of Industrial and Systems Engineering at Georgia Tech focuses on core disciplines for both Industrial Engineering (such as manufacturing and quality control) and Systems Engineering (such as global logistics and system optimization). U.S. News & World Report consistently ranks the program at number one. [ citation needed ]
The "industrial option" for mechanical engineering was first offered at then Georgia School of Technology in 1924. [ 1 ] The Department of Industrial Engineering was created in 1945 with Frank Groseclose as its first director and professor. In 1948, the department was elevated to its current status of School of Industrial Engineering with Groseclose serving as its first dean. [ 2 ]
This article about a university or other tertiary education institution in the U.S. state of Georgia is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H._Milton_Stewart_School_of_Industrial_and_Systems_Engineering |
Nortilidine [ 1 ] is the major active metabolite of tilidine . It is formed from tilidine by demethylation in the liver . The racemate has opioid analgesic effects roughly equivalent in potency to that of morphine. [ 2 ] The (1 R ,2 S )-isomer has NMDA antagonist activity. The drug also acts as a dopamine reuptake inhibitor . [ 3 ] The reversed-ester of nortilidine is also known, as is the corresponding analogue with the cyclohexene ring replaced by cyclopentane, [ 4 ] which have almost identical properties to nortilidine. [ 5 ]
Nortilidine has been sold as a designer drug , first being identified in Poland in May 2020. [ 6 ] [ 7 ] | https://en.wikipedia.org/wiki/H21NO2 |
Diazomethane is an organic chemical compound with the formula CH 2 N 2 , discovered by German chemist Hans von Pechmann in 1894. It is the simplest diazo compound . In the pure form at room temperature, it is an extremely sensitive explosive yellow gas ; thus, it is almost universally used as a solution in diethyl ether . The compound is a popular methylating agent in the laboratory, but it is too hazardous to be employed on an industrial scale without special precautions. [ 4 ] Use of diazomethane has been significantly reduced by the introduction of the safer and equivalent reagent trimethylsilyldiazomethane . [ 5 ]
For safety and convenience diazomethane is always prepared as needed as a solution in ether and used as such. It converts carboxylic acids to methyl esters and phenols into their methyl ethers . The reaction is thought to proceed via proton transfer from carboxylic acid to diazomethane to give a methyldiazonium cation, which reacts with the carboxylate ion to give the methyl ester and nitrogen gas. Labeling studies indicate that the initial proton transfer is faster than the methyl transfer step. [ 6 ] Since proton transfer is required for the reaction to proceed, this reaction is selective for the more acidic carboxylic acids (p K a ~ 5) and phenols (p K a ~ 10) over aliphatic alcohols (p K a ~ 15). [ 7 ]
In more specialized applications, diazomethane and other diazoalkyl reagents are used in the Arndt–Eistert reaction and the Büchner–Curtius–Schlotterbeck reaction for homologation of various compounds. [ 8 ] [ 9 ]
Diazomethane reacts with alcohols or phenols in presence of boron trifluoride (BF 3 ) to give methyl ethers .
Diazomethane is also frequently used as a carbene source. It readily takes part in 1,3-dipolar cycloadditions .
A wide variety of routes have been developed for the laboratory production of diazomethane. [ 10 ] In general, the synthesis of these all involves the addition of methylamine to an electron-deficient species, before treatment with nitrite and mineral acid ( nitrous acid ) to form an N -methyl nitrosamide. Diazomethane is prepared by hydrolysis of an ethereal solution of these N -methyl nitrosamides with aqueous base. Examples include:
Diazomethane reacts with alkaline solutions of D 2 O to give the deuterated derivative CD 2 N 2 . [ 20 ] This can be used for isotopic labeling studies.
The ease with which diazomethane explodes makes it too hazardous to handle in large quantities. Despite this, it can be used on an industrial scale using on-demand flow chemistry . In these processes the rate of production is matched by the rate of consumption, such that the amount of diazomethane present at any one time is very low. [ 21 ] [ 4 ]
The concentration of CH 2 N 2 can be determined in either of two convenient ways. It can be treated with an excess of benzoic acid in cold Et 2 O. Unreacted benzoic acid is then back-titrated with standard NaOH. Alternatively, the concentration of CH 2 N 2 in Et 2 O can be determined spectrophotometrically at 410 nm where its extinction coefficient , ε, is 7.2. [ citation needed ] The gas-phase concentration of diazomethane can be determined using photoacoustic spectroscopy . [ 4 ]
Diazomethane is both isomeric and isoelectronic with the more stable cyanamide , but they do not interconvert.
Many substituted derivatives of diazomethane have been prepared:
Diazomethane is toxic by inhalation or by contact with the skin or eyes (TLV 0.2 ppm). Symptoms include chest discomfort, headache, weakness and, in severe cases, collapse. [ 26 ] Symptoms may be delayed. Deaths from diazomethane poisoning have been reported. In one instance a laboratory worker consumed a hamburger near a fumehood where he was generating a large quantity of diazomethane, and died four days later from fulminating pneumonia. [ 27 ] Like any other alkylating agent it is expected to be carcinogenic, but such concerns are overshadowed by its serious acute toxicity.
CH 2 N 2 may explode in contact with sharp edges, such as ground-glass joints, even scratches in glassware. [ 28 ] Glassware should be inspected before use and preparation should take place behind a blast shield. Specialized kits to prepare diazomethane with flame-polished joints are commercially available.
The compound explodes when heated beyond 100 °C, exposed to intense light, alkali metals, or calcium sulfate. Use of a blast shield is highly recommended while using this compound.
Proof-of-concept work has been done with microfluidics , in which continuous point-of-use synthesis from N -methyl- N -nitrosourea and 0.93 M potassium hydroxide in water was followed by point-of-use conversion with benzoic acid , resulting in a 65% yield of the methyl benzoate ester within seconds at temperatures ranging from 0 to 50 °C. The yield was better than under capillary conditions; the microfluidics were credited with "suppression of hot spots, low holdup, isothermal conditions, and intensive mixing." [ 29 ]
The stable compound cyanamide , whose minor tautomer is carbodiimide , is an isomer of diazomethane. Less stable but still isolable isomers of diazomethane include the cyclic 3 H -diazirine and isocyanoamine ( isodiazomethane ). [ 30 ] [ 31 ] In addition, the parent nitrilimine has been observed under matrix isolation conditions. [ 32 ] | https://en.wikipedia.org/wiki/H2C-N2 |
Diazomethane is an organic chemical compound with the formula CH 2 N 2 , discovered by German chemist Hans von Pechmann in 1894. It is the simplest diazo compound . In the pure form at room temperature, it is an extremely sensitive explosive yellow gas ; thus, it is almost universally used as a solution in diethyl ether . The compound is a popular methylating agent in the laboratory, but it is too hazardous to be employed on an industrial scale without special precautions. [ 4 ] Use of diazomethane has been significantly reduced by the introduction of the safer and equivalent reagent trimethylsilyldiazomethane . [ 5 ]
For safety and convenience diazomethane is always prepared as needed as a solution in ether and used as such. It converts carboxylic acids to methyl esters and phenols into their methyl ethers . The reaction is thought to proceed via proton transfer from carboxylic acid to diazomethane to give a methyldiazonium cation, which reacts with the carboxylate ion to give the methyl ester and nitrogen gas. Labeling studies indicate that the initial proton transfer is faster than the methyl transfer step. [ 6 ] Since proton transfer is required for the reaction to proceed, this reaction is selective for the more acidic carboxylic acids (p K a ~ 5) and phenols (p K a ~ 10) over aliphatic alcohols (p K a ~ 15). [ 7 ]
In more specialized applications, diazomethane and other diazoalkyl reagents are used in the Arndt–Eistert reaction and the Büchner–Curtius–Schlotterbeck reaction for homologation of various compounds. [ 8 ] [ 9 ]
Diazomethane reacts with alcohols or phenols in presence of boron trifluoride (BF 3 ) to give methyl ethers .
Diazomethane is also frequently used as a carbene source. It readily takes part in 1,3-dipolar cycloadditions .
A wide variety of routes have been developed for the laboratory production of diazomethane. [ 10 ] In general, the synthesis of these all involves the addition of methylamine to an electron-deficient species, before treatment with nitrite and mineral acid ( nitrous acid ) to form an N -methyl nitrosamide. Diazomethane is prepared by hydrolysis of an ethereal solution of these N -methyl nitrosamides with aqueous base. Examples include:
Diazomethane reacts with alkaline solutions of D 2 O to give the deuterated derivative CD 2 N 2 . [ 20 ] This can be used for isotopic labeling studies.
The ease with which diazomethane explodes makes it too hazardous to handle in large quantities. Despite this, it can be used on an industrial scale using on-demand flow chemistry . In these processes the rate of production is matched by the rate of consumption, such that the amount of diazomethane present at any one time is very low. [ 21 ] [ 4 ]
The concentration of CH 2 N 2 can be determined in either of two convenient ways. It can be treated with an excess of benzoic acid in cold Et 2 O. Unreacted benzoic acid is then back-titrated with standard NaOH. Alternatively, the concentration of CH 2 N 2 in Et 2 O can be determined spectrophotometrically at 410 nm where its extinction coefficient , ε, is 7.2. [ citation needed ] The gas-phase concentration of diazomethane can be determined using photoacoustic spectroscopy . [ 4 ]
Diazomethane is both isomeric and isoelectronic with the more stable cyanamide , but they do not interconvert.
Many substituted derivatives of diazomethane have been prepared:
Diazomethane is toxic by inhalation or by contact with the skin or eyes (TLV 0.2 ppm). Symptoms include chest discomfort, headache, weakness and, in severe cases, collapse. [ 26 ] Symptoms may be delayed. Deaths from diazomethane poisoning have been reported. In one instance a laboratory worker consumed a hamburger near a fumehood where he was generating a large quantity of diazomethane, and died four days later from fulminating pneumonia. [ 27 ] Like any other alkylating agent it is expected to be carcinogenic, but such concerns are overshadowed by its serious acute toxicity.
CH 2 N 2 may explode in contact with sharp edges, such as ground-glass joints, even scratches in glassware. [ 28 ] Glassware should be inspected before use and preparation should take place behind a blast shield. Specialized kits to prepare diazomethane with flame-polished joints are commercially available.
The compound explodes when heated beyond 100 °C, exposed to intense light, alkali metals, or calcium sulfate. Use of a blast shield is highly recommended while using this compound.
Proof-of-concept work has been done with microfluidics , in which continuous point-of-use synthesis from N -methyl- N -nitrosourea and 0.93 M potassium hydroxide in water was followed by point-of-use conversion with benzoic acid , resulting in a 65% yield of the methyl benzoate ester within seconds at temperatures ranging from 0 to 50 °C. The yield was better than under capillary conditions; the microfluidics were credited with "suppression of hot spots, low holdup, isothermal conditions, and intensive mixing." [ 29 ]
The stable compound cyanamide , whose minor tautomer is carbodiimide , is an isomer of diazomethane. Less stable but still isolable isomers of diazomethane include the cyclic 3 H -diazirine and isocyanoamine ( isodiazomethane ). [ 30 ] [ 31 ] In addition, the parent nitrilimine has been observed under matrix isolation conditions. [ 32 ] | https://en.wikipedia.org/wiki/H2C-N≡N |
Acetylene ( systematic name : ethyne ) is a chemical compound with the formula C 2 H 2 and structure HC≡CH . It is a hydrocarbon and the simplest alkyne . [ 8 ] This colorless gas is widely used as a fuel and a chemical building block. It is unstable in its pure form and thus is usually handled as a solution. [ 9 ] Pure acetylene is odorless, but commercial grades usually have a marked odor due to impurities such as divinyl sulfide and phosphine . [ 9 ] [ 10 ]
As an alkyne, acetylene is unsaturated because its two carbon atoms are bonded together in a triple bond . The carbon–carbon triple bond places all four atoms in the same straight line, with CCH bond angles of 180°. [ 11 ] The triple bond in acetylene results in a high energy content that is released when acetylene is burned. [ 12 ]
Acetylene was discovered in 1836 by Edmund Davy , who identified it as a "new carburet of hydrogen". [ 13 ] [ 14 ] It was an accidental discovery while attempting to isolate potassium metal. By heating potassium carbonate with carbon at very high temperatures, he produced a residue of what is now known as potassium carbide , (K 2 C 2 ), which reacted with water to release the new gas. [ 12 ] It was rediscovered in 1860 by French chemist Marcellin Berthelot , who coined the name acétylène . [ 15 ] Berthelot's empirical formula for acetylene (C 4 H 2 ), as well as the alternative name " quadricarbure d'hydrogène " (hydrogen quadricarbide), were incorrect because many chemists at that time used the wrong atomic mass for carbon (6 instead of 12). [ 16 ] Berthelot was able to prepare this gas by passing vapours of organic compounds (methanol, ethanol, etc.) through a red hot tube and collecting the effluent . He also found that acetylene was formed by sparking electricity through mixed cyanogen and hydrogen gases. Berthelot later obtained acetylene directly by passing hydrogen between the poles of a carbon arc . [ 17 ] [ 18 ]
Since the 1950s, acetylene has mainly been manufactured by the partial combustion of methane in the US, much of the EU, and many other countries: [ 9 ] [ 19 ] [ 20 ]
It is a recovered side product in production of ethylene by cracking of hydrocarbons . Approximately 400,000 tonnes were produced by this method in 1983. [ 9 ] Its presence in ethylene is usually undesirable because of its explosive character and its ability to poison Ziegler–Natta catalysts . It is selectively hydrogenated into ethylene, usually using Pd – Ag catalysts. [ 21 ]
The heaviest alkanes in petroleum and natural gas are cracked into lighter molecules which are dehydrogenated at high temperature:
This last reaction is implemented in the process of anaerobic decomposition of methane by microwave plasma. [ 22 ]
The first acetylene produced was by Edmund Davy in 1836, via potassium carbide. [ 23 ] Acetylene was historically produced by hydrolysis (reaction with water) of calcium carbide: [ 12 ]
This reaction was discovered by Friedrich Wöhler in 1862, [ 24 ] but a suitable commercial scale production method which allowed acetylene to be put into wider scale use was not found until 1892 by the Canadian inventor Thomas Willson while searching for a viable commercial production method for aluminum. [ 25 ]
As late as the early 21st century, China, Japan, and Eastern Europe produced acetylene primarily by this method. [ 26 ]
The use of this technology has since declined worldwide with the notable exception of China, with its emphasis on coal-based chemical industry, as of 2013. Otherwise oil has increasingly supplanted coal as the chief source of reduced carbon. [ 27 ]
Calcium carbide production requires high temperatures, ~2000 °C, necessitating the use of an electric arc furnace . In the US, this process was an important part of the late-19th century revolution in chemistry enabled by the massive hydroelectric power project at Niagara Falls . [ 28 ]
In terms of valence bond theory , in each carbon atom the 2s orbital hybridizes with one 2p orbital thus forming an sp hybrid. The other two 2p orbitals remain unhybridized. The two ends of the two sp hybrid orbital overlap to form a strong σ valence bond between the carbons, while on each of the other two ends hydrogen atoms attach also by σ bonds. The two unchanged 2p orbitals form a pair of weaker π bonds . [ 29 ]
Since acetylene is a linear symmetrical molecule , it possesses the D ∞h point group . [ 30 ]
At atmospheric pressure, acetylene cannot exist as a liquid and does not have a melting point. The triple point on the phase diagram corresponds to the melting point (−80.8 °C) at the minimal pressure at which liquid acetylene can exist (1.27 atm). At temperatures below the triple point, solid acetylene can change directly to the vapour (gas) by sublimation . The sublimation point at atmospheric pressure is −84.0 °C. [ 31 ]
At room temperature, the solubility of acetylene in acetone is 27.9 g per kg. For the same amount of dimethylformamide (DMF), the solubility is 51 g. At
20.26 bar, the solubility increases to 689.0 and 628.0 g for acetone and DMF, respectively. These solvents are used in pressurized gas cylinders. [ 32 ]
Approximately 20% of acetylene is supplied by the industrial gases industry for oxyacetylene gas welding and cutting due to the high temperature of the flame. Combustion of acetylene with oxygen produces a flame of over 3,600 K (3,330 °C; 6,020 °F), releasing 11.8 kJ /g. Oxygen with acetylene is the hottest burning common gas mixture. [ 33 ] Acetylene is the third-hottest natural chemical flame after dicyanoacetylene 's 5,260 K (4,990 °C; 9,010 °F) and cyanogen at 4,798 K (4,525 °C; 8,177 °F). Oxy-acetylene welding was a popular welding process in previous decades. The development and advantages of arc-based welding processes have made oxy-fuel welding nearly extinct for many applications. Acetylene usage for welding has dropped significantly. On the other hand, oxy-acetylene welding equipment is quite versatile – not only because the torch is preferred for some sorts of iron or steel welding (as in certain artistic applications), but also because it lends itself easily to brazing, braze-welding, metal heating (for annealing or tempering, bending or forming), the loosening of corroded nuts and bolts, and other applications. Bell Canada cable-repair technicians still use portable acetylene-fuelled torch kits as a soldering tool for sealing lead sleeve splices in manholes and in some aerial locations. Oxyacetylene welding may also be used in areas where electricity is not readily accessible. Oxyacetylene cutting is used in many metal fabrication shops. For use in welding and cutting, the working pressures must be controlled by a regulator, since above 15 psi (100 kPa), if subjected to a shockwave (caused, for example, by a flashback ), acetylene decomposes explosively into hydrogen and carbon . [ 34 ]
Acetylene is useful for many processes, but few are conducted on a commercial scale. [ 35 ]
One of the major chemical applications is ethynylation of formaldehyde. [ 9 ] Acetylene adds to aldehydes and ketones to form α-ethynyl alcohols:
The reaction gives butynediol , with propargyl alcohol as the by-product. Copper acetylide is used as the catalyst. [ 36 ] [ 37 ]
In addition to ethynylation, acetylene reacts with carbon monoxide , acetylene reacts to give acrylic acid , or acrylic esters. Metal catalysts are required. These derivatives form products such as acrylic fibers , glasses , paints , resins , and polymers . Except in China, use of acetylene as a chemical feedstock has declined by 70% from 1965 to 2007 owing to cost and environmental considerations. [ 38 ] In China, acetylene is a major precursor to vinyl chloride . [ 35 ]
Prior to the widespread use of petrochemicals, coal-derived acetylene was a building block for several industrial chemicals. Thus acetylene can be hydrated to give acetaldehyde , which in turn can be oxidized to acetic acid. Processes leading to acrylates were also commercialized. Almost all of these processes became obsolete with the availability of petroleum-derived ethylene and propylene. [ 39 ]
In 1881, the Russian chemist Mikhail Kucherov [ 40 ] described the hydration of acetylene to acetaldehyde using catalysts such as mercury(II) bromide . Before the advent of the Wacker process , this reaction was conducted on an industrial scale. [ 41 ]
The polymerization of acetylene with Ziegler–Natta catalysts produces polyacetylene films. Polyacetylene, a chain of CH centres with alternating single and double bonds, was one of the first discovered organic semiconductors . Its reaction with iodine produces a highly electrically conducting material. Although such materials are not useful, these discoveries led to the developments of organic semiconductors , as recognized by the Nobel Prize in Chemistry in 2000 to Alan J. Heeger , Alan G MacDiarmid , and Hideki Shirakawa . [ 9 ]
In the 1920s, pure acetylene was experimentally used as an inhalation anesthetic . [ 42 ]
Acetylene is sometimes used for carburization (that is, hardening) of steel when the object is too large to fit into a furnace. [ 43 ]
Acetylene is used to volatilize carbon in radiocarbon dating . The carbonaceous material in an archeological sample is treated with lithium metal in a small specialized research furnace to form lithium carbide (also known as lithium acetylide). The carbide can then be reacted with water, as usual, to form acetylene gas to feed into a mass spectrometer to measure the isotopic ratio of carbon-14 to carbon-12. [ 44 ]
Acetylene combustion produces a strong, bright light and the ubiquity of carbide lamps drove much acetylene commercialization in the early 20th century. Common applications included coastal lighthouses , [ 45 ] street lights , [ 12 ] and automobile [ 46 ] and mining headlamps . [ 47 ] In most of these applications, direct combustion is a fire hazard , and so acetylene has been replaced, first by incandescent lighting and many years later by low-power/high-lumen LEDs. Nevertheless, acetylene lamps remain in limited use in remote or otherwise inaccessible areas and in countries with a weak or unreliable central electric grid . [ 47 ]
The energy richness of the C≡C triple bond and the rather high solubility of acetylene in water make it a suitable substrate for bacteria, provided an adequate source is available. [ 48 ] A number of bacteria living on acetylene have been identified. The enzyme acetylene hydratase catalyzes the hydration of acetylene to give acetaldehyde : [ 49 ]
Acetylene is a moderately common chemical in the universe, often associated with the atmospheres of gas giants . [ 50 ] One curious discovery of acetylene is on Enceladus , a moon of Saturn . Natural acetylene is believed to form from catalytic decomposition of long-chain hydrocarbons at temperatures of 1,700 K (1,430 °C; 2,600 °F) and above. Since such temperatures are highly unlikely on such a small distant body, this discovery is potentially suggestive of catalytic reactions within that moon, making it a promising site to search for prebiotic chemistry. [ 51 ] [ 52 ]
In vinylation reactions, H−X compounds add across the triple bond. Alcohols and phenols add to acetylene to give vinyl ethers . Thiols give vinyl thioethers. Similarly, vinylpyrrolidone and vinylcarbazole are produced industrially by vinylation of 2-pyrrolidone and carbazole . [ 32 ] [ 9 ]
The hydration of acetylene is a vinylation reaction, but the resulting vinyl alcohol isomerizes to acetaldehyde . The reaction is catalyzed by mercury salts. This reaction once was the dominant technology for acetaldehyde production, but it has been displaced by the Wacker process , which affords acetaldehyde by oxidation of ethylene , a cheaper feedstock. A similar situation applies to the conversion of acetylene to the valuable vinyl chloride by hydrochlorination vs the oxychlorination of ethylene.
Vinyl acetate is used instead of acetylene for some vinylations, which are more accurately described as transvinylations . [ 53 ] Higher esters of vinyl acetate have been used in the synthesis of vinyl formate .
Acetylene and its derivatives (2-butyne, diphenylacetylene, etc.) form complexes with transition metals . Its bonding to the metal is somewhat similar to that of ethylene complexes. These complexes are intermediates in many catalytic reactions such as alkyne trimerisation to benzene, tetramerization to cyclooctatetraene , [ 9 ] and carbonylation to hydroquinone : [ 54 ]
Metal acetylides , species of the formula L n M−C 2 R , are also common. Copper(I) acetylide and silver acetylide can be formed in aqueous solutions with ease due to a favorable solubility equilibrium . [ 55 ]
Acetylene has a p K a of 25, acetylene can be deprotonated by a superbase to form an acetylide : [ 55 ]
Various organometallic [ 56 ] and inorganic [ 57 ] reagents are effective.
Acetylene can be semihydrogenated to ethylene , providing a feedstock for a variety of polyethylene plastics. Halogens add to the triple bond.
Acetylene is not especially toxic, but when generated from calcium carbide , or CaC 2 , it can contain toxic impurities such as traces of phosphine and arsine , which gives it a distinct garlic -like smell. It is also highly flammable, as are most light hydrocarbons, hence its use in welding. Its most singular hazard is associated with its intrinsic instability, especially when it is pressurized: under certain conditions acetylene can react in an exothermic addition-type reaction to form a number of products, typically benzene and/or vinylacetylene , possibly in addition to carbon and hydrogen . [ citation needed ] Although it is stable at normal pressures and temperatures, if it is subjected to pressures as low as 15 psig it can explode. [ 12 ] The safe limit for acetylene therefore is 101 kPa gage , or 15 psig. [ 58 ] [ 59 ] Additionally, if acetylene is initiated by intense heat or a shockwave, it can decompose explosively if the absolute pressure of the gas exceeds about 200 kilopascals (29 psi). It is therefore supplied and stored dissolved in acetone or dimethylformamide (DMF), [ 59 ] [ 60 ] [ 61 ] contained in a gas cylinder with a porous filling , which renders it safe to transport and use, given proper handling. Acetylene cylinders should be used in the upright position to avoid withdrawing acetone during use. [ 62 ]
Information on safe storage of acetylene in upright cylinders is provided by the OSHA, [ 63 ] [ 64 ] Compressed Gas Association, [ 59 ] United States Mine Safety and Health Administration (MSHA), [ 65 ] EIGA, [ 62 ] and other agencies.
Copper catalyses the decomposition of acetylene, and as a result acetylene should not be transported in copper pipes. [ 66 ]
Cylinders should be stored in an area segregated from oxidizers to avoid exacerbated reaction in case of fire/leakage. [ 59 ] [ 64 ] Acetylene cylinders should not be stored in confined spaces, enclosed vehicles, garages, and buildings, to avoid unintended leakage leading to explosive atmosphere. [ 59 ] [ 64 ] In the US, National Electric Code (NEC) requires consideration for hazardous areas including those where acetylene may be released during accidents or leaks. [ 67 ] Consideration may include electrical classification and use of listed Group A electrical components in US. [ 67 ] Further information on determining the areas requiring special consideration is in NFPA 497. [ 68 ] In Europe, ATEX also requires consideration for hazardous areas where flammable gases may be released during accidents or leaks. [ 62 ] | https://en.wikipedia.org/wiki/H2C2 |
Glyoxal is an organic compound with the chemical formula OCHCHO. It is the smallest dialdehyde (a compound with two aldehyde groups). It is a crystalline solid, white at low temperatures and yellow near the melting point (15 °C). The liquid is yellow, and the vapor is green. [ 2 ]
Pure glyoxal is not commonly encountered because glyoxal is usually handled as a 40% aqueous solution (density near 1.24 g/mL). It forms a series of hydrates , including oligomers . For many purposes, these hydrated oligomers behave equivalently to glyoxal. Glyoxal is produced industrially as a precursor to many products. [ 3 ]
Glyoxal was first prepared and named by the German-British chemist Heinrich Debus (1824–1915) by reacting ethanol with nitric acid . [ 4 ] [ 5 ]
Commercial glyoxal is prepared either by the gas-phase oxidation of ethylene glycol in the presence of a silver or copper catalyst (the Laporte process) or by the liquid-phase oxidation of acetaldehyde with nitric acid . [ 3 ]
The first commercial glyoxal source was in Lamotte , France, started in 1960. The single largest commercial source is BASF in Ludwigshafen , Germany , at around 60,000 tons per year. Other production sites exist also in the US and China. Commercial bulk glyoxal is made and reported as a 40% solution in water by weight [ 3 ] (approx. 1:5 molar ratio of glyoxal to water).
Glyoxal may be synthesized in the laboratory by oxidation of acetaldehyde with selenious acid [ 6 ] or by ozonolysis of benzene . [ 7 ]
Anhydrous glyoxal is prepared by heating solid glyoxal hydrate(s) with phosphorus pentoxide and condensing the vapors in a cold trap . [ 8 ]
The experimentally determined Henry's law constant of glyoxal is:
Advanced glycation end-products (AGEs) are proteins or lipids that become glycated as the result of a high-sugar diet. [ 10 ] They are a bio-marker implicated in aging and the development, or worsening, of many degenerative diseases , such as diabetes , atherosclerosis , chronic kidney disease , and Alzheimer's disease . [ 11 ]
Guanine bases in DNA can undergo non-enzymatic glycation by glyoxal to form glyoxal-guanine adducts. [ 12 ] These adducts may then produce DNA crosslinks . Glycation of DNA may also lead to mutation , breaks in DNA and cytotoxicity . [ 13 ] In humans, glyoxal-glycated nucleotides can be repaired by the protein DJ-1 also known as Park7. [ 13 ]
Coated paper and textile finishes use large amounts of glyoxal as a crosslinker for starch -based formulations. It condenses with urea to afford 4,5-dihydroxy-2-imidazolidinone, which further reacts with formaldehyde to give the bis(hydroxymethyl) derivative dimethylol ethylene urea , which is used for wrinkle-resistant chemical treatments of clothing, i.e. permanent press. [ 3 ]
Glyoxal is used as a solubilizer and cross-linking agent in polymer chemistry .
Glyoxal is a valuable building block in organic synthesis , especially in the synthesis of heterocycles such as imidazoles . [ 14 ] A convenient form of the reagent for use in the laboratory is its bis(hemiacetal) with ethylene glycol , 1,4-dioxane-2,3-diol. This compound is commercially available.
Glyoxal solutions can also be used as a fixative for histology , that is, a method of preserving cells for examining them under a microscope.
Glyoxal is supplied typically as a 40% aqueous solution. [ 3 ] Like other small aldehydes , glyoxal forms hydrates. Furthermore, the hydrates condense to give a series of oligomers, some of which remain of uncertain structure. For most applications, the exact nature of the species in solution is inconsequential. At least one hydrate of glyoxal is sold commercially, glyoxal trimer dihydrate: [(CHO) 2 ] 3 (H 2 O) 2 (CAS 4405-13-4).
Other glyoxal equivalents are available, such as the ethylene glycol hemiacetal 1,4-dioxane- trans -2,3-diol ( CAS 4845-50-5, m.p. 91–95 °C).
It is estimated that, at concentrations less than 1 M , glyoxal exists predominantly as the monomer or hydrates thereof, i.e., OCHCHO, OCHCH(OH) 2 , or (HO) 2 CHCH(OH) 2 . At concentrations above 1 M, dimers predominate. These dimers are probably dioxolanes , with the formula [(HO)CH] 2 O 2 CHCHO. Dimer and trimers precipitate as solids from cold solutions. [ 15 ]
Glyoxal has been observed as a trace gas in the atmosphere, e.g. as an oxidation product of hydrocarbons. [ 16 ] Tropospheric concentrations of 0–200 ppt by volume have been reported, in polluted regions up to 1 ppb by volume. [ 17 ]
The LD 50 (oral, rats) is 3.3 g/kg, [ 3 ] when that of common salt is 3 g/kg. [ 18 ] | https://en.wikipedia.org/wiki/H2C2O2 |
Oxalic acid is an organic acid with the systematic name ethanedioic acid and chemical formula HO−C(=O)−C(=O)−OH , also written as (COOH) 2 or (CO 2 H) 2 or H 2 C 2 O 4 . It is the simplest dicarboxylic acid . It is a white crystalline solid that forms a colorless solution in water. Its name is derived from early investigators who isolated oxalic acid from flowering plants of the genus Oxalis , commonly known as wood-sorrels. It occurs naturally in many foods. Excessive ingestion of oxalic acid or prolonged skin contact can be dangerous.
Oxalic acid is a much stronger acid than acetic acid . It is a reducing agent [ 9 ] and its conjugate bases hydrogen oxalate ( HC 2 O − 4 ) and oxalate ( C 2 O 2− 4 ) are chelating agents for metal cations. It is used as a cleaning agent, especially for the removal of rust , because it forms a water-soluble ferric iron complex, the ferrioxalate ion. Oxalic acid typically occurs as the dihydrate with the formula H 2 C 2 O 4 ·2H 2 O .
The preparation of salts of oxalic acid from plants had been known since at least 1745, when the Dutch botanist and physician Herman Boerhaave isolated a salt from wood sorrel , akin to kraft process . [ 10 ] By 1773, François Pierre Savary of Fribourg, Switzerland had isolated oxalic acid from its salt in sorrel. [ 11 ]
In 1776, Swedish chemists Carl Wilhelm Scheele and Torbern Olof Bergman [ 12 ] produced oxalic acid by reacting sugar with concentrated nitric acid ; Scheele called the acid that resulted socker-syra or såcker-syra (sugar acid). By 1784, Scheele had shown that "sugar acid" and oxalic acid from natural sources were identical. [ 13 ] The modern name was introduced along with many other acid names by de Morveau , Lavoisier and coauthors in 1787. [ 14 ]
In 1824, the German chemist Friedrich Wöhler obtained oxalic acid by reacting cyanogen with ammonia in aqueous solution. [ 15 ] This experiment may represent the first synthesis of a natural product . [ 16 ]
Oxalic acid is mainly manufactured by the oxidation of carbohydrates or glucose using nitric acid or air in the presence of vanadium pentoxide . Another process uses oxygen to regenerate the nitric acid, using a variety of precursors including glycolic acid and ethylene glycol . [ 17 ] As of 2011, this process was only used by Mitsubishi in Japan. [ 18 ] A newer method entails oxidative carbonylation of alcohols to give the diesters of oxalic acid:
These diesters are subsequently hydrolyzed to oxalic acid. Approximately 120,000 tonnes are produced annually. [ 16 ]
Historically oxalic acid was obtained exclusively by using caustics, such as sodium or potassium hydroxide , on sawdust , followed by acidification of the oxalate by mineral acids, such as sulfuric acid . [ 19 ] Oxalic acid can also be formed by the heating of sodium formate in the presence of an alkaline catalyst. [ 20 ]
Although it can be readily purchased, oxalic acid can be prepared in the laboratory by oxidizing sucrose using nitric acid in the presence of a small amount of vanadium pentoxide as a catalyst . [ 21 ]
The dihydrate can be converted to the anhydrous form by heating or azeotropic distillation . [ 22 ]
Anhydrous oxalic acid exists as two polymorphs ; in one the hydrogen-bonding results in a chain-like structure, whereas the hydrogen bonding pattern in the other form defines a sheet-like structure. [ 23 ] Because the anhydrous material is both acidic and hydrophilic (water seeking), it is used in esterifications .
The dihydrate H 2 C 2 O 4 ·2 H 2 O has space group C 5 2 h – P 2 1 / n , with lattice parameters a = 611.9 pm , b = 360.7 pm , c = 1205.7 pm , β = 106°19′ , Z = 2 . [ 24 ] The main inter-atomic distances are: C−C 153 pm, C−O 1 129 pm, C−O 2 119 pm. [ 25 ]
Oxalic acid's p K a values vary in the literature from 1.25 to 1.46 and from 3.81 to 4.40. [ 26 ] [ 27 ] [ 28 ] The 100th ed of the CRC, released in 2019, has values of 1.25 and 3.81. [ 29 ] Oxalic acid is relatively strong compared to other carboxylic acids :
Oxalic acid undergoes many of the reactions characteristic for other carboxylic acids. It forms esters such as dimethyl oxalate ( m.p. 52.5 to 53.5 °C, 126.5 to 128.3 °F). [ 30 ] It forms an acid chloride called oxalyl chloride .
Transition metal oxalate complexes are numerous, e.g. the drug oxaliplatin . Oxalic acid has been shown to reduce manganese dioxide MnO 2 in manganese ores to allow the leaching of the metal by sulfuric acid . [ 31 ]
Oxalic acid is an important reagent in lanthanide chemistry. Hydrated lanthanide oxalates form readily in very strongly acidic solutions as a densely crystalline , easily filtered form, largely free of contamination by nonlanthanide elements:
Thermal decomposition of these oxalates gives the oxides , which is the most commonly marketed form of these elements. [ 32 ]
Oxalic acid and oxalates can be oxidized by permanganate in an autocatalytic reaction. [ 33 ]
Oxalic acid vapor decomposes at 125–175 °C into carbon dioxide CO 2 and formic acid HCOOH. Photolysis with 237–313 nm UV light also produces carbon monoxide CO and water. [ 34 ]
Evaporation of a solution of urea and oxalic acid in 2:1 molar ratio yields a solid crystalline compound H 2 C 2 O 4 ·2CO(NH 2 ) 2 , consisting of stacked two-dimensional networks of the neutral molecules held together by hydrogen bonds with the oxygen atoms. [ 35 ]
At least two pathways exist for the enzyme-mediated formation of oxalate. In one pathway, oxaloacetate , a component of the Krebs citric acid cycle , is hydrolyzed to oxalate and acetic acid by the enzyme oxaloacetase : [ 36 ]
It also arises from the dehydrogenation of glycolic acid , which is produced by the metabolism of ethylene glycol .
Early investigators isolated oxalic acid from wood-sorrel ( Oxalis ). Members of the spinach family and the brassicas ( cabbage , broccoli , brussels sprouts ) are high in oxalates, as are sorrel and umbellifers like parsley . [ 37 ] The leaves and stems of all species of the genus Chenopodium and related genera of the family Amaranthaceae , which includes quinoa , contain high levels of oxalic acid. [ 38 ] Rhubarb leaves contain about 0.5% oxalic acid, and jack-in-the-pulpit ( Arisaema triphyllum ) contains calcium oxalate crystals. Similarly, the Virginia creeper , a common decorative vine, produces oxalic acid in its berries as well as oxalate crystals in the sap, in the form of raphides . Bacteria produce oxalates from oxidation of carbohydrates . [ 16 ]
Plants of the genus Fenestraria produce optical fibers made from crystalline oxalic acid to transmit light to subterranean photosynthetic sites. [ 39 ]
Carambola , also known as starfruit, also contains oxalic acid along with caramboxin . Citrus juice contains small amounts of oxalic acid.
The formation of naturally occurring calcium oxalate patinas on certain limestone and marble statues and monuments has been proposed to be caused by the chemical reaction of the carbonate stone with oxalic acid secreted by lichen or other microorganisms . [ 40 ] [ 41 ]
Many soil fungus species secrete oxalic acid, which results in greater solubility of metal cations and increased availability of certain soil nutrients, and can lead to the formation of calcium oxalate crystals. [ 42 ] [ 43 ] Some fungi such as Aspergillus niger have been extensively studied for the industrial production of oxalic acid; [ 44 ] however, those processes are not yet economically competitive with production from oil and gas. [ 45 ] Cryphonectria parasitica may excrete oxalic acid containing solutions at the advancing edge of its chestnut cambium infection. The lower pH (<2.5) of more concentrated oxalic acid excretions may degrade cambium cell walls and have a toxic effect on chestnut cambium cells. Cambium cells that burst provide nutrients for a blight infection advance. [ 46 ] [ 47 ]
The conjugate base of oxalic acid is the hydrogenoxalate anion, and its conjugate base ( oxalate ) is a competitive inhibitor of the lactate dehydrogenase (LDH) enzyme. [ 48 ] LDH catalyses the conversion of pyruvate to lactic acid (end product of the fermentation (anaerobic) process) oxidising the coenzyme NADH to NAD + and H + concurrently. Restoring NAD + levels is essential to the continuation of anaerobic energy metabolism through glycolysis . As cancer cells preferentially use anaerobic metabolism (see Warburg effect ) inhibition of LDH has been shown to inhibit tumor formation and growth, [ 49 ] thus is an interesting potential course of cancer treatment.
Oxalic acid plays a key role in the interaction between pathogenic fungi and plants. Small amounts of oxalic acid enhances plant resistance to fungi, but higher amounts cause widespread programmed cell death of the plant and help with fungi infection. Plants normally produce it in small amounts, but some pathogenic fungi such as Sclerotinia sclerotiorum cause a toxic accumulation. [ 50 ]
Oxalate, besides being biosynthesised, may also be biodegraded. Oxalobacter formigenes is an important gut bacterium that helps animals (including humans) degrade oxalate. [ 51 ]
Oxalic acid's main applications include cleaning or bleaching, especially for the removal of rust (iron complexing agent). Its utility in rust removal agents is due to its forming a stable, water-soluble salt with ferric iron, ferrioxalate ion. Oxalic acid is an ingredient in some tooth whitening products. About 25% of produced oxalic acid is used as a mordant in dyeing processes. It is also used in bleaches , especially for pulpwood , cork, straw, cane, feathers, and for rust removal and other cleaning, in baking powder, and as a third reagent in silica analysis instruments.
Oxalic acid is used by some beekeepers as a miticide against the parasitic varroa mite . [ 52 ]
Dilute solutions (0.05–0.15 M ) of oxalic acid can be used to remove iron from clays such as kaolinite to produce light-colored ceramics . [ 53 ]
Oxalic acid can be used to clean minerals like many other acids. Two such examples are quartz crystals and pyrite. [ 54 ] [ 55 ] [ 56 ]
Oxalic acid is sometimes used in the aluminum anodizing process, with or without sulfuric acid. [ 57 ] Compared to sulfuric-acid anodizing, the coatings obtained are thinner and exhibit lower surface roughness.
Oxalic acid is also widely used as a wood bleach, most often in its crystalline form to be mixed with water to its proper dilution for use. [ citation needed ]
Oxalic acid is also used in electronic and semiconductor industries. In 2006 it was reported being used in electrochemical–mechanical planarization of copper layers in the semiconductor devices fabrication process. [ 58 ]
Reduction of carbon dioxide to oxalic acid by various methods, such as electrocatalysis using a copper complex, [ 59 ] is under study as a proposed chemical intermediate for carbon capture and utilization . [ 60 ]
[ 61 ] [ clarification needed ]
Oxalic acid has an oral LD Lo (lowest published lethal dose) of 600 mg/kg. [ 65 ] It has been reported that the lethal oral dose is 15 to 30 grams. [ 66 ] The toxicity of oxalic acid is due to kidney failure caused by precipitation of solid calcium oxalate . [ 67 ]
Oxalate is known to cause mitochondrial dysfunction . [ 68 ]
Ingestion of ethylene glycol results in oxalic acid as a metabolite which can also cause acute kidney failure.
Most kidney stones , 76%, are composed of calcium oxalate . [ 69 ]
^a Unless otherwise cited, all measurements are based on raw vegetable weights with original moisture content. | https://en.wikipedia.org/wiki/H2C2O4 |
Propadiene ( / p r oʊ p ə ˈ d aɪ iː n / ) or allene ( / ˈ æ l iː n / ) is the organic compound with the formula H 2 C=C=CH 2 . It is the simplest allene , i.e. a compound with two adjacent carbon double bonds . [ 3 ] As a constituent of MAPP gas , it has been used as a fuel for specialized welding .
Propadiene exists in equilibrium with methylacetylene (propyne) and the mixture is sometimes called MAPD for m ethyl a cetylene- p ropa d iene:
for which K eq = 0.22 at 270 °C or 0.1 at 5 °C.
MAPD is produced as a side product, often an undesirable one, of dehydrogenation of propane to produce propene , an important feedstock in the chemical industry . MAPD interferes with the catalytic polymerization of propene. [ 4 ]
In 2019 it was announced that propadiene had been detected in the atmosphere of Saturn's moon Titan using the NASA Infrared Telescope Facility . [ 5 ] This was the first time that propadiene had been detected in space, and the second structural isomeric pair (paired with propyne ) detected in Titan's atmosphere, after HCN - HNC . [ 6 ] [ 7 ] | https://en.wikipedia.org/wiki/H2C=C=CH2 |
Vinyl alcohol , also called ethenol (IUPAC name; not ethanol) or ethylenol , is the simplest enol . With the formula C H 2 CH O H , it is a labile compound that converts to acetaldehyde immediately upon isolation near room temperature. [ 1 ] It is not a practical precursor to any compound.
Vinyl alcohol can be formed by the pyrolytic elimination of water from ethylene glycol at a temperature of 900 °C and low pressure. Such processes are of no practical importance. [ 2 ]
Under normal conditions, vinyl alcohol converts ( tautomerizes ) to acetaldehyde :
At room temperature, acetaldehyde ( H 3 CC(O)H ) is more stable than vinyl alcohol ( H 2 C=CHOH ) by 42.7 kJ/mol. [ 3 ] Vinyl alcohol gas isomerizes to the aldehyde with a half-life of 30 min at room temperature. [ 1 ]
The uncatalyzed keto–enol tautomerism by a 1,3-hydrogen migration is forbidden by the Woodward–Hoffmann rules and therefore has a high activation barrier and is not a significant pathway at or near room temperature. However, even trace amounts of acids or bases (including water) can catalyze the reaction. Even with rigorous precautions to minimize adventitious moisture or proton sources, vinyl alcohol can only be stored for minutes to hours before it isomerizes to acetaldehyde. ( Carbonic acid is another example of a substance that is stable when rigorously pure, but decomposes rapidly due to catalysis by trace moisture.)
The tautomerization can also be catalyzed via photochemical process. These findings suggest that the keto–enol tautomerization is a viable route under atmospheric and stratospheric conditions, relevant to a role for vinyl alcohol in the production of organic acids in the atmosphere. [ 5 ] [ 6 ]
Vinyl alcohol can be stabilized by controlling the water concentration in the system and utilizing the kinetic favorability of the deuterium -produced kinetic isotope effect ( k H + / k D + = 4.75, k H 2 O / k D 2 O = 12). Deuterium stabilization can be accomplished through hydrolysis of a ketene precursor in the presence of a slight stoichiometric excess of heavy water (D 2 O). Studies show that the tautomerization process is significantly inhibited at ambient temperatures ( k t ≈ 10 −6 M/s), and the half-life of the enol form can easily be increased to t 1/2 = 42 minutes for first-order hydrolysis kinetics. [ 7 ]
Because of the instability of vinyl alcohol, the thermoplastic polyvinyl alcohol (PVA or PVOH) is made indirectly by polymerization of vinyl acetate followed by hydrolysis of the ester bonds (Ac = acetyl; HOAc = acetic acid): [ 8 ]
Several metal complexes are known that contain vinyl alcohol as a ligand . One example is Pt(acac)(η 2 -C 2 H 3 OH)Cl. [ 9 ]
Vinyl alcohol was detected in the molecular cloud Sagittarius B in 2001, the last of the three stable isomers of C 2 H 4 O (after acetaldehyde and ethylene oxide ) to be detected in space. [ 10 ] [ 11 ] Its stability in the (dilute) interstellar medium shows that its tautomerization does not happen unimolecularly , [ 11 ] a fact attributed to the size of the activation energy barrier to the rearrangement being insurmountable at temperatures present in interstellar space. [ 12 ] The vinyl alcohol to acetaldehyde rearrangement is the only keto-enol tautomerisation to have been detected in deep space, induced by the provision of secondary electrons from galactic cosmic rays . [ 12 ] | https://en.wikipedia.org/wiki/H2C=CH-O-H |
Vinyl alcohol , also called ethenol (IUPAC name; not ethanol) or ethylenol , is the simplest enol . With the formula C H 2 CH O H , it is a labile compound that converts to acetaldehyde immediately upon isolation near room temperature. [ 1 ] It is not a practical precursor to any compound.
Vinyl alcohol can be formed by the pyrolytic elimination of water from ethylene glycol at a temperature of 900 °C and low pressure. Such processes are of no practical importance. [ 2 ]
Under normal conditions, vinyl alcohol converts ( tautomerizes ) to acetaldehyde :
At room temperature, acetaldehyde ( H 3 CC(O)H ) is more stable than vinyl alcohol ( H 2 C=CHOH ) by 42.7 kJ/mol. [ 3 ] Vinyl alcohol gas isomerizes to the aldehyde with a half-life of 30 min at room temperature. [ 1 ]
The uncatalyzed keto–enol tautomerism by a 1,3-hydrogen migration is forbidden by the Woodward–Hoffmann rules and therefore has a high activation barrier and is not a significant pathway at or near room temperature. However, even trace amounts of acids or bases (including water) can catalyze the reaction. Even with rigorous precautions to minimize adventitious moisture or proton sources, vinyl alcohol can only be stored for minutes to hours before it isomerizes to acetaldehyde. ( Carbonic acid is another example of a substance that is stable when rigorously pure, but decomposes rapidly due to catalysis by trace moisture.)
The tautomerization can also be catalyzed via photochemical process. These findings suggest that the keto–enol tautomerization is a viable route under atmospheric and stratospheric conditions, relevant to a role for vinyl alcohol in the production of organic acids in the atmosphere. [ 5 ] [ 6 ]
Vinyl alcohol can be stabilized by controlling the water concentration in the system and utilizing the kinetic favorability of the deuterium -produced kinetic isotope effect ( k H + / k D + = 4.75, k H 2 O / k D 2 O = 12). Deuterium stabilization can be accomplished through hydrolysis of a ketene precursor in the presence of a slight stoichiometric excess of heavy water (D 2 O). Studies show that the tautomerization process is significantly inhibited at ambient temperatures ( k t ≈ 10 −6 M/s), and the half-life of the enol form can easily be increased to t 1/2 = 42 minutes for first-order hydrolysis kinetics. [ 7 ]
Because of the instability of vinyl alcohol, the thermoplastic polyvinyl alcohol (PVA or PVOH) is made indirectly by polymerization of vinyl acetate followed by hydrolysis of the ester bonds (Ac = acetyl; HOAc = acetic acid): [ 8 ]
Several metal complexes are known that contain vinyl alcohol as a ligand . One example is Pt(acac)(η 2 -C 2 H 3 OH)Cl. [ 9 ]
Vinyl alcohol was detected in the molecular cloud Sagittarius B in 2001, the last of the three stable isomers of C 2 H 4 O (after acetaldehyde and ethylene oxide ) to be detected in space. [ 10 ] [ 11 ] Its stability in the (dilute) interstellar medium shows that its tautomerization does not happen unimolecularly , [ 11 ] a fact attributed to the size of the activation energy barrier to the rearrangement being insurmountable at temperatures present in interstellar space. [ 12 ] The vinyl alcohol to acetaldehyde rearrangement is the only keto-enol tautomerisation to have been detected in deep space, induced by the provision of secondary electrons from galactic cosmic rays . [ 12 ] | https://en.wikipedia.org/wiki/H2C=CH-OH |
Ethylene ( IUPAC name: ethene ) is a hydrocarbon which has the formula C 2 H 4 or H 2 C=CH 2 . It is a colourless, flammable gas with a faint "sweet and musky " odour when pure. [ 7 ] It is the simplest alkene (a hydrocarbon with carbon–carbon double bonds ).
Ethylene is widely used in the chemical industry, and its worldwide production (over 150 million tonnes in 2016 [ 8 ] ) exceeds that of any other organic compound . [ 9 ] [ 10 ] Much of this production goes toward creating polythene , which is a widely used plastic containing polymer chains of ethylene units in various chain lengths. Production emits greenhouse gases , including methane from feedstock production and carbon dioxide from any non- sustainable energy used.
Ethylene is also an important natural plant hormone and is used in agriculture to induce ripening of fruits . [ 11 ] The hydrate of ethylene is ethanol .
This hydrocarbon has four hydrogen atoms bound to a pair of carbon atoms that are connected by a double bond . All six atoms that comprise ethylene are coplanar . The H-C-H angle is 117.4°, close to the 120° for ideal sp² hybridized carbon. The molecule is also relatively weak: rotation about the C-C bond is a very low energy process that requires breaking the π-bond by supplying heat at 50 °C. [ citation needed ]
The π-bond in the ethylene molecule is responsible for its useful reactivity. The double bond is a region of high electron density , thus it is susceptible to attack by electrophiles . Many reactions of ethylene are catalyzed by transition metals, which bind transiently to the ethylene using both the π and π* orbitals. [ citation needed ]
Being a simple molecule, ethylene is spectroscopically simple. Its UV-vis spectrum is still used as a test of theoretical methods. [ 12 ]
Major industrial reactions of ethylene include in order of scale: 1) polymerization , 2) oxidation , 3) halogenation and hydrohalogenation , 4) alkylation , 5) hydration , 6) oligomerization , and 7) hydroformylation . In the United States and Europe , approximately 90% of ethylene is used to produce ethylene oxide , ethylene dichloride , ethylbenzene and polyethylene . [ 13 ] Most of the reactions with ethylene are electrophilic addition . [ citation needed ]
Polyethylene production uses more than half of the world's ethylene supply. Polyethylene, also called polyethene and polythene , is the world's most widely used plastic. It is primarily used to make films in packaging , carrier bags and trash liners . Linear alpha-olefins , produced by oligomerization (formation of short-chain molecules) are used as precursors , detergents , plasticisers , synthetic lubricants , additives, and also as co-monomers in the production of polyethylenes. [ 13 ]
Ethylene is oxidized to produce ethylene oxide , a key raw material in the production of surfactants and detergents by ethoxylation . Ethylene oxide is also hydrolyzed to produce ethylene glycol , widely used as an automotive antifreeze as well as higher molecular weight glycols, glycol ethers , and polyethylene terephthalate . [ 14 ] [ 15 ]
Ethylene oxidation in the presence of a palladium catalyst can form acetaldehyde . This conversion remains a major industrial process (10M kg/y). [ 16 ] The process proceeds via the initial complexation of ethylene to a Pd(II) center. [ citation needed ]
Major intermediates from the halogenation and hydrohalogenation of ethylene include ethylene dichloride , ethyl chloride , and ethylene dibromide . The addition of chlorine entails " oxychlorination ", i.e. chlorine itself is not used. Some products derived from this group are polyvinyl chloride , trichloroethylene , perchloroethylene , methyl chloroform , polyvinylidene chloride and copolymers , and ethyl bromide . [ 17 ]
Major chemical intermediates from the alkylation with ethylene is ethylbenzene , precursor to styrene . Styrene is used principally in polystyrene for packaging and insulation, as well as in styrene-butadiene rubber for tires and footwear. On a smaller scale, ethyltoluene , ethylanilines, 1,4-hexadiene, and aluminium alkyls. Products of these intermediates include polystyrene , unsaturated polyesters and ethylene-propylene terpolymers . [ 17 ]
The hydroformylation (oxo reaction) of ethylene results in propionaldehyde , a precursor to propionic acid and n-propyl alcohol . [ 17 ]
Ethylene has long represented the major nonfermentative precursor to ethanol . The original method entailed its conversion to diethyl sulfate , followed by hydrolysis. The main method practiced since the mid-1990s is the direct hydration of ethylene catalyzed by solid acid catalysts : [ 18 ]
Ethylene is dimerized by hydrovinylation to give n -butenes using processes licensed by Lummus or IFP . The Lummus process produces mixed n -butenes (primarily 2-butenes ) while the IFP process produces 1-butene . 1-Butene is used as a comonomer in the production of certain kinds of polyethylene . [ 19 ]
Ethylene is a hormone that affects the ripening and flowering of many plants. It is widely used to control freshness in horticulture and fruits . [ 20 ] The scrubbing of naturally occurring ethylene delays ripening. [ 21 ] Adsorption of ethylene by nets coated in titanium dioxide gel has also been shown to be effective. [ 22 ]
An example of a niche use is as an anesthetic agent (in an 85% ethylene/15% oxygen ratio). [ 23 ] It is also used as a refrigerant gas for low temperature applications under the name R-1150. [ 24 ]
Global ethylene production was 107 million tonnes in 2005, [ 9 ] 109 million tonnes in 2006, [ 25 ] 138 million tonnes in 2010, and 141 million tonnes in 2011. [ 26 ] By 2013, ethylene was produced by at least 117 companies in 32 countries. To meet the ever-increasing demand for ethylene, sharp increases in production facilities are added globally, particularly in the Mideast and in China . [ 27 ] Production emits greenhouse gas , namely significant amounts of carbon dioxide. [ 28 ]
Ethylene is produced by several methods in the petrochemical industry . A primary method is steam cracking (SC) where hydrocarbons and steam are heated to 750–950 °C. This process converts large hydrocarbons into smaller ones and introduces unsaturation. When ethane is the feedstock, ethylene is the product. Ethylene is separated from the resulting mixture by repeated compression and distillation . [ 17 ] In Europe and Asia, ethylene is obtained mainly from cracking naphtha, gasoil and condensates with the coproduction of propylene, C4 olefins and aromatics (pyrolysis gasoline). [ 29 ] Other procedures employed for the production of ethylene include Fischer-Tropsch synthesis and methanol-to-olefins (MTO). [ 30 ]
Although of great value industrially, ethylene is rarely synthesized in the laboratory and is ordinarily purchased. [ 31 ] It can be produced via dehydration of ethanol with sulfuric acid or in the gas phase with aluminium oxide or activated alumina . [ 32 ]
Ethylene is produced from methionine in nature. The immediate precursor is 1-aminocyclopropane-1-carboxylic acid . [ 33 ]
Ethylene is a fundamental ligand in transition metal alkene complexes . One of the first organometallic compounds, Zeise's salt is a complex of ethylene. Useful reagents containing ethylene include Pt(PPh 3 ) 2 (C 2 H 4 ) and Rh 2 Cl 2 (C 2 H 4 ) 4 . The Rh-catalysed hydroformylation of ethylene is conducted on an industrial scale to provide propionaldehyde . [ 35 ]
Some geologists and scholars believe that the famous Greek Oracle at Delphi (the Pythia ) went into her trance-like state as an effect of ethylene rising from ground faults. [ 36 ]
Ethylene appears to have been discovered by Johann Joachim Becher , who obtained it by heating ethanol with sulfuric acid; [ 37 ] he mentioned the gas in his Physica Subterranea (1669). [ 38 ] Joseph Priestley also mentions the gas in his Experiments and observations relating to the various branches of natural philosophy: with a continuation of the observations on air (1779), where he reports that Jan Ingenhousz saw ethylene synthesized in the same way by a Mr. Enée in Amsterdam in 1777 and that Ingenhousz subsequently produced the gas himself. [ 39 ] The properties of ethylene were studied in 1795 by four Dutch chemists, Johann Rudolph Deimann, Adrien Paets van Troostwyck, Anthoni Lauwerenburgh and Nicolas Bondt, who found that it differed from hydrogen gas and that it contained both carbon and hydrogen. [ 40 ] This group also discovered that ethylene could be combined with chlorine to produce the Dutch oil , 1,2-dichloroethane ; this discovery gave ethylene the name used for it at that time, olefiant gas (oil-making gas.) [ 41 ] The term olefiant gas is in turn the etymological origin of the modern word "olefin", the class of hydrocarbons in which ethylene is the first member. [ citation needed ]
In the mid-19th century, the suffix -ene (an Ancient Greek root added to the end of female names meaning "daughter of") was widely used to refer to a molecule or part thereof that contained one fewer hydrogen atoms than the molecule being modified. Thus, ethylene ( C 2 H 4 ) was the "daughter of ethyl " ( C 2 H 5 ). The name ethylene was used in this sense as early as 1852. [ 42 ]
In 1866, the German chemist August Wilhelm von Hofmann proposed a system of hydrocarbon nomenclature in which the suffixes -ane, -ene, -ine, -one, and -une were used to denote the hydrocarbons with 0, 2, 4, 6, and 8 fewer hydrogens than their parent alkane . [ 43 ] In this system, ethylene became ethene . Hofmann's system eventually became the basis for the Geneva nomenclature approved by the International Congress of Chemists in 1892, which remains at the core of the IUPAC nomenclature. However, by that time, the name ethylene was deeply entrenched, and it remains in wide use today, especially in the chemical industry.
Following experimentation by Luckhardt, Crocker, and Carter at the University of Chicago, [ 44 ] ethylene was used as an anesthetic. [ 45 ] [ 7 ] It remained in use through the 1940s, even while chloroform was being phased out. Its pungent odor and its explosive nature limit its use today. [ 46 ]
The 1979 IUPAC nomenclature rules made an exception for retaining the non-systematic name ethylene ; [ 47 ] however, this decision was reversed in the 1993 rules, [ 48 ] and it remains unchanged in the newest 2013 recommendations, [ 49 ] so the IUPAC name is now ethene . In the IUPAC system, the name ethylene is reserved for the divalent group -CH 2 CH 2 -. Hence, names like ethylene oxide and ethylene dibromide are permitted, but the use of the name ethylene for the two-carbon alkene is not. Nevertheless, use of the name ethylene for H 2 C=CH 2 (and propylene for H 2 C=CHCH 3 ) is still prevalent among chemists in North America. [ 50 ]
"A key factor affecting petrochemicals life-cycle emissions is the methane intensity of feedstocks, especially in the production segment." [ 51 ] Emissions from cracking of naptha and natural gas (common in the US as gas is cheap there) depend a lot on the source of energy (for example gas burnt to provide high temperatures [ 52 ] ) but that from naptha is certainly more per kg of feedstock. [ 53 ] Both steam cracking and production from natural gas via ethane are estimated to emit 1.8 to 2kg of CO2 per kg ethylene produced, [ 54 ] totalling over 260 million tonnes a year. [ 55 ] This is more than all other manufactured chemicals except cement and ammonia. [ 56 ] According to a 2022 report using renewable or nuclear energy could cut emissions by almost half. [ 53 ]
Like all hydrocarbons, ethylene is a combustible asphyxiant . It is listed as an IARC group 3 agent , since there is no current evidence that it causes cancer in humans. [ 57 ] | https://en.wikipedia.org/wiki/H2C=CH2 |
Methylene imine is an organic compound with the chemical formula H 2 C=NH . The simplest imine , it is a stable, colorless gas that has been detected throughout the universe. [ 1 ] Structural parameters determined by microwave spectroscopy include a C=N bond length of 1.27 Å , an N–H bond length of 1.02 Å and an H−N=C bond angle of 110.5°. [ 2 ] Because unhindered imines polymerize or oligomerize when concentrated, methylene imine has not been isolated as a liquid or bulk solid. Attempted synthesis of methylene imine from the reaction of ammonia and formaldehyde produces hexamethylenetetramine . [ 3 ] | https://en.wikipedia.org/wiki/H2C=N-H |
Diazomethane is an organic chemical compound with the formula CH 2 N 2 , discovered by German chemist Hans von Pechmann in 1894. It is the simplest diazo compound . In the pure form at room temperature, it is an extremely sensitive explosive yellow gas ; thus, it is almost universally used as a solution in diethyl ether . The compound is a popular methylating agent in the laboratory, but it is too hazardous to be employed on an industrial scale without special precautions. [ 4 ] Use of diazomethane has been significantly reduced by the introduction of the safer and equivalent reagent trimethylsilyldiazomethane . [ 5 ]
For safety and convenience diazomethane is always prepared as needed as a solution in ether and used as such. It converts carboxylic acids to methyl esters and phenols into their methyl ethers . The reaction is thought to proceed via proton transfer from carboxylic acid to diazomethane to give a methyldiazonium cation, which reacts with the carboxylate ion to give the methyl ester and nitrogen gas. Labeling studies indicate that the initial proton transfer is faster than the methyl transfer step. [ 6 ] Since proton transfer is required for the reaction to proceed, this reaction is selective for the more acidic carboxylic acids (p K a ~ 5) and phenols (p K a ~ 10) over aliphatic alcohols (p K a ~ 15). [ 7 ]
In more specialized applications, diazomethane and other diazoalkyl reagents are used in the Arndt–Eistert reaction and the Büchner–Curtius–Schlotterbeck reaction for homologation of various compounds. [ 8 ] [ 9 ]
Diazomethane reacts with alcohols or phenols in presence of boron trifluoride (BF 3 ) to give methyl ethers .
Diazomethane is also frequently used as a carbene source. It readily takes part in 1,3-dipolar cycloadditions .
A wide variety of routes have been developed for the laboratory production of diazomethane. [ 10 ] In general, the synthesis of these all involves the addition of methylamine to an electron-deficient species, before treatment with nitrite and mineral acid ( nitrous acid ) to form an N -methyl nitrosamide. Diazomethane is prepared by hydrolysis of an ethereal solution of these N -methyl nitrosamides with aqueous base. Examples include:
Diazomethane reacts with alkaline solutions of D 2 O to give the deuterated derivative CD 2 N 2 . [ 20 ] This can be used for isotopic labeling studies.
The ease with which diazomethane explodes makes it too hazardous to handle in large quantities. Despite this, it can be used on an industrial scale using on-demand flow chemistry . In these processes the rate of production is matched by the rate of consumption, such that the amount of diazomethane present at any one time is very low. [ 21 ] [ 4 ]
The concentration of CH 2 N 2 can be determined in either of two convenient ways. It can be treated with an excess of benzoic acid in cold Et 2 O. Unreacted benzoic acid is then back-titrated with standard NaOH. Alternatively, the concentration of CH 2 N 2 in Et 2 O can be determined spectrophotometrically at 410 nm where its extinction coefficient , ε, is 7.2. [ citation needed ] The gas-phase concentration of diazomethane can be determined using photoacoustic spectroscopy . [ 4 ]
Diazomethane is both isomeric and isoelectronic with the more stable cyanamide , but they do not interconvert.
Many substituted derivatives of diazomethane have been prepared:
Diazomethane is toxic by inhalation or by contact with the skin or eyes (TLV 0.2 ppm). Symptoms include chest discomfort, headache, weakness and, in severe cases, collapse. [ 26 ] Symptoms may be delayed. Deaths from diazomethane poisoning have been reported. In one instance a laboratory worker consumed a hamburger near a fumehood where he was generating a large quantity of diazomethane, and died four days later from fulminating pneumonia. [ 27 ] Like any other alkylating agent it is expected to be carcinogenic, but such concerns are overshadowed by its serious acute toxicity.
CH 2 N 2 may explode in contact with sharp edges, such as ground-glass joints, even scratches in glassware. [ 28 ] Glassware should be inspected before use and preparation should take place behind a blast shield. Specialized kits to prepare diazomethane with flame-polished joints are commercially available.
The compound explodes when heated beyond 100 °C, exposed to intense light, alkali metals, or calcium sulfate. Use of a blast shield is highly recommended while using this compound.
Proof-of-concept work has been done with microfluidics , in which continuous point-of-use synthesis from N -methyl- N -nitrosourea and 0.93 M potassium hydroxide in water was followed by point-of-use conversion with benzoic acid , resulting in a 65% yield of the methyl benzoate ester within seconds at temperatures ranging from 0 to 50 °C. The yield was better than under capillary conditions; the microfluidics were credited with "suppression of hot spots, low holdup, isothermal conditions, and intensive mixing." [ 29 ]
The stable compound cyanamide , whose minor tautomer is carbodiimide , is an isomer of diazomethane. Less stable but still isolable isomers of diazomethane include the cyclic 3 H -diazirine and isocyanoamine ( isodiazomethane ). [ 30 ] [ 31 ] In addition, the parent nitrilimine has been observed under matrix isolation conditions. [ 32 ] | https://en.wikipedia.org/wiki/H2C=N2 |
Diazomethane is an organic chemical compound with the formula CH 2 N 2 , discovered by German chemist Hans von Pechmann in 1894. It is the simplest diazo compound . In the pure form at room temperature, it is an extremely sensitive explosive yellow gas ; thus, it is almost universally used as a solution in diethyl ether . The compound is a popular methylating agent in the laboratory, but it is too hazardous to be employed on an industrial scale without special precautions. [ 4 ] Use of diazomethane has been significantly reduced by the introduction of the safer and equivalent reagent trimethylsilyldiazomethane . [ 5 ]
For safety and convenience diazomethane is always prepared as needed as a solution in ether and used as such. It converts carboxylic acids to methyl esters and phenols into their methyl ethers . The reaction is thought to proceed via proton transfer from carboxylic acid to diazomethane to give a methyldiazonium cation, which reacts with the carboxylate ion to give the methyl ester and nitrogen gas. Labeling studies indicate that the initial proton transfer is faster than the methyl transfer step. [ 6 ] Since proton transfer is required for the reaction to proceed, this reaction is selective for the more acidic carboxylic acids (p K a ~ 5) and phenols (p K a ~ 10) over aliphatic alcohols (p K a ~ 15). [ 7 ]
In more specialized applications, diazomethane and other diazoalkyl reagents are used in the Arndt–Eistert reaction and the Büchner–Curtius–Schlotterbeck reaction for homologation of various compounds. [ 8 ] [ 9 ]
Diazomethane reacts with alcohols or phenols in presence of boron trifluoride (BF 3 ) to give methyl ethers .
Diazomethane is also frequently used as a carbene source. It readily takes part in 1,3-dipolar cycloadditions .
A wide variety of routes have been developed for the laboratory production of diazomethane. [ 10 ] In general, the synthesis of these all involves the addition of methylamine to an electron-deficient species, before treatment with nitrite and mineral acid ( nitrous acid ) to form an N -methyl nitrosamide. Diazomethane is prepared by hydrolysis of an ethereal solution of these N -methyl nitrosamides with aqueous base. Examples include:
Diazomethane reacts with alkaline solutions of D 2 O to give the deuterated derivative CD 2 N 2 . [ 20 ] This can be used for isotopic labeling studies.
The ease with which diazomethane explodes makes it too hazardous to handle in large quantities. Despite this, it can be used on an industrial scale using on-demand flow chemistry . In these processes the rate of production is matched by the rate of consumption, such that the amount of diazomethane present at any one time is very low. [ 21 ] [ 4 ]
The concentration of CH 2 N 2 can be determined in either of two convenient ways. It can be treated with an excess of benzoic acid in cold Et 2 O. Unreacted benzoic acid is then back-titrated with standard NaOH. Alternatively, the concentration of CH 2 N 2 in Et 2 O can be determined spectrophotometrically at 410 nm where its extinction coefficient , ε, is 7.2. [ citation needed ] The gas-phase concentration of diazomethane can be determined using photoacoustic spectroscopy . [ 4 ]
Diazomethane is both isomeric and isoelectronic with the more stable cyanamide , but they do not interconvert.
Many substituted derivatives of diazomethane have been prepared:
Diazomethane is toxic by inhalation or by contact with the skin or eyes (TLV 0.2 ppm). Symptoms include chest discomfort, headache, weakness and, in severe cases, collapse. [ 26 ] Symptoms may be delayed. Deaths from diazomethane poisoning have been reported. In one instance a laboratory worker consumed a hamburger near a fumehood where he was generating a large quantity of diazomethane, and died four days later from fulminating pneumonia. [ 27 ] Like any other alkylating agent it is expected to be carcinogenic, but such concerns are overshadowed by its serious acute toxicity.
CH 2 N 2 may explode in contact with sharp edges, such as ground-glass joints, even scratches in glassware. [ 28 ] Glassware should be inspected before use and preparation should take place behind a blast shield. Specialized kits to prepare diazomethane with flame-polished joints are commercially available.
The compound explodes when heated beyond 100 °C, exposed to intense light, alkali metals, or calcium sulfate. Use of a blast shield is highly recommended while using this compound.
Proof-of-concept work has been done with microfluidics , in which continuous point-of-use synthesis from N -methyl- N -nitrosourea and 0.93 M potassium hydroxide in water was followed by point-of-use conversion with benzoic acid , resulting in a 65% yield of the methyl benzoate ester within seconds at temperatures ranging from 0 to 50 °C. The yield was better than under capillary conditions; the microfluidics were credited with "suppression of hot spots, low holdup, isothermal conditions, and intensive mixing." [ 29 ]
The stable compound cyanamide , whose minor tautomer is carbodiimide , is an isomer of diazomethane. Less stable but still isolable isomers of diazomethane include the cyclic 3 H -diazirine and isocyanoamine ( isodiazomethane ). [ 30 ] [ 31 ] In addition, the parent nitrilimine has been observed under matrix isolation conditions. [ 32 ] | https://en.wikipedia.org/wiki/H2C=N=N |
Methylene imine is an organic compound with the chemical formula H 2 C=NH . The simplest imine , it is a stable, colorless gas that has been detected throughout the universe. [ 1 ] Structural parameters determined by microwave spectroscopy include a C=N bond length of 1.27 Å , an N–H bond length of 1.02 Å and an H−N=C bond angle of 110.5°. [ 2 ] Because unhindered imines polymerize or oligomerize when concentrated, methylene imine has not been isolated as a liquid or bulk solid. Attempted synthesis of methylene imine from the reaction of ammonia and formaldehyde produces hexamethylenetetramine . [ 3 ] | https://en.wikipedia.org/wiki/H2C=NH |
Formaldehyde ( / f ɔːr ˈ m æ l d ɪ h aɪ d / ⓘ for- MAL -di-hide , US also / f ə r -/ ⓘ fər- ) ( systematic name methanal ) is an organic compound with the chemical formula CH 2 O and structure H−CHO , more precisely H 2 C=O . The compound is a pungent, colourless gas that polymerises spontaneously into paraformaldehyde . It is stored as aqueous solutions ( formalin ), which consists mainly of the hydrate CH 2 (OH) 2 . It is the simplest of the aldehydes ( R−CHO ). As a precursor to many other materials and chemical compounds, in 2006 the global production of formaldehyde was estimated at 12 million tons per year. [ 14 ] It is mainly used in the production of industrial resins , e.g., for particle board and coatings .
Formaldehyde also occurs naturally. It is derived from the degradation of serine , dimethylglycine , and lipids . Demethylases act by converting N-methyl groups to formaldehyde. [ 15 ]
Formaldehyde is classified as a group 1 carcinogen [ note 1 ] [ 17 ] and can cause respiratory and skin irritation upon exposure. [ 16 ]
Formaldehyde is more complicated than many simple carbon compounds in that it adopts several diverse forms. These compounds can often be used interchangeably and can be interconverted. [ citation needed ]
A small amount of stabilizer , such as methanol , is usually added to suppress oxidation and polymerization . A typical commercial-grade formalin may contain 10–12% methanol in addition to various metallic impurities.
"Formaldehyde" was first used as a generic trademark in 1893 following a previous trade name, "formalin". [ 18 ]
Molecular formaldehyde contains a central carbon atom with a double bond to the oxygen atom and a single bond to each hydrogen atom . This structure is summarised by the condensed formula H 2 C=O. [ 19 ] The molecule is planar, Y-shaped and its molecular symmetry belongs to the C 2v point group . [ 20 ] The precise molecular geometry of gaseous formaldehyde has been determined by gas electron diffraction [ 19 ] [ 21 ] and microwave spectroscopy . [ 22 ] [ 23 ] The bond lengths are 1.21 Å for the carbon–oxygen bond [ 19 ] [ 21 ] [ 22 ] [ 23 ] [ 24 ] and around 1.11 Å for the carbon–hydrogen bond , [ 19 ] [ 21 ] [ 22 ] [ 23 ] while the H–C–H bond angle is 117°, [ 22 ] [ 23 ] close to the 120° angle found in an ideal trigonal planar molecule . [ 19 ] Some excited electronic states of formaldehyde are pyramidal rather than planar as in the ground state . [ 24 ]
Processes in the upper atmosphere contribute more than 80% of the total formaldehyde in the environment. [ 25 ] Formaldehyde is an intermediate in the oxidation (or combustion ) of methane , as well as of other carbon compounds, e.g. in forest fires , automobile exhaust, and tobacco smoke . When produced in the atmosphere by the action of sunlight and oxygen on atmospheric methane and other hydrocarbons , it becomes part of smog . Formaldehyde has also been detected in outer space.
Formaldehyde and its adducts are ubiquitous in nature. Food may contain formaldehyde at levels 1–100 mg/kg. [ 26 ] Formaldehyde, formed in the metabolism of the amino acids serine and threonine , is found in the bloodstream of humans and other primates at concentrations of approximately 50 micromolar . [ 27 ] Experiments in which animals are exposed to an atmosphere containing isotopically labeled formaldehyde have demonstrated that even in deliberately exposed animals, the majority of formaldehyde-DNA adducts found in non-respiratory tissues are derived from endogenously produced formaldehyde. [ 28 ]
Formaldehyde does not accumulate in the environment, because it is broken down within a few hours by sunlight or by bacteria present in soil or water. Humans metabolize formaldehyde quickly, converting it to formic acid . [ 29 ] [ 30 ] It nonetheless presents significant health concerns , as a contaminant .
Formaldehyde appears to be a useful probe in astrochemistry due to prominence of the 1 10 ←1 11 and 2 11 ←2 12 K -doublet transitions. It was the first polyatomic organic molecule detected in the interstellar medium . [ 31 ] Since its initial detection in 1969, it has been observed in many regions of the galaxy . Because of the widespread interest in interstellar formaldehyde, it has been extensively studied, yielding new extragalactic sources. [ 32 ] A proposed mechanism for the formation is the hydrogenation of CO ice: [ 33 ]
HCN , HNC , H 2 CO, and dust have also been observed inside the comae of comets C/2012 F6 (Lemmon) and C/2012 S1 (ISON) . [ 34 ] [ 35 ]
Formaldehyde was discovered in 1859 by the Russian chemist Aleksandr Butlerov (1828–1886) when he attempted to synthesize methanediol ("methylene glycol") from iodomethane and silver oxalate . [ 36 ] In his paper, Butlerov referred to formaldehyde as "dioxymethylen" (methylene dioxide) because his empirical formula for it was incorrect, as atomic weights were not precisely determined until the Karlsruhe Congress .
The compound was identified as an aldehyde by August Wilhelm von Hofmann , who first announced its production by passing methanol vapor in air over hot platinum wire. [ 37 ] [ 38 ] With modifications, Hofmann's method remains the basis of the present day industrial route.
Solution routes to formaldehyde also entail oxidation of methanol or iodomethane . [ 39 ]
Formaldehyde is produced industrially by the catalytic oxidation of methanol . The most common catalysts are silver metal (i.e. the FASIL process ), iron(III) oxide , [ 40 ] iron molybdenum oxides (e.g. iron(III) molybdate ) with a molybdenum -enriched surface, [ 41 ] or vanadium oxides . In the commonly used formox process , methanol and oxygen react at c. 250–400 °C in presence of iron oxide in combination with molybdenum and/or vanadium to produce formaldehyde according to the chemical equation : [ 42 ]
The silver-based catalyst usually operates at a higher temperature, about 650 °C. Two chemical reactions on it simultaneously produce formaldehyde: that shown above and the dehydrogenation reaction:
In principle, formaldehyde could be generated by oxidation of methane , but this route is not industrially viable because the methanol is more easily oxidized than methane. [ 42 ]
Formaldehyde is produced via several enzyme-catalyzed routes. [ 43 ] Living beings, including humans, produce formaldehyde as part of their metabolism. Formaldehyde is key to several bodily functions (e.g. epigenetics [ 27 ] ), but its amount must also be tightly controlled to avoid self-poisoning. [ 44 ]
Formaldehyde is catabolized by alcohol dehydrogenase ADH5 and aldehyde dehydrogenase ALDH2 . [ 45 ]
Formaldehyde is a building block in the synthesis of many other compounds of specialised and industrial significance. It exhibits most of the chemical properties of other aldehydes but is more reactive. [ 46 ]
Monomeric CH 2 O is a gas and is rarely encountered in the laboratory. Aqueous formaldehyde, unlike some other small aldehydes (which need specific conditions to oligomerize through aldol condensation ) oligomerizes spontaneously at a common state. The trimer 1,3,5-trioxane, (CH 2 O) 3 , is a typical oligomer. Many cyclic oligomers of other sizes have been isolated. Similarly, formaldehyde hydrates to give the geminal diol methanediol , which condenses further to form hydroxy-terminated oligomers HO(CH 2 O) n H. The polymer is called paraformaldehyde . The higher concentration of formaldehyde—the more equilibrium shifts towards polymerization. Diluting with water or increasing the solution temperature, as well as adding alcohols (such as methanol or ethanol) lowers that tendency.
Gaseous formaldehyde polymerizes at active sites on vessel walls, but the mechanism of the reaction is unknown. [ 47 ] Small amounts of hydrogen chloride , boron trifluoride , or stannic chloride present in gaseous formaldehyde provide the catalytic effect and make the polymerization rapid. [ 48 ]
Formaldehyde forms cross-links by first combining with a protein to form methylol , which loses a water molecule to form a Schiff base . [ 49 ] The Schiff base can then react with DNA or protein to create a cross-linked product. [ 49 ] This reaction is the basis for the most common process of chemical fixation .
Formaldehyde is readily oxidized by atmospheric oxygen into formic acid . For this reason, commercial formaldehyde is typically contaminated with formic acid. Formaldehyde can be hydrogenated into methanol .
In the Cannizzaro reaction , formaldehyde and base react to produce formic acid and methanol, a disproportionation reaction .
Formaldehyde reacts with many compounds, resulting in hydroxymethylation :
The resulting hydroxymethyl derivatives typically react further. Thus, amines give hexahydro-1,3,5-triazines :
Similarly, when combined with hydrogen sulfide , it forms trithiane : [ 50 ]
In the presence of acids, it participates in electrophilic aromatic substitution reactions with aromatic compounds resulting in hydroxymethylated derivatives:
When conducted in the presence of hydrogen chloride, the product is the chloromethyl compound, as described in the Blanc chloromethylation . If the arene is electron-rich, as in phenols, elaborate condensations ensue. With 4-substituted phenols one obtains calixarenes . [ 51 ] Phenol results in polymers.
Many amino acids react with formaldehyde. [ 43 ] Cysteine converts to thioproline .
Formaldehyde is a common precursor to more complex compounds and materials. In approximate order of decreasing consumption, products generated from formaldehyde include urea formaldehyde resin , melamine resin , phenol formaldehyde resin , polyoxymethylene plastics , 1,4-butanediol , and methylene diphenyl diisocyanate . [ 42 ] The textile industry uses formaldehyde-based resins as finishers to make fabrics crease-resistant. [ 52 ]
When condensed with phenol , urea , or melamine , formaldehyde produces, respectively, hard thermoset phenol formaldehyde resin, urea formaldehyde resin, and melamine resin. These polymers are permanent adhesives used in plywood and carpeting . They are also foamed to make insulation , or cast into moulded products. Production of formaldehyde resins accounts for more than half of formaldehyde consumption.
Formaldehyde is also a precursor to polyfunctional alcohols such as pentaerythritol , which is used to make paints and explosives . Other formaldehyde derivatives include methylene diphenyl diisocyanate, an important component in polyurethane paints and foams, and hexamine , which is used in phenol-formaldehyde resins as well as the explosive RDX .
Condensation with acetaldehyde affords pentaerythritol , a chemical necessary in synthesizing PETN , a high explosive: [ 53 ]
An aqueous solution of formaldehyde can be useful as a disinfectant as it kills most bacteria and fungi (including their spores). It is used as an additive in vaccine manufacturing to inactivate toxins and pathogens. [ 54 ] Formaldehyde releasers are used as biocides in personal care products such as cosmetics. Although present at levels not normally considered harmful, they are known to cause allergic contact dermatitis in certain sensitised individuals. [ 55 ]
Aquarists use formaldehyde as a treatment for the parasites Ichthyophthirius multifiliis and Cryptocaryon irritans . [ 56 ] Formaldehyde is one of the main disinfectants recommended for destroying anthrax . [ 57 ]
Formaldehyde is also approved for use in the manufacture of animal feeds in the US. It is an antimicrobial agent used to maintain complete animal feeds or feed ingredients Salmonella negative for up to 21 days. [ 58 ]
Formaldehyde preserves or fixes tissue or cells. The process involves cross-linking of primary amino groups . The European Union has banned the use of formaldehyde as a biocide (including embalming ) under the Biocidal Products Directive (98/8/EC) due to its carcinogenic properties. [ 59 ] [ 60 ] Countries with a strong tradition of embalming corpses, such as Ireland and other colder-weather countries, have raised concerns. Despite reports to the contrary, [ 61 ] no decision on the inclusion of formaldehyde on Annex I of the Biocidal Products Directive for product-type 22 (embalming and taxidermist fluids) had been made as of September 2009 [update] . [ 62 ]
Formaldehyde-based crosslinking is exploited in ChIP-on-chip or ChIP-sequencing genomics experiments, where DNA-binding proteins are cross-linked to their cognate binding sites on the chromosome and analyzed to determine what genes are regulated by the proteins. Formaldehyde is also used as a denaturing agent in RNA gel electrophoresis , preventing RNA from forming secondary structures. A solution of 4% formaldehyde fixes pathology tissue specimens at about one mm per hour at room temperature.
Formaldehyde and 18 M (concentrated) sulfuric acid makes Marquis reagent —which can identify alkaloids and other compounds.
In photography, formaldehyde is used in low concentrations for the process C-41 (color negative film) stabilizer in the final wash step, [ 63 ] as well as in the process E-6 pre-bleach step, to make it unnecessary in the final wash. Due to improvements in dye coupler chemistry, more modern (2006 or later) E-6 and C-41 films do not need formaldehyde, as their dyes are already stable.
In view of its widespread use, toxicity, and volatility, formaldehyde poses a significant danger to human health. [ 64 ] [ 65 ] In 2011, the US National Toxicology Program described formaldehyde as "known to be a human carcinogen". [ 66 ] [ 67 ] [ 68 ]
Concerns are associated with chronic (long-term) exposure by inhalation as may happen from thermal or chemical decomposition of formaldehyde-based resins and the production of formaldehyde resulting from the combustion of a variety of organic compounds (for example, exhaust gases). As formaldehyde resins are used in many construction materials , it is one of the more common indoor air pollutants . [ 69 ] [ 70 ] At concentrations above 0.1 ppm in air, formaldehyde can irritate the eyes and mucous membranes . [ 71 ] Formaldehyde inhaled at this concentration may cause headaches, a burning sensation in the throat, and difficulty breathing, and can trigger or aggravate asthma symptoms. [ 72 ] [ 73 ]
The CDC considers formaldehyde as a systemic poison. Formaldehyde poisoning can cause permanent changes in the nervous system 's functions. [ 74 ]
A 1988 Canadian study of houses with urea-formaldehyde foam insulation found that formaldehyde levels as low as 0.046 ppm were positively correlated with eye and nasal irritation. [ 75 ] A 2009 review of studies has shown a strong association between exposure to formaldehyde and the development of childhood asthma . [ 76 ]
A theory was proposed for the carcinogenesis of formaldehyde in 1978. [ 77 ] In 1987 the United States Environmental Protection Agency (EPA) classified it as a probable human carcinogen , and after more studies the WHO International Agency for Research on Cancer (IARC) in 1995 also classified it as a probable human carcinogen . Further information and evaluation of all known data led the IARC to reclassify formaldehyde as a known human carcinogen [ 78 ] associated with nasal sinus cancer and nasopharyngeal cancer . [ 79 ] Studies in 2009 and 2010 have also shown a positive correlation between exposure to formaldehyde and the development of leukemia , particularly myeloid leukemia . [ 80 ] [ 81 ] Nasopharyngeal and sinonasal cancers are relatively rare, with a combined annual incidence in the United States of < 4,000 cases. [ 82 ] [ 83 ] About 30,000 cases of myeloid leukemia occur in the United States each year. [ 84 ] [ 85 ] Some evidence suggests that workplace exposure to formaldehyde contributes to sinonasal cancers. [ 86 ] Professionals exposed to formaldehyde in their occupation, such as funeral industry workers and embalmers , showed an increased risk of leukemia and brain cancer compared with the general population. [ 87 ] Other factors are important in determining individual risk for the development of leukemia or nasopharyngeal cancer. [ 86 ] [ 88 ] [ 89 ] In yeast, formaldehyde is found to perturb pathways for DNA repair and cell cycle. [ 90 ]
In the residential environment, formaldehyde exposure comes from a number of routes; formaldehyde can be emitted by treated wood products, such as plywood or particle board , but it is produced by paints, varnishes , floor finishes, and cigarette smoking as well. [ 91 ] In July 2016, the U.S. EPA released a prepublication version of its final rule on Formaldehyde Emission Standards for Composite Wood Products. [ 92 ] These new rules impact manufacturers, importers, distributors, and retailers of products containing composite wood, including fiberboard, particleboard, and various laminated products, who must comply with more stringent record-keeping and labeling requirements. [ 93 ]
The U.S. EPA allows no more than 0.016 ppm formaldehyde in the air in new buildings constructed for that agency. [ 94 ] [ failed verification ] A U.S. EPA study found a new home measured 0.076 ppm when brand new and 0.045 ppm after 30 days. [ 95 ] The Federal Emergency Management Agency (FEMA) has also announced limits on the formaldehyde levels in trailers purchased by that agency. [ 96 ] The EPA recommends the use of "exterior-grade" pressed-wood products with phenol instead of urea resin to limit formaldehyde exposure, since pressed-wood products containing formaldehyde resins are often a significant source of formaldehyde in homes. [ 79 ]
The eyes are most sensitive to formaldehyde exposure: The lowest level at which many people can begin to smell formaldehyde ranges between 0.05 and 1 ppm. The maximum concentration value at the workplace is 0.3 ppm. [ 97 ] [ need quotation to verify ] In controlled chamber studies, individuals begin to sense eye irritation at about 0.5 ppm; 5 to 20 percent report eye irritation at 0.5 to 1 ppm; and greater certainty for sensory irritation occurred at 1 ppm and above. While some agencies have used a level as low as 0.1 ppm as a threshold for irritation, the expert panel found that a level of 0.3 ppm would protect against nearly all irritation. In fact, the expert panel found that a level of 1.0 ppm would avoid eye irritation—the most sensitive endpoint—in 75–95% of all people exposed. [ 98 ]
Formaldehyde levels in building environments are affected by a number of factors. These include the potency of formaldehyde-emitting products present, the ratio of the surface area of emitting materials to volume of space, environmental factors, product age, interactions with other materials, and ventilation conditions. Formaldehyde emits from a variety of construction materials, furnishings, and consumer products. The three products that emit the highest concentrations are medium density fiberboard , hardwood plywood, and particle board. Environmental factors such as temperature and relative humidity can elevate levels because formaldehyde has a high vapor pressure . Formaldehyde levels from building materials are the highest when a building first opens because materials would have less time to off-gas. Formaldehyde levels decrease over time as the sources suppress.
In operating rooms , formaldehyde is produced as a byproduct of electrosurgery and is present in surgical smoke, exposing surgeons and healthcare workers to potentially unsafe concentrations. [ 99 ]
Formaldehyde levels in air can be sampled and tested in several ways, including impinger, treated sorbent, and passive monitors. [ 100 ] The National Institute for Occupational Safety and Health (NIOSH) has measurement methods numbered 2016, 2541, 3500, and 3800. [ 101 ]
In June 2011, the twelfth edition of the National Toxicology Program (NTP) Report on Carcinogens (RoC) changed the listing status of formaldehyde from "reasonably anticipated to be a human carcinogen" to "known to be a human carcinogen." [ 66 ] [ 67 ] [ 68 ] Concurrently, a National Academy of Sciences (NAS) committee was convened and issued an independent review of the draft U.S. EPA IRIS assessment of formaldehyde, providing a comprehensive health effects assessment and quantitative estimates of human risks of adverse effects. [ 102 ]
For most people, irritation from formaldehyde is temporary and reversible, although formaldehyde can cause allergies and is part of the standard patch test series. In 2005–06, it was the seventh-most-prevalent allergen in patch tests (9.0%). [ 103 ] People with formaldehyde allergy are advised to avoid formaldehyde releasers as well (e.g., Quaternium-15 , imidazolidinyl urea , and diazolidinyl urea ). [ 104 ] People who suffer allergic reactions to formaldehyde tend to display lesions on the skin in the areas that have had direct contact with the substance, such as the neck or thighs (often due to formaldehyde released from permanent press finished clothing) or dermatitis on the face (typically from cosmetics). [ 55 ] Formaldehyde has been banned in cosmetics in both Sweden [ 105 ] and Japan . [ 106 ]
In humans, ingestion of as little as 30 millilitres (1.0 US fl oz) of 37% formaldehyde solution can cause death. Other symptoms associated with ingesting such a solution include gastrointestinal damage (vomiting, abdominal pain), and systematic damage (dizziness). [ 74 ] Testing for formaldehyde is by blood and/or urine by gas chromatography–mass spectrometry . Other methods to detect formaldehyde include infrared detection, gas detector tubes, gas detectors using electrochemical sensors, and high-performance liquid chromatography (HPLC). HPLC is the most sensitive. [ 107 ]
The fifteenth edition (2021) of the U.S. National Toxicology Program Report on Carcinogens notes that currently in the U.S. “The general population can be exposed to formaldehyde primarily from breathing indoor or outdoor air, from tobacco smoke, from use of cosmetic products containing formaldehyde, and, to a more limited extent, from ingestion of food and water.” Affected water includes groundwater, surface water, and bottled water. It also notes that occupational exposure can be significant. [ 108 ]
Formaldehyde in food can be present naturally, added as an inadvertent contaminant, or intentionally added as a preservative, disinfectant, or bacteriostatic agent . Cooking and smoking food can also result in formaldehyde being produced in food. Foods that the U.S. National Toxicology Program has reported to have higher levels compared to other foods are fish, seafood, and smoked ham. It also notes formaldehyde in food generally occurs in a bound form and that formaldehyde is unstable in an aqueous solution . [ 109 ]
Scandals have broken in both the 2005 Indonesia food scare and 2007 Vietnam food scare regarding the addition of formaldehyde to foods to extend shelf life. In 2011, after a four-year absence, Indonesian authorities found foods with formaldehyde being sold in markets in a number of regions across the country. [ 110 ] In August 2011, at least at two Carrefour supermarkets, the Central Jakarta Livestock and Fishery Sub-Department found cendol containing 10 parts per million of formaldehyde. [ 111 ] In 2014, the owner of two noodle factories in Bogor , Indonesia, was arrested for using formaldehyde in noodles. [ 112 ] Foods known to be contaminated included noodles, salted fish, and tofu. Chicken and beer were also rumored to be contaminated. In some places, such as China, manufacturers still use formaldehyde illegally as a preservative in foods, which exposes people to formaldehyde ingestion. [ 113 ]
In 2011 in Nakhon Ratchasima , Thailand, truckloads of rotten chicken were treated with formaldehyde for sale in which "a large network", including 11 slaughterhouses run by a criminal gang, were implicated. [ 114 ] In 2012, 1 billion rupiah (almost US$100,000) of fish imported from Pakistan to Batam , Indonesia, were found laced with formaldehyde. [ 115 ]
Formalin contamination of foods has been reported in Bangladesh , with stores and supermarkets selling fruits, fishes, and vegetables that have been treated with formalin to keep them fresh. [ 116 ] However, in 2015, a Formalin Control Bill was passed in the Parliament of Bangladesh with a provision of life-term imprisonment as the maximum punishment as well as a maximum fine of 2,000,000 BDT but not less than 500,000 BDT for importing, producing, or hoarding formalin without a license. [ 117 ]
In the early 1900s, formaldehyde was frequently added by US milk plants to milk bottles as a method of pasteurization due to the lack of knowledge and concern [ 118 ] regarding formaldehyde's toxicity. [ 119 ] [ 120 ]
Formaldehyde was one of the chemicals used in 19th century industrialised food production that was investigated by Dr. Harvey W. Wiley with his famous 'Poison Squad' as part of the US Department of Agriculture . This led to the 1906 Pure Food and Drug Act , a landmark event in the early history of food regulation in the United States . [ 121 ]
Formaldehyde is banned from use in certain applications (preservatives for liquid-cooling and processing systems, slimicides , metalworking-fluid preservatives, and antifouling products) under the Biocidal Products Directive. [ 122 ] [ 123 ] In the EU, the maximum allowed concentration of formaldehyde in finished products is 0.2%, and any product that exceeds 0.05% has to include a warning that the product contains formaldehyde. [ 55 ]
In the United States, Congress passed a bill July 7, 2010, regarding the use of formaldehyde in hardwood plywood , particle board , and medium density fiberboard . The bill limited the allowable amount of formaldehyde emissions from these wood products to 0.09 ppm, and required companies to meet this standard by January 2013. [ 124 ] The final U.S. EPA rule specified maximum emissions of "0.05 ppm formaldehyde for hardwood plywood, 0.09 ppm formaldehyde for particleboard, 0.11 ppm formaldehyde for medium-density fiberboard, and 0.13 ppm formaldehyde for thin medium-density fiberboard." [ 125 ]
Formaldehyde was declared a toxic substance by the 1999 Canadian Environmental Protection Act . [ 126 ]
The FDA is proposing a ban on hair relaxers with formaldehyde due to cancer concerns. [ 127 ] | https://en.wikipedia.org/wiki/H2C=O |
HCNH + , also known as protonated hydrogen cyanide , is a molecular ion of astrophysical interest. It also exists in the condensed state when formed by superacids .
In the ground state , HC + N H is a simple linear molecule, whereas its excited triplet state is expected to have cis and trans isomeric forms . The higher-energy structural isomers H 2 CN + and C + N H 2 have also been studied theoretically. [ 5 ]
As a relatively simple molecular ion, HCNH + has been extensively studied in the laboratory. The very first spectrum taken at any wavelength focused on the ν 2 (C−H stretch) ro-vibrational band in the infrared . [ 6 ] Soon afterward, the same authors reported on their investigation of the ν 1 (N−H stretch) band. [ 7 ] Following these initial studies, several groups published manuscripts on the various ro-vibrational spectra of HCNH + , including studies of the ν 3 band (C≡N stretch), [ 8 ] the ν 4 band (H−C≡N bend), [ 9 ] and the ν 5 band (H−N≡C bend)
. [ 10 ]
While all of these studies focused on ro-vibrational spectra in the infrared , it was not until 1998 that technology advanced far enough for an investigation of the pure rotational spectrum of HCNH + in the microwave region to take place. At that time, microwave spectra for HCNH + and its isotopomers HCND + and DCND + were published. [ 11 ] Recently, the pure rotational spectrum of HCNH + was measured again in order to more precisely determine the molecular rotational constants B and D . [ 12 ]
According to the database at astrochemistry.net , the most advanced chemical models of HCNH + include 71 total formation reactions and 21 total destruction reactions. Of these, however, only a handful dominate the overall formation and destruction. [ 13 ] In the case of formation, the 7 dominant reactions are:
HCNH + was first detected in interstellar space in 1986 toward the dense cloud Sgr B2 using the NRAO 12 m dish and the Texas Millimeter Wave Observatory . [ 14 ] These observations utilized the J = 1–0, 2–1, and 3–2 pure rotational transitions at 74, 148, and 222 GHz, respectively.
Since the initial detection, HCNH + has also been observed in TMC-1 [ 15 ] [ 16 ] as well as DR 21(OH) [ 15 ] . [ 17 ] The initial detection toward Sgr B2 has also been confirmed. [ 15 ] [ 18 ] All 3 of these sources are dense molecular clouds, and to date HCNH + has not been detected in diffuse interstellar material.
While not directly detected via spectroscopy, the existence of HCNH + has been inferred to exist in the atmosphere of Saturn 's largest moon, Titan , [ 19 ] based on data from the Ion and Neutral Mass Spectrometer (INMS) instrument aboard the Cassini space probe. Models of Titan's atmosphere had predicted that HCNH + would be the dominant ion present, and a strong peak in the mass spectrum at m / z = 28 seems to support this theory.
In 1997, observations were made of the long-period comet Hale–Bopp in an attempt to find HCNH + , [ 20 ] but it was not detected. However, the upper limit derived from these observations, along with the detections of HCN , HNC , and CN , is important in understanding the chemistry associated with comets . | https://en.wikipedia.org/wiki/H2CN+ |
Carbonic acid is a chemical compound with the chemical formula H 2 C O 3 . The molecule rapidly converts to water and carbon dioxide in the presence of water. However, in the absence of water, it is quite stable at room temperature . [ 5 ] [ 6 ] The interconversion of carbon dioxide and carbonic acid is related to the breathing cycle of animals and the acidification of natural waters . [ 4 ]
In biochemistry and physiology, the name "carbonic acid" is sometimes applied to aqueous solutions of carbon dioxide . These chemical species play an important role in the bicarbonate buffer system , used to maintain acid–base homeostasis . [ 7 ]
In chemistry , the term "carbonic acid" strictly refers to the chemical compound with the formula H 2 CO 3 . Some biochemistry literature effaces the distinction between carbonic acid and carbon dioxide dissolved in extracellular fluid.
In physiology , carbon dioxide excreted by the lungs may be called volatile acid or respiratory acid .
At ambient temperatures, pure carbonic acid is a stable gas. [ 6 ] There are two main methods to produce anhydrous carbonic acid: reaction of hydrogen chloride and potassium bicarbonate at 100 K in methanol and proton irradiation of pure solid carbon dioxide . [ 3 ] Chemically, it behaves as a diprotic Brønsted acid . [ 8 ] [ 9 ]
Carbonic acid monomers exhibit three conformational isomers : cis–cis, cis–trans, and trans–trans. [ 10 ]
At low temperatures and atmospheric pressure , solid carbonic acid is amorphous and lacks Bragg peaks in X-ray diffraction . [ 11 ] But at high pressure, carbonic acid crystallizes, and modern analytical spectroscopy can measure its geometry.
According to neutron diffraction of dideuterated carbonic acid ( D 2 CO 3 ) in a hybrid clamped cell ( Russian alloy / copper-beryllium ) at 1.85 GPa, the molecules are planar and form dimers joined by pairs of hydrogen bonds . All three C-O bonds are nearly equidistant at 1.34 Å , intermediate between typical C-O and C=O distances (respectively 1.43 and 1.23 Å). The unusual C-O bond lengths are attributed to delocalized π bonding in the molecule's center and extraordinarily strong hydrogen bonds. The same effects also induce a very short O—O separation (2.13 Å), through the 136° O-H-O angle imposed by the doubly hydrogen-bonded 8-membered rings. [ 4 ] Longer O—O distances are observed in strong intramolecular hydrogen bonds, e.g. in oxalic acid , where the distances exceed 2.4 Å. [ 11 ]
In even a slight presence of water, carbonic acid dehydrates to carbon dioxide and water , which then catalyzes further decomposition. [ 6 ] For this reason, carbon dioxide can be considered the carbonic acid anhydride .
The hydration equilibrium constant at 25 °C is [ H 2 CO 3 ]/[CO 2 ] ≈ 1.7×10 −3 in pure water [ 12 ] and ≈ 1.2×10 −3 in seawater . [ 13 ] Hence the majority of carbon dioxide at geophysical or biological air-water interfaces does not convert to carbonic acid, remaining dissolved CO 2 gas. However, the uncatalyzed equilibrium is reached quite slowly: the rate constants are 0.039 s −1 for hydration and 23 s −1 for dehydration.
In the presence of the enzyme carbonic anhydrase , equilibrium is instead reached rapidly, and the following reaction takes precedence: [ 14 ] HCO 3 − + H + ↽ − − ⇀ CO 2 + H 2 O {\displaystyle {\ce {HCO3^- {+}H^+ <=> CO2 {+}H2O}}}
When the created carbon dioxide exceeds its solubility, gas evolves and a third equilibrium CO 2 ( soln ) ↽ − − ⇀ CO 2 ( g ) {\displaystyle {\ce {CO_2 (soln) <=> CO_2 (g)}}} must also be taken into consideration. The equilibrium constant for this reaction is defined by Henry's law .
The two reactions can be combined for the equilibrium in solution: HCO 3 − + H + ↽ − − ⇀ CO 2 ( soln ) + H 2 O K 3 = [ H + ] [ HCO 3 − ] [ CO 2 ( soln ) ] {\displaystyle {\begin{aligned}{\ce {HCO3^{-}{}+ H+{}<=> CO2(soln){}+ H2O}}&&K_{3}={\frac {[{\ce {H+}}][{\ce {HCO3^-}}]}{[{\ce {CO2(soln)}}]}}\end{aligned}}} When Henry's law is used to calculate the denominator care is needed with regard to units since Henry's law constant can be commonly expressed with 8 different dimensionalities. [ 15 ]
In wastewater treatment and agriculture irrigation, carbonic acid is used to acidify the water similar to sulfuric acid and sulfurous acid produced by sulfur burners. [ 16 ]
In the beverage industry , sparkling or "fizzy water" is usually referred to as carbonated water . It is made by dissolving carbon dioxide under a small positive pressure in water. Many soft drinks treated the same way effervesce .
Significant amounts of molecular H 2 CO 3 exist in aqueous solutions subjected to pressures of multiple gigapascals (tens of thousands of atmospheres) in planetary interiors. [ 17 ] [ 18 ] Pressures of 0.6–1.6 GPa at 100 K , and 0.75–1.75 GPa at 300 K are attained in the cores of large icy satellites such as Ganymede , Callisto , and Titan , where water and carbon dioxide are present. Pure carbonic acid, being denser, is expected to have sunk under the ice layers and separate them from the rocky cores of these moons. [ 19 ]
Carbonic acid is the formal Brønsted–Lowry conjugate acid of the bicarbonate anion, stable in alkaline solution . The protonation constants have been measured to great precision, but depend on overall ionic strength I . The two equilibria most easily measured are as follows: CO 3 2 − + H + ↽ − − ⇀ HCO 3 − β 1 = [ HCO 3 − ] [ H + ] [ CO 3 2 − ] CO 3 2 − + 2 H + ↽ − − ⇀ H 2 CO 3 β 2 = [ H 2 CO 3 ] [ H + ] 2 [ CO 3 2 − ] {\displaystyle {\begin{aligned}{\ce {CO3^{2-}{}+ H+{}<=> HCO3^-}}&&\beta _{1}={\frac {[{\ce {HCO3^-}}]}{[{\ce {H+}}][{\ce {CO3^{2-}}}]}}\\{\ce {CO3^{2-}{}+ 2H+{}<=> H2CO3}}&&\beta _{2}={\frac {[{\ce {H2CO3}}]}{[{\ce {H+}}]^{2}[{\ce {CO3^{2-}}}]}}\end{aligned}}} where brackets indicate the concentration of species . At 25 °C, these equilibria empirically satisfy [ 20 ] log ( β 1 ) = 0 .54 I 2 − 0 .96 I + 9 .93 log ( β 2 ) = − 2 .5 I 2 − 0 .043 I + 16 .07 {\displaystyle {\begin{alignedat}{6}\log(\beta _{1})=&&0&.54&I^{2}-0&.96&I+&&9&.93\\\log(\beta _{2})=&&-2&.5&I^{2}-0&.043&I+&&16&.07\end{alignedat}}} log( β 1 ) decreases with increasing I , as does log( β 2 ) . In a solution absent other ions (e.g. I = 0 ), these curves imply the following stepwise dissociation constants : p K 1 = log ( β 2 ) − log ( β 1 ) = 6.77 p K 2 = log ( β 1 ) = 9.93 {\displaystyle {\begin{alignedat}{3}p{\text{K}}_{1}&=\log(\beta _{2})-\log(\beta _{1})&=6.77\\p{\text{K}}_{2}&=\log(\beta _{1})&=9.93\end{alignedat}}} Direct values for these constants in the literature include p K 1 = 6.35 and p K 2 - p K 1 = 3.49 . [ 21 ]
To interpret these numbers, note that two chemical species in an acid equilibrium are equiconcentrated when p K = p H . In particular, the extracellular fluid ( cytosol ) in biological systems exhibits p H ≈ 7.2 , so that carbonic acid will be almost 50%-dissociated at equilibrium.
The Bjerrum plot shows typical equilibrium concentrations, in solution, in seawater , of carbon dioxide and the various species derived from it, as a function of pH . [ 8 ] [ 9 ] As human industrialization has increased the proportion of carbon dioxide in Earth's atmosphere , the proportion of carbon dioxide dissolved in sea- and freshwater as carbonic acid is also expected to increase. This rise in dissolved acid is also expected to acidify those waters, generating a decrease in pH. [ 22 ] [ 23 ] It has been estimated that the increase in dissolved carbon dioxide has already caused the ocean's average surface pH to decrease by about 0.1 from pre-industrial levels. | https://en.wikipedia.org/wiki/H2CO3 |
Thiocarbonic acid is an acid with the chemical formula H 2 CS 3 (or S=C(SH) 2 ). It is an analog of carbonic acid H 2 CO 3 (or O=C(OH) 2 ), in which all oxygen atoms are replaced with sulfur atoms. It is an unstable hydrophobic red oily liquid. [ 1 ]
It is often referred to as trithiocarbonic acid so as to differentiate it from other carbonic acids containing sulfur, such as monothiocarbonic O , O -acid S=C(OH) 2 , monothiocarbonic O , S -acid O=C(OH)(SH) , dithiocarbonic O , S -acid S=C(OH)(SH) and dithiocarbonic S , S -acid O=C(SH) 2 (see thiocarbonates ).
It was first reported in brief by Zeise in 1824 and later in more detail by Berzelius in 1826, [ 2 ] in both cases it was produced by the action of carbon disulfide on a hydrosulfide salt (e.g. potassium hydrosulfide ). [ 3 ]
Treatment with acids liberates the thiocarbonic acid as a red oil:
Both the acid and many of its salts are unstable and decompose via the release of carbon disulfide, particularly upon heating:
An improved synthesis involves addition of barium trithiocarbonate to hydrochloric acid at 0 °C. This method provided samples with which many measurement have been made. [ 1 ]
Despite its lability, crystals of thiocarbonic acid have been examined by X-ray crystallography , which confirms the anticipated molecular structure of a trigonal planar molecular geometry at the central carbon atom. The C-S bond lengths range from 1.69 to 1.77 Å . [ 4 ]
Thiocarbonic acid is acidic , with the first p K a being around 2. The second p K a is near 7. It dissolves S 8 , but does not react with it. [ 1 ]
Salts and esters of trithiocarbonic acid are called trithiocarbonates , and they are sometimes called thioxanthates .
Thiocarbonic acid reacts with bifunctional reagents to give rings . 1,2-Dichloroethane gives ethylene trithiocarbonate ( S=CS 2 (CH 2 ) 2 ). Oxalyl chloride gives oxalyl trithiocarbonate ( S=CS 2 (C=O) 2 ).
Thiocarbonic acid currently has no significant applications. Its esters find use in RAFT polymerization . | https://en.wikipedia.org/wiki/H2CS3 |
Chromic acid is a chemical compound with the chemical formula H 2 Cr O 4 . It is also a jargon for a solution formed by the addition of sulfuric acid to aqueous solutions of dichromate . It consists at least in part of chromium trioxide . [ 3 ]
The term "chromic acid" is usually used for a mixture made by adding concentrated sulfuric acid to a dichromate , which may contain a variety of compounds, including solid chromium trioxide . This kind of chromic acid may be used as a cleaning mixture for glass. Chromic acid may also refer to the molecular species, H 2 CrO 4 of which the trioxide is the anhydride . Chromic acid features chromium in an oxidation state of +6 (and a valence of VI or 6). It is a strong and corrosive oxidizing agent and a moderate carcinogen .
Molecular chromic acid, H 2 CrO 4 , in principle, resembles sulfuric acid , H 2 SO 4 . It would ionize accordingly:
The p K a for the equilibrium is not well characterized. Reported values vary between about −0.8 to 1.6. [ 4 ] The structure of the mono anion has been determined by X-ray crystallography . In this tetrahedral oxyanion, three Cr-O bond lengths are 156 pm and the Cr-OH bond is 201 pm [ 5 ]
[HCrO 4 ] − condenses to form dichromate:
Furthermore, the dichromate can be protonated:
Loss of the second proton occurs in the pH range 4–8, making the ion [HCrO 4 ] − a weak acid . [ citation needed ]
Molecular chromic acid could in principle be made by adding chromium trioxide to water ( cf. manufacture of sulfuric acid ).
In practice, the reverse reaction occurs: molecular chromic acid dehydrates . Some insights can be gleaned from observations on the reaction of dichromate solutions with sulfuric acid. The first colour change from orange to red signals the conversion of dichromate to chromic acid. Under these conditions deep red crystals of chromium trioxide precipitate from the mixture, without further colour change.
Chromium trioxide is the anhydride of molecular chromic acid. It is a Lewis acid and can react with a Lewis base, such as pyridine in a non-aqueous medium such as dichloromethane ( Collins reagent ).
Higher chromic acids with the formula H 2 Cr n O (3 n +1) are probable components of concentrated solutions of chromic acid.
Chromic acid is an intermediate in chromium plating, and is also used in ceramic glazes, and colored glass. Because a solution of chromic acid in sulfuric acid (also known as a sulfochromic mixture or chromosulfuric acid ) is a powerful oxidizing agent , it can be used to clean laboratory glassware , particularly of otherwise insoluble organic residues. This application has declined due to environmental concerns. [ 8 ] Furthermore, the acid leaves trace amounts of paramagnetic chromic ions ( Cr 3+ ) that can interfere with certain applications, such as NMR spectroscopy . This is especially the case for NMR tubes . [ 9 ] Piranha solution can be used for the same task, without leaving metallic residues behind.
Chromic acid was widely used in the musical instrument repair industry, due to its ability to "brighten" raw brass . A chromic acid dip leaves behind a bright yellow patina on the brass. Due to growing health and environmental concerns, many have discontinued use of this chemical in their repair shops.
It was used in hair dye in the 1940s, under the name Melereon . [ 10 ]
It is used as a bleach in processing black and white photographic reversal film . [ 11 ]
Chromic acid is capable of oxidizing many kinds of organic compounds and many variations on this reagent have been developed:
In organic chemistry , dilute solutions of chromic acid can be used to oxidize primary or secondary alcohols to the corresponding aldehydes and ketones . Similarly, it can also be used to oxidize an aldehyde to its corresponding carboxylic acid . Tertiary alcohols and ketones are unaffected. Because the oxidation is signaled by a color change from orange to brownish green (indicating chromium being reduced from oxidation state +6 to +3), chromic acid is commonly used as a lab reagent in high school or undergraduate college chemistry as a qualitative analytical test for the presence of primary or secondary alcohols , or aldehydes. [ 12 ]
In oxidations of alcohols or aldehydes into carboxylic acids , chromic acid is one of several reagents, including several that are catalytic. For example, nickel(II) salts catalyze oxidations by bleach (hypochlorite). [ 17 ] Aldehydes are relatively easily oxidized to carboxylic acids, and mild oxidizing agents are sufficient. Silver(I) compounds have been used for this purpose. Each oxidant offers advantages and disadvantages. Instead of using chemical oxidants, electrochemical oxidation is often possible.
Hexavalent chromium compounds (including chromium trioxide, chromic acids, chromates, chlorochromates) are toxic and carcinogenic . Chromium trioxide and chromic acids are strong oxidizers and may react violently if mixed with easily oxidizable organic substances.
Chromic acid burns are treated with a dilute sodium thiosulfate solution. [ 18 ] | https://en.wikipedia.org/wiki/H2Cr2O7 |
Chromic acid is a chemical compound with the chemical formula H 2 Cr O 4 . It is also a jargon for a solution formed by the addition of sulfuric acid to aqueous solutions of dichromate . It consists at least in part of chromium trioxide . [ 3 ]
The term "chromic acid" is usually used for a mixture made by adding concentrated sulfuric acid to a dichromate , which may contain a variety of compounds, including solid chromium trioxide . This kind of chromic acid may be used as a cleaning mixture for glass. Chromic acid may also refer to the molecular species, H 2 CrO 4 of which the trioxide is the anhydride . Chromic acid features chromium in an oxidation state of +6 (and a valence of VI or 6). It is a strong and corrosive oxidizing agent and a moderate carcinogen .
Molecular chromic acid, H 2 CrO 4 , in principle, resembles sulfuric acid , H 2 SO 4 . It would ionize accordingly:
The p K a for the equilibrium is not well characterized. Reported values vary between about −0.8 to 1.6. [ 4 ] The structure of the mono anion has been determined by X-ray crystallography . In this tetrahedral oxyanion, three Cr-O bond lengths are 156 pm and the Cr-OH bond is 201 pm [ 5 ]
[HCrO 4 ] − condenses to form dichromate:
Furthermore, the dichromate can be protonated:
Loss of the second proton occurs in the pH range 4–8, making the ion [HCrO 4 ] − a weak acid . [ citation needed ]
Molecular chromic acid could in principle be made by adding chromium trioxide to water ( cf. manufacture of sulfuric acid ).
In practice, the reverse reaction occurs: molecular chromic acid dehydrates . Some insights can be gleaned from observations on the reaction of dichromate solutions with sulfuric acid. The first colour change from orange to red signals the conversion of dichromate to chromic acid. Under these conditions deep red crystals of chromium trioxide precipitate from the mixture, without further colour change.
Chromium trioxide is the anhydride of molecular chromic acid. It is a Lewis acid and can react with a Lewis base, such as pyridine in a non-aqueous medium such as dichloromethane ( Collins reagent ).
Higher chromic acids with the formula H 2 Cr n O (3 n +1) are probable components of concentrated solutions of chromic acid.
Chromic acid is an intermediate in chromium plating, and is also used in ceramic glazes, and colored glass. Because a solution of chromic acid in sulfuric acid (also known as a sulfochromic mixture or chromosulfuric acid ) is a powerful oxidizing agent , it can be used to clean laboratory glassware , particularly of otherwise insoluble organic residues. This application has declined due to environmental concerns. [ 8 ] Furthermore, the acid leaves trace amounts of paramagnetic chromic ions ( Cr 3+ ) that can interfere with certain applications, such as NMR spectroscopy . This is especially the case for NMR tubes . [ 9 ] Piranha solution can be used for the same task, without leaving metallic residues behind.
Chromic acid was widely used in the musical instrument repair industry, due to its ability to "brighten" raw brass . A chromic acid dip leaves behind a bright yellow patina on the brass. Due to growing health and environmental concerns, many have discontinued use of this chemical in their repair shops.
It was used in hair dye in the 1940s, under the name Melereon . [ 10 ]
It is used as a bleach in processing black and white photographic reversal film . [ 11 ]
Chromic acid is capable of oxidizing many kinds of organic compounds and many variations on this reagent have been developed:
In organic chemistry , dilute solutions of chromic acid can be used to oxidize primary or secondary alcohols to the corresponding aldehydes and ketones . Similarly, it can also be used to oxidize an aldehyde to its corresponding carboxylic acid . Tertiary alcohols and ketones are unaffected. Because the oxidation is signaled by a color change from orange to brownish green (indicating chromium being reduced from oxidation state +6 to +3), chromic acid is commonly used as a lab reagent in high school or undergraduate college chemistry as a qualitative analytical test for the presence of primary or secondary alcohols , or aldehydes. [ 12 ]
In oxidations of alcohols or aldehydes into carboxylic acids , chromic acid is one of several reagents, including several that are catalytic. For example, nickel(II) salts catalyze oxidations by bleach (hypochlorite). [ 17 ] Aldehydes are relatively easily oxidized to carboxylic acids, and mild oxidizing agents are sufficient. Silver(I) compounds have been used for this purpose. Each oxidant offers advantages and disadvantages. Instead of using chemical oxidants, electrochemical oxidation is often possible.
Hexavalent chromium compounds (including chromium trioxide, chromic acids, chromates, chlorochromates) are toxic and carcinogenic . Chromium trioxide and chromic acids are strong oxidizers and may react violently if mixed with easily oxidizable organic substances.
Chromic acid burns are treated with a dilute sodium thiosulfate solution. [ 18 ] | https://en.wikipedia.org/wiki/H2CrO4 |
Ethylene ( IUPAC name: ethene ) is a hydrocarbon which has the formula C 2 H 4 or H 2 C=CH 2 . It is a colourless, flammable gas with a faint "sweet and musky " odour when pure. [ 7 ] It is the simplest alkene (a hydrocarbon with carbon–carbon double bonds ).
Ethylene is widely used in the chemical industry, and its worldwide production (over 150 million tonnes in 2016 [ 8 ] ) exceeds that of any other organic compound . [ 9 ] [ 10 ] Much of this production goes toward creating polythene , which is a widely used plastic containing polymer chains of ethylene units in various chain lengths. Production emits greenhouse gases , including methane from feedstock production and carbon dioxide from any non- sustainable energy used.
Ethylene is also an important natural plant hormone and is used in agriculture to induce ripening of fruits . [ 11 ] The hydrate of ethylene is ethanol .
This hydrocarbon has four hydrogen atoms bound to a pair of carbon atoms that are connected by a double bond . All six atoms that comprise ethylene are coplanar . The H-C-H angle is 117.4°, close to the 120° for ideal sp² hybridized carbon. The molecule is also relatively weak: rotation about the C-C bond is a very low energy process that requires breaking the π-bond by supplying heat at 50 °C. [ citation needed ]
The π-bond in the ethylene molecule is responsible for its useful reactivity. The double bond is a region of high electron density , thus it is susceptible to attack by electrophiles . Many reactions of ethylene are catalyzed by transition metals, which bind transiently to the ethylene using both the π and π* orbitals. [ citation needed ]
Being a simple molecule, ethylene is spectroscopically simple. Its UV-vis spectrum is still used as a test of theoretical methods. [ 12 ]
Major industrial reactions of ethylene include in order of scale: 1) polymerization , 2) oxidation , 3) halogenation and hydrohalogenation , 4) alkylation , 5) hydration , 6) oligomerization , and 7) hydroformylation . In the United States and Europe , approximately 90% of ethylene is used to produce ethylene oxide , ethylene dichloride , ethylbenzene and polyethylene . [ 13 ] Most of the reactions with ethylene are electrophilic addition . [ citation needed ]
Polyethylene production uses more than half of the world's ethylene supply. Polyethylene, also called polyethene and polythene , is the world's most widely used plastic. It is primarily used to make films in packaging , carrier bags and trash liners . Linear alpha-olefins , produced by oligomerization (formation of short-chain molecules) are used as precursors , detergents , plasticisers , synthetic lubricants , additives, and also as co-monomers in the production of polyethylenes. [ 13 ]
Ethylene is oxidized to produce ethylene oxide , a key raw material in the production of surfactants and detergents by ethoxylation . Ethylene oxide is also hydrolyzed to produce ethylene glycol , widely used as an automotive antifreeze as well as higher molecular weight glycols, glycol ethers , and polyethylene terephthalate . [ 14 ] [ 15 ]
Ethylene oxidation in the presence of a palladium catalyst can form acetaldehyde . This conversion remains a major industrial process (10M kg/y). [ 16 ] The process proceeds via the initial complexation of ethylene to a Pd(II) center. [ citation needed ]
Major intermediates from the halogenation and hydrohalogenation of ethylene include ethylene dichloride , ethyl chloride , and ethylene dibromide . The addition of chlorine entails " oxychlorination ", i.e. chlorine itself is not used. Some products derived from this group are polyvinyl chloride , trichloroethylene , perchloroethylene , methyl chloroform , polyvinylidene chloride and copolymers , and ethyl bromide . [ 17 ]
Major chemical intermediates from the alkylation with ethylene is ethylbenzene , precursor to styrene . Styrene is used principally in polystyrene for packaging and insulation, as well as in styrene-butadiene rubber for tires and footwear. On a smaller scale, ethyltoluene , ethylanilines, 1,4-hexadiene, and aluminium alkyls. Products of these intermediates include polystyrene , unsaturated polyesters and ethylene-propylene terpolymers . [ 17 ]
The hydroformylation (oxo reaction) of ethylene results in propionaldehyde , a precursor to propionic acid and n-propyl alcohol . [ 17 ]
Ethylene has long represented the major nonfermentative precursor to ethanol . The original method entailed its conversion to diethyl sulfate , followed by hydrolysis. The main method practiced since the mid-1990s is the direct hydration of ethylene catalyzed by solid acid catalysts : [ 18 ]
Ethylene is dimerized by hydrovinylation to give n -butenes using processes licensed by Lummus or IFP . The Lummus process produces mixed n -butenes (primarily 2-butenes ) while the IFP process produces 1-butene . 1-Butene is used as a comonomer in the production of certain kinds of polyethylene . [ 19 ]
Ethylene is a hormone that affects the ripening and flowering of many plants. It is widely used to control freshness in horticulture and fruits . [ 20 ] The scrubbing of naturally occurring ethylene delays ripening. [ 21 ] Adsorption of ethylene by nets coated in titanium dioxide gel has also been shown to be effective. [ 22 ]
An example of a niche use is as an anesthetic agent (in an 85% ethylene/15% oxygen ratio). [ 23 ] It is also used as a refrigerant gas for low temperature applications under the name R-1150. [ 24 ]
Global ethylene production was 107 million tonnes in 2005, [ 9 ] 109 million tonnes in 2006, [ 25 ] 138 million tonnes in 2010, and 141 million tonnes in 2011. [ 26 ] By 2013, ethylene was produced by at least 117 companies in 32 countries. To meet the ever-increasing demand for ethylene, sharp increases in production facilities are added globally, particularly in the Mideast and in China . [ 27 ] Production emits greenhouse gas , namely significant amounts of carbon dioxide. [ 28 ]
Ethylene is produced by several methods in the petrochemical industry . A primary method is steam cracking (SC) where hydrocarbons and steam are heated to 750–950 °C. This process converts large hydrocarbons into smaller ones and introduces unsaturation. When ethane is the feedstock, ethylene is the product. Ethylene is separated from the resulting mixture by repeated compression and distillation . [ 17 ] In Europe and Asia, ethylene is obtained mainly from cracking naphtha, gasoil and condensates with the coproduction of propylene, C4 olefins and aromatics (pyrolysis gasoline). [ 29 ] Other procedures employed for the production of ethylene include Fischer-Tropsch synthesis and methanol-to-olefins (MTO). [ 30 ]
Although of great value industrially, ethylene is rarely synthesized in the laboratory and is ordinarily purchased. [ 31 ] It can be produced via dehydration of ethanol with sulfuric acid or in the gas phase with aluminium oxide or activated alumina . [ 32 ]
Ethylene is produced from methionine in nature. The immediate precursor is 1-aminocyclopropane-1-carboxylic acid . [ 33 ]
Ethylene is a fundamental ligand in transition metal alkene complexes . One of the first organometallic compounds, Zeise's salt is a complex of ethylene. Useful reagents containing ethylene include Pt(PPh 3 ) 2 (C 2 H 4 ) and Rh 2 Cl 2 (C 2 H 4 ) 4 . The Rh-catalysed hydroformylation of ethylene is conducted on an industrial scale to provide propionaldehyde . [ 35 ]
Some geologists and scholars believe that the famous Greek Oracle at Delphi (the Pythia ) went into her trance-like state as an effect of ethylene rising from ground faults. [ 36 ]
Ethylene appears to have been discovered by Johann Joachim Becher , who obtained it by heating ethanol with sulfuric acid; [ 37 ] he mentioned the gas in his Physica Subterranea (1669). [ 38 ] Joseph Priestley also mentions the gas in his Experiments and observations relating to the various branches of natural philosophy: with a continuation of the observations on air (1779), where he reports that Jan Ingenhousz saw ethylene synthesized in the same way by a Mr. Enée in Amsterdam in 1777 and that Ingenhousz subsequently produced the gas himself. [ 39 ] The properties of ethylene were studied in 1795 by four Dutch chemists, Johann Rudolph Deimann, Adrien Paets van Troostwyck, Anthoni Lauwerenburgh and Nicolas Bondt, who found that it differed from hydrogen gas and that it contained both carbon and hydrogen. [ 40 ] This group also discovered that ethylene could be combined with chlorine to produce the Dutch oil , 1,2-dichloroethane ; this discovery gave ethylene the name used for it at that time, olefiant gas (oil-making gas.) [ 41 ] The term olefiant gas is in turn the etymological origin of the modern word "olefin", the class of hydrocarbons in which ethylene is the first member. [ citation needed ]
In the mid-19th century, the suffix -ene (an Ancient Greek root added to the end of female names meaning "daughter of") was widely used to refer to a molecule or part thereof that contained one fewer hydrogen atoms than the molecule being modified. Thus, ethylene ( C 2 H 4 ) was the "daughter of ethyl " ( C 2 H 5 ). The name ethylene was used in this sense as early as 1852. [ 42 ]
In 1866, the German chemist August Wilhelm von Hofmann proposed a system of hydrocarbon nomenclature in which the suffixes -ane, -ene, -ine, -one, and -une were used to denote the hydrocarbons with 0, 2, 4, 6, and 8 fewer hydrogens than their parent alkane . [ 43 ] In this system, ethylene became ethene . Hofmann's system eventually became the basis for the Geneva nomenclature approved by the International Congress of Chemists in 1892, which remains at the core of the IUPAC nomenclature. However, by that time, the name ethylene was deeply entrenched, and it remains in wide use today, especially in the chemical industry.
Following experimentation by Luckhardt, Crocker, and Carter at the University of Chicago, [ 44 ] ethylene was used as an anesthetic. [ 45 ] [ 7 ] It remained in use through the 1940s, even while chloroform was being phased out. Its pungent odor and its explosive nature limit its use today. [ 46 ]
The 1979 IUPAC nomenclature rules made an exception for retaining the non-systematic name ethylene ; [ 47 ] however, this decision was reversed in the 1993 rules, [ 48 ] and it remains unchanged in the newest 2013 recommendations, [ 49 ] so the IUPAC name is now ethene . In the IUPAC system, the name ethylene is reserved for the divalent group -CH 2 CH 2 -. Hence, names like ethylene oxide and ethylene dibromide are permitted, but the use of the name ethylene for the two-carbon alkene is not. Nevertheless, use of the name ethylene for H 2 C=CH 2 (and propylene for H 2 C=CHCH 3 ) is still prevalent among chemists in North America. [ 50 ]
"A key factor affecting petrochemicals life-cycle emissions is the methane intensity of feedstocks, especially in the production segment." [ 51 ] Emissions from cracking of naptha and natural gas (common in the US as gas is cheap there) depend a lot on the source of energy (for example gas burnt to provide high temperatures [ 52 ] ) but that from naptha is certainly more per kg of feedstock. [ 53 ] Both steam cracking and production from natural gas via ethane are estimated to emit 1.8 to 2kg of CO2 per kg ethylene produced, [ 54 ] totalling over 260 million tonnes a year. [ 55 ] This is more than all other manufactured chemicals except cement and ammonia. [ 56 ] According to a 2022 report using renewable or nuclear energy could cut emissions by almost half. [ 53 ]
Like all hydrocarbons, ethylene is a combustible asphyxiant . It is listed as an IARC group 3 agent , since there is no current evidence that it causes cancer in humans. [ 57 ] | https://en.wikipedia.org/wiki/H2C⚌CH2 |
Formaldehyde ( / f ɔːr ˈ m æ l d ɪ h aɪ d / ⓘ for- MAL -di-hide , US also / f ə r -/ ⓘ fər- ) ( systematic name methanal ) is an organic compound with the chemical formula CH 2 O and structure H−CHO , more precisely H 2 C=O . The compound is a pungent, colourless gas that polymerises spontaneously into paraformaldehyde . It is stored as aqueous solutions ( formalin ), which consists mainly of the hydrate CH 2 (OH) 2 . It is the simplest of the aldehydes ( R−CHO ). As a precursor to many other materials and chemical compounds, in 2006 the global production of formaldehyde was estimated at 12 million tons per year. [ 14 ] It is mainly used in the production of industrial resins , e.g., for particle board and coatings .
Formaldehyde also occurs naturally. It is derived from the degradation of serine , dimethylglycine , and lipids . Demethylases act by converting N-methyl groups to formaldehyde. [ 15 ]
Formaldehyde is classified as a group 1 carcinogen [ note 1 ] [ 17 ] and can cause respiratory and skin irritation upon exposure. [ 16 ]
Formaldehyde is more complicated than many simple carbon compounds in that it adopts several diverse forms. These compounds can often be used interchangeably and can be interconverted. [ citation needed ]
A small amount of stabilizer , such as methanol , is usually added to suppress oxidation and polymerization . A typical commercial-grade formalin may contain 10–12% methanol in addition to various metallic impurities.
"Formaldehyde" was first used as a generic trademark in 1893 following a previous trade name, "formalin". [ 18 ]
Molecular formaldehyde contains a central carbon atom with a double bond to the oxygen atom and a single bond to each hydrogen atom . This structure is summarised by the condensed formula H 2 C=O. [ 19 ] The molecule is planar, Y-shaped and its molecular symmetry belongs to the C 2v point group . [ 20 ] The precise molecular geometry of gaseous formaldehyde has been determined by gas electron diffraction [ 19 ] [ 21 ] and microwave spectroscopy . [ 22 ] [ 23 ] The bond lengths are 1.21 Å for the carbon–oxygen bond [ 19 ] [ 21 ] [ 22 ] [ 23 ] [ 24 ] and around 1.11 Å for the carbon–hydrogen bond , [ 19 ] [ 21 ] [ 22 ] [ 23 ] while the H–C–H bond angle is 117°, [ 22 ] [ 23 ] close to the 120° angle found in an ideal trigonal planar molecule . [ 19 ] Some excited electronic states of formaldehyde are pyramidal rather than planar as in the ground state . [ 24 ]
Processes in the upper atmosphere contribute more than 80% of the total formaldehyde in the environment. [ 25 ] Formaldehyde is an intermediate in the oxidation (or combustion ) of methane , as well as of other carbon compounds, e.g. in forest fires , automobile exhaust, and tobacco smoke . When produced in the atmosphere by the action of sunlight and oxygen on atmospheric methane and other hydrocarbons , it becomes part of smog . Formaldehyde has also been detected in outer space.
Formaldehyde and its adducts are ubiquitous in nature. Food may contain formaldehyde at levels 1–100 mg/kg. [ 26 ] Formaldehyde, formed in the metabolism of the amino acids serine and threonine , is found in the bloodstream of humans and other primates at concentrations of approximately 50 micromolar . [ 27 ] Experiments in which animals are exposed to an atmosphere containing isotopically labeled formaldehyde have demonstrated that even in deliberately exposed animals, the majority of formaldehyde-DNA adducts found in non-respiratory tissues are derived from endogenously produced formaldehyde. [ 28 ]
Formaldehyde does not accumulate in the environment, because it is broken down within a few hours by sunlight or by bacteria present in soil or water. Humans metabolize formaldehyde quickly, converting it to formic acid . [ 29 ] [ 30 ] It nonetheless presents significant health concerns , as a contaminant .
Formaldehyde appears to be a useful probe in astrochemistry due to prominence of the 1 10 ←1 11 and 2 11 ←2 12 K -doublet transitions. It was the first polyatomic organic molecule detected in the interstellar medium . [ 31 ] Since its initial detection in 1969, it has been observed in many regions of the galaxy . Because of the widespread interest in interstellar formaldehyde, it has been extensively studied, yielding new extragalactic sources. [ 32 ] A proposed mechanism for the formation is the hydrogenation of CO ice: [ 33 ]
HCN , HNC , H 2 CO, and dust have also been observed inside the comae of comets C/2012 F6 (Lemmon) and C/2012 S1 (ISON) . [ 34 ] [ 35 ]
Formaldehyde was discovered in 1859 by the Russian chemist Aleksandr Butlerov (1828–1886) when he attempted to synthesize methanediol ("methylene glycol") from iodomethane and silver oxalate . [ 36 ] In his paper, Butlerov referred to formaldehyde as "dioxymethylen" (methylene dioxide) because his empirical formula for it was incorrect, as atomic weights were not precisely determined until the Karlsruhe Congress .
The compound was identified as an aldehyde by August Wilhelm von Hofmann , who first announced its production by passing methanol vapor in air over hot platinum wire. [ 37 ] [ 38 ] With modifications, Hofmann's method remains the basis of the present day industrial route.
Solution routes to formaldehyde also entail oxidation of methanol or iodomethane . [ 39 ]
Formaldehyde is produced industrially by the catalytic oxidation of methanol . The most common catalysts are silver metal (i.e. the FASIL process ), iron(III) oxide , [ 40 ] iron molybdenum oxides (e.g. iron(III) molybdate ) with a molybdenum -enriched surface, [ 41 ] or vanadium oxides . In the commonly used formox process , methanol and oxygen react at c. 250–400 °C in presence of iron oxide in combination with molybdenum and/or vanadium to produce formaldehyde according to the chemical equation : [ 42 ]
The silver-based catalyst usually operates at a higher temperature, about 650 °C. Two chemical reactions on it simultaneously produce formaldehyde: that shown above and the dehydrogenation reaction:
In principle, formaldehyde could be generated by oxidation of methane , but this route is not industrially viable because the methanol is more easily oxidized than methane. [ 42 ]
Formaldehyde is produced via several enzyme-catalyzed routes. [ 43 ] Living beings, including humans, produce formaldehyde as part of their metabolism. Formaldehyde is key to several bodily functions (e.g. epigenetics [ 27 ] ), but its amount must also be tightly controlled to avoid self-poisoning. [ 44 ]
Formaldehyde is catabolized by alcohol dehydrogenase ADH5 and aldehyde dehydrogenase ALDH2 . [ 45 ]
Formaldehyde is a building block in the synthesis of many other compounds of specialised and industrial significance. It exhibits most of the chemical properties of other aldehydes but is more reactive. [ 46 ]
Monomeric CH 2 O is a gas and is rarely encountered in the laboratory. Aqueous formaldehyde, unlike some other small aldehydes (which need specific conditions to oligomerize through aldol condensation ) oligomerizes spontaneously at a common state. The trimer 1,3,5-trioxane, (CH 2 O) 3 , is a typical oligomer. Many cyclic oligomers of other sizes have been isolated. Similarly, formaldehyde hydrates to give the geminal diol methanediol , which condenses further to form hydroxy-terminated oligomers HO(CH 2 O) n H. The polymer is called paraformaldehyde . The higher concentration of formaldehyde—the more equilibrium shifts towards polymerization. Diluting with water or increasing the solution temperature, as well as adding alcohols (such as methanol or ethanol) lowers that tendency.
Gaseous formaldehyde polymerizes at active sites on vessel walls, but the mechanism of the reaction is unknown. [ 47 ] Small amounts of hydrogen chloride , boron trifluoride , or stannic chloride present in gaseous formaldehyde provide the catalytic effect and make the polymerization rapid. [ 48 ]
Formaldehyde forms cross-links by first combining with a protein to form methylol , which loses a water molecule to form a Schiff base . [ 49 ] The Schiff base can then react with DNA or protein to create a cross-linked product. [ 49 ] This reaction is the basis for the most common process of chemical fixation .
Formaldehyde is readily oxidized by atmospheric oxygen into formic acid . For this reason, commercial formaldehyde is typically contaminated with formic acid. Formaldehyde can be hydrogenated into methanol .
In the Cannizzaro reaction , formaldehyde and base react to produce formic acid and methanol, a disproportionation reaction .
Formaldehyde reacts with many compounds, resulting in hydroxymethylation :
The resulting hydroxymethyl derivatives typically react further. Thus, amines give hexahydro-1,3,5-triazines :
Similarly, when combined with hydrogen sulfide , it forms trithiane : [ 50 ]
In the presence of acids, it participates in electrophilic aromatic substitution reactions with aromatic compounds resulting in hydroxymethylated derivatives:
When conducted in the presence of hydrogen chloride, the product is the chloromethyl compound, as described in the Blanc chloromethylation . If the arene is electron-rich, as in phenols, elaborate condensations ensue. With 4-substituted phenols one obtains calixarenes . [ 51 ] Phenol results in polymers.
Many amino acids react with formaldehyde. [ 43 ] Cysteine converts to thioproline .
Formaldehyde is a common precursor to more complex compounds and materials. In approximate order of decreasing consumption, products generated from formaldehyde include urea formaldehyde resin , melamine resin , phenol formaldehyde resin , polyoxymethylene plastics , 1,4-butanediol , and methylene diphenyl diisocyanate . [ 42 ] The textile industry uses formaldehyde-based resins as finishers to make fabrics crease-resistant. [ 52 ]
When condensed with phenol , urea , or melamine , formaldehyde produces, respectively, hard thermoset phenol formaldehyde resin, urea formaldehyde resin, and melamine resin. These polymers are permanent adhesives used in plywood and carpeting . They are also foamed to make insulation , or cast into moulded products. Production of formaldehyde resins accounts for more than half of formaldehyde consumption.
Formaldehyde is also a precursor to polyfunctional alcohols such as pentaerythritol , which is used to make paints and explosives . Other formaldehyde derivatives include methylene diphenyl diisocyanate, an important component in polyurethane paints and foams, and hexamine , which is used in phenol-formaldehyde resins as well as the explosive RDX .
Condensation with acetaldehyde affords pentaerythritol , a chemical necessary in synthesizing PETN , a high explosive: [ 53 ]
An aqueous solution of formaldehyde can be useful as a disinfectant as it kills most bacteria and fungi (including their spores). It is used as an additive in vaccine manufacturing to inactivate toxins and pathogens. [ 54 ] Formaldehyde releasers are used as biocides in personal care products such as cosmetics. Although present at levels not normally considered harmful, they are known to cause allergic contact dermatitis in certain sensitised individuals. [ 55 ]
Aquarists use formaldehyde as a treatment for the parasites Ichthyophthirius multifiliis and Cryptocaryon irritans . [ 56 ] Formaldehyde is one of the main disinfectants recommended for destroying anthrax . [ 57 ]
Formaldehyde is also approved for use in the manufacture of animal feeds in the US. It is an antimicrobial agent used to maintain complete animal feeds or feed ingredients Salmonella negative for up to 21 days. [ 58 ]
Formaldehyde preserves or fixes tissue or cells. The process involves cross-linking of primary amino groups . The European Union has banned the use of formaldehyde as a biocide (including embalming ) under the Biocidal Products Directive (98/8/EC) due to its carcinogenic properties. [ 59 ] [ 60 ] Countries with a strong tradition of embalming corpses, such as Ireland and other colder-weather countries, have raised concerns. Despite reports to the contrary, [ 61 ] no decision on the inclusion of formaldehyde on Annex I of the Biocidal Products Directive for product-type 22 (embalming and taxidermist fluids) had been made as of September 2009 [update] . [ 62 ]
Formaldehyde-based crosslinking is exploited in ChIP-on-chip or ChIP-sequencing genomics experiments, where DNA-binding proteins are cross-linked to their cognate binding sites on the chromosome and analyzed to determine what genes are regulated by the proteins. Formaldehyde is also used as a denaturing agent in RNA gel electrophoresis , preventing RNA from forming secondary structures. A solution of 4% formaldehyde fixes pathology tissue specimens at about one mm per hour at room temperature.
Formaldehyde and 18 M (concentrated) sulfuric acid makes Marquis reagent —which can identify alkaloids and other compounds.
In photography, formaldehyde is used in low concentrations for the process C-41 (color negative film) stabilizer in the final wash step, [ 63 ] as well as in the process E-6 pre-bleach step, to make it unnecessary in the final wash. Due to improvements in dye coupler chemistry, more modern (2006 or later) E-6 and C-41 films do not need formaldehyde, as their dyes are already stable.
In view of its widespread use, toxicity, and volatility, formaldehyde poses a significant danger to human health. [ 64 ] [ 65 ] In 2011, the US National Toxicology Program described formaldehyde as "known to be a human carcinogen". [ 66 ] [ 67 ] [ 68 ]
Concerns are associated with chronic (long-term) exposure by inhalation as may happen from thermal or chemical decomposition of formaldehyde-based resins and the production of formaldehyde resulting from the combustion of a variety of organic compounds (for example, exhaust gases). As formaldehyde resins are used in many construction materials , it is one of the more common indoor air pollutants . [ 69 ] [ 70 ] At concentrations above 0.1 ppm in air, formaldehyde can irritate the eyes and mucous membranes . [ 71 ] Formaldehyde inhaled at this concentration may cause headaches, a burning sensation in the throat, and difficulty breathing, and can trigger or aggravate asthma symptoms. [ 72 ] [ 73 ]
The CDC considers formaldehyde as a systemic poison. Formaldehyde poisoning can cause permanent changes in the nervous system 's functions. [ 74 ]
A 1988 Canadian study of houses with urea-formaldehyde foam insulation found that formaldehyde levels as low as 0.046 ppm were positively correlated with eye and nasal irritation. [ 75 ] A 2009 review of studies has shown a strong association between exposure to formaldehyde and the development of childhood asthma . [ 76 ]
A theory was proposed for the carcinogenesis of formaldehyde in 1978. [ 77 ] In 1987 the United States Environmental Protection Agency (EPA) classified it as a probable human carcinogen , and after more studies the WHO International Agency for Research on Cancer (IARC) in 1995 also classified it as a probable human carcinogen . Further information and evaluation of all known data led the IARC to reclassify formaldehyde as a known human carcinogen [ 78 ] associated with nasal sinus cancer and nasopharyngeal cancer . [ 79 ] Studies in 2009 and 2010 have also shown a positive correlation between exposure to formaldehyde and the development of leukemia , particularly myeloid leukemia . [ 80 ] [ 81 ] Nasopharyngeal and sinonasal cancers are relatively rare, with a combined annual incidence in the United States of < 4,000 cases. [ 82 ] [ 83 ] About 30,000 cases of myeloid leukemia occur in the United States each year. [ 84 ] [ 85 ] Some evidence suggests that workplace exposure to formaldehyde contributes to sinonasal cancers. [ 86 ] Professionals exposed to formaldehyde in their occupation, such as funeral industry workers and embalmers , showed an increased risk of leukemia and brain cancer compared with the general population. [ 87 ] Other factors are important in determining individual risk for the development of leukemia or nasopharyngeal cancer. [ 86 ] [ 88 ] [ 89 ] In yeast, formaldehyde is found to perturb pathways for DNA repair and cell cycle. [ 90 ]
In the residential environment, formaldehyde exposure comes from a number of routes; formaldehyde can be emitted by treated wood products, such as plywood or particle board , but it is produced by paints, varnishes , floor finishes, and cigarette smoking as well. [ 91 ] In July 2016, the U.S. EPA released a prepublication version of its final rule on Formaldehyde Emission Standards for Composite Wood Products. [ 92 ] These new rules impact manufacturers, importers, distributors, and retailers of products containing composite wood, including fiberboard, particleboard, and various laminated products, who must comply with more stringent record-keeping and labeling requirements. [ 93 ]
The U.S. EPA allows no more than 0.016 ppm formaldehyde in the air in new buildings constructed for that agency. [ 94 ] [ failed verification ] A U.S. EPA study found a new home measured 0.076 ppm when brand new and 0.045 ppm after 30 days. [ 95 ] The Federal Emergency Management Agency (FEMA) has also announced limits on the formaldehyde levels in trailers purchased by that agency. [ 96 ] The EPA recommends the use of "exterior-grade" pressed-wood products with phenol instead of urea resin to limit formaldehyde exposure, since pressed-wood products containing formaldehyde resins are often a significant source of formaldehyde in homes. [ 79 ]
The eyes are most sensitive to formaldehyde exposure: The lowest level at which many people can begin to smell formaldehyde ranges between 0.05 and 1 ppm. The maximum concentration value at the workplace is 0.3 ppm. [ 97 ] [ need quotation to verify ] In controlled chamber studies, individuals begin to sense eye irritation at about 0.5 ppm; 5 to 20 percent report eye irritation at 0.5 to 1 ppm; and greater certainty for sensory irritation occurred at 1 ppm and above. While some agencies have used a level as low as 0.1 ppm as a threshold for irritation, the expert panel found that a level of 0.3 ppm would protect against nearly all irritation. In fact, the expert panel found that a level of 1.0 ppm would avoid eye irritation—the most sensitive endpoint—in 75–95% of all people exposed. [ 98 ]
Formaldehyde levels in building environments are affected by a number of factors. These include the potency of formaldehyde-emitting products present, the ratio of the surface area of emitting materials to volume of space, environmental factors, product age, interactions with other materials, and ventilation conditions. Formaldehyde emits from a variety of construction materials, furnishings, and consumer products. The three products that emit the highest concentrations are medium density fiberboard , hardwood plywood, and particle board. Environmental factors such as temperature and relative humidity can elevate levels because formaldehyde has a high vapor pressure . Formaldehyde levels from building materials are the highest when a building first opens because materials would have less time to off-gas. Formaldehyde levels decrease over time as the sources suppress.
In operating rooms , formaldehyde is produced as a byproduct of electrosurgery and is present in surgical smoke, exposing surgeons and healthcare workers to potentially unsafe concentrations. [ 99 ]
Formaldehyde levels in air can be sampled and tested in several ways, including impinger, treated sorbent, and passive monitors. [ 100 ] The National Institute for Occupational Safety and Health (NIOSH) has measurement methods numbered 2016, 2541, 3500, and 3800. [ 101 ]
In June 2011, the twelfth edition of the National Toxicology Program (NTP) Report on Carcinogens (RoC) changed the listing status of formaldehyde from "reasonably anticipated to be a human carcinogen" to "known to be a human carcinogen." [ 66 ] [ 67 ] [ 68 ] Concurrently, a National Academy of Sciences (NAS) committee was convened and issued an independent review of the draft U.S. EPA IRIS assessment of formaldehyde, providing a comprehensive health effects assessment and quantitative estimates of human risks of adverse effects. [ 102 ]
For most people, irritation from formaldehyde is temporary and reversible, although formaldehyde can cause allergies and is part of the standard patch test series. In 2005–06, it was the seventh-most-prevalent allergen in patch tests (9.0%). [ 103 ] People with formaldehyde allergy are advised to avoid formaldehyde releasers as well (e.g., Quaternium-15 , imidazolidinyl urea , and diazolidinyl urea ). [ 104 ] People who suffer allergic reactions to formaldehyde tend to display lesions on the skin in the areas that have had direct contact with the substance, such as the neck or thighs (often due to formaldehyde released from permanent press finished clothing) or dermatitis on the face (typically from cosmetics). [ 55 ] Formaldehyde has been banned in cosmetics in both Sweden [ 105 ] and Japan . [ 106 ]
In humans, ingestion of as little as 30 millilitres (1.0 US fl oz) of 37% formaldehyde solution can cause death. Other symptoms associated with ingesting such a solution include gastrointestinal damage (vomiting, abdominal pain), and systematic damage (dizziness). [ 74 ] Testing for formaldehyde is by blood and/or urine by gas chromatography–mass spectrometry . Other methods to detect formaldehyde include infrared detection, gas detector tubes, gas detectors using electrochemical sensors, and high-performance liquid chromatography (HPLC). HPLC is the most sensitive. [ 107 ]
The fifteenth edition (2021) of the U.S. National Toxicology Program Report on Carcinogens notes that currently in the U.S. “The general population can be exposed to formaldehyde primarily from breathing indoor or outdoor air, from tobacco smoke, from use of cosmetic products containing formaldehyde, and, to a more limited extent, from ingestion of food and water.” Affected water includes groundwater, surface water, and bottled water. It also notes that occupational exposure can be significant. [ 108 ]
Formaldehyde in food can be present naturally, added as an inadvertent contaminant, or intentionally added as a preservative, disinfectant, or bacteriostatic agent . Cooking and smoking food can also result in formaldehyde being produced in food. Foods that the U.S. National Toxicology Program has reported to have higher levels compared to other foods are fish, seafood, and smoked ham. It also notes formaldehyde in food generally occurs in a bound form and that formaldehyde is unstable in an aqueous solution . [ 109 ]
Scandals have broken in both the 2005 Indonesia food scare and 2007 Vietnam food scare regarding the addition of formaldehyde to foods to extend shelf life. In 2011, after a four-year absence, Indonesian authorities found foods with formaldehyde being sold in markets in a number of regions across the country. [ 110 ] In August 2011, at least at two Carrefour supermarkets, the Central Jakarta Livestock and Fishery Sub-Department found cendol containing 10 parts per million of formaldehyde. [ 111 ] In 2014, the owner of two noodle factories in Bogor , Indonesia, was arrested for using formaldehyde in noodles. [ 112 ] Foods known to be contaminated included noodles, salted fish, and tofu. Chicken and beer were also rumored to be contaminated. In some places, such as China, manufacturers still use formaldehyde illegally as a preservative in foods, which exposes people to formaldehyde ingestion. [ 113 ]
In 2011 in Nakhon Ratchasima , Thailand, truckloads of rotten chicken were treated with formaldehyde for sale in which "a large network", including 11 slaughterhouses run by a criminal gang, were implicated. [ 114 ] In 2012, 1 billion rupiah (almost US$100,000) of fish imported from Pakistan to Batam , Indonesia, were found laced with formaldehyde. [ 115 ]
Formalin contamination of foods has been reported in Bangladesh , with stores and supermarkets selling fruits, fishes, and vegetables that have been treated with formalin to keep them fresh. [ 116 ] However, in 2015, a Formalin Control Bill was passed in the Parliament of Bangladesh with a provision of life-term imprisonment as the maximum punishment as well as a maximum fine of 2,000,000 BDT but not less than 500,000 BDT for importing, producing, or hoarding formalin without a license. [ 117 ]
In the early 1900s, formaldehyde was frequently added by US milk plants to milk bottles as a method of pasteurization due to the lack of knowledge and concern [ 118 ] regarding formaldehyde's toxicity. [ 119 ] [ 120 ]
Formaldehyde was one of the chemicals used in 19th century industrialised food production that was investigated by Dr. Harvey W. Wiley with his famous 'Poison Squad' as part of the US Department of Agriculture . This led to the 1906 Pure Food and Drug Act , a landmark event in the early history of food regulation in the United States . [ 121 ]
Formaldehyde is banned from use in certain applications (preservatives for liquid-cooling and processing systems, slimicides , metalworking-fluid preservatives, and antifouling products) under the Biocidal Products Directive. [ 122 ] [ 123 ] In the EU, the maximum allowed concentration of formaldehyde in finished products is 0.2%, and any product that exceeds 0.05% has to include a warning that the product contains formaldehyde. [ 55 ]
In the United States, Congress passed a bill July 7, 2010, regarding the use of formaldehyde in hardwood plywood , particle board , and medium density fiberboard . The bill limited the allowable amount of formaldehyde emissions from these wood products to 0.09 ppm, and required companies to meet this standard by January 2013. [ 124 ] The final U.S. EPA rule specified maximum emissions of "0.05 ppm formaldehyde for hardwood plywood, 0.09 ppm formaldehyde for particleboard, 0.11 ppm formaldehyde for medium-density fiberboard, and 0.13 ppm formaldehyde for thin medium-density fiberboard." [ 125 ]
Formaldehyde was declared a toxic substance by the 1999 Canadian Environmental Protection Act . [ 126 ]
The FDA is proposing a ban on hair relaxers with formaldehyde due to cancer concerns. [ 127 ] | https://en.wikipedia.org/wiki/H2C⚌O |
The fluoronium ion is an inorganic cation with the chemical formula H 2 F + . It is one of the cations found in fluoroantimonic acid . [ 1 ] The structure of the salt with the Sb 2 F − 11 anion, has been determined. [ 2 ] [ 3 ] The fluoronium ion is isoelectronic with the water molecule and the azanide ion .
The term can also refer to organyl substituted species of type H– + F –R, R– + F –R, or R 2 C=F + . In contrast to the heavier halogens, which have long been known to form open-chain halonium ions (such as [Me 2 Cl] + [Al(OTeF 5 ) 4 ] – ) [ 4 ] as well as cyclic haliranium ions, fluorine was not believed to form fluoronium ions of type R– + F –R until the recent characterization of a fluoronium ion locked in a designed cage structure by Lectka and coworkers. [ 5 ] Recent solvolysis experiments and NMR spectroscopic studies on a metastable [C–F–C] + fluoronium ion strongly support the dicoordinated fluoronium structure over the alternative rapidly equilibrating classical carbocation. Definitive structural proof of the symmetrical [C–F–C] + was reported by Riedel, Lectka, and coworkers by single crystal X-ray diffraction analysis. Besides its synthesis and crystallographic characterization as the [Sb 2 F 11 ] − salt, vibrational spectra could be recorded and a detailed analysis concerning the nature of the bonding situation in this fluoronium ion and its heavier halonium homologues was reported. [ 6 ] | https://en.wikipedia.org/wiki/H2F+ |
Sodium fluorosilicate
Hexafluorosilicic acid is an inorganic compound with the chemical formula H 2 SiF 6 . Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.
Hexafluorosilicic acid is produced naturally on a large scale in volcanoes. [ 2 ] [ 3 ] It is manufactured as a coproduct in the production of phosphate fertilizers . The resulting hexafluorosilicic acid is almost exclusively consumed as a precursor to aluminum trifluoride and synthetic cryolite , which are used in aluminium processing. Salts derived from hexafluorosilicic acid are called hexafluorosilicates .
Hexafluorosilicic acid has been crystallized as various hydrates. These include ( H 5 O 2 ) 2 SiF 6 , the more complicated (H 5 O 2 ) 2 SiF 6 ·2H 2 O, and (H 5 O 2 )(H 7 O 3 )SiF 6 ·4.5H 2 O. In all of these salts, the octahedral hexafluorosilicate anion is hydrogen bonded to the cations. [ 4 ]
Aqueous solutions of hexafluorosilicic acid are often described as H 2 SiF 6 .
Hexafluorosilicic acid is produced commercially from fluoride-containing minerals that also contain silicates. Specifically, apatite and fluorapatite are treated with sulfuric acid to give phosphoric acid , a precursor to several water-soluble fertilizers. This is called the wet phosphoric acid process . [ 5 ] As a by-product, approximately 50 kg of hexafluorosilicic acid is produced per tonne of HF owing to reactions involving silica-containing mineral impurities. [ 6 ] : 3
Some of the hydrogen fluoride (HF) produced during this process in turn reacts with silicon dioxide (SiO 2 ) impurities, which are unavoidable constituents of the mineral feedstock, to give silicon tetrafluoride . Thus formed, the silicon tetrafluoride reacts further with HF. [ citation needed ] The net process can be described as: [ 7 ] [ page needed ]
Hexafluorosilicic acid can also be produced by treating silicon tetrafluoride with hydrofluoric acid. [ 7 ]
Hexafluorosilic acid is only stable in hydrogen fluoride or acidic aqueous solutions. In any other circumstance, it acts as a source of hydrofluoric acid . Thus, for example, hexafluorosilicic acid pure or in oleum solution evolves silicon tetrafluoride until the residual hydrogen fluoride re-establishes equilibrium: [ 7 ]
In alkaline-to-neutral aqueous solutions, hexafluorosilicic acid readily hydrolyzes to fluoride anions and amorphous, hydrated silica ("SiO 2 "). Strong bases give fluorosilicate salts at first, but any stoichiometric excess begins hydrolysis. [ 7 ] At the concentrations usually used for water fluoridation , 99% hydrolysis occurs: [ 6 ] [ 8 ]
Neutralization of solutions of hexafluorosilicic acid with alkali metal bases produces the corresponding alkali metal fluorosilicate salts:
The resulting salt Na 2 SiF 6 is mainly used in water fluoridation . Related ammonium and barium salts are produced similarly for other applications. At room temperature 15–30% concentrated hexafluorosilicic acid undergoes similar reactions with chlorides , hydroxides , and carbonates of alkali and alkaline earth metals . [ 9 ]
Sodium hexafluorosilicate for instance may be produced by treating sodium chloride ( NaCl ) by hexafluorosilicic acid: [ 6 ] : 3 [ 10 ] : 7
Heating sodium hexafluorosilicate gives silicon tetrafluoride : [ 10 ] : 8
The majority of the hexafluorosilicic acid is converted to aluminium fluoride and synthetic cryolite . These materials are central to the conversion of aluminium ore into aluminium metal. The conversion to aluminium trifluoride is described as: [ 7 ]
Hexafluorosilicic acid is also converted to a variety of useful hexafluorosilicate salts. The potassium salt, Potassium fluorosilicate , is used in the production of porcelains, the magnesium salt for hardened concretes and as an insecticide, and the barium salts for phosphors.
Hexafluorosilicic acid and the salts are used as wood preservation agents. [ 11 ]
Hexafluorosilicic acid is also used as an electrolyte in the Betts electrolytic process for refining lead.
Hexafluorosilicic acid (identified as hydrofluorosilicic acid on the label) along with oxalic acid are the active ingredients used in Iron Out rust-removing cleaning products, which are essentially varieties of laundry sour .
H 2 SiF 6 is a specialized reagent in organic synthesis for cleaving Si–O bonds of silyl ethers . It is more reactive for this purpose than HF. It reacts faster with t - butyldimethysilyl ( TBDMS ) ethers than triisopropylsilyl ( TIPS ) ethers. [ 12 ]
The application of hexafluorosilicic acid to a calcium-rich surface such as concrete will give that surface some resistance to acid attack. [ 13 ]
Calcium fluoride (CaF 2 ) is an insoluble solid that is acid resistant.
Some rare minerals, encountered either within volcanic or coal-fire fumaroles, are salts of the hexafluorosilicic acid. Examples include ammonium hexafluorosilicate that naturally occurs as two polymorphs: cryptohalite and bararite . [ 14 ] [ 15 ] [ 16 ]
Hexafluorosilicic acid can release hydrogen fluoride (HF) when evaporated, so it has similar risks. Inhalation of the vapors may cause lung edema . Like hydrogen fluoride, it attacks glass and stoneware . [ 17 ] The LD 50 value of hexafluorosilicic acid is 430 mg/kg. [ 6 ] | https://en.wikipedia.org/wiki/H2F6Si |
Iron tetracarbonyl dihydride is the organometallic compound with the formula H 2 Fe(CO) 4 . This compound was the first transition metal hydride discovered. The complex is stable at low temperatures but decomposes rapidly at temperatures above –20 °C. [ 1 ]
Iron tetracarbonyl dihydride was first produced by Hieber and Leutert from iron pentacarbonyl , which is first converted to HFe(CO) − 4 : [ 2 ] [ 3 ]
Since the compound is thermally labile and sensitive to light, ideal conditions in 1930's Munich called for winter nights. The early method was called the "polar night synthesis."
As recommended by Hieber and Leutert, the compound can be purified by trap-to-trap distillation. [ 1 ] [ 4 ]
In iron tetracarbonyl hydride the Fe(CO) 4 group has C 2v molecular symmetry with a geometry intermediate between octahedral and tetrahedral . Viewed as an octahedral complex, the hydride ligands are cis . Viewed as a tetrahedral Fe(CO) 4 complex, the hydrides occupy adjacent faces of the tetrahedron. [ 5 ] Although the structure of tetracarbonyliron with the hydrogen atoms bound as a single H 2 ligand has been proposed as an intermediate in some rearrangement reactions, [ 6 ] the stable state for the compound has the two atoms as independent ligands. [ 7 ]
Iron tetracarbonyl dihydride undergoes rapid ligand substitutions by phosphorus ligands:
The substitution mechanism is proposed to entail transient formation of a 16e − formyl intermediate. [ 8 ]
H 2 Fe(CO) 4 has p K 1 of 6.8 and p K 2 of 15. [ 9 ] The monoanion [HFe(CO) 4 ] − has more extensive reaction chemistry because it is more stable than the dihydride. [ 10 ] [ 11 ] The monoanion is an intermediate in the homogeneous iron-carbonyl-catalyzed water-gas shift reaction (WGSR). The slow step in the WGSR is the proton transfer from water to the iron hydride anion. [ 12 ] | https://en.wikipedia.org/wiki/H2Fe(CO)4 |
Sinkankasite , mineral formula: H 2 MnAl(PO 4 ) 2 (OH)·6H 2 O , was named after John Sinkankas (1915–2002), noted author and mineral collector, Scripps Institute of Oceanography . [ 4 ] It is triclinic; as colorless, bladed to prismatic crystals up to 4 mm in length, often as divergent, radial aggregates and as pseudomorphs after triphlyte crystals; occurs in the Barker pegmatite (formerly Ferguson pegmatite ), east of Keystone, South Dakota , and in the Palermo pegmatite , North Groton, New Hampshire . [ 5 ]
This article about a specific phosphate mineral is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2MnAl(PO4)2(OH)·6H2O |
Molybdic acid refers to hydrated forms of molybdenum trioxide and related species. The monohydrate (MoO 3 ·H 2 O) and the dihydrate (MoO 3 ·2H 2 O) are well characterized. They are yellow diamagnetic solids.
Solid forms of molybdic acid are coordination polymers . The monohydrate MoO 3 ·H 2 O consists of layers of octahedrally coordinated MoO 5 ·(H 2 O) units where 4 vertices are shared. [ 3 ] The dihydrate (image shown above) has the same layer structure with the "extra" H 2 O molecule intercalated between the layers. [ 4 ]
In acidified aqueous solutions of molybdic acid, the complex MoO 3 (H 2 O) 3 is observed. Once again, molybdenum adopts octahedral molecular geometry , probably with three oxo ligands and three aquo ligands . [ 5 ]
The salts of molybdic acid are called molybdates . They arise by adding base to solutions of molybdic acid.
Many molybdenum oxides are used as heterogeneous catalysts , e.g. for oxidations . Molybdic acid and its salts are used to make the Froehde reagent for the presumptive identification of alkaloids. | https://en.wikipedia.org/wiki/H2MoO4 |
Tiomolibdic acid (trade name Decuprate ) is a chelating agent under investigation for the treatment of cancer and of Wilson's disease , [ 1 ] a rare and potentially fatal disease in which the body cannot regulate copper . It is developed by Wilson Therapeutics and used in form of the salt bis- choline tetrathiomolybdate.
Wilson's disease is an autosomal recessive genetic disorder that is manifested by serious hepatic , neurologic or psychiatric symptoms . The disease is fatal if left untreated. It is estimated that 1 individual in every 30,000 to 100,000 worldwide has Wilson's disease. [ 2 ]
Bis-choline tetrathiomolybdate has been evaluated in clinical trials in patients with various forms of cancer [ 3 ] [ 4 ] [ 5 ] and has received orphan designation in the US and EU as a potential therapy against Wilson's disease. [ 6 ] [ 7 ]
Tiomolibdic acid selectively forms highly stable complexes with copper and proteins. These complexes are then believed to be primarily excreted via the bile , restoring the normal excretion route of copper that is impaired in patients with Wilson's disease. [ 8 ] [ 9 ] [ 10 ]
The binding and excretion mechanism is stable; whereas many de-coppering agents form unstable complexes that are excreted via urine. [ 11 ]
As of November 2014, a Phase 2, multi-centre, open-label study was recruiting newly diagnosed Wilson's disease patients 18 and older to evaluate the efficacy and safety of bis-choline tetrathiomolybdate administration over a 24-week period. [ 12 ] [ 13 ]
As of 2016, tetrathiomolybdate had been tested in over 500 patients for up to seven years, primarily in oncology [ 3 ] [ 4 ] [ 5 ] [ 14 ] [ 15 ] [ 16 ] [ 17 ] [ 18 ] [ 19 ] [ 20 ] and Wilson's disease , [ 21 ] [ 22 ] [ 23 ] [ 24 ] as well as some other clinical pathologies. [ 25 ] [ 26 ]
The data suggest that bis-choline tetrathiomolybdate can rapidly lower and control toxic free copper levels and improve clinical symptoms in Wilson's disease patients. The data also suggest that it is generally well tolerated, with the potential for a reduced risk of neurological worsening after initiation of therapy compared to existing therapies. [ 22 ] [ 23 ] [ 24 ]
Previous clinical studies with bis-choline tetrathiomolybdate in oncology patients have shown that it can lower and maintain copper levels using a once or twice daily oral dosing. [ 4 ] [ 5 ] This may be helpful since untreated Wilson's disease may lead to death within several years of the onset of symptoms, [ 27 ] and medication use should continue throughout the patient's lifespan. Patient compliance is crucial for clinical improvement, and it is a particular challenge for Wilson's disease patients taking de-coppering treatments. [ 28 ]
Tiomolibdic acid is the recommended International nonproprietary name (INN). [ 29 ] | https://en.wikipedia.org/wiki/H2MoS4 |
A nitrenium ion (also called: aminylium ion or imidonium ion (obsolete)) in organic chemistry is a reactive intermediate based on nitrogen with both an electron lone pair and a positive charge and with two substituents ( R 2 N + ). [ 1 ] [ 2 ] Nitrenium ions are isoelectronic with carbenes , and can exist in either a singlet or a triplet state . The parent nitrenium ion, NH + 2 , is a ground state triplet species with a gap of 30 kcal/mol (130 kJ/mol) to the lowest energy singlet state. Conversely, most arylnitrenium ions are ground state singlets. Certain substituted arylnitrenium ions can be ground state triplets, however. Nitrenium ions can have microsecond or longer lifetimes in water. [ 3 ]
Aryl nitrenium ions are of biological interest because of their involvement in certain DNA damaging processes. They are generated upon in vivo oxidation of arylamines. The regiochemistry and energetics of the reaction of phenylnitrenium ion with guanine has been investigated using density functional theory computations. [ 4 ]
Nitrenium species have been exploited as intermediates in organic reactions. [ 5 ] They are typically generated via heterolysis of N–X (X = N, O, Halogen ) bonds. For instance, they are formed upon treatment of chloramine derivatives with silver salts or by activation of aryl hydroxylamine derivatives or aryl azides with Brønsted or Lewis acids. [ 6 ] The Bamberger rearrangement is an early example of a reaction that is now thought to proceed via an aryl nitrenium intermediate. They can also act as electrophiles in electrophilic aromatic substitution . [ 7 ] | https://en.wikipedia.org/wiki/H2N+ |
Glycine (symbol Gly or G ; [ 6 ] / ˈ ɡ l aɪ s iː n / ⓘ ) [ 7 ] is an amino acid that has a single hydrogen atom as its side chain . It is the simplest stable amino acid. Glycine is one of the proteinogenic amino acids . It is encoded by all the codons starting with GG (GGU, GGC, GGA, GGG). [ 8 ] Glycine disrupts the formation of alpha-helices in secondary protein structure . Its small side chain causes it to favor random coils instead. [ 9 ] Glycine is also an inhibitory neurotransmitter [ 10 ] – interference with its release within the spinal cord (such as during a Clostridium tetani infection) can cause spastic paralysis due to uninhibited muscle contraction. [ 11 ]
It is the only achiral proteinogenic amino acid . [ 12 ] It can fit into both hydrophilic and hydrophobic environments, due to its minimal side chain of only one hydrogen atom. [ 13 ]
Glycine was discovered in 1820 by French chemist Henri Braconnot when he hydrolyzed gelatin by boiling it with sulfuric acid . [ 14 ] He originally called it "sugar of gelatin", [ 15 ] [ 16 ] but French chemist Jean-Baptiste Boussingault showed in 1838 that it contained nitrogen. [ 17 ] In 1847 American scientist Eben Norton Horsford , then a student of the German chemist Justus von Liebig , proposed the name "glycocoll"; [ 18 ] [ 19 ] however, the Swedish chemist Berzelius suggested the simpler current name a year later. [ 20 ] [ 21 ] The name comes from the Greek word γλυκύς "sweet tasting" [ 22 ] (which is also related to the prefixes glyco- and gluco- , as in glycoprotein and glucose ). In 1858, the French chemist Auguste Cahours determined that glycine was an amine of acetic acid . [ 23 ]
Although glycine can be isolated from hydrolyzed proteins , this route is not used for industrial production, as it can be manufactured more conveniently by chemical synthesis. [ 24 ] The two main processes are amination of chloroacetic acid with ammonia , giving glycine and hydrochloric acid , [ 25 ] and the Strecker amino acid synthesis , [ 26 ] which is the main synthetic method in the United States and Japan. [ 27 ] About 15 thousand tonnes are produced annually in this way. [ 28 ]
Glycine is also co-generated as an impurity in the synthesis of EDTA , arising from reactions of the ammonia co-product. [ 29 ]
Its acid–base properties are most important. In aqueous solution, glycine is amphoteric : below pH = 2.4, it converts to the ammonium cation called glycinium. Above about pH 9.6, it converts to glycinate.
Glycine functions as a bidentate ligand for many metal ions, forming amino acid complexes . [ 30 ] A typical complex is Cu(glycinate) 2 , i.e. Cu(H 2 NCH 2 CO 2 ) 2 , which exists both in cis and trans isomers. [ 31 ] [ 32 ]
With acid chlorides, glycine converts to the amidocarboxylic acid, such as hippuric acid [ 33 ] and acetylglycine . [ 34 ] With nitrous acid , one obtains glycolic acid ( van Slyke determination ). With methyl iodide , the amine becomes quaternized to give trimethylglycine , a natural product:
Glycine condenses with itself to give peptides, beginning with the formation of glycylglycine : [ 35 ]
Pyrolysis of glycine or glycylglycine gives 2,5-diketopiperazine , the cyclic diamide. [ 36 ]
Glycine forms esters with alcohols . They are often isolated as their hydrochloride , such as glycine methyl ester hydrochloride . Otherwise, the free ester tends to convert to diketopiperazine .
As a bifunctional molecule, glycine reacts with many reagents. These can be classified into N-centered and carboxylate-center reactions.
Glycine is not essential to the human diet , as it is biosynthesized in the body from the amino acid serine , which is in turn derived from 3-phosphoglycerate . In most organisms, the enzyme serine hydroxymethyltransferase catalyses this transformation via the cofactor pyridoxal phosphate : [ 37 ]
In E. coli , antibiotics that target folate depletes the supply of active tetrahydrofolates, halting glycine biosynthesis as a consequence. [ 38 ]
In the liver of vertebrates , glycine synthesis is catalyzed by glycine synthase (also called glycine cleavage enzyme). This conversion is readily reversible : [ 37 ]
In addition to being synthesized from serine, glycine can also be derived from threonine , choline or hydroxyproline via inter-organ metabolism of the liver and kidneys. [ 39 ]
Glycine is degraded via three pathways. The predominant pathway in animals and plants is the reverse of the glycine synthase pathway mentioned above. In this context, the enzyme system involved is usually called the glycine cleavage system : [ 37 ]
In the second pathway, glycine is degraded in two steps. The first step is the reverse of glycine biosynthesis from serine with serine hydroxymethyl transferase. Serine is then converted to pyruvate by serine dehydratase . [ 37 ]
In the third pathway of its degradation, glycine is converted to glyoxylate by D-amino acid oxidase . Glyoxylate is then oxidized by hepatic lactate dehydrogenase to oxalate in an NAD + -dependent reaction. [ 37 ]
The half-life of glycine and its elimination from the body varies significantly based on dose. [ 40 ] In one study, the half-life varied between 0.5 and 4.0 hours. [ 40 ]
The principal function of glycine is it acts as a precursor to proteins . Most proteins incorporate only small quantities of glycine, a notable exception being collagen , which contains about 35% glycine due to its periodically repeated role in the formation of collagen's helix structure in conjunction with hydroxyproline . [ 37 ] [ 41 ] In the genetic code , glycine is coded by all codons starting with GG, namely GGU, GGC, GGA and GGG. [ 8 ]
In higher eukaryotes , δ-aminolevulinic acid , the key precursor to porphyrins , is biosynthesized from glycine and succinyl-CoA by the enzyme ALA synthase . Glycine provides the central C 2 N subunit of all purines . [ 37 ]
Glycine is an inhibitory neurotransmitter in the central nervous system , especially in the spinal cord , brainstem , and retina . When glycine receptors are activated, chloride enters the neuron via ionotropic receptors, causing an inhibitory postsynaptic potential (IPSP). Strychnine is a strong antagonist at ionotropic glycine receptors, whereas bicuculline is a weak one. Glycine is a required co-agonist along with glutamate for NMDA receptors . In contrast to the inhibitory role of glycine in the spinal cord, this behaviour is facilitated at the ( NMDA ) glutamatergic receptors which are excitatory. [ 42 ] The LD 50 of glycine is 7930 mg/kg in rats (oral), [ 43 ] and it usually causes death by hyperexcitability. [ citation needed ]
Glycine conjugation pathway has not been fully investigated. [ 44 ] Glycine is thought to be a hepatic detoxifier of a number endogenous and xenobiotic organic acids. [ 45 ] Bile acids are normally conjugated to glycine in order to increase their solubility in water. [ 46 ]
The human body rapidly clears sodium benzoate by combining it with glycine to form hippuric acid which is then excreted. [ 47 ] The metabolic pathway for this begins with the conversion of benzoate by butyrate-CoA ligase into an intermediate product, benzoyl-CoA , [ 48 ] which is then metabolized by glycine N -acyltransferase into hippuric acid. [ 49 ]
In the US, glycine is typically sold in two grades: United States Pharmacopeia ("USP"), and technical grade. USP grade sales account for approximately 80 to 85 percent of the U.S. market for glycine. If purity greater than the USP standard is needed, for example for intravenous injections, a more expensive pharmaceutical grade glycine can be used. Technical grade glycine, which may or may not meet USP grade standards, is sold at a lower price for use in industrial applications, e.g., as an agent in metal complexing and finishing. [ 50 ]
Glycine is not widely used in foods for its nutritional value, except in infusions. Instead, glycine's role in food chemistry is as a flavorant. It is mildly sweet, and it counters the aftertaste of saccharine . It also has preservative properties, perhaps owing to its complexation to metal ions. Metal glycinate complexes, e.g. copper(II) glycinate are used as supplements for animal feeds. [ 28 ]
As of 1971 [update] , the U.S. Food and Drug Administration "no longer regards glycine and its salts as generally recognized as safe for use in human food", [ 52 ] and only permits food uses of glycine under certain conditions. [ 53 ]
Glycine has been researched for its potential to extend life . [ 54 ] [ 55 ] The proposed mechanisms of this effect are its ability to clear methionine from the body, and activating autophagy . [ 54 ]
Glycine is an intermediate in the synthesis of a variety of chemical products. It is used in the manufacture of the herbicides glyphosate , [ 56 ] iprodione , glyphosine, imiprothrin , and eglinazine. [ 28 ] It is used as an intermediate of antibiotics such as thiamphenicol . [ citation needed ]
Glycine is a significant component of some solutions used in the SDS-PAGE method of protein analysis. It serves as a buffering agent, maintaining pH and preventing sample damage during electrophoresis. [ 57 ] Glycine is also used to remove protein-labeling antibodies from Western blot membranes to enable the probing of numerous proteins of interest from SDS-PAGE gel. This allows more data to be drawn from the same specimen, increasing the reliability of the data, reducing the amount of sample processing, and number of samples required. [ 58 ] This process is known as stripping.
The presence of glycine outside the Earth was confirmed in 2009, based on the analysis of samples that had been taken in 2004 by the NASA spacecraft Stardust from comet Wild 2 and subsequently returned to Earth. Glycine had previously been identified in the Murchison meteorite in 1970. [ 59 ] The discovery of glycine in outer space bolstered the hypothesis of so-called soft-panspermia , which claims that the "building blocks" of life are widespread throughout the universe. [ 60 ] In 2016, detection of glycine within Comet 67P/Churyumov–Gerasimenko by the Rosetta spacecraft was announced. [ 61 ]
The detection of glycine outside the Solar System in the interstellar medium has been debated. [ 62 ]
Glycine is proposed to be defined by early genetic codes. [ 63 ] [ 64 ] [ 65 ] [ 66 ] For example, low complexity regions (in proteins), that may resemble the proto-peptides of the early genetic code are highly enriched in glycine. [ 66 ] | https://en.wikipedia.org/wiki/H2N-CH2-COOH |
Formamide is an amide derived from formic acid . It is a colorless liquid which is miscible with water and has an ammonia -like odor. It is chemical feedstock for the manufacture of sulfa drugs and other pharmaceuticals , herbicides and pesticides , and in the manufacture of hydrocyanic acid . It has been used as a softener for paper and fiber. It is a solvent for many ionic compounds . It has also been used as a solvent for resins and plasticizers . [ 4 ] Some astrobiologists suggest that it may be an alternative to water as the main solvent in other forms of life. [ 5 ]
Formamides are compounds of the type RR′NCHO. One important formamide is dimethylformamide , (CH 3 ) 2 NCHO.
In the past, formamide was produced by treating formic acid with ammonia , which produces ammonium formate , which in turn yields formamide upon heating: [ 6 ]
Formamide is also generated by aminolysis of ethyl formate : [ 7 ]
The current industrial process for the manufacture of formamide involves the carbonylation of ammonia: [ 4 ]
An alternative two-stage process involves the ammonolysis of methyl formate , which is formed from carbon monoxide and methanol :
Formamide is used in the industrial production of hydrogen cyanide . It is also used as a solvent for processing various polymers such as polyacrylonitrile . [ 8 ]
Formamide decomposes into carbon monoxide and ammonia when heated above 100 °C.
The reaction is slow below 160 °C, but accelerates thereafter. At very high temperatures, the reaction products shift to hydrogen cyanide (HCN) and water instead:
The same effect occurs in the presence of solid acid catalysts. [ 8 ]
Formamide is a constituent of cryoprotectant vitrification mixtures used for cryopreservation of tissues and organs .
Formamide is also used as an RNA stabiliser in gel electrophoresis by deionizing RNA. In capillary electrophoresis, it is used for stabilizing (single) strands of denatured DNA.
Another use is to add it in sol-gel solutions in order to avoid cracking during sintering .
Formamide, in its pure state, has been used as an alternative solvent for the electrostatic self-assembly of polymer nanofilms. [ 9 ]
Formamide is used to prepare primary amines directly from ketones via their N-formyl derivatives, using the Leuckart reaction .
Formamides are intermediates in the methanogenesis cycle.
Formamide has been proposed as an alternative solvent to water, perhaps with the ability to support life with alternative biochemistries to that currently found on Earth. It forms by the hydrolysis of hydrogen cyanide. With a large dipole moment, its solvation properties are similar to those of water. [ 11 ]
Formamide has been shown to convert to traces of guanine upon heating in the presence of ultraviolet light. [ 12 ]
Several prebiotic chemical reactions producing amino acid derivatives have been shown to take place in formamide. [ 13 ]
Contact with skin and eyes is not recommended. With an LD50 of grams per kg, formamide is of low acute toxicity. It also has low mutagenicity. [ 8 ]
Formamide is classified as toxic to reproductive health. [ 14 ] | https://en.wikipedia.org/wiki/H2N-CHO |
Carbamic acid , which might also be called aminoformic acid or aminocarboxylic acid , [ 2 ] is the chemical compound with the formula H 2 NCOOH . It can be obtained by the reaction of ammonia NH 3 and carbon dioxide CO 2 at very low temperatures, which also yields ammonium carbamate [NH 4 ] + [NH 2 CO 2 ] − . The compound is stable only up to about 250 K (−23 °C); at higher temperatures it decomposes into those two gases. [ 3 ] The solid apparently consists of dimers , with the two molecules connected by hydrogen bonds between the two carboxyl groups –COOH. [ 4 ]
Carbamic acid could be seen as both an amine and carboxylic acid , and therefore an amino acid ; [ 3 ] however, the attachment of the carboxyl group –COOH directly to the nitrogen atom (without any intermediate carbon chain) makes it behave very differently from the amino acids with intermediate carbon chain. ( Glycine NH 2 CH 2 COOH is generally considered to be the simplest amino acid.) The hydroxyl group –OH attached to the carbon also excludes it from the amide class.
The term "carbamic acid" is also used generically for any compounds of the form RR′NCOOH, where R and R′ are organic groups or hydrogen. [ 5 ]
Deprotonation of a carbamic acid yields a carbamate anion RR′NCOO − , the salts of which can be relatively stable. Carbamate is also a term used for esters of carbamic acids, such as methyl carbamate H 2 N−C(=O)−OCH 3 . The carbamoyl functional group RR′N–C(=O)– (often denoted by Cbm ) is the carbamic acid molecule minus the OH part of the carboxyl.
Carbamic acid is a planar molecule. [ 3 ]
The H 2 N− group of carbamic acid, unlike that of most amines, cannot be protonated to an ammonium group H 3 N + − . The zwitterionic form H 3 N + −COO − is very unstable and promptly decomposes into ammonia and carbon dioxide, [ 6 ] yet there is a report of its detection in ices irradiated with high-energy protons . [ 3 ]
Carbamic acid is formally the parent compound of several important families of organic compounds:
Many substituted carbamic acids (RHNCOOH or RR′NCOOH), can be readily synthesized by bubbling carbon dioxide through solutions of the corresponding amine ( RNH 2 or RR′NH, respectively) in an appropriate solvent, such as DMSO or supercritical carbon dioxide. [ 5 ] These carbamic acids are generally unstable at room temperature, reverting to the parent amine and carbon dioxide. [ 7 ]
Unlike carbamic acids, carbamate esters are generally stable at room temperature as a higher state. They are prepared by reaction of carbamoyl chlorides with alcohols, the addition of alcohols to isocyanates , and the reaction of carbonate esters with ammonia. [ 8 ] Methyl carbamate and ethyl carbamate are among the simplest examples and have historically been used in the textile industry, both are now suspected carcinogens. Benzyl carbamate is also known.
The enzyme class carbamate kinase , involved in several metabolic pathways of living organisms, catalyzes the formation of carbamoyl phosphate H 2 N−C(=O)−O−PO 2− 3 :
An important example of an enzyme with this activity is carbamoyl phosphate synthetase , e.g. carbamoyl phosphate synthetase I carrying out the first step of the urea cycle in order to dispose of waste ammonia.
One hemoglobin molecule can carry four molecules of carbon dioxide to the lungs as carbamate groups formed by reaction of CO 2 with four terminal amine groups of the deoxy form . The resulting compound is called carbaminohaemoglobin .
Carbamic acid is an intermediate in the industrial production of urea , which involves the reaction of carbon dioxide and ammonia. [ 9 ]
Some carbamate esters have use as muscle relaxants , including Emylcamate , Phenprobamate , Styramate and other members of ATC code M03BA . These bind to the barbiturate site of the GABA A receptor. [ 10 ]
Several carbamic acid based insecticides have been developed; for example aldicarb , carbaryl , carbofuran . [ 11 ]
An amine functional group −NH 2 can be protected from unwanted reactions by being formed as carbamate ester residue –NHC(=O)–OR. Hydrolysis of the ester bond then produces a carbamic acid –NHC(=O)OH, which then loses carbon dioxide yielding the desired amine. | https://en.wikipedia.org/wiki/H2N-COOH |
The molecular formula H 2 N 2 O 2 (molar mass: 62.03 g/mol) may refer to: | https://en.wikipedia.org/wiki/H2N2O2 |
Glycine (symbol Gly or G ; [ 6 ] / ˈ ɡ l aɪ s iː n / ⓘ ) [ 7 ] is an amino acid that has a single hydrogen atom as its side chain . It is the simplest stable amino acid. Glycine is one of the proteinogenic amino acids . It is encoded by all the codons starting with GG (GGU, GGC, GGA, GGG). [ 8 ] Glycine disrupts the formation of alpha-helices in secondary protein structure . Its small side chain causes it to favor random coils instead. [ 9 ] Glycine is also an inhibitory neurotransmitter [ 10 ] – interference with its release within the spinal cord (such as during a Clostridium tetani infection) can cause spastic paralysis due to uninhibited muscle contraction. [ 11 ]
It is the only achiral proteinogenic amino acid . [ 12 ] It can fit into both hydrophilic and hydrophobic environments, due to its minimal side chain of only one hydrogen atom. [ 13 ]
Glycine was discovered in 1820 by French chemist Henri Braconnot when he hydrolyzed gelatin by boiling it with sulfuric acid . [ 14 ] He originally called it "sugar of gelatin", [ 15 ] [ 16 ] but French chemist Jean-Baptiste Boussingault showed in 1838 that it contained nitrogen. [ 17 ] In 1847 American scientist Eben Norton Horsford , then a student of the German chemist Justus von Liebig , proposed the name "glycocoll"; [ 18 ] [ 19 ] however, the Swedish chemist Berzelius suggested the simpler current name a year later. [ 20 ] [ 21 ] The name comes from the Greek word γλυκύς "sweet tasting" [ 22 ] (which is also related to the prefixes glyco- and gluco- , as in glycoprotein and glucose ). In 1858, the French chemist Auguste Cahours determined that glycine was an amine of acetic acid . [ 23 ]
Although glycine can be isolated from hydrolyzed proteins , this route is not used for industrial production, as it can be manufactured more conveniently by chemical synthesis. [ 24 ] The two main processes are amination of chloroacetic acid with ammonia , giving glycine and hydrochloric acid , [ 25 ] and the Strecker amino acid synthesis , [ 26 ] which is the main synthetic method in the United States and Japan. [ 27 ] About 15 thousand tonnes are produced annually in this way. [ 28 ]
Glycine is also co-generated as an impurity in the synthesis of EDTA , arising from reactions of the ammonia co-product. [ 29 ]
Its acid–base properties are most important. In aqueous solution, glycine is amphoteric : below pH = 2.4, it converts to the ammonium cation called glycinium. Above about pH 9.6, it converts to glycinate.
Glycine functions as a bidentate ligand for many metal ions, forming amino acid complexes . [ 30 ] A typical complex is Cu(glycinate) 2 , i.e. Cu(H 2 NCH 2 CO 2 ) 2 , which exists both in cis and trans isomers. [ 31 ] [ 32 ]
With acid chlorides, glycine converts to the amidocarboxylic acid, such as hippuric acid [ 33 ] and acetylglycine . [ 34 ] With nitrous acid , one obtains glycolic acid ( van Slyke determination ). With methyl iodide , the amine becomes quaternized to give trimethylglycine , a natural product:
Glycine condenses with itself to give peptides, beginning with the formation of glycylglycine : [ 35 ]
Pyrolysis of glycine or glycylglycine gives 2,5-diketopiperazine , the cyclic diamide. [ 36 ]
Glycine forms esters with alcohols . They are often isolated as their hydrochloride , such as glycine methyl ester hydrochloride . Otherwise, the free ester tends to convert to diketopiperazine .
As a bifunctional molecule, glycine reacts with many reagents. These can be classified into N-centered and carboxylate-center reactions.
Glycine is not essential to the human diet , as it is biosynthesized in the body from the amino acid serine , which is in turn derived from 3-phosphoglycerate . In most organisms, the enzyme serine hydroxymethyltransferase catalyses this transformation via the cofactor pyridoxal phosphate : [ 37 ]
In E. coli , antibiotics that target folate depletes the supply of active tetrahydrofolates, halting glycine biosynthesis as a consequence. [ 38 ]
In the liver of vertebrates , glycine synthesis is catalyzed by glycine synthase (also called glycine cleavage enzyme). This conversion is readily reversible : [ 37 ]
In addition to being synthesized from serine, glycine can also be derived from threonine , choline or hydroxyproline via inter-organ metabolism of the liver and kidneys. [ 39 ]
Glycine is degraded via three pathways. The predominant pathway in animals and plants is the reverse of the glycine synthase pathway mentioned above. In this context, the enzyme system involved is usually called the glycine cleavage system : [ 37 ]
In the second pathway, glycine is degraded in two steps. The first step is the reverse of glycine biosynthesis from serine with serine hydroxymethyl transferase. Serine is then converted to pyruvate by serine dehydratase . [ 37 ]
In the third pathway of its degradation, glycine is converted to glyoxylate by D-amino acid oxidase . Glyoxylate is then oxidized by hepatic lactate dehydrogenase to oxalate in an NAD + -dependent reaction. [ 37 ]
The half-life of glycine and its elimination from the body varies significantly based on dose. [ 40 ] In one study, the half-life varied between 0.5 and 4.0 hours. [ 40 ]
The principal function of glycine is it acts as a precursor to proteins . Most proteins incorporate only small quantities of glycine, a notable exception being collagen , which contains about 35% glycine due to its periodically repeated role in the formation of collagen's helix structure in conjunction with hydroxyproline . [ 37 ] [ 41 ] In the genetic code , glycine is coded by all codons starting with GG, namely GGU, GGC, GGA and GGG. [ 8 ]
In higher eukaryotes , δ-aminolevulinic acid , the key precursor to porphyrins , is biosynthesized from glycine and succinyl-CoA by the enzyme ALA synthase . Glycine provides the central C 2 N subunit of all purines . [ 37 ]
Glycine is an inhibitory neurotransmitter in the central nervous system , especially in the spinal cord , brainstem , and retina . When glycine receptors are activated, chloride enters the neuron via ionotropic receptors, causing an inhibitory postsynaptic potential (IPSP). Strychnine is a strong antagonist at ionotropic glycine receptors, whereas bicuculline is a weak one. Glycine is a required co-agonist along with glutamate for NMDA receptors . In contrast to the inhibitory role of glycine in the spinal cord, this behaviour is facilitated at the ( NMDA ) glutamatergic receptors which are excitatory. [ 42 ] The LD 50 of glycine is 7930 mg/kg in rats (oral), [ 43 ] and it usually causes death by hyperexcitability. [ citation needed ]
Glycine conjugation pathway has not been fully investigated. [ 44 ] Glycine is thought to be a hepatic detoxifier of a number endogenous and xenobiotic organic acids. [ 45 ] Bile acids are normally conjugated to glycine in order to increase their solubility in water. [ 46 ]
The human body rapidly clears sodium benzoate by combining it with glycine to form hippuric acid which is then excreted. [ 47 ] The metabolic pathway for this begins with the conversion of benzoate by butyrate-CoA ligase into an intermediate product, benzoyl-CoA , [ 48 ] which is then metabolized by glycine N -acyltransferase into hippuric acid. [ 49 ]
In the US, glycine is typically sold in two grades: United States Pharmacopeia ("USP"), and technical grade. USP grade sales account for approximately 80 to 85 percent of the U.S. market for glycine. If purity greater than the USP standard is needed, for example for intravenous injections, a more expensive pharmaceutical grade glycine can be used. Technical grade glycine, which may or may not meet USP grade standards, is sold at a lower price for use in industrial applications, e.g., as an agent in metal complexing and finishing. [ 50 ]
Glycine is not widely used in foods for its nutritional value, except in infusions. Instead, glycine's role in food chemistry is as a flavorant. It is mildly sweet, and it counters the aftertaste of saccharine . It also has preservative properties, perhaps owing to its complexation to metal ions. Metal glycinate complexes, e.g. copper(II) glycinate are used as supplements for animal feeds. [ 28 ]
As of 1971 [update] , the U.S. Food and Drug Administration "no longer regards glycine and its salts as generally recognized as safe for use in human food", [ 52 ] and only permits food uses of glycine under certain conditions. [ 53 ]
Glycine has been researched for its potential to extend life . [ 54 ] [ 55 ] The proposed mechanisms of this effect are its ability to clear methionine from the body, and activating autophagy . [ 54 ]
Glycine is an intermediate in the synthesis of a variety of chemical products. It is used in the manufacture of the herbicides glyphosate , [ 56 ] iprodione , glyphosine, imiprothrin , and eglinazine. [ 28 ] It is used as an intermediate of antibiotics such as thiamphenicol . [ citation needed ]
Glycine is a significant component of some solutions used in the SDS-PAGE method of protein analysis. It serves as a buffering agent, maintaining pH and preventing sample damage during electrophoresis. [ 57 ] Glycine is also used to remove protein-labeling antibodies from Western blot membranes to enable the probing of numerous proteins of interest from SDS-PAGE gel. This allows more data to be drawn from the same specimen, increasing the reliability of the data, reducing the amount of sample processing, and number of samples required. [ 58 ] This process is known as stripping.
The presence of glycine outside the Earth was confirmed in 2009, based on the analysis of samples that had been taken in 2004 by the NASA spacecraft Stardust from comet Wild 2 and subsequently returned to Earth. Glycine had previously been identified in the Murchison meteorite in 1970. [ 59 ] The discovery of glycine in outer space bolstered the hypothesis of so-called soft-panspermia , which claims that the "building blocks" of life are widespread throughout the universe. [ 60 ] In 2016, detection of glycine within Comet 67P/Churyumov–Gerasimenko by the Rosetta spacecraft was announced. [ 61 ]
The detection of glycine outside the Solar System in the interstellar medium has been debated. [ 62 ]
Glycine is proposed to be defined by early genetic codes. [ 63 ] [ 64 ] [ 65 ] [ 66 ] For example, low complexity regions (in proteins), that may resemble the proto-peptides of the early genetic code are highly enriched in glycine. [ 66 ] | https://en.wikipedia.org/wiki/H2NCH2COOH |
50 g/L ethanol ~4 g/L acetonitrile [ 4 ]
Urea , also called carbamide (because it is a diamide of carbonic acid ), is an organic compound with chemical formula CO(NH 2 ) 2 . This amide has two amino groups (– NH 2 ) joined by a carbonyl functional group (–C(=O)–). It is thus the simplest amide of carbamic acid . [ 6 ]
Urea serves an important role in the cellular metabolism of nitrogen -containing compounds by animals and is the main nitrogen-containing substance in the urine of mammals . Urea is Neo-Latin , from French urée , from Ancient Greek οὖρον ( oûron ) ' urine ' , itself from Proto-Indo-European *h₂worsom .
It is a colorless, odorless solid, highly soluble in water, and practically non-toxic ( LD 50 is 15 g/kg for rats). [ 7 ] Dissolved in water, it is neither acidic nor alkaline . The body uses it in many processes, most notably nitrogen excretion . The liver forms it by combining two ammonia molecules ( NH 3 ) with a carbon dioxide ( CO 2 ) molecule in the urea cycle . Urea is widely used in fertilizers as a source of nitrogen (N) and is an important raw material for the chemical industry .
In 1828, Friedrich Wöhler discovered that urea can be produced from inorganic starting materials, which was an important conceptual milestone in chemistry. This showed for the first time that a substance previously known only as a byproduct of life could be synthesized in the laboratory without biological starting materials, thereby contradicting the widely held doctrine of vitalism , which stated that only living organisms could produce the chemicals of life.
The structure of the molecule of urea is O=C(−NH 2 ) 2 . The urea molecule is planar when in a solid crystal because of sp 2 hybridization of the N orbitals. [ 8 ] [ 9 ] It is non-planar with C 2 symmetry when in the gas phase [ 10 ] or in aqueous solution, [ 9 ] with C–N–H and H–N–H bond angles that are intermediate between the trigonal planar angle of 120° and the tetrahedral angle of 109.5°. In solid urea, the oxygen center is engaged in two N–H–O hydrogen bonds . The resulting hydrogen-bond network is probably established at the cost of efficient molecular packing: The structure is quite open, the ribbons forming tunnels with square cross-section. The carbon in urea is described as sp 2 hybridized, the C-N bonds have significant double bond character, and the carbonyl oxygen is relatively basic. Urea's high aqueous solubility reflects its ability to engage in extensive hydrogen bonding with water.
By virtue of its tendency to form porous frameworks, urea has the ability to trap many organic compounds. In these so-called clathrates , the organic "guest" molecules are held in channels formed by interpenetrating helices composed of hydrogen-bonded urea molecules. In this way, urea-clathrates have been well investigated for separations. [ 11 ]
Urea is a weak base, with a p K b of 13.9. [ 5 ] When combined with strong acids, it undergoes protonation at oxygen to form uronium salts. [ 13 ] [ 14 ] It is also a Lewis base , forming metal complexes of the type [M(urea) 6 ] n + . [ 15 ]
Urea reacts with malonic esters to make barbituric acids .
Molten urea decomposes into ammonium cyanate at about 152 °C, and into ammonia and isocyanic acid above 160 °C: [ 16 ]
Heating above 160 °C yields biuret NH 2 CONHCONH 2 and triuret NH 2 CONHCONHCONH 2 via reaction with isocyanic acid: [ 17 ] [ 16 ]
At higher temperatures it converts to a range of condensation products , including cyanuric acid (CNOH) 3 , guanidine HNC(NH 2 ) 2 , and melamine . [ 17 ] [ 16 ]
In aqueous solution, urea slowly equilibrates with ammonium cyanate. This elimination reaction [ 18 ] cogenerates isocyanic acid , which can carbamylate proteins, in particular the N-terminal amino group, the side chain amino of lysine , and to a lesser extent the side chains of arginine and cysteine . [ 19 ] [ 20 ] Each carbamylation event adds 43 daltons to the mass of the protein, which can be observed in protein mass spectrometry . [ 20 ] For this reason, pure urea solutions should be freshly prepared and used, as aged solutions may develop a significant concentration of cyanate (20 mM in 8 M urea). [ 20 ] Dissolving urea in ultrapure water followed by removing ions (i.e. cyanate) with a mixed-bed ion-exchange resin and storing that solution at 4 °C is a recommended preparation procedure. [ 21 ] However, cyanate will build back up to significant levels within a few days. [ 20 ] Alternatively, adding 25–50 mM ammonium chloride to a concentrated urea solution decreases formation of cyanate because of the common ion effect . [ 20 ] [ 22 ]
Urea is readily quantified by a number of different methods, such as the diacetyl monoxime colorimetric method, and the Berthelot reaction (after initial conversion of urea to ammonia via urease). These methods are amenable to high throughput instrumentation, such as automated flow injection analyzers [ 23 ] and 96-well micro-plate spectrophotometers. [ 24 ]
Ureas describes a class of chemical compounds that share the same functional group, a carbonyl group attached to two organic amine residues: R 1 R 2 N−C(=O)−NR 3 R 4 , where R 1 , R 2 , R 3 and R 4 groups are hydrogen (–H), organyl or other groups. Examples include carbamide peroxide , allantoin , and hydantoin . Ureas are closely related to biurets and related in structure to amides , carbamates , carbodiimides , and thiocarbamides .
More than 90% of world industrial production of urea is destined for use as a nitrogen-release fertilizer . [ 17 ] Urea has the highest nitrogen content of all solid nitrogenous fertilizers in common use. Therefore, it has a low transportation cost per unit of nitrogen nutrient . The most common impurity of synthetic urea is biuret , which impairs plant growth. Urea breaks down in the soil to give ammonium ions ( NH + 4 ). The ammonium is taken up by the plant through its roots. In some soils, the ammonium is oxidized by bacteria to give nitrate ( NO − 3 ), which is also a nitrogen-rich plant nutrient. The loss of nitrogenous compounds to the atmosphere and runoff is wasteful and environmentally damaging so urea is sometimes modified to enhance the efficiency of its agricultural use. Techniques to make controlled-release fertilizers that slow the release of nitrogen include the encapsulation of urea in an inert sealant, and conversion of urea into derivatives such as urea-formaldehyde compounds, which degrade into ammonia at a pace matching plants' nutritional requirements.
Urea is a raw material for the manufacture of formaldehyde based resins , such as UF, MUF, and MUPF, used mainly in wood-based panels, for instance, particleboard , fiberboard , OSB, and plywood . [ 25 ]
Urea can be used in a reaction with nitric acid to make urea nitrate , a high explosive that is used industrially and as part of some improvised explosive devices .
Urea is used in Selective Non-Catalytic Reduction (SNCR) and Selective Catalytic Reduction (SCR) reactions to reduce the NO x pollutants in exhaust gases from combustion from diesel , dual fuel, and lean-burn natural gas engines. The BlueTec system, for example, injects a water-based urea solution into the exhaust system. Ammonia ( NH 3 ) produced by the hydrolysis of urea reacts with nitrogen oxides ( NO x ) and is converted into nitrogen gas ( N 2 ) and water within the catalytic converter. The conversion of noxious NO x to innocuous N 2 is described by the following simplified global equation: [ 26 ]
When urea is used, a pre-reaction (hydrolysis) occurs to first convert it to ammonia:
Being a solid highly soluble in water (545 g/L at 25 °C), [ 2 ] urea is much easier and safer to handle and store than the more irritant , caustic and hazardous ammonia ( NH 3 ), so it is the reactant of choice. Trucks and cars using these catalytic converters need to carry a supply of diesel exhaust fluid , also sold as AdBlue , a solution of urea in water.
Urea in concentrations up to 10 M is a powerful protein denaturant as it disrupts the noncovalent bonds in the proteins. This property can be exploited to increase the solubility of some proteins. A mixture of urea and choline chloride is used as a deep eutectic solvent (DES), a substance similar to ionic liquid . When used in a deep eutectic solvent, urea gradually denatures the proteins that are solubilized. [ 27 ]
Urea in concentrations up to 8 M can be used to make fixed brain tissue transparent to visible light while still preserving fluorescent signals from labeled cells. This allows for much deeper imaging of neuronal processes than previously obtainable using conventional one photon or two photon confocal microscopes. [ 28 ]
Urea-containing creams are used as topical dermatological products to promote rehydration of the skin . Urea 40% is indicated for psoriasis , xerosis , onychomycosis , ichthyosis , eczema , keratosis , keratoderma , corns, and calluses . If covered by an occlusive dressing , 40% urea preparations may also be used for nonsurgical debridement of nails . Urea 40% "dissolves the intercellular matrix" [ 29 ] [ 30 ] of the nail plate. Only diseased or dystrophic nails are removed, as there is no effect on healthy portions of the nail. [ 31 ] This drug (as carbamide peroxide ) is also used as an earwax removal aid. [ 32 ]
Urea has also been studied as a diuretic . It was first used by Dr. W. Friedrich in 1892. [ 33 ] In a 2010 study of ICU patients, urea was used to treat euvolemic hyponatremia and was found safe, inexpensive, and simple. [ 34 ]
Like saline , urea has been injected into the uterus to induce abortion , although this method is no longer in widespread use. [ 35 ]
The blood urea nitrogen (BUN) test is a measure of the amount of nitrogen in the blood that comes from urea. It is used as a marker of renal function , though it is inferior to other markers such as creatinine because blood urea levels are influenced by other factors such as diet, dehydration, [ 36 ] and liver function.
Urea has also been studied as an excipient in drug-coated balloon (DCB) coating formulations to enhance local drug delivery to stenotic blood vessels. [ 37 ] [ 38 ] Urea, when used as an excipient in small doses (~3 μg/mm 2 ) to coat DCB surface was found to form crystals that increase drug transfer without adverse toxic effects on vascular endothelial cells . [ 39 ]
Urea labeled with carbon-14 or carbon-13 is used in the urea breath test , which is used to detect the presence of the bacterium Helicobacter pylori ( H. pylori ) in the stomach and duodenum of humans, associated with peptic ulcers . The test detects the characteristic enzyme urease , produced by H. pylori , by a reaction that produces ammonia from urea. This increases the pH (reduces the acidity) of the stomach environment around the bacteria. Similar bacteria species to H. pylori can be identified by the same test in animals such as apes , dogs , and cats (including big cats ).
Amino acids from ingested food (or produced from catabolism of muscle protein) that are used for the synthesis of proteins and other biological substances can be oxidized by the body as an alternative source of energy, yielding urea and carbon dioxide . [ 47 ] The oxidation pathway starts with the removal of the amino group by a transaminase ; the amino group is then fed into the urea cycle . The first step in the conversion of amino acids into metabolic waste in the liver is removal of the alpha-amino nitrogen, which produces ammonia . Because ammonia is toxic, it is excreted immediately by fish, converted into uric acid by birds, and converted into urea by mammals. [ 48 ]
Ammonia ( NH 3 ) is a common byproduct of the metabolism of nitrogenous compounds. Ammonia is smaller, more volatile, and more mobile than urea. If allowed to accumulate, ammonia would raise the pH in cells to toxic levels. Therefore, many organisms convert ammonia to urea, even though this synthesis has a net energy cost. Being practically neutral and highly soluble in water, urea is a safe vehicle for the body to transport and excrete excess nitrogen.
Urea is synthesized in the body of many organisms as part of the urea cycle , either from the oxidation of amino acids or from ammonia . In this cycle, amino groups donated by ammonia and L - aspartate are converted to urea, while L - ornithine , citrulline , L - argininosuccinate , and L - arginine act as intermediates. Urea production occurs in the liver and is regulated by N -acetylglutamate . Urea is then dissolved into the blood (in the reference range of 2.5 to 6.7 mmol/L) and further transported and excreted by the kidney as a component of urine . In addition, a small amount of urea is excreted (along with sodium chloride and water) in sweat .
In water, the amine groups undergo slow displacement by water molecules, producing ammonia, ammonium ions , and bicarbonate ions . For this reason, old, stale urine has a stronger odor than fresh urine.
The cycling of and excretion of urea by the kidneys is a vital part of mammalian metabolism. Besides its role as carrier of waste nitrogen, urea also plays a role in the countercurrent exchange system of the nephrons , that allows for reabsorption of water and critical ions from the excreted urine . Urea is reabsorbed in the inner medullary collecting ducts of the nephrons, [ 49 ] thus raising the osmolarity in the medullary interstitium surrounding the thin descending limb of the loop of Henle , which makes the water reabsorb.
By action of the urea transporter 2 , some of this reabsorbed urea eventually flows back into the thin descending limb of the tubule, [ 50 ] through the collecting ducts, and into the excreted urine. The body uses this mechanism, which is controlled by the antidiuretic hormone , to create hyperosmotic urine — i.e., urine with a higher concentration of dissolved substances than the blood plasma . This mechanism is important to prevent the loss of water, maintain blood pressure , and maintain a suitable concentration of sodium ions in the blood plasma.
The equivalent nitrogen content (in grams ) of urea (in mmol ) can be estimated by the conversion factor 0.028 g/mmol. [ 51 ] Furthermore, 1 gram of nitrogen is roughly equivalent to 6.25 grams of protein , and 1 gram of protein is roughly equivalent to 5 grams of muscle tissue. In situations such as muscle wasting , 1 mmol of excessive urea in the urine (as measured by urine volume in litres multiplied by urea concentration in mmol/L) roughly corresponds to a muscle loss of 0.67 gram.
In aquatic organisms the most common form of nitrogen waste is ammonia, whereas land-dwelling organisms convert the toxic ammonia to either urea or uric acid . Urea is found in the urine of mammals and amphibians , as well as some fish. Birds and saurian reptiles have a different form of nitrogen metabolism that requires less water, and leads to nitrogen excretion in the form of uric acid. Tadpoles excrete ammonia, but shift to urea production during metamorphosis . Despite the generalization above, the urea pathway has been documented not only in mammals and amphibians, but in many other organisms as well, including birds, invertebrates , insects, plants, yeast , fungi , and even microorganisms . [ 52 ]
Urea can be irritating to skin, eyes, and the respiratory tract. Repeated or prolonged contact with urea in fertilizer form on the skin may cause dermatitis . [ 53 ]
High concentrations in the blood can be damaging. Ingestion of low concentrations of urea, such as are found in typical human urine , are not dangerous with additional water ingestion within a reasonable time-frame. Many animals (e.g. camels , rodents or dogs) have a much more concentrated urine which may contain a higher urea amount than normal human urine.
Urea can cause algal blooms to produce toxins, and its presence in the runoff from fertilized land may play a role in the increase of toxic blooms. [ 54 ]
The substance decomposes on heating above melting point, producing toxic gases, and reacts violently with strong oxidants, nitrites, inorganic chlorides, chlorites and perchlorates, causing fire and explosion. [ 55 ]
Urea was first discovered in urine in 1727 by the Dutch scientist Herman Boerhaave , [ 56 ] although this discovery is often attributed to the French chemist Hilaire Rouelle as well as William Cruickshank . [ 57 ]
Boerhaave used the following steps to isolate urea: [ 58 ] [ 59 ]
In 1828, the German chemist Friedrich Wöhler obtained urea artificially by treating silver cyanate with ammonium chloride . [ 60 ] [ 61 ] [ 62 ]
This was the first time an organic compound was artificially synthesized from inorganic starting materials, without the involvement of living organisms. The results of this experiment implicitly discredited vitalism , the theory that the chemicals of living organisms are fundamentally different from those of inanimate matter. This insight was important for the development of organic chemistry . His discovery prompted Wöhler to write triumphantly to Jöns Jakob Berzelius :
I must tell you that I can make urea without the use of kidneys, either man or dog. Ammonium cyanate is urea.
In fact, his second sentence was incorrect. Ammonium cyanate [NH 4 ] + [OCN] − and urea CO(NH 2 ) 2 are two different chemicals with the same empirical formula CON 2 H 4 , which are in chemical equilibrium heavily favoring urea under standard conditions . [ 63 ] Regardless, with his discovery, Wöhler secured a place among the pioneers of organic chemistry.
Uremic frost was first described in 1865 by Harald Hirschsprung , the first Danish pediatrician in 1870 who also described the disease that carries his name in 1886. Uremic frost has become rare since the advent of dialysis . It is the classical pre-dialysis era description of crystallized urea deposits over the skin of patients with prolonged kidney failure and severe uremia. [ 64 ]
Urea was first noticed by Herman Boerhaave in the early 18th century from evaporates of urine. In 1773, Hilaire Rouelle obtained crystals containing urea from human urine by evaporating it and treating it with alcohol in successive filtrations. [ 65 ] This method was aided by Carl Wilhelm Scheele 's discovery that urine treated by concentrated nitric acid precipitated crystals. Antoine François, comte de Fourcroy and Louis Nicolas Vauquelin discovered in 1799 that the nitrated crystals were identical to Rouelle's substance and invented the term "urea." [ 66 ] [ 67 ] Berzelius made further improvements to its purification [ 68 ] and finally William Prout , in 1817, succeeded in obtaining and determining the chemical composition of the pure substance. [ 69 ] In the evolved procedure, urea was precipitated as urea nitrate by adding strong nitric acid to urine. To purify the resulting crystals, they were dissolved in boiling water with charcoal and filtered. After cooling, pure crystals of urea nitrate form. To reconstitute the urea from the nitrate, the crystals are dissolved in warm water, and barium carbonate added. The water is then evaporated and anhydrous alcohol added to extract the urea. This solution is drained off and evaporated, leaving pure urea.
Ureas in the more general sense can be accessed in the laboratory by reaction of phosgene with primary or secondary amines :
These reactions proceed through an isocyanate intermediate. Non-symmetric ureas can be accessed by the reaction of primary or secondary amines with an isocyanate.
Urea can also be produced by heating ammonium cyanate to 60 °C.
In 2020, worldwide production capacity was approximately 180 million tonnes. [ 70 ]
For use in industry, urea is produced from synthetic ammonia and carbon dioxide . As large quantities of carbon dioxide are produced during the ammonia manufacturing process as a byproduct of burning hydrocarbons to generate heat (predominantly natural gas, and less often petroleum derivatives or coal), urea production plants are almost always located adjacent to the site where the ammonia is manufactured.
The basic process, patented in 1922, is called the Bosch–Meiser urea process after its discoverers Carl Bosch and Wilhelm Meiser. [ 71 ] The process consists of two main equilibrium reactions , with incomplete conversion of the reactants. The first is carbamate formation : the fast exothermic reaction of liquid ammonia with gaseous carbon dioxide ( CO 2 ) at high temperature and pressure to form ammonium carbamate ( [NH 4 ] + [NH 2 COO] − ): [ 17 ]
The second is urea conversion : the slower endothermic decomposition of ammonium carbamate into urea and water:
The overall conversion of NH 3 and CO 2 to urea is exothermic, with the reaction heat from the first reaction driving the second. The conditions that favor urea formation (high temperature) have an unfavorable effect on the carbamate formation equilibrium. The process conditions are a compromise: the ill-effect on the first reaction of the high temperature (around 190 °C) needed for the second is compensated for by conducting the process under high pressure (140–175 bar), which favors the first reaction. Although it is necessary to compress gaseous carbon dioxide to this pressure, the ammonia is available from the ammonia production plant in liquid form, which can be pumped into the system much more economically. To allow the slow urea formation reaction time to reach equilibrium, a large reaction space is needed, so the synthesis reactor in a large urea plant tends to be a massive pressure vessel.
Because the urea conversion is incomplete, the urea must be separated from the unconverted reactants, including the ammonium carbamate. Various commercial urea processes are characterized by the conditions under which urea forms and the way that unconverted reactants are further processed.
In early "straight-through" urea plants, reactant recovery (the first step in "recycling") was done by letting down the system pressure to atmospheric to let the carbamate decompose back to ammonia and carbon dioxide. Originally, because it was not economic to recompress the ammonia and carbon dioxide for recycle, the ammonia at least would be used for the manufacture of other products such as ammonium nitrate or ammonium sulfate , and the carbon dioxide was usually wasted. Later process schemes made recycling unused ammonia and carbon dioxide practical. This was accomplished by the "total recycle process", developed in the 1940s to 1960s and now called the "conventional recycle process". It proceeds by depressurizing the reaction solution in stages (first to 18–25 bar and then to 2–5 bar) and passing it at each stage through a steam-heated carbamate decomposer , then recombining the resulting carbon dioxide and ammonia in a falling-film carbamate condenser and pumping the carbamate solution back into the urea reaction vessel. [ 17 ]
The "conventional recycle process" for recovering and reusing the reactants has largely been supplanted by a stripping process, developed in the early 1960s by Stamicarbon in The Netherlands, that operates at or near the full pressure of the reaction vessel. It reduces the complexity of the multi-stage recycle scheme, and it reduces the amount of water recycled in the carbamate solution, which has an adverse effect on the equilibrium in the urea conversion reaction and thus on overall plant efficiency. Effectively all new urea plants use the stripper, and many total recycle urea plants have converted to a stripping process. [ 17 ] [ 73 ]
In the conventional recycle processes, carbamate decomposition is promoted by reducing the overall pressure, which reduces the partial pressure of both ammonia and carbon dioxide, allowing these gasses to be separated from the urea product solution. The stripping process achieves a similar effect without lowering the overall pressure, by suppressing the partial pressure of just one of the reactants in order to promote carbamate decomposition. Instead of feeding carbon dioxide gas directly to the urea synthesis reactor with the ammonia, as in the conventional process, the stripping process first routes the carbon dioxide through the stripper. The stripper is a carbamate decomposer that provides a large amount of gas-liquid contact. This flushes out free ammonia, reducing its partial pressure over the liquid surface and carrying it directly to a carbamate condenser (also under full system pressure). From there, reconstituted ammonium carbamate liquor is passed to the urea production reactor. That eliminates the medium-pressure stage of the conventional recycle process. [ 17 ] [ 73 ]
The three main side reactions that produce impurities have in common that they decompose urea.
Urea hydrolyzes back to ammonium carbamate in the hottest stages of the synthesis plant, especially in the stripper, so residence times in these stages are designed to be short. [ 17 ]
Biuret is formed when two molecules of urea combine with the loss of a molecule of ammonia.
Normally this reaction is suppressed in the synthesis reactor by maintaining an excess of ammonia, but after the stripper, it occurs until the temperature is reduced. [ 17 ] Biuret is undesirable in urea fertilizer because it is toxic to crop plants to varying degrees, [ 74 ] but it is sometimes desirable as a nitrogen source when used in animal feed. [ 75 ]
Isocyanic acid HNCO and ammonia NH 3 results from the thermal decomposition of ammonium cyanate [NH 4 ] + [OCN] − , which is in chemical equilibrium with urea:
This decomposition is at its worst when the urea solution is heated at low pressure, which happens when the solution is concentrated for prilling or granulation (see below). The reaction products mostly volatilize into the overhead vapours, and recombine when these condense to form urea again, which contaminates the process condensate. [ 17 ]
Ammonium carbamate solutions are highly corrosive to metallic construction materials – even to resistant forms of stainless steel – especially in the hottest parts of the synthesis plant such as the stripper. Historically corrosion has been minimized (although not eliminated) by continuous injection of a small amount of oxygen (as air) into the plant to establish and maintain a passive oxide layer on exposed stainless steel surfaces. Highly corrosion resistant materials have been introduced to reduce the need for passivation oxygen, such as specialized duplex stainless steels in the 1990s, and zirconium or zirconium-clad titanium tubing in the 2000s. [ 17 ]
Urea can be produced in solid forms ( prills , granules , pellets or crystals) or as solutions.
For its main use as a fertilizer urea is mostly marketed in solid form, either as prills or granules. Prills are solidified droplets, whose production predates satisfactory urea granulation processes. Prills can be produced more cheaply than granules, but the limited size of prills (up to about 2.1 mm in diameter), their low crushing strength, and the caking or crushing of prills during bulk storage and handling make them inferior to granules. Granules are produced by acretion onto urea seed particles by spraying liquid urea in a succession of layers. Formaldehyde is added during the production of both prills and granules in order to increase crushing strength and suppress caking. Other shaping techniques such as pastillization (depositing uniform-sized liquid droplets onto a cooling conveyor belt) are also used. [ 17 ]
Solutions of urea and ammonium nitrate in water (UAN) are commonly used as a liquid fertilizer. In admixture, the combined solubility of ammonium nitrate and urea is so much higher than that of either component alone that it gives a stable solution with a total nitrogen content (32%) approaching that of solid ammonium nitrate (33.5%), though not, of course, that of urea itself (46%). UAN allows use of ammonium nitrate without the explosion hazard. [ 17 ] UAN accounts for 80% of the liquid fertilizers in the US. [ 76 ] | https://en.wikipedia.org/wiki/H2NCONH2 |
Carbamic acid , which might also be called aminoformic acid or aminocarboxylic acid , [ 2 ] is the chemical compound with the formula H 2 NCOOH . It can be obtained by the reaction of ammonia NH 3 and carbon dioxide CO 2 at very low temperatures, which also yields ammonium carbamate [NH 4 ] + [NH 2 CO 2 ] − . The compound is stable only up to about 250 K (−23 °C); at higher temperatures it decomposes into those two gases. [ 3 ] The solid apparently consists of dimers , with the two molecules connected by hydrogen bonds between the two carboxyl groups –COOH. [ 4 ]
Carbamic acid could be seen as both an amine and carboxylic acid , and therefore an amino acid ; [ 3 ] however, the attachment of the carboxyl group –COOH directly to the nitrogen atom (without any intermediate carbon chain) makes it behave very differently from the amino acids with intermediate carbon chain. ( Glycine NH 2 CH 2 COOH is generally considered to be the simplest amino acid.) The hydroxyl group –OH attached to the carbon also excludes it from the amide class.
The term "carbamic acid" is also used generically for any compounds of the form RR′NCOOH, where R and R′ are organic groups or hydrogen. [ 5 ]
Deprotonation of a carbamic acid yields a carbamate anion RR′NCOO − , the salts of which can be relatively stable. Carbamate is also a term used for esters of carbamic acids, such as methyl carbamate H 2 N−C(=O)−OCH 3 . The carbamoyl functional group RR′N–C(=O)– (often denoted by Cbm ) is the carbamic acid molecule minus the OH part of the carboxyl.
Carbamic acid is a planar molecule. [ 3 ]
The H 2 N− group of carbamic acid, unlike that of most amines, cannot be protonated to an ammonium group H 3 N + − . The zwitterionic form H 3 N + −COO − is very unstable and promptly decomposes into ammonia and carbon dioxide, [ 6 ] yet there is a report of its detection in ices irradiated with high-energy protons . [ 3 ]
Carbamic acid is formally the parent compound of several important families of organic compounds:
Many substituted carbamic acids (RHNCOOH or RR′NCOOH), can be readily synthesized by bubbling carbon dioxide through solutions of the corresponding amine ( RNH 2 or RR′NH, respectively) in an appropriate solvent, such as DMSO or supercritical carbon dioxide. [ 5 ] These carbamic acids are generally unstable at room temperature, reverting to the parent amine and carbon dioxide. [ 7 ]
Unlike carbamic acids, carbamate esters are generally stable at room temperature as a higher state. They are prepared by reaction of carbamoyl chlorides with alcohols, the addition of alcohols to isocyanates , and the reaction of carbonate esters with ammonia. [ 8 ] Methyl carbamate and ethyl carbamate are among the simplest examples and have historically been used in the textile industry, both are now suspected carcinogens. Benzyl carbamate is also known.
The enzyme class carbamate kinase , involved in several metabolic pathways of living organisms, catalyzes the formation of carbamoyl phosphate H 2 N−C(=O)−O−PO 2− 3 :
An important example of an enzyme with this activity is carbamoyl phosphate synthetase , e.g. carbamoyl phosphate synthetase I carrying out the first step of the urea cycle in order to dispose of waste ammonia.
One hemoglobin molecule can carry four molecules of carbon dioxide to the lungs as carbamate groups formed by reaction of CO 2 with four terminal amine groups of the deoxy form . The resulting compound is called carbaminohaemoglobin .
Carbamic acid is an intermediate in the industrial production of urea , which involves the reaction of carbon dioxide and ammonia. [ 9 ]
Some carbamate esters have use as muscle relaxants , including Emylcamate , Phenprobamate , Styramate and other members of ATC code M03BA . These bind to the barbiturate site of the GABA A receptor. [ 10 ]
Several carbamic acid based insecticides have been developed; for example aldicarb , carbaryl , carbofuran . [ 11 ]
An amine functional group −NH 2 can be protected from unwanted reactions by being formed as carbamate ester residue –NHC(=O)–OR. Hydrolysis of the ester bond then produces a carbamic acid –NHC(=O)OH, which then loses carbon dioxide yielding the desired amine. | https://en.wikipedia.org/wiki/H2NCOOH |
Nickel(II) hydroxide is the inorganic compound with the formula Ni(OH) 2 . It is a lime-green solid that dissolves with decomposition in ammonia and amines and is attacked by acids. It is electroactive, being converted to the Ni(III) oxy-hydroxide , leading to widespread applications in rechargeable batteries . [ 6 ]
Nickel(II) hydroxide has two well-characterized polymorphs , α and β. The α structure consists of Ni(OH) 2 layers with intercalated anions or water. [ 7 ] [ 8 ] The β form adopts a hexagonal close-packed structure of Ni 2+ and OH − ions. [ 7 ] [ 8 ] In the presence of water, the α polymorph typically recrystallizes to the β form. [ 7 ] [ 9 ] In addition to the α and β polymorphs, several γ nickel hydroxides have been described, distinguished by crystal structures with much larger inter-sheet distances. [ 7 ]
The mineral form of Ni(OH) 2 , theophrastite , was first identified in the Vermion region of northern Greece, in 1980. It is found naturally as a translucent emerald-green crystal formed in thin sheets near the boundaries of idocrase or chlorite crystals. [ 10 ] A nickel-magnesium variant of the mineral, (Ni,Mg)(OH) 2 had been previously discovered at Hagdale on the island of Unst in Scotland. [ 11 ]
Nickel(II) hydroxide is frequently used in electrical car batteries. [ 8 ] Specifically, Ni(OH) 2 readily oxidizes to nickel oxyhydroxide, NiOOH, in combination with a reduction reaction, often of a metal hydride (reaction 1 and 2). [ 12 ] [ 13 ]
Reaction 1 Ni(OH) 2 + OH − → NiO(OH) + H 2 O + e −
Reaction 2 M + H 2 O + e − → MH + OH −
Net Reaction (in H 2 O) Ni(OH) 2 + M → NiOOH + MH
Of the two polymorphs, α-Ni(OH) 2 has a higher theoretical capacity and thus is generally considered to be preferable in electrochemical applications. However, it transforms to β-Ni(OH) 2 in alkaline solutions, leading to many investigations into the possibility of stabilized α-Ni(OH) 2 electrodes for industrial applications. [ 9 ]
The synthesis entails treating aqueous solutions of nickel(II) salts with potassium hydroxide. When the same reaction is conducted in the presence of bromine, the product is Ni 3 O 2 (OH) 4 . [ 14 ]
The Ni 2+ ion is a carcinogen when inhaled. | https://en.wikipedia.org/wiki/H2NiO2 |
Hydrogen peroxide is a chemical compound with the formula H 2 O 2 . In its pure form, it is a very pale blue [ 5 ] liquid that is slightly more viscous than water . It is used as an oxidizer , bleaching agent, and antiseptic , usually as a dilute solution (3%–6% by weight) in water for consumer use and in higher concentrations for industrial use. Concentrated hydrogen peroxide, or " high-test peroxide ", decomposes explosively when heated and has been used as both a monopropellant and an oxidizer in rocketry . [ 6 ]
Hydrogen peroxide is a reactive oxygen species and the simplest peroxide , a compound having an oxygen–oxygen single bond . It decomposes slowly into water and elemental oxygen when exposed to light, and rapidly in the presence of organic or reactive compounds. It is typically stored with a stabilizer in a weakly acidic solution in an opaque bottle. Hydrogen peroxide is found in biological systems including the human body. Enzymes that use or decompose hydrogen peroxide are classified as peroxidases .
The boiling point of H 2 O 2 has been extrapolated as being 150.2 °C (302.4 °F), approximately 50 °C (90 °F) higher than water. In practice, hydrogen peroxide will undergo potentially explosive thermal decomposition if heated to this temperature. It may be safely distilled at lower temperatures under reduced pressure. [ 7 ]
Hydrogen peroxide forms stable adducts with urea ( hydrogen peroxide–urea ), sodium carbonate ( sodium percarbonate ) and other compounds. [ 8 ] An acid-base adduct with triphenylphosphine oxide is a useful "carrier" for H 2 O 2 in some reactions.
Hydrogen peroxide ( H 2 O 2 ) is a nonplanar molecule with (twisted) C 2 symmetry ; this was first shown by Paul-Antoine Giguère in 1950 using infrared spectroscopy . [ 9 ] [ 10 ] Although the O−O bond is a single bond , the molecule has a relatively high rotational barrier of 386 cm −1 (4.62 kJ / mol ) for rotation between enantiomers via the trans configuration, and 2460 cm −1 (29.4 kJ/mol) via the cis configuration. [ 11 ] These barriers are proposed to be due to repulsion between the lone pairs of the adjacent oxygen atoms and dipolar effects between the two O–H bonds. For comparison, the rotational barrier for ethane is 1040 cm −1 (12.4 kJ/mol).
The approximately 100° dihedral angle between the two O–H bonds makes the molecule chiral . It is the smallest and simplest molecule to exhibit enantiomerism . It has been proposed that the enantiospecific interactions of one rather than the other may have led to amplification of one enantiomeric form of ribonucleic acids and therefore an origin of homochirality in an RNA world . [ 12 ]
The molecular structures of gaseous and crystalline H 2 O 2 are significantly different. This difference is attributed to the effects of hydrogen bonding , which is absent in the gaseous state. [ 13 ] Crystals of H 2 O 2 are tetragonal with the space group D 4 4 or P 4 1 2 1 2. [ 14 ]
In aqueous solutions , hydrogen peroxide forms a eutectic mixture, exhibiting freezing-point depression down as low as −56 °C; pure water has a freezing point of 0 °C and pure hydrogen peroxide of −0.43 °C. The boiling point of the same mixtures is also depressed in relation with the mean of both boiling points (125.1 °C). It occurs at 114 °C. This boiling point is 14 °C greater than that of pure water and 36.2 °C less than that of pure hydrogen peroxide. [ 15 ]
Hydrogen peroxide is most commonly available as a solution in water. For consumers, it is usually available from pharmacies at 3 and 6 wt% concentrations. The concentrations are sometimes described in terms of the volume of oxygen gas generated; one milliliter of a 20-volume solution generates twenty milliliters of oxygen gas when completely decomposed. For laboratory use, 30 wt% solutions are most common. Commercial grades from 70% to 98% are also available, but due to the potential of solutions of more than 68% hydrogen peroxide to be converted entirely to steam and oxygen (with the temperature of the steam increasing as the concentration increases above 68%) these grades are potentially far more hazardous and require special care in dedicated storage areas. Buyers must typically allow inspection by commercial manufacturers.
Hydrogen peroxide has several structural analogues with H m X−XH n bonding arrangements (water also shown for comparison). It has the highest (theoretical) boiling point of this series (X = O, S, N, P). Its melting point is also fairly high, being comparable to that of hydrazine and water, with only hydroxylamine crystallising significantly more readily, indicative of particularly strong hydrogen bonding. Diphosphane and hydrogen disulfide exhibit only weak hydrogen bonding and have little chemical similarity to hydrogen peroxide. Structurally, the analogues all adopt similar skewed structures, due to repulsion between adjacent lone pairs .
Hydrogen peroxide is produced by various biological processes mediated by enzymes .
Hydrogen peroxide has been detected in surface water, in groundwater, and in the atmosphere . It can also form when water is exposed to UV light. [ 16 ] Sea water contains 0.5 to 14 μg/L of hydrogen peroxide, and freshwater contains 1 to 30 μg/L. [ 17 ] Concentrations in air are about 0.4 to 4 μg/m 3 , varying over several orders of magnitude depending in conditions such as season, altitude, daylight and water vapor content. In rural nighttime air it is less than 0.014 μg/m 3 , and in moderate photochemical smog it is 14 to 42 μg/m 3 . [ 18 ]
The amount of hydrogen peroxide in biological systems can be assayed using a fluorometric assay . [ 19 ]
Alexander von Humboldt is sometimes said to have been the first to report the first synthetic peroxide, barium peroxide , in 1799 as a by-product of his attempts to decompose air, although this is disputed due to von Humboldt's ambiguous wording. [ 20 ] Nineteen years later Louis Jacques Thénard recognized that this compound could be used for the preparation of a previously unknown compound, which he described as eau oxygénée ("oxygenated water") — subsequently known as hydrogen peroxide. [ 21 ] [ 22 ] [ 23 ]
An improved version of Thénard's process used hydrochloric acid , followed by addition of sulfuric acid to precipitate the barium sulfate byproduct. This process was used from the end of the 19th century until the middle of the 20th century. [ 24 ]
The bleaching effect of peroxides and their salts on natural dyes had been known since Thénard's experiments in the 1820s, but early attempts of industrial production of peroxides failed. The first plant producing hydrogen peroxide was built in 1873 in Berlin . The discovery of the synthesis of hydrogen peroxide by electrolysis with sulfuric acid introduced the more efficient electrochemical method. It was first commercialized in 1908 in Weißenstein , Carinthia , Austria. The anthraquinone process , which is still used, was developed during the 1930s by the German chemical manufacturer IG Farben in Ludwigshafen . The increased demand and improvements in the synthesis methods resulted in the rise of the annual production of hydrogen peroxide from 35,000 tonnes in 1950, to over 100,000 tonnes in 1960, to 300,000 tonnes by 1970; by 1998 it reached 2.7 million tonnes. [ 17 ]
Early attempts failed to produce neat hydrogen peroxide. Anhydrous hydrogen peroxide was first obtained by vacuum distillation . [ 25 ]
Determination of the molecular structure of hydrogen peroxide proved to be very difficult. In 1892, the Italian physical chemist Giacomo Carrara (1864–1925) determined its molecular mass by freezing-point depression , which confirmed that its molecular formula is H 2 O 2 . [ 26 ] H 2 O=O seemed to be just as possible as the modern structure, and as late as in the middle of the 20th century at least half a dozen hypothetical isomeric variants of two main options seemed to be consistent with the available evidence. [ 27 ] In 1934, the English mathematical physicist William Penney and the Scottish physicist Gordon Sutherland proposed a molecular structure for hydrogen peroxide that was very similar to the presently accepted one. [ 28 ] [ 29 ]
In 1994, world production of H 2 O 2 was around 1.9 million tonnes and grew to 2.2 million in 2006, [ 30 ] most of which was at a concentration of 70% or less. In that year, bulk 30% H 2 O 2 sold for around 0.54 USD / kg , equivalent to US$1.50/kg (US$0.68/ lb ) on a "100% basis". [ 31 ] [ clarification needed ]
Today, hydrogen peroxide is manufactured almost exclusively by the anthraquinone process , which was originally developed by BASF in 1939. It begins with the reduction of an anthraquinone (such as 2-ethylanthraquinone or the 2-amyl derivative) to the corresponding anthrahydroquinone, typically by hydrogenation on a palladium catalyst . In the presence of oxygen , the anthrahydroquinone then undergoes autoxidation : the labile hydrogen atoms of the hydroxy groups transfer to the oxygen molecule, to give hydrogen peroxide and regenerating the anthraquinone. Most commercial processes achieve oxidation by bubbling compressed air through a solution of the anthrahydroquinone, with the hydrogen peroxide then extracted from the solution and the anthraquinone recycled back for successive cycles of hydrogenation and oxidation. [ 31 ] [ 32 ]
The net reaction for the anthraquinone-catalyzed process is: [ 31 ]
The economics of the process depend heavily on effective recycling of the extraction solvents, the hydrogenation catalyst and the expensive quinone .
Hydrogen peroxide was once prepared industrially by hydrolysis of ammonium persulfate :
[NH 4 ] 2 S 2 O 8 was itself obtained by the electrolysis of a solution of ammonium bisulfate ( [NH 4 ]HSO 4 ) in sulfuric acid . [ 33 ]
Small amounts are formed by electrolysis, photochemistry , electric arc , and related methods. [ 34 ]
A commercially viable route for hydrogen peroxide via the reaction of hydrogen with oxygen favours production of water but can be stopped at the peroxide stage. [ 35 ] [ 36 ] One economic obstacle has been that direct processes give a dilute solution uneconomic for transportation. None of these has yet reached a point where it can be used for industrial-scale synthesis.
Hydrogen peroxide is about 1000 times stronger as an acid than water. [ 37 ]
Hydrogen peroxide disproportionates to form water and oxygen with a Δ H o of −2884.5 kJ / kg [ 38 ] and a Δ S of 70.5 J/(mol·K):
The rate of decomposition increases with rise in temperature, concentration, and pH . H 2 O 2 is unstable under alkaline conditions. Decomposition is catalysed by various redox-active ions or compounds, including most transition metals and their compounds (e.g. manganese dioxide ( MnO 2 ), silver , and platinum ). [ 39 ]
The redox properties of hydrogen peroxide depend on pH. In acidic solutions, H 2 O 2 is a powerful oxidizer .
Sulfite ( SO 2− 3 ) is oxidized to sulfate ( SO 2− 4 ).
Under alkaline conditions, hydrogen peroxide is a reductant. When H 2 O 2 acts as a reducing agent, oxygen gas is also produced. For example, hydrogen peroxide will reduce sodium hypochlorite and potassium permanganate , which is a convenient method for preparing oxygen in the laboratory:
The oxygen produced from hydrogen peroxide and sodium hypochlorite is in the singlet state .
Hydrogen peroxide also reduces silver oxide to silver :
Although usually a reductant, alkaline hydrogen peroxide converts Mn(II) to the dioxide:
In a related reaction, potassium permanganate is reduced to Mn 2+ by acidic H 2 O 2 : [ 5 ]
Hydrogen peroxide is frequently used as an oxidizing agent . Illustrative is oxidation of thioethers to form sulfoxides , such as conversion of thioanisole to methyl phenyl sulfoxide : [ 40 ] [ 41 ]
Alkaline hydrogen peroxide is used for epoxidation of electron-deficient alkenes such as acrylic acid derivatives, [ 42 ] and for the oxidation of alkylboranes to alcohols , the second step of hydroboration-oxidation . It is also the principal reagent in the Dakin oxidation process.
Hydrogen peroxide is a weak acid, forming hydroperoxide or peroxide salts with many metals.
It also converts metal oxides into the corresponding peroxides. For example, upon treatment with hydrogen peroxide, chromic acid ( CrO 3 and H 2 SO 4 ) forms a blue peroxide CrO(O 2 ) 2 .
The aerobic oxidation of glucose in the presence of the enzyme glucose oxidase produces hydrogen peroxide. The conversion affords gluconolactone : [ 43 ]
Superoxide dismutases (SOD)s are enzymes that promote the disproportionation of superoxide into oxygen and hydrogen peroxide. [ 44 ]
Peroxisomes are organelles found in virtually all eukaryotic cells. [ 45 ] They are involved in the catabolism of very long chain fatty acids , branched chain fatty acids , D -amino acids , polyamines , and biosynthesis of plasmalogens and ether phospholipids , which are found in mammalian brains and lungs. [ 46 ] They produce hydrogen peroxide in a process catalyzed by flavin adenine dinucleotide (FAD): [ 47 ]
Hydrogen peroxide arises by the degradation of adenosine monophosphate , which yields hypoxanthine . Hypoxanthine is then oxidatively catabolized first to xanthine and then to uric acid , and the reaction is catalyzed by the enzyme xanthine oxidase : [ 48 ]
Hypoxanthine
Xanthine oxidase
Xanthine
Xanthine oxidase
Uric acid
The degradation of guanosine monophosphate yields xanthine as an intermediate product which is then converted in the same way to uric acid with the formation of hydrogen peroxide. [ 48 ]
Catalase , another peroxisomal enzyme, uses this H 2 O 2 to oxidize other substrates, including phenols , formic acid , formaldehyde , and alcohol , by means of a peroxidation reaction:
thus eliminating the poisonous hydrogen peroxide in the process.
This reaction is important in liver and kidney cells, where the peroxisomes neutralize various toxic substances that enter the blood. Some of the ethanol humans drink is oxidized to acetaldehyde in this way. [ 49 ] In addition, when excess H 2 O 2 accumulates in the cell, catalase converts it to H 2 O through this reaction:
Glutathione peroxidase , a selenoenzyme , also catalyzes the disproportionation of hydrogen peroxide.
The reaction of Fe 2+ and hydrogen peroxide is the basis of the Fenton reaction , which generates hydroxyl radicals , which are of significance in biology:
The Fenton reaction explains the toxicity of hydrogen peroxides because the hydroxyl radicals rapidly and irreversibly oxidize all organic compounds, including proteins , membrane lipids , and DNA . [ 50 ] Hydrogen peroxide is a significant source of oxidative DNA damage in living cells. DNA damage includes formation of 8-Oxo-2'-deoxyguanosine among many other altered bases, as well as strand breaks, inter-strand crosslinks, and deoxyribose damage. [ 51 ] By interacting with Cl − , hydrogen peroxide also leads to chlorinated DNA bases. [ 51 ] Hydroxyl radicals readily damage vital cellular components, especially those of the mitochondria . [ 52 ] [ 53 ] [ 54 ] The compound is a major factor implicated in the free-radical theory of aging , based on its ready conversion into a hydroxyl radical .
Eggs of sea urchin , shortly after fertilization by a sperm, produce hydrogen peroxide. It is then converted to hydroxyl radicals (HO•), which initiate radical polymerization , which surrounds the eggs with a protective layer of polymer .
The bombardier beetle combines hydroquinone and hydrogen peroxide, leading to a violent exothermic chemical reaction to produce boiling, foul-smelling liquid that partially becomes a gas ( flash evaporation ) and is expelled through an outlet valve with a loud popping sound. [ 55 ] [ 56 ] [ 57 ]
As a proposed signaling molecule , hydrogen peroxide may regulate a wide variety of biological processes. [ 58 ] [ 59 ] At least one study has tried to link hydrogen peroxide production to cancer. [ 60 ]
About 60% of the world's production of hydrogen peroxide is used for pulp- and paper-bleaching . [ 30 ] The second major industrial application is the manufacture of sodium percarbonate and sodium perborate , which are used as mild bleaches in laundry detergents . A representative conversion is:
Sodium percarbonate, which is an adduct of sodium carbonate and hydrogen peroxide, is the active ingredient in such laundry products as OxiClean and Tide laundry detergent . When dissolved in water, it releases hydrogen peroxide and sodium carbonate. [ 24 ] By themselves these bleaching agents are only effective at wash temperatures of 60 °C (140 °F) or above and so, often are used in conjunction with bleach activators , which facilitate cleaning at lower temperatures.
Hydrogen peroxide has also been used as a flour bleaching agent and a tooth and bone whitening agent.
It is used in the production of various organic peroxides with dibenzoyl peroxide being a high volume example. [ 61 ] Peroxy acids , such as peracetic acid and meta -chloroperoxybenzoic acid also are produced using hydrogen peroxide. Hydrogen peroxide has been used for creating organic peroxide -based explosives, such as acetone peroxide . It is used as an initiator in polymerizations . Hydrogen peroxide reacts with certain di- esters , such as phenyl oxalate ester (cyalume), to produce chemiluminescence ; this application is most commonly encountered in the form of glow sticks .
The reaction with borax leads to sodium perborate , a bleach used in laundry detergents:
Hydrogen peroxide is used in certain waste-water treatment processes to remove organic impurities. In advanced oxidation processing , the Fenton reaction [ 62 ] [ 63 ] gives the highly reactive hydroxyl radical (•OH). This degrades organic compounds, including those that are ordinarily robust, such as aromatic or halogenated compounds . [ 64 ] It can also oxidize sulfur -based compounds present in the waste; which is beneficial as it generally reduces their odour. [ 65 ]
Hydrogen peroxide may be used for the sterilization of various surfaces, [ 66 ] including surgical instruments, [ 67 ] and may be deployed as a vapour ( VHP ) for room sterilization. [ 68 ] H 2 O 2 demonstrates broad-spectrum efficacy against viruses, bacteria, yeasts, and bacterial spores. [ 69 ] [ 70 ] In general, greater activity is seen against Gram-positive than Gram-negative bacteria; however, the presence of catalase or other peroxidases in these organisms may increase tolerance in the presence of lower concentrations. [ 71 ] Lower levels of concentration (3%) will work against most spores; higher concentrations (7 to 30%) and longer contact times will improve sporicidal activity. [ 70 ] [ 72 ]
Hydrogen peroxide is seen as an environmentally safe alternative to chlorine -based bleaches, as it degrades to form oxygen and water and it is generally recognized as safe as an antimicrobial agent by the U.S. Food and Drug Administration (FDA). [ 73 ]
High-concentration H 2 O 2 is referred to as "high-test peroxide" (HTP). It can be used as either a monopropellant (not mixed with fuel) or the oxidizer component of a bipropellant rocket . Use as a monopropellant takes advantage of the decomposition of 70–98% concentration hydrogen peroxide into steam and oxygen. The propellant is pumped into a reaction chamber, where a catalyst, usually a silver or platinum screen, triggers decomposition, producing steam at over 600 °C (1,100 °F), which is expelled through a nozzle , generating thrust . H 2 O 2 monopropellant produces a maximal specific impulse ( I sp ) of 161 s (1.6 kN·s /kg). Peroxide was the first major monopropellant adopted for use in rocket applications. Hydrazine eventually replaced hydrogen peroxide monopropellant thruster applications primarily because of a 25% increase in the vacuum specific impulse. [ 74 ] Hydrazine (toxic) and hydrogen peroxide (less toxic [ACGIH TLV 0.01 and 1 ppm respectively]) are the only two monopropellants (other than cold gases) to have been widely adopted and utilized for propulsion and power applications. [ citation needed ] The Bell Rocket Belt , reaction control systems for X-1 , X-15 , Centaur , Mercury , Little Joe , as well as the turbo-pump gas generators for X-1, X-15, Jupiter, Redstone and Viking used hydrogen peroxide as a monopropellant. [ 75 ] The RD-107 engines (used from 1957 to present) in the R-7 series of rockets decompose hydrogen peroxide to power the turbopumps.
In bipropellant applications, H 2 O 2 is decomposed to oxidize a burning fuel. Specific impulses as high as 350 s (3.5 kN·s/kg) can be achieved, depending on the fuel. Peroxide used as an oxidizer gives a somewhat lower I sp than liquid oxygen but is dense, storable, and non-cryogenic and can be more easily used to drive gas turbines to give high pressures using an efficient closed cycle . It may also be used for regenerative cooling of rocket engines. Peroxide was used very successfully as an oxidizer in World War II German rocket motors (e.g., T-Stoff , containing oxyquinoline stabilizer, for both the Walter HWK 109-500 Starthilfe RATO externally podded monopropellant booster system and the Walter HWK 109-509 rocket motor series used for the Me 163 B), most often used with C-Stoff in a self-igniting hypergolic combination, and for the low-cost British Black Knight and Black Arrow launchers. Presently, HTP is used on ILR-33 AMBER [ 76 ] and Nucleus [ 77 ] suborbital rockets.
In the 1940s and 1950s, the Hellmuth Walter KG –conceived turbine used hydrogen peroxide for use in submarines while submerged; it was found to be too noisy and require too much maintenance compared to diesel-electric power systems. Some torpedoes used hydrogen peroxide as oxidizer or propellant. Operator error in the use of hydrogen peroxide torpedoes was named as possible causes for the sinking of HMS Sidon and the Russian submarine Kursk . [ 78 ] SAAB Underwater Systems is manufacturing the Torpedo 2000. This torpedo, used by the Swedish Navy , is powered by a piston engine propelled by HTP as an oxidizer and kerosene as a fuel in a bipropellant system. [ 79 ] [ 80 ]
Hydrogen peroxide has various domestic uses, primarily as a cleaning and disinfecting agent.
Diluted H 2 O 2 (between 1.9% and 12%) mixed with aqueous ammonia has been used to bleach human hair . The chemical's bleaching property lends its name to the phrase " peroxide blonde ". [ 81 ] Hydrogen peroxide is also used for tooth whitening . It may be found in most whitening toothpastes. Hydrogen peroxide has shown positive results involving teeth lightness and chroma shade parameters. [ 82 ] It works by oxidizing colored pigments onto the enamel where the shade of the tooth may become lighter. [ further explanation needed ] Hydrogen peroxide may be mixed with baking soda and salt to make a homemade toothpaste. [ 83 ]
Hydrogen peroxide reacts with blood as a bleaching agent, and so if a blood stain is fresh, or not too old, liberal application of hydrogen peroxide, if necessary in more than single application, will bleach the stain fully out. After about two minutes of the application, the blood should be firmly blotted out. [ 84 ] [ 85 ]
Hydrogen peroxide may be used to treat acne , [ 86 ] although benzoyl peroxide is a more common treatment.
The use of dilute hydrogen peroxide as an oral cleansing agent has been reviewed academically to determine its usefulness in treating gingivitis and plaque . Although there is a positive effect when compared with a placebo, it was concluded that chlorhexidine is a much more effective treatment. [ 87 ]
Some horticulturists and users of hydroponics advocate the use of weak hydrogen peroxide solution in watering solutions. Its spontaneous decomposition releases oxygen that enhances a plant's root development and helps to treat root rot (cellular root death due to lack of oxygen) and a variety of other pests. [ 88 ] [ 89 ]
For general watering concentrations, around 0.1% is in use. This can be increased up to one percent for antifungal actions. [ 90 ] Tests show that plant foliage can safely tolerate concentrations up to 3%. [ 91 ]
Hydrogen peroxide is used in aquaculture for controlling mortality caused by various microbes. In 2019, the U.S. FDA approved it for control of Saprolegniasis in all coldwater finfish and all fingerling and adult coolwater and warmwater finfish, for control of external columnaris disease in warm-water finfish, and for control of Gyrodactylus spp. in freshwater-reared salmonids. [ 92 ] Laboratory tests conducted by fish culturists have demonstrated that common household hydrogen peroxide may be used safely to provide oxygen for small fish. The hydrogen peroxide releases oxygen by decomposition when it is exposed to catalysts such as manganese dioxide .
Hydrogen peroxide may be used in combination with a UV-light source to remove yellowing from white or light grey acrylonitrile butadiene styrene (ABS) plastics to partially or fully restore the original color. In the retrocomputing scene, this process is commonly referred to as retrobright .
Regulations vary, but low concentrations, such as 5%, are widely available and legal to buy for medical use. Most over-the-counter peroxide solutions are not suitable for ingestion. Higher concentrations may be considered hazardous and typically are accompanied by a safety data sheet (SDS). In high concentrations, hydrogen peroxide is an aggressive oxidizer and will corrode many materials, including human skin. In the presence of a reducing agent , high concentrations of H 2 O 2 will react violently. [ 93 ] While concentrations up to 35% produce only "white" oxygen bubbles in the skin (and some biting pain) that disappear with the blood within 30–45 minutes, concentrations of 98% dissolve paper. However, concentrations as low as 3% can be dangerous for the eye because of oxygen evolution within the eye. [ 94 ]
High-concentration hydrogen peroxide streams, typically above 40%, should be considered hazardous due to concentrated hydrogen peroxide's meeting the definition of a DOT oxidizer according to U.S. regulations if released into the environment. The EPA Reportable Quantity (RQ) for D001 hazardous wastes is 100 pounds (45 kg), or approximately 10 US gallons (38 L), of concentrated hydrogen peroxide.
Hydrogen peroxide should be stored in a cool, dry, well-ventilated area and away from any flammable or combustible substances. It should be stored in a container composed of non-reactive materials such as stainless steel or glass (other materials including some plastics and aluminium alloys may also be suitable). [ 95 ] As it breaks down quickly when exposed to light, it should be stored in an opaque container, and pharmaceutical formulations typically come in brown bottles that block light. [ 96 ]
Hydrogen peroxide, either in pure or diluted form, may pose several risks, the main one being that it forms explosive mixtures upon contact with organic compounds. [ 97 ] Distillation of hydrogen peroxide at normal pressures is highly dangerous. It is corrosive, especially when concentrated, but even domestic-strength solutions may cause irritation to the eyes, mucous membranes , and skin. [ 98 ] Swallowing hydrogen peroxide solutions is particularly dangerous, as decomposition in the stomach releases large quantities of gas (ten times the volume of a 3% solution), leading to internal bloating. Inhaling over 10% can cause severe pulmonary irritation. [ 99 ]
With a significant vapour pressure (1.2 kPa at 50 °C), [ 100 ] hydrogen peroxide vapour is potentially hazardous. According to U.S. NIOSH, the immediately dangerous to life and health (IDLH) limit is only 75 ppm. [ 101 ] The U.S. Occupational Safety and Health Administration (OSHA) has established a permissible exposure limit of 1.0 ppm calculated as an 8-hour time-weighted average (29 CFR 1910.1000, Table Z-1). [ 97 ] Hydrogen peroxide has been classified by the American Conference of Governmental Industrial Hygienists (ACGIH) as a "known animal carcinogen, with unknown relevance on humans". [ 102 ] For workplaces where there is a risk of exposure to the hazardous concentrations of the vapours, continuous monitors for hydrogen peroxide should be used. Information on the hazards of hydrogen peroxide is available from OSHA [ 97 ] and from the ATSDR. [ 103 ]
Historically, hydrogen peroxide was used for disinfecting wounds, partly because of its low cost and prompt availability compared to other antiseptics . [ 104 ]
There is conflicting evidence on hydrogen peroxide's effect on wound healing. Some research finds benefit, while other research find delays and healing inhibition. [ 105 ] Its use for home treatment of wounds is generally not recommended. [ 106 ] 1.5–3% hydrogen peroxide is used as a disinfectant in dentistry, especially in endodotic treatments together with hypochlorite and chlorhexidine and 1–1.5% is also useful for treatment of inflammation of third molars (wisdom teeth). [ 107 ]
Practitioners of alternative medicine have advocated the use of hydrogen peroxide for various conditions, including emphysema , influenza , AIDS , and in particular cancer . [ 108 ] There is no evidence of effectiveness and in some cases it has proved fatal. [ 109 ] [ 110 ] [ 111 ] [ 112 ]
Both the effectiveness and safety of hydrogen peroxide therapy is scientifically questionable. Hydrogen peroxide is produced by the immune system, but in a carefully controlled manner. Cells called phagocytes engulf pathogens and then use hydrogen peroxide to destroy them. The peroxide is toxic to both the cell and the pathogen and so is kept within a special compartment, called a phagosome . Free hydrogen peroxide will damage any tissue it encounters via oxidative stress , a process that also has been proposed as a cause of cancer. [ 113 ] Claims that hydrogen peroxide therapy increases cellular levels of oxygen have not been supported. The quantities administered would be expected to provide very little additional oxygen compared to that available from normal respiration. It is also difficult to raise the level of oxygen around cancer cells within a tumour, as the blood supply tends to be poor, a situation known as tumor hypoxia .
Large oral doses of hydrogen peroxide at a 3% concentration may cause irritation and blistering to the mouth, throat, and abdomen as well as abdominal pain, vomiting, and diarrhea. [ 109 ] Ingestion of hydrogen peroxide at concentrations of 35% or higher has been implicated as the cause of numerous gas embolism events resulting in hospitalisation. In these cases, hyperbaric oxygen therapy was used to treat the embolisms. [ 114 ]
Intravenous injection of hydrogen peroxide has been linked to several deaths. [ 115 ] [ 111 ] [ 112 ] The American Cancer Society states that "there is no scientific evidence that hydrogen peroxide is a safe, effective, or useful cancer treatment." [ 110 ] Furthermore, the therapy is not approved by the U.S. FDA.
Bibliography | https://en.wikipedia.org/wiki/H2O2 |
Trioxidane (systematically named dihydrogen trioxide , [ 2 ] [ 3 ] ), also called hydrogen trioxide [ 4 ] [ 5 ] is an inorganic compound with the chemical formula H[O] 3 H (can be written as [H( μ -O 3 )H] or [H 2 O 3 ] ). It is one of the unstable hydrogen polyoxides . [ 4 ] In aqueous solutions, trioxidane decomposes to form water and singlet oxygen :
The reverse reaction, the addition of singlet oxygen to water, typically does not occur in part due to the scarcity of singlet oxygen. In biological systems, however, ozone is known to be generated from singlet oxygen, and the presumed mechanism is an antibody-catalyzed production of trioxidane from singlet oxygen. [ 2 ]
Trioxidane can be obtained in small, but detectable, amounts in reactions of ozone and hydrogen peroxide , or by the electrolysis of water . Larger quantities have been prepared by the reaction of ozone with organic reducing agents at low temperatures in a variety of organic solvents, such as the anthraquinone process . It is also formed during the decomposition of organic hydrotrioxides (ROOOH). [ 3 ] Alternatively, trioxidane can be prepared by reduction of ozone with 1,2-diphenylhydrazine at low temperature. Using a resin-bound version of the latter, relatively pure trioxidane can be isolated as a solution in organic solvent. Preparation of high purity solutions is possible using the methyltrioxorhenium(VII) catalyst. [ 5 ] In acetone- d 6 at −20 °C, the characteristic 1 H NMR signal of trioxidane could be observed at a chemical shift of 13.1 ppm. [ 3 ] Solutions of hydrogen trioxide in diethyl ether can be safely stored at −20 °C for as long as a week. [ 5 ]
The reaction of ozone with hydrogen peroxide is known as the "peroxone process". This mixture has been used for some time for treating groundwater contaminated with organic compounds. The reaction produces H 2 O 3 and H 2 O 5 . [ 6 ]
In 1970-75, Giguère et al. observed infrared and Raman spectra of dilute aqueous solutions of trioxidane. [ 4 ] In 2005, trioxidane was observed experimentally by microwave spectroscopy in a supersonic jet. The molecule exists in a skewed structure, with an oxygen–oxygen–oxygen–hydrogen dihedral angle of 81.8°. The oxygen–oxygen bond lengths of 142.8 picometer are slightly shorter than the 146.4 pm oxygen–oxygen bonds in hydrogen peroxide . [ 7 ] Various dimeric and trimeric forms also seem to exist.
There is a trend of increasing gas-phase acidity and corresponding p K a as the number of oxygen atoms in the chain increases in HO n H structures ( n =1,2,3). [ 8 ]
Trioxidane readily decomposes into water and singlet oxygen, with a half-life of about 16 minutes in organic solvents at room temperature, but only milliseconds in water. It reacts with organic sulfides to form sulfoxides , but little else is known of its reactivity.
Recent research found that trioxidane is the active ingredient responsible for the antimicrobial properties of the well known ozone / hydrogen peroxide mix. Because these two compounds are present in biological systems as well it is argued that an antibody in the human body can generate trioxidane as a powerful oxidant against invading bacteria. [ 2 ] [ 9 ] The source of the compound in biological systems is the reaction between singlet oxygen and water (which proceeds in either direction, of course, according to concentrations), with the singlet oxygen being produced by immune cells. [ 3 ] [ 10 ]
Computational chemistry predicts that more oxygen chain molecules or hydrogen polyoxides exist and that even indefinitely long oxygen chains can exist in a low-temperature gas. With this spectroscopic evidence a search for these type of molecules can start in interstellar space . [ 7 ] A 2022 publication suggested the possibility of the presence of detectable concentrations of polyoxides in the atmosphere. [ 11 ] | https://en.wikipedia.org/wiki/H2O3 |
Selenous acid (or selenious acid ) is the chemical compound with the formula H 2 SeO 3 . Structurally, it is more accurately described by O=Se(OH) 2 . It is the principal oxoacid of selenium ; the other being selenic acid .
Selenous acid is analogous to sulfurous acid , but it is more readily isolated. Selenous acid is easily formed upon the addition of selenium dioxide to water. As a crystalline solid, the compound can be seen as pyramidal molecules that are interconnected with hydrogen bonds. In solution it is a diprotic acid: [ 3 ]
It is moderately oxidizing in nature, but kinetically slow. In 1 M H + :
In 1 M OH − :
Selenous acid is hygroscopic . [ 4 ] [ 5 ]
The major use is in protecting and changing the color of steel, especially steel parts on firearms. [ 6 ] The so-called cold-bluing process uses selenous acid, copper(II) nitrate , and nitric acid to change the color of the steel from silver-grey to blue-grey or black. Alternative procedures use copper sulfate and phosphoric acid instead. This process deposits a coating of copper selenide and is fundamentally different from other bluing processes which generate black iron oxide . Some older razor blades were also made of blued steel. [ 6 ]
Another use for selenious acid is the chemical darkening and patination of copper, brass and bronze, producing a rich dark brown color that can be further enhanced with mechanical abrasion. [ citation needed ]
It is used in organic synthesis as an oxidizing agent for the synthesis of 1,2-dicarbonyl compounds, e.g. in laboratory preparation of glyoxal (oxaldehyde) from acetaldehyde . [ 7 ]
Selenious acid is a key component of the Mecke reagent used for drug checking. [ 8 ] [ 9 ]
Selenous acid can supply the trace element indicated in people as a source of selenium. [ 10 ] [ 11 ]
Like many selenium compounds, selenous acid is highly toxic in excessive quantities, and ingestion of any significant quantity of selenous acid is usually fatal, however it is an approved dietary source in proper amounts. Symptoms of selenium poisoning can occur several hours after exposure, and may include stupor , nausea , severe hypotension and death. | https://en.wikipedia.org/wiki/H2O3Se |
Tellurous acid is an inorganic compound with the formula H 2 TeO 3 . It is the oxoacid of tellurium(IV). [ 2 ] This compound is not well characterized. An alternative way of writing its formula is (HO) 2 TeO. In principle, tellurous acid would form by treatment of tellurium dioxide with water, that is by hydrolysis. The related conjugate base is well known in the form of several salts such as potassium hydrogen tellurite, KHTeO 3 .
In contrast to the analogous compound selenous acid , tellurous acid is only metastable. Most tellurite salts contain the TeO 2− 3 ion. Oxidation of its aqueous solution with hydrogen peroxide gives the tellurate ion. It is usually prepared as an aqueous solution where it acts as a weak acid. [ 1 ] [ 3 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2O3Te |
Tetraoxidane is an inorganic compound of hydrogen and oxygen with the chemical formula H 2 O 4 . [ 1 ] [ 2 ] [ 3 ] This is one of the unstable hydrogen polyoxides . [ 4 ]
The compound is prepared by a chemical reaction between hydroperoxyl radicals ( HOO• ) at low temperatures: [ 5 ] [ 6 ]
This is the fourth member of the polyoxidanes. The first three are water [(mon)oxidane], hydrogen peroxide (dioxidane), and trioxidane . Tetraoxidane is more unstable than the previous compounds. The term "tetraoxidane" extends beyond the parent compound to several daughter compounds of the general formula R 2 O 4 , where R can be hydrogen, halogen atoms, or various inorganic and organic monovalent radicals. The two Rs together can be replaced by a divalent radical, so heterocyclic tetraoxidanes also exist. [ 7 ]
Tetraoxidane autoionizes when in liquid form: | https://en.wikipedia.org/wiki/H2O4 |
Selenic acid is the inorganic compound with the formula H 2 SeO 4 . It is an oxoacid of selenium , and its structure is more accurately described as O 2 Se(OH) 2 . It is a colorless compound. Although it has few uses, one of its salts, sodium selenate is used in the production of glass and animal feeds. [ 3 ]
The molecule is tetrahedral, as predicted by VSEPR theory . The Se–O bond length is 161 pm . [ 4 ] In the solid state, it crystallizes in an orthorhombic structure. [ 5 ]
It is prepared by oxidising selenium compounds in lower oxidation states. One method involves the oxidation of selenium dioxide with hydrogen peroxide :
Unlike the production sulfuric acid by hydration of sulfur trioxide , the hydration of selenium trioxide is an impractical method. [ 4 ] Instead, selenic acid may also be prepared by the oxidation of selenous acid ( H 2 SeO 3 ) with halogens, such as chlorine or bromine , or with potassium permanganate . [ 6 ] Using chlorine or bromine as the oxidising agents also produces hydrochloric or hydrobromic acid as a side-product, which needs to be removed from the solution since they can reduce the selenic acid to selenous acid. [ 7 ]
To obtain the anhydrous acid as a crystalline solid, the resulting solution is evaporated at temperatures below 140 °C (413 K; 284 °F) in a vacuum. [ 8 ]
Like sulfuric acid , selenic acid is a strong acid that is hygroscopic and extremely soluble in water. Concentrated solutions are viscous. Crystalline mono- and di- hydrates are known. [ 6 ] The monohydrate melts at 26 °C, and the dihydrate melts at −51.7 °C. [ 4 ]
Selenic acid is a stronger oxidizer than sulfuric acid , [ 9 ] capable of liberating chlorine from chloride ions , being reduced to selenous acid in the process:
It decomposes above 200 °C, liberating oxygen gas and being reduced to selenous acid: [ 6 ]
Selenic acid reacts with barium salts to precipitate solid BaSeO 4 , analogous to the sulfate. In general, selenate salts resemble sulfate salts, but are more soluble. Many selenate salts have the same crystal structure as the corresponding sulfate salts. [ 4 ]
Treatment with fluorosulfuric acid gives selenoyl fluoride : [ 8 ]
Hot, concentrated selenic acid reacts with gold , forming a reddish-yellow solution of gold(III) selenate: [ 10 ]
Selenic acid is used as a specialized oxidizing agent. | https://en.wikipedia.org/wiki/H2O4Se |
Uranyl hydroxide is a hydroxide of uranium with the chemical formula UO 2 (OH) 2 in the monomeric form and [(UO 2 ) 2 (OH) 4 ] 2- in the dimeric; both forms may exist in normal aqueous media. In aerobic conditions, up to 5 hydroxides can bind to uranyl ([(UO 2 ) 2 (OH) 5 ] 3- ). Uranyl hydroxide hydrate is precipitated as a colloidal yellowcake from oxidized uranium liquors near neutral pH.
Uranyl hydroxide was once used in glassmaking and ceramics in the colouring of the vitreous phases and the preparation of pigments for high temperature firing. The introduction of alkaline di uranates (like sodium diuranate ) into glasses leads to yellow by transmission, green by reflection; moreover these glasses become dichroic and fluorescent under ultraviolet rays.
Uranyl hydroxide is teratogenic and radioactive .
The formation of uranyl hydroxide hydrate can occur via hydrated uranyl fluoride [(UO 2 F 2 )(H 2 O)] 7 ·4H 2 O which is not stable at an elevated water vapor pressure. A complete loss of fluorine is undergone and the formation of uranyl hydroxide hydrate ([(UO 2 ) 4 O(OH) 6 ]·5H 2 O) occurs. This uranyl hydroxide species is structurally similar to the uranyl hydroxide hydrate minerals schoepite and metaschoepite. X-ray diffraction data was gathered and found that this species has expanded interlayer spacing suggesting there may be additional water molecules in between uranyl layers. Unlike metaschoepite, however, this species does not form UO 2 (OH) 2 upon dehydration. [ 1 ]
UO 2 (OH) 2 reacts with water in a hydration reaction to form [(UO 2 (OH) 2 )(H 2 O)] + and the monohydrate form also reacted with water to form dihydrates [(UO 2 OH)(H 2 O) 2 ] + and trihydrates [(UO 2 OH)(H 2 O) 3 ] + . The hydration reaction to form the monohydrate was significantly slower than if the hydroxide were replaced with acetate or nitrate. This could be due to the strongly basic (OH) − reducing the Lewis acidity of U or because the more complex acetate and nitrate anions provide more degrees of freedom. However, it was found that the formation of the dihydrate uranyl hydroxide hydrate (2) was nearly three times faster than the monohydrate (1). [ 2 ]
A mechanism for oxygen exchange between the UO 2 2+ cations in a highly alkaline solution was proposed and investigated by Shamov et al. in the Journal of the American Chemical Society . [ 3 ] An equilibrium between [UO 2 (OH) 4 ] 2- and [UO 2 (OH) 5 ] 3- was observed followed by the formation of the stable [UO 3 (OH) 3 ·H 2 O] 3- intermediate that formed from [UO 2 (OH) 5 ] 3- via intramolecular water elimination. | https://en.wikipedia.org/wiki/H2O4U |
Tungstic acid refers to hydrated forms of tungsten trioxide , WO 3 . Both a monohydrate (WO 3 ·H 2 O) and hemihydrate (WO 3 · 1 / 2 H 2 O) [ 1 ] are known. Molecular species akin to sulfuric acid , i.e. (HO) 2 WO 2 are not observed.
The solid-state structure of WO 3 ·H 2 O consists of layers of octahedrally coordinated WO 5 (H 2 O) units where 4 vertices are shared. [ 2 ] The dihydrate has the same layer structure with the extra H 2 O molecule intercalated . [ 2 ] The monohydrate is a yellow solid and insoluble in water. The classical name for this acid is 'acid of wolfram'. Salts of tungstic acid are tungstates .
The acid was discovered by Carl Wilhelm Scheele in 1781. [ 3 ]
Tungstic acid is obtained by the action of strong acids on solutions of alkali metallic tungstates. It may also be prepared from the reaction between hydrogen carbonate and sodium tungstate . It can also be obtained from pure tungsten by reaction with hydrogen peroxide . [ 4 ]
It is used as a mordant and a dye in textiles. | https://en.wikipedia.org/wiki/H2O4W |
Xenic acid is a proposed noble gas compound with the chemical formula H 2 XeO 4 or XeO 2 (OH) 2 . It has not been isolated, and the published characterization data are ambiguous. [ 2 ]
Salts of xenic acid are called xenates , containing the HXeO − 4 anion, such as monosodium xenate . They tend to disproportionate into xenon gas and perxenates : [ 3 ]
The energy given off is sufficient to form ozone from diatomic oxygen:
Salts containing the deprotonated anion XeO 2− 4 are presently unknown. [ 3 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2O4Xe |
Pentaoxidane is an inorganic compound of hydrogen and oxygen with the chemical formula H 2 O 5 . [ 1 ] This is one of the most unstable hydrogen polyoxides . [ 2 ] [ 3 ] [ 4 ] | https://en.wikipedia.org/wiki/H2O5 |
Peroxymonosulfuric acid , also known as persulfuric acid , peroxysulfuric acid is the inorganic compound with the formula H 2 SO 5 . It is a white solid. It is a component of Caro's acid , which is a solution of peroxymonosulfuric acid in sulfuric acid containing small amounts of water. [ 4 ] Peroxymonosulfuric acid is a very strong oxidant ( E 0 = +2.51 V).
In peroxymonosulfuric acid, the S(VI) center adopts its characteristic tetrahedral geometry; the connectivity is indicated by the formula HO–O–S(O) 2 –OH. The S-O- H proton is more acidic. [ 4 ]
The German chemist Heinrich Caro first reported investigations of mixtures of hydrogen peroxide and sulfuric acid. [ 5 ]
One laboratory scale preparation of Caro's acid involves the combination of chlorosulfuric acid and hydrogen peroxide : [ 6 ]
Patents include more than one reaction for preparation of Caro's acid, usually as an intermediate for the production of potassium monopersulfate (PMPS) , a bleaching and oxidizing agent. One route employs the following reaction: [ 7 ]
This reaction is related to " piranha solution ".
H 2 SO 5 and Caro's acid have been used for a variety of disinfectant and cleaning applications, e.g., swimming pool treatment and denture cleaning. It is used in gold mining to destroy the cyanide in the waste stream (" Tailings ").
Alkali metal salts of H 2 SO 5 , especially oxone , are widely investigated.
These peroxy acids can be explosive. Explosions have been reported at Brown University [ 8 ] and Sun Oil . As with all strong oxidizing agents, peroxysulfuric acid is incompatible with organic compounds . | https://en.wikipedia.org/wiki/H2O5S |
Disulfuric acid (alternative spelling disulphuric acid ) or pyrosulfuric acid (alternative spelling pyrosulphuric acid ), also named oleum , is a sulfur oxoacid . [ 3 ] It is a major constituent of fuming sulfuric acid, oleum , and this is how most chemists encounter it. As confirmed by X-ray crystallography , the molecule consists of a pair of SO 2 (OH) groups joined by an oxide. [ 4 ]
It is also a minor constituent of liquid anhydrous sulfuric acid due to the equilibria:
The acid is prepared by reacting excess sulfur trioxide (SO 3 ) with sulfuric acid:
Disulfuric acid can be seen as the sulfuric acid analog of an acid anhydride . The mutual electron-withdrawing effects of each sulfuric acid unit on its neighbour causes a marked increase in acidity. Disulfuric acid is strong enough to protonate "normal" sulfuric acid in the (anhydrous) sulfuric acid solvent system. There are salts of disulfuric acid, commonly called pyrosulfates , e.g. potassium pyrosulfate .
There are other related acids with the general formula H 2 O·(SO 3 ) x though none can be isolated. | https://en.wikipedia.org/wiki/H2O7S2 |
This page provides supplementary data to the article properties of water .
Further comprehensive authoritative data can be found at the NIST Chemistry WebBook page on thermophysical properties of fluids. [ 1 ]
88.00 at 0 °C 86.04 at 5 °C 84.11 at 10 °C 82.22 at 15 °C 80.36 at 20 °C 78.54 at 25 °C 76.75 at 30 °C 75.00 at 35 °C 73.28 at 40 °C 71.59 at 45 °C 69.94 at 50 °C 66.74 at 60 °C 63.68 at 70 °C 60.76 at 80 °C 57.98 at 90 °C 55.33 at 100 °C
Vapor pressure formula for steam in equilibrium with liquid water: [ 14 ]
where P is equilibrium vapor pressure in k Pa , and T is temperature in kelvins .
For T = 273 K to 333 K: A = 7.2326; B = 1750.286; C = 38.1.
For T = 333 K to 423 K: A = 7.0917; B = 1668.21; C = 45.1.
Data in the table above is given for water–steam equilibria at various temperatures over the entire temperature range at which liquid water can exist. Pressure of the equilibrium is given in the second column in k Pa . The third column is the heat content of each gram of the liquid phase relative to water at 0 °C. The fourth column is the heat of vaporization of each gram of liquid that changes to vapor. The fifth column is the work P Δ V done by each gram of liquid that changes to vapor. The sixth column is the density of the vapor.
Data obtained from CRC Handbook of Chemistry and Physics 44th ed., p. 2390.
‡ Ice XI triple point is theoretical and has never been obtained
Note: ρ is density, n is refractive index at 589 nm, [ clarification needed ] and η is viscosity, all at 20 °C; T eq is the equilibrium temperature between two phases: ice/liquid solution for T eq < 0–0.1 °C and NaCl/liquid solution for T eq above 0.1 °C.
The data that follows was copied and translated from the German language Wikipedia version of this page (which has moved to here ). It provides supplementary physical, thermodynamic, and vapor pressure data, some of which is redundant with data in the tables above, and some of which is additional.
In the following tables, values are temperature-dependent and to a lesser degree pressure-dependent, and are arranged by state of aggregation (s = solid, lq = liquid, g = gas), which are clearly a function of temperature and pressure. All of the data were computed from data given in "Formulation of the Thermodynamic Properties of Ordinary Water Substance for Scientific and General Use" (IAPWS , 1984) (obsolete as of 1995). [ 22 ] This applies to:
In the following table, material data are given for standard pressure of 0.1 M Pa (equivalent to 1 bar). Up to 99.63 °C (the boiling point of water at 0.1 MPa), at this pressure water exists as a liquid. Above that, it exists as water vapor. Note that the boiling point of 100.0 °C is at a pressure of 0.101325 MPa (1 atm ), which is the average atmospheric pressure.
In the following table, material data are given with a pressure of 611.7 Pa (equivalent to 0.006117 bar). Up to a temperature of 0.01 °C, the triple point of water, water normally exists as ice, except for supercooled water, for which one data point is tabulated here. At the triple point, ice can exist together with both liquid water and vapor. At higher temperatures, the data are for water vapor only.
The following table is based on different, complementary sources and approximation formulas, whose values are of various quality and accuracy. The values in the temperature range of −100 °C to 100 °C were inferred from D. Sunday (1982) and are quite uniform and exact. The values in the temperature range of the boiling point of water up to the critical point (100 °C to 374 °C) are drawn from different sources and are substantially less accurate; hence they should be used only as approximate values. [ 23 ] [ 24 ] [ 25 ] [ 26 ]
To use the values correctly, consider the following points:
The table values for −100 °C to 100 °C were computed by the following formulas, where T is in kelvins and vapor pressures, P w and P i , are in pascals .
Over liquid water
For temperature range: 173.15 K to 373.15 K or equivalently −100 °C to 100 °C
Over ice
For temperature range: 173.15 K to 273.15 K or equivalently −100 °C to 0 °C
At triple point
An important basic value, which is not registered in the table, is the saturated vapor pressure at the triple point of water. The internationally accepted value according to measurements of Guildner, Johnson and Jones (1976) amounts to:
Accepted standardized value of the magnetic susceptibility of water at 20 °C (room temperature) is −12.97 cm 3 /mol. [ 27 ]
Accepted standardized value of the magnetic susceptibility of water at 20 °C (room temperature) is −0.702 cm 3 /g. [ 27 ] | https://en.wikipedia.org/wiki/H2O_(data_page) |
Diphosphene is a compound having the formula (PH) 2 . It exists as two geometric isomers , E and Z . [ 1 ] Diphosphene is also the parent member of the entire class of diphosphene compounds with the formula (PR) 2 , where R is an organyl group. [ 2 ]
Visible radiation induces cis-trans isomerization, [ 3 ] although further irradiation can excite the molecule to a triplet diradical state. In triplet trans-HPPH, the P-P bond length is predicted to be 2.291 Å. It is not only longer than the P-P double bond in ground state trans-bis(2,4,6-tri-tert-butylphenyl)diphosphene, but also longer than that of P-P single bond in H 2 P−PH 2 . Calculation of the dihedral angle of trans-HPPH suggests that it is almost 90 degree, which means the formation of π {\displaystyle \pi } and π ∗ {\displaystyle \pi ^{*}} P-P bonds is forbidden and σ bond is enhanced. [ 4 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2P2 |
Polonium hydride (also known as polonium dihydride , hydrogen polonide , or polane ) is a chemical compound with the formula Po H 2 . It is a liquid at room temperature, the second hydrogen chalcogenide with this property after water . It is very unstable chemically and tends to decompose into elemental polonium and hydrogen . It is a volatile and very labile compound, from which many polonides can be derived. Additionally, it is radioactive. [ 2 ]
Polonium hydride cannot be produced by direct reaction from the elements upon heating. Other unsuccessful routes to synthesis include the reaction of polonium tetrachloride (PoCl 4 ) with lithium aluminium hydride (LiAlH 4 ), which only produces elemental polonium, and the reaction of hydrochloric acid with magnesium polonide (MgPo). The fact that these synthesis routes do not work may be caused by the radiolysis of polonium hydride upon formation. [ 3 ]
Trace quantities of polonium hydride may be prepared by reacting hydrochloric acid with polonium-plated magnesium foil. In addition, the diffusion of trace quantities of polonium in palladium or platinum that is saturated with hydrogen (see palladium hydride ) may be due to the formation and migration of polonium hydride. [ 3 ]
Polonium hydride is a more covalent compound than most metal hydrides because polonium straddles the border between metals and metalloids and has some nonmetallic properties. It is intermediate between a hydrogen halide like hydrogen chloride and a metal hydride like stannane .
It should have properties similar to that of hydrogen selenide and hydrogen telluride , other borderline hydrides . It is expected to be an endothermic compound, like the lighter hydrogen telluride and hydrogen selenide, and therefore would decompose into its constituent elements, releasing heat in the process. The amount of heat given off in the decomposition of polonium hydride is over 100 kJ/mol , the largest of all the hydrogen chalcogenides .
It is predicted that, like the other hydrogen chalcogenides, polonium may form two types of salts : polonide (containing the Po 2− anion ) and one from polonium hydride (containing –PoH, which would be the polonium analogue of thiol , selenol and tellurol ). However, no salts from polonium hydride are known. An example of a polonide is lead polonide (PbPo), which occurs naturally as lead is formed in the alpha decay of polonium. [ 4 ]
Polonium hydride is difficult to work with due to the extreme radioactivity of polonium and its compounds and has only been prepared in very dilute tracer quantities. As a result, its physical properties are not definitely known. [ 3 ] It is also unknown if polonium hydride forms an acidic solution in water like its lighter homologues, or if it behaves more like a metal hydride (see also hydrogen astatide ). | https://en.wikipedia.org/wiki/H2Po |
Chloroplatinic acid (also known as hexachloroplatinic acid ) is an inorganic compound with the formula [H 3 O] 2 [PtCl 6 ](H 2 O) x (0 ≤ x ≤ 6). A red solid, it is an important commercial source of platinum , usually as an aqueous solution . Although often written in shorthand as H 2 PtCl 6 , it is the hydronium (H 3 O + ) salt of the hexachloroplatinate anion ( PtCl 2− 6 ). [ 1 ] [ 2 ] [ 3 ] Hexachloroplatinic acid is highly hygroscopic .
Hexachloroplatinic acid may be produced via a variety of methods. The most common of these methods involves dissolution of platinum in aqua regia . Other methods include exposing an aqueous suspension of platinum particles to chlorine gas, or via electrolysis.
When produced by the aqua regia route, hexachloroplatinic acid is thought to arise by the following equation: [ 4 ] [ 5 ]
The resulting orange/red solution can be evaporated to produce brownish red crystals. Some authors suggest that hexachloroplatinic acid produced using this method is contaminated with nitrosonium hexachloroplatinate. Newer literature indicates that this is not the case, and that once the nitric acid has been driven off, samples prepared via this method contain no detectable nitrogen.
Alternative methods have been investigated and described, often motivated by the avoidance of nitrogen contamination. [ 6 ]
According to X-ray crystallography , hexachloroplatinic acid consists of octahedral PtCl 2− 6 ions linked by hydrogen bonding. The cubic array of these octahedra encase water molecules. [ citation needed ]
When heated, hexachloroplatinic acid decomposes to platinum(IV) chloride . [ 1 ]
Chloroplatinic acid was popularized for the quantitative analysis of potassium. The potassium is selectively precipitated from solution as potassium hexachloroplatinate. Determinations were done in 85% (v/v) alcohol solutions with excess platinate ions, and the precipitated product was weighed. Potassium could be detected for solutions as dilute as 0.02 to 0.2% (m/v). [ 7 ]
This method for determination of potassium was advantageous compared to the sodium cobaltinitrite method used previously, since it required a single precipitation reaction. [ 7 ] Gravimetric analysis of precipitated products has been supplanted by modern instrumental analysis methods such as ion selective electrodes , flame photometry , ICP-AES or ICP-MS .
Upon treatment with an ammonium salt, such as ammonium chloride , chloroplatinic acid converts to ammonium hexachloroplatinate , which precipitates as a solid. [ 4 ] Upon heating in an atmosphere of hydrogen , the ammonium salt converts to elemental platinum. Platinum is often isolated from ores or recycled from residues using this method. [ 8 ]
Like many platinum compounds, chloroplatinic acid is a catalyst (or precatalyst) for hydrogenation and related reactions. As first reported by John Speier and colleagues from Dow Corning , it catalyzes the addition of hydrosilanes to olefins, i.e. hydrosilylation . Early demonstration reactions used isopropanol solutions of trichlorosilane (SiHCl 3 ) with pentenes . Prior work on the addition of silanes to alkenes required radical reactions that were inefficient. [ 9 ] [ 10 ] As well as with Karstedt's catalyst , Speier's catalyst enjoys widespread use for hydrosilylation, the main drawback is the deliquescent properties of the catalyst. [ 11 ]
It is generally agreed that chloroplatinic acid is a precursor to the actual catalyst. A possible role for colloidal platinum or zero-valent complexes has also been considered. [ 12 ]
Chloroplatinic acid prepared from aqua regia is proposed to contain nitrosonium hexachloroplatinate, (NO) 2 PtCl 6 . Nitrosonium hexachloroplatinate is obtained by the reaction of nitrosyl chloride (NOCl) and platinum metal. [ 13 ] Nitrosonium hexachloroplatinate has been found to react vigorously with water and hydrochloric acid, making contamination of chloroplatinic acid prepared with aqua regia with nitrosonium hexachloroplatinate unlikely. [ citation needed ] | https://en.wikipedia.org/wiki/H2PtCl6 |
Hydrogen sulfide is a chemical compound with the formula H 2 S . It is a colorless chalcogen-hydride gas , and is toxic, corrosive, and flammable. Trace amounts in ambient atmosphere have a characteristic foul odor of rotten eggs. [ 11 ] Swedish chemist Carl Wilhelm Scheele is credited with having discovered the chemical composition of purified hydrogen sulfide in 1777. [ 12 ]
Hydrogen sulfide is toxic to humans and most other animals by inhibiting cellular respiration in a manner similar to hydrogen cyanide . When it is inhaled or its salts are ingested in high amounts, damage to organs occurs rapidly with symptoms ranging from breathing difficulties to convulsions and death. [ 13 ] [ 14 ] Despite this, the human body produces small amounts of this sulfide and its mineral salts, and uses it as a signalling molecule . [ 15 ]
Hydrogen sulfide is often produced from the microbial breakdown of organic matter in the absence of oxygen, such as in swamps and sewers; this process is commonly known as anaerobic digestion , which is done by sulfate-reducing microorganisms . It also occurs in volcanic gases , natural gas deposits, and sometimes in well-drawn water.
Hydrogen sulfide is slightly denser than air. A mixture of H 2 S and air can be explosive.
In general, hydrogen sulfide acts as a reducing agent , as indicated by its ability to reduce sulfur dioxide in the Claus process . Hydrogen sulfide burns in oxygen with a blue flame to form sulfur dioxide ( SO 2 ) and water :
If an excess of oxygen is present, sulfur trioxide ( SO 3 ) is formed, which quickly hydrates to sulfuric acid :
It is slightly soluble in water and acts as a weak acid ( p K a = 6.9 in 0.01–0.1 mol/litre solutions at 18 °C), giving the hydrosulfide ion HS − . Hydrogen sulfide and its solutions are colorless. When exposed to air, it slowly oxidizes to form elemental sulfur, which is not soluble in water. The sulfide anion S 2− is not formed in aqueous solution. [ 16 ]
H 2 S and H 2 O exchange protons rapidly. This behavior is the basis of technologies for the purification of deuterium oxide ("heavy water" or D 2 O ), which exploits the easy distillation of these compounds. [ 17 ]
At pressures above 90 GPa ( gigapascal ), hydrogen sulfide becomes a metallic conductor of electricity. When cooled below a critical temperature this high-pressure phase exhibits superconductivity . The critical temperature increases with pressure, ranging from 23 K at 100 GPa to 150 K at 200 GPa. [ 18 ] If hydrogen sulfide is pressurized at higher temperatures, then cooled, the critical temperature reaches 203 K (−70 °C), which was the highest accepted superconducting critical temperature until the discovery of Lanthanum decahydride in 2019. By substituting a small part of sulfur with phosphorus and using even higher pressures, it has been predicted that it may be possible to raise the critical temperature to above 0 °C (273 K) and achieve room-temperature superconductivity . [ 19 ]
Hydrogen sulfide decomposes without a presence of a catalyst under atmospheric pressure around 1200 °C into hydrogen and sulfur. [ 20 ]
Hydrogen sulfide reacts with metal ions to form metal sulfides, which are insoluble, often dark colored solids. This behavior is the basis of the use of hydrogen sulfide as a reagent in the qualitative inorganic analysis of metal ions. In these analyses, heavy metal (and nonmetal ) ions (e.g., Pb(II), Cu(II), Hg(II), As(III)) are precipitated from solution upon exposure to H 2 S . The components of the resulting solid are then identified by their reactivity. Lead(II) acetate paper is used to detect hydrogen sulfide because it readily converts to lead(II) sulfide , which is black. [ 21 ] [ 22 ]
Hydrogen sulfide is also responsible for tarnishing on various metals including copper and silver ; the chemical responsible for black toning found on silver coins is silver sulfide ( Ag 2 S ), which is produced when the silver on the surface of the coin reacts with atmospheric hydrogen sulfide. [ 23 ] Coins that have been subject to toning by hydrogen sulfide and other sulfur-containing compounds may have the toning add to the numismatic value of a coin based on aesthetics, as the toning may produce thin-film interference , resulting in the coin taking on an attractive coloration. [ 24 ] Coins can also be intentionally treated with hydrogen sulfide to induce toning, though artificial toning can be distinguished from natural toning, and is generally criticised among collectors. [ 25 ]
Hydrogen sulfide is most commonly obtained by its separation from sour gas , which is natural gas with a high content of H 2 S . It can also be produced by treating hydrogen with molten elemental sulfur at about 450 °C. Hydrocarbons can serve as a source of hydrogen in this process. [ 26 ]
The very favorable thermodynamics for the hydrogenation of sulfur implies that the dehydrogenation (or cracking ) of hydrogen sulfide would require very high temperatures. [ 27 ]
A standard lab preparation is to treat ferrous sulfide with a strong acid in a Kipp generator :
For use in qualitative inorganic analysis , thioacetamide is used to generate H 2 S :
Many metal and nonmetal sulfides, e.g. aluminium sulfide , phosphorus pentasulfide , silicon disulfide liberate hydrogen sulfide upon exposure to water: [ 28 ]
This gas is also produced by heating sulfur with solid organic compounds and by reducing sulfurated organic compounds with hydrogen.
It can also be produced by mixing ammonium thiocyanate to concentrated sulphuric acid and adding water to it.
Hydrogen sulfide can be generated in cells via enzymatic or non-enzymatic pathways. Three enzymes catalyze formation of H 2 S : cystathionine γ-lyase (CSE), cystathionine β-synthetase (CBS), and 3-mercaptopyruvate sulfurtransferase (3-MST). [ 29 ] CBS and CSE are the main proponents of H 2 S biogenesis, which follows the trans-sulfuration pathway. [ 30 ] These enzymes have been identified in a breadth of biological cells and tissues, and their activity is induced by a number of disease states. [ 31 ] These enzymes are characterized by the transfer of a sulfur atom from methionine to serine to form a cysteine molecule. [ 30 ] 3-MST also contributes to hydrogen sulfide production by way of the cysteine catabolic pathway. [ 31 ] [ 30 ] Dietary amino acids, such as methionine and cysteine serve as the primary substrates for the transulfuration pathways and in the production of hydrogen sulfide. Hydrogen sulfide can also be derived from proteins such as ferredoxins and Rieske proteins . [ 31 ]
Sulfate-reducing (resp. sulfur-reducing ) bacteria generate usable energy under low-oxygen conditions by using sulfates (resp. elemental sulfur) to oxidize organic compounds or hydrogen; this produces hydrogen sulfide as a waste product. [ 32 ]
H 2 S in the body acts as a gaseous signaling molecule with implications for health and in diseases. [ 29 ] [ 33 ] [ 34 ] [ 35 ]
Hydrogen sulfide is involved in vasodilation in animals, as well as in increasing seed germination and stress responses in plants. [ 36 ] Hydrogen sulfide signaling is moderated by reactive oxygen species (ROS) and reactive nitrogen species (RNS). [ 36 ] H 2 S has been shown to interact with the NO pathway resulting in several different cellular effects, including the inhibition of cGMP phosphodiesterases, [ 37 ] as well as the formation of another signal called nitrosothiol. [ 36 ] Hydrogen sulfide is also known to increase the levels of glutathione, which acts to reduce or disrupt ROS levels in cells. [ 36 ]
The field of H 2 S biology has advanced from environmental toxicology to investigate the roles of endogenously produced H 2 S in physiological conditions and in various pathophysiological states. [ 38 ] H 2 S has been implicated in cancer, in Down syndrome and in vascular disease. [ 39 ] [ 40 ] [ 41 ] [ 42 ]
At lower concentrations, it stimulates mitochondrial function via multiple mechanisms including direct electron donation. [ 43 ] [ 44 ] However, at higher concentrations, it inhibits Complex IV of the mitochondrial electron transport chain, which effectively reduces ATP generation and biochemical activity within cells. [ 36 ]
Hydrogen sulfide is mainly consumed as a precursor to elemental sulfur. This conversion, called the Claus process , involves partial oxidation to sulfur dioxide. The latter reacts with hydrogen sulfide to give elemental sulfur. The conversion is catalyzed by alumina. [ 45 ]
Many fundamental organosulfur compounds are produced using hydrogen sulfide. These include methanethiol , ethanethiol , and thioglycolic acid . [ 26 ] Hydrosulfides can be used in the production of thiophenol . [ 46 ]
Upon combining with alkali metal bases, hydrogen sulfide converts to alkali hydrosulfides such as sodium hydrosulfide and sodium sulfide :
Sodium sulfides are used in the paper making industry. Specifically, salts of SH − break bonds between lignin and cellulose components of pulp in the Kraft process . [ 26 ]
As indicated above, many metal ions react with hydrogen sulfide to give the corresponding metal sulfides. Oxidic ores are sometimes treated with hydrogen sulfide to give the corresponding metal sulfides which are more readily purified by flotation .Metal parts are sometimes passivated with hydrogen sulfide. Catalysts used in hydrodesulfurization are routinely activated with hydrogen sulfide. [ 26 ]
Volcanoes and some hot springs (as well as cold springs ) emit some H 2 S . Hydrogen sulfide can be present naturally in well water, often as a result of the action of sulfate-reducing bacteria . [ 47 ] [ better source needed ] Hydrogen sulfide is produced by the human body in small quantities through bacterial breakdown of proteins containing sulfur in the intestinal tract; it therefore contributes to the characteristic odor of flatulence. It is also produced in the mouth ( halitosis ). [ 48 ]
A portion of global H 2 S emissions are due to human activity. By far the largest industrial source of H 2 S is petroleum refineries : The hydrodesulfurization process liberates sulfur from petroleum by the action of hydrogen. The resulting H 2 S is converted to elemental sulfur by partial combustion via the Claus process , which is a major source of elemental sulfur. Other anthropogenic sources of hydrogen sulfide include coke ovens, paper mills (using the Kraft process), tanneries and sewerage . H 2 S arises from virtually anywhere where elemental sulfur comes in contact with organic material, especially at high temperatures. Depending on environmental conditions, it is responsible for deterioration of material through the action of some sulfur oxidizing microorganisms. It is called biogenic sulfide corrosion . [ citation needed ]
In 2011 it was reported that increased concentrations of H 2 S were observed in the Bakken formation crude, possibly due to oil field practices, and presented challenges such as "health and environmental risks, corrosion of wellbore, added expense with regard to materials handling and pipeline equipment, and additional refinement requirements". [ 49 ]
Besides living near gas and oil drilling operations, ordinary citizens can be exposed to hydrogen sulfide by being near waste water treatment facilities, landfills and farms with manure storage. Exposure occurs through breathing contaminated air or drinking contaminated water. [ 50 ]
In municipal waste landfill sites , the burial of organic material rapidly leads to the production of anaerobic digestion within the waste mass and, with the humid atmosphere and relatively high temperature that accompanies biodegradation , biogas is produced as soon as the air within the waste mass has been reduced. If there is a source of sulfate bearing material, such as plasterboard or natural gypsum (calcium sulfate dihydrate), under anaerobic conditions sulfate reducing bacteria converts this to hydrogen sulfide. These bacteria cannot survive in air but the moist, warm, anaerobic conditions of buried waste that contains a high source of carbon – in inert landfills, paper and glue used in the fabrication of products such as plasterboard can provide a rich source of carbon [ 51 ] – is an excellent environment for the formation of hydrogen sulfide.
In industrial anaerobic digestion processes, such as waste water treatment or the digestion of organic waste from agriculture , hydrogen sulfide can be formed from the reduction of sulfate and the degradation of amino acids and proteins within organic compounds. [ 52 ] Sulfates are relatively non-inhibitory to methane forming bacteria but can be reduced to H 2 S by sulfate reducing bacteria , of which there are several genera. [ 53 ]
A number of processes have been designed to remove hydrogen sulfide from drinking water . [ 54 ]
Hydrogen sulfide is commonly found in raw natural gas and biogas. It is typically removed by amine gas treating technologies. In such processes, the hydrogen sulfide is first converted to an ammonium salt, whereas the natural gas is unaffected. [ 57 ] [ 58 ]
The bisulfide anion is subsequently regenerated by heating of the amine sulfide solution. Hydrogen sulfide generated in this process is typically converted to elemental sulfur using the Claus Process .
The underground mine gas term for foul-smelling hydrogen sulfide-rich gas mixtures is stinkdamp . Hydrogen sulfide is a highly toxic and flammable gas ( flammable range : 4.3–46%). It can poison several systems in the body, although the nervous system is most affected. [ citation needed ] The toxicity of H 2 S is comparable with that of carbon monoxide . [ 59 ] It binds with iron in the mitochondrial cytochrome enzymes , thus preventing cellular respiration . Its toxic properties were described in detail in 1843 by Justus von Liebig . [ 60 ]
Even before hydrogen sulfide was discovered, Italian physician Bernardino Ramazzini hypothesized in his 1713 book De Morbis Artificum Diatriba that occupational diseases of sewer-workers and blackening of coins in their clothes may be caused by an unknown invisible volatile acid (moreover, in late 18th century toxic gas emanation from Paris sewers became a problem for the citizens and authorities). [ 61 ]
Although very pungent at first (it smells like rotten eggs [ 62 ] ), it quickly deadens the sense of smell, creating temporary anosmia , [ 63 ] so victims may be unaware of its presence until it is too late. Safe handling procedures are provided by its safety data sheet (SDS) . [ 64 ]
Since hydrogen sulfide occurs naturally in the body, the environment, and the gut, enzymes exist to metabolize it. At some threshold level, believed to average around 300–350 ppm, the oxidative enzymes become overwhelmed. Many personal safety gas detectors, such as those used by utility, sewage and petrochemical workers, are set to alarm at as low as 5 to 10 ppm and to go into high alarm at 15 ppm. Metabolism causes oxidation to sulfate, which is harmless. [ 65 ] Hence, low levels of hydrogen sulfide may be tolerated indefinitely. [ citation needed ]
Exposure to lower concentrations can result in eye irritation, a sore throat and cough , nausea, shortness of breath, and fluid in the lungs . [ 59 ] These effects are believed to be due to hydrogen sulfide combining with alkali present in moist surface tissues to form sodium sulfide , a caustic . [ 66 ] These symptoms usually subside in a few weeks.
Long-term, low-level exposure may result in fatigue , loss of appetite, headaches , irritability, poor memory, and dizziness . Chronic exposure to low level H 2 S (around 2 ppm ) has been implicated in increased miscarriage and reproductive health issues among Russian and Finnish wood pulp workers, [ 67 ] but the reports have not (as of 1995) been replicated.
Short-term, high-level exposure can induce immediate collapse, with loss of breathing and a high probability of death. If death does not occur, high exposure to hydrogen sulfide can lead to cortical pseudolaminar necrosis , degeneration of the basal ganglia and cerebral edema . [ 59 ] Although respiratory paralysis may be immediate, it can also be delayed up to 72 hours. [ 68 ]
Inhalation of H 2 S resulted in about 7 workplace deaths per year in the U.S. (2011–2017 data), second only to carbon monoxide (17 deaths per year) for workplace chemical inhalation deaths. [ 69 ]
Treatment involves immediate inhalation of amyl nitrite , injections of sodium nitrite , or administration of 4-dimethylaminophenol in combination with inhalation of pure oxygen, administration of bronchodilators to overcome eventual bronchospasm , and in some cases hyperbaric oxygen therapy (HBOT). [ 59 ] HBOT has clinical and anecdotal support. [ 74 ] [ 75 ] [ 76 ]
Hydrogen sulfide was used by the British Army as a chemical weapon during World War I . It was not considered to be an ideal war gas, partially due to its flammability and because the distinctive smell could be detected from even a small leak, alerting the enemy to the presence of the gas. It was nevertheless used on two occasions in 1916 when other gases were in short supply. [ 77 ]
On September 2, 2005, a leak in the propeller room of a Royal Caribbean Cruise Liner docked in Los Angeles resulted in the deaths of 3 crewmen due to a sewage line leak. As a result, all such compartments are now required to have a ventilation system. [ 78 ] [ 79 ]
A dump of toxic waste containing hydrogen sulfide is believed to have caused 17 deaths and thousands of illnesses in Abidjan , on the West African coast, in the 2006 Côte d'Ivoire toxic waste dump . [ 80 ]
In September 2008, three workers were killed and two suffered serious injury, including long term brain damage, at a mushroom growing company in Langley , British Columbia . A valve to a pipe that carried chicken manure , straw and gypsum to the compost fuel for the mushroom growing operation became clogged, and as workers unclogged the valve in a confined space without proper ventilation the hydrogen sulfide that had built up due to anaerobic decomposition of the material was released, poisoning the workers in the surrounding area. [ 81 ] An investigator said there could have been more fatalities if the pipe had been fully cleared and/or if the wind had changed directions. [ 82 ]
In 2014, levels of hydrogen sulfide as high as 83 ppm were detected at a recently built mall in Thailand called Siam Square One at the Siam Square area. Shop tenants at the mall reported health complications such as sinus inflammation, breathing difficulties and eye irritation. After investigation it was determined that the large amount of gas originated from imperfect treatment and disposal of waste water in the building. [ 83 ]
In 2014, hydrogen sulfide gas killed workers at the Promenade shopping center in North Scottsdale, Arizona , USA [ 84 ] after climbing into 15 ft deep chamber without wearing personal protective gear . "Arriving crews recorded high levels of hydrogen cyanide and hydrogen sulfide coming out of the sewer."
In November 2014, a substantial amount of hydrogen sulfide gas shrouded the central, eastern and southeastern parts of Moscow . Residents living in the area were urged to stay indoors by the emergencies ministry. Although the exact source of the gas was not known, blame had been placed on a Moscow oil refinery. [ 85 ]
In June 2016, a mother and her daughter were found dead in their still-running 2006 Porsche Cayenne SUV against a guardrail on Florida's Turnpike , initially thought to be victims of carbon monoxide poisoning . [ 86 ] [ 87 ] Their deaths remained unexplained as the medical examiner waited for results of toxicology tests on the victims, [ 88 ] until urine tests revealed that hydrogen sulfide was the cause of death. A report from the Orange-Osceola Medical Examiner's Office indicated that toxic fumes came from the Porsche's starter battery , located under the front passenger seat. [ 89 ] [ 90 ]
In January 2017, three utility workers in Key Largo, Florida , died one by one within seconds of descending into a narrow space beneath a manhole cover to check a section of paved street. [ 91 ] In an attempt to save the men, a firefighter who entered the hole without his air tank (because he could not fit through the hole with it) collapsed within seconds and had to be rescued by a colleague. [ 92 ] The firefighter was airlifted to Jackson Memorial Hospital and later recovered. [ 93 ] [ 94 ] A Monroe County Sheriff officer initially determined that the space contained hydrogen sulfide and methane gas produced by decomposing vegetation. [ 95 ]
On May 24, 2018, two workers were killed, another seriously injured, and 14 others hospitalized by hydrogen sulfide inhalation at a Norske Skog paper mill in Albury, New South Wales . [ 96 ] [ 97 ] An investigation by SafeWork NSW found that the gas was released from a tank used to hold process water . The workers were exposed at the end of a 3-day maintenance period. Hydrogen sulfide had built up in an upstream tank, which had been left stagnant and untreated with biocide during the maintenance period. These conditions allowed sulfate-reducing bacteria to grow in the upstream tank, as the water contained small quantities of wood pulp and fiber . The high rate of pumping from this tank into the tank involved in the incident caused hydrogen sulfide gas to escape from various openings around its top when pumping was resumed at the end of the maintenance period. The area above it was sufficiently enclosed for the gas to pool there, despite not being identified as a confined space by Norske Skog. One of the workers who was killed was exposed while investigating an apparent fluid leak in the tank, while the other who was killed and the worker who was badly injured were attempting to rescue the first after he collapsed on top of it. In a resulting criminal case , Norske Skog was accused of failing to ensure the health and safety of its workforce at the plant to a reasonably practicable extent. It pleaded guilty, and was fined AU$1,012,500 and ordered to fund the production of an anonymized educational video about the incident. [ 98 ] [ 99 ] [ 96 ] [ 100 ]
In October 2019, an Odessa, Texas employee of Aghorn Operating Inc. and his wife were killed due to a water pump failure. Produced water with a high concentration of hydrogen sulfide was released by the pump. The worker died while responding to an automated phone call he had received alerting him to a mechanical failure in the pump, while his wife died after driving to the facility to check on him. [ 101 ] A CSB investigation cited lax safety practices at the facility, such as an informal lockout-tagout procedure and a nonfunctioning hydrogen sulfide alert system. [ 102 ]
The gas, produced by mixing certain household ingredients, was used in a suicide wave in 2008 in Japan. [ 103 ] The wave prompted staff at Tokyo's suicide prevention center to set up a special hotline during " Golden Week ", as they received an increase in calls from people wanting to kill themselves during the annual May holiday. [ 104 ]
As of 2010, this phenomenon has occurred in a number of US cities, prompting warnings to those arriving at the site of the suicide. [ 105 ] [ 106 ] [ 107 ] [ 108 ] [ 109 ]
In 2020, H 2 S ingestion was used as a suicide method by Japanese pro wrestler Hana Kimura . [ 110 ]
In 2024, Lucy-Bleu Knight, stepdaughter of famed musician Slash , also used H 2 S ingestion to commit suicide. [ 111 ]
Hydrogen sulfide is a central participant in the sulfur cycle , the biogeochemical cycle of sulfur on Earth. [ 112 ]
In the absence of oxygen , sulfur-reducing and sulfate-reducing bacteria derive energy from oxidizing hydrogen or organic molecules by reducing elemental sulfur or sulfate to hydrogen sulfide. Other bacteria liberate hydrogen sulfide from sulfur-containing amino acids ; this gives rise to the odor of rotten eggs and contributes to the odor of flatulence .
As organic matter decays under low-oxygen (or hypoxic ) conditions (such as in swamps, eutrophic lakes or dead zones of oceans), sulfate-reducing bacteria will use the sulfates present in the water to oxidize the organic matter, producing hydrogen sulfide as waste. Some of the hydrogen sulfide will react with metal ions in the water to produce metal sulfides, which are not water-soluble. These metal sulfides, such as ferrous sulfide FeS, are often black or brown, leading to the dark color of sludge .
Several groups of bacteria can use hydrogen sulfide as fuel, oxidizing it to elemental sulfur or to sulfate by using dissolved oxygen, metal oxides (e.g., iron oxyhydroxides and manganese oxides ), or nitrate as electron acceptors. [ 113 ]
The purple sulfur bacteria and the green sulfur bacteria use hydrogen sulfide as an electron donor in photosynthesis , thereby producing elemental sulfur. This mode of photosynthesis is older than the mode of cyanobacteria , algae , and plants , which uses water as electron donor and liberates oxygen.
The biochemistry of hydrogen sulfide is a key part of the chemistry of the iron-sulfur world . In this model of the origin of life on Earth, geologically produced hydrogen sulfide is postulated as an electron donor driving the reduction of carbon dioxide. [ 114 ]
Hydrogen sulfide is lethal to most animals, but a few highly specialized species ( extremophiles ) do thrive in habitats that are rich in this compound. [ 115 ]
In the deep sea, hydrothermal vents and cold seeps with high levels of hydrogen sulfide are home to a number of extremely specialized lifeforms, ranging from bacteria to fish. [ which? ] [ 116 ] Because of the absence of sunlight at these depths, these ecosystems rely on chemosynthesis rather than photosynthesis . [ 117 ]
Freshwater springs rich in hydrogen sulfide are mainly home to invertebrates, but also include a small number of fish: Cyprinodon bobmilleri (a pupfish from Mexico), Limia sulphurophila (a poeciliid from the Dominican Republic ), Gambusia eurystoma (a poeciliid from Mexico), and a few Poecilia (poeciliids from Mexico). [ 115 ] [ 118 ] Invertebrates and microorganisms in some cave systems, such as Movile Cave , are adapted to high levels of hydrogen sulfide. [ 119 ]
Hydrogen sulfide has often been detected in the interstellar medium. [ 120 ] It also occurs in the clouds of planets in our solar system. [ 121 ] [ 122 ]
Hydrogen sulfide has been implicated in several mass extinctions that have occurred in the Earth's past. In particular, a buildup of hydrogen sulfide in the atmosphere may have caused, or at least contributed to, the Permian-Triassic extinction event 252 million years ago. [ 123 ] [ 124 ] [ 125 ]
Organic residues from these extinction boundaries indicate that the oceans were anoxic (oxygen-depleted) and had species of shallow plankton that metabolized H 2 S . The formation of H 2 S may have been initiated by massive volcanic eruptions, which emitted carbon dioxide and methane into the atmosphere, which warmed the oceans, lowering their capacity to absorb oxygen that would otherwise oxidize H 2 S . The increased levels of hydrogen sulfide could have killed oxygen-generating plants as well as depleted the ozone layer, causing further stress. Small H 2 S blooms have been detected in modern times in the Dead Sea and in the Atlantic Ocean off the coast of Namibia . [ 123 ] | https://en.wikipedia.org/wiki/H2S |
Sulfuric acid ( American spelling and the preferred IUPAC name ) or sulphuric acid ( Commonwealth spelling ), known in antiquity as oil of vitriol , is a mineral acid composed of the elements sulfur , oxygen , and hydrogen , with the molecular formula H 2 SO 4 . It is a colorless, odorless, and viscous liquid that is miscible with water. [ 7 ]
Pure sulfuric acid does not occur naturally due to its strong affinity to water vapor ; it is hygroscopic and readily absorbs water vapor from the air . [ 7 ] Concentrated sulfuric acid is a strong oxidant with powerful dehydrating properties, making it highly corrosive towards other materials, from rocks to metals. Phosphorus pentoxide is a notable exception in that it is not dehydrated by sulfuric acid but, to the contrary, dehydrates sulfuric acid to sulfur trioxide . Upon addition of sulfuric acid to water, a considerable amount of heat is released; thus, the reverse procedure of adding water to the acid is generally avoided since the heat released may boil the solution, spraying droplets of hot acid during the process. Upon contact with body tissue, sulfuric acid can cause severe acidic chemical burns and secondary thermal burns due to dehydration. [ 8 ] [ 9 ] Dilute sulfuric acid is substantially less hazardous without the oxidative and dehydrating properties; though, it is handled with care for its acidity.
Many methods for its production are known, including the contact process , the wet sulfuric acid process , and the lead chamber process . [ 10 ] Sulfuric acid is also a key substance in the chemical industry . It is most commonly used in fertilizer manufacture [ 11 ] but is also important in mineral processing , oil refining , wastewater treating , and chemical synthesis . It has a wide range of end applications, including in domestic acidic drain cleaners , [ 12 ] as an electrolyte in lead-acid batteries , as a dehydrating compound, and in various cleaning agents .
Sulfuric acid can be obtained by dissolving sulfur trioxide in water.
Although nearly 100% sulfuric acid solutions can be made, the subsequent loss of SO 3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade, which is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are: [ 13 ] [ 14 ]
"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process , chamber acid being the acid produced in the lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid ) and tower acid being the acid recovered from the bottom of the Glover tower. [ 13 ] [ 14 ] They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10 M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations ), is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. [ 14 ]
Sulfuric acid contains not only H 2 SO 4 molecules, but is actually an equilibrium of many other chemical species, as it is shown in the table below.
Sulfuric acid is a colorless oily liquid, and has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, [ 16 ] and 98% sulfuric acid has a vapor pressure of <1 mmHg at 40 °C. [ 17 ]
In the solid state, sulfuric acid is a molecular solid that forms monoclinic crystals with nearly trigonal lattice parameters. The structure consists of layers parallel to the (010) plane, in which each molecule is connected by hydrogen bonds to two others. [ 3 ] Hydrates H 2 SO 4 · n H 2 O are known for n = 1, 2, 3, 4, 6.5, and 8, although most intermediate hydrates are stable against disproportionation . [ 18 ]
Anhydrous H 2 SO 4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity , a consequence of autoprotolysis , i.e. self- protonation : [ 15 ]
The equilibrium constant for autoprotolysis (25 °C) is: [ 15 ]
The corresponding equilibrium constant for water , K w is 10 −14 , a factor of 10 10 (10 billion) smaller.
In spite of the viscosity of the acid, the effective conductivities of the H 3 SO + 4 and HSO − 4 ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.
The hydration reaction of sulfuric acid is highly exothermic . [ 19 ]
As indicated by its acid dissociation constant , sulfuric acid is a strong acid:
The product of this ionization is HSO − 4 , the bisulfate anion. Bisulfate is a far weaker acid:
The product of this second dissociation is SO 2− 4 , the sulfate anion.
Concentrated sulfuric acid has a powerful dehydrating property, removing water ( H 2 O ) from other chemical compounds such as table sugar ( sucrose ) and other carbohydrates , to produce carbon , steam , and heat. Dehydration of table sugar (sucrose) is a common laboratory demonstration. [ 21 ] The sugar darkens as carbon is formed, and a rigid column of black, porous carbon called a carbon snake may emerge. [ 22 ]
Similarly, mixing starch into concentrated sulfuric acid gives elemental carbon and water. The effect of this can also be seen when concentrated sulfuric acid is spilled on paper. Paper is composed of cellulose , a polysaccharide related to starch. The cellulose reacts to give a burnt appearance in which the carbon appears much like soot that results from fire.
Although less dramatic, the action of the acid on cotton , even in diluted form, destroys the fabric.
The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystals change into white powder as water is removed.
Sulfuric acid reacts with most bases to give the corresponding sulfate or bisulfate.
Aluminium sulfate , also known as paper maker's alum, is made by treating bauxite with sulfuric acid:
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate , for example, displaces acetic acid , CH 3 COOH , and forms sodium bisulfate :
Similarly, treating potassium nitrate with sulfuric acid produces nitric acid . Sulfuric acid reacts with sodium chloride , and gives hydrogen chloride gas and sodium bisulfate :
When combined with nitric acid , sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO + 2 , which is important in nitration reactions involving electrophilic aromatic substitution . This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols.
When allowed to react with superacids , sulfuric acid can act as a base and can be protonated, forming the [H 3 SO 4 ] + ion. Salts of [H 3 SO 4 ] + have been prepared (e.g. trihydroxyoxosulfonium hexafluoroantimonate(V) [H 3 SO 4 ] + [SbF 6 ] − ) using the following reaction in liquid HF :
The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply fluoroantimonic acid , however, has met with failure, as pure sulfuric acid undergoes self-ionization to give [H 3 O] + ions:
which prevents the conversion of H 2 SO 4 to [H 3 SO 4 ] + by the HF/ SbF 5 system. [ 23 ]
Even diluted sulfuric acid reacts with many metals via a single displacement reaction, like other typical acids , producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series ) such as iron , aluminium , zinc , manganese , magnesium , and nickel .
Concentrated sulfuric acid can serve as an oxidizing agent , releasing sulfur dioxide: [ 8 ]
Lead and tungsten , however, are resistant to sulfuric acid.
Hot concentrated sulfuric acid oxidizes carbon [ 24 ] (as bituminous coal ) and sulfur :
Benzene and many derivatives undergo electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids : [ 25 ]
Sulfuric acid can be used to produce hydrogen from water :
The compounds of sulfur and iodine are recovered and reused, hence the process is called the sulfur–iodine cycle . This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur–iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy . It is an alternative to electrolysis , and does not require hydrocarbons like current methods of steam reforming . But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it. [ 26 ] [ 27 ]
Sulfuric acid is rarely encountered naturally on Earth in anhydrous form, due to its great affinity for water . Dilute sulfuric acid is a constituent of acid rain , which is formed by atmospheric oxidation of sulfur dioxide in the presence of water —i.e. oxidation of sulfurous acid . When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water).
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as pyrite :
The resulting highly acidic water is called acid mine drainage (AMD) or acid rock drainage (ARD).
The Fe 2+ can be further oxidized to Fe 3+ :
The Fe 3+ produced can be precipitated as the hydroxide or hydrous iron oxide :
The iron(III) ion (" ferric iron ") can also oxidize pyrite:
FeS 2 (s) + 14 Fe 3+ + 8 H 2 O → 15 Fe 2+ + 2 SO 2− 4 + 16 H +
When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process.
ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.
Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales ) concentrates sulfuric acid in cell vacuoles. [ 28 ]
In the stratosphere , the atmosphere's second layer that is generally between 10–50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical : [ 29 ]
Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer. [ 29 ]
The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. [ 30 ] Sulfuric acid ice has been detected on Jupiter 's moon Europa , where it forms when sulfur ions from Jupiter's magnetosphere implant into the icy surface. [ 31 ]
Sulfuric acid is produced from sulfur , oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).
In the first step, sulfur is burned to produce sulfur dioxide.
The sulfur dioxide is oxidized to sulfur trioxide by oxygen in the presence of a vanadium(V) oxide catalyst . This reaction is reversible and the formation of the sulfur trioxide is exothermic.
The sulfur trioxide is absorbed into 97–98% H 2 SO 4 to form oleum ( H 2 S 2 O 7 ), also known as fuming sulfuric acid or pyrosulphuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.
Directly dissolving SO 3 in water, called the " wet sulfuric acid process ", is rarely practiced because the reaction is extremely exothermic, resulting in a hot aerosol of sulfuric acid that requires condensation and separation.
In the first step, sulfur is burned to produce sulfur dioxide:
or, alternatively, hydrogen sulfide ( H 2 S ) gas is incinerated to SO 2 gas:
The sulfur dioxide then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst .
The sulfur trioxide is hydrated into sulfuric acid H 2 SO 4 :
The last step is the condensation of the sulfuric acid to liquid 97–98% H 2 SO 4 :
Burning sulfur together with saltpeter ( potassium nitrate , KNO 3 ), in the presence of steam, has been used historically. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid.
Prior to 1900, most sulfuric acid was manufactured by the lead chamber process . [ 32 ] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.
A wide variety of laboratory syntheses are known, and typically begin from sulfur dioxide or an equivalent salt . In the metabisulfite method, hydrochloric acid reacts with metabisulfite to produce sulfur dioxide vapors. The gas is bubbled through nitric acid , which will release brown/red vapors of nitrogen dioxide as the reaction proceeds. The completion of the reaction is indicated by the ceasing of the fumes. This method conveniently does not produce an inseparable mist. [ citation needed ]
Alternatively, dissolving sulfur dioxide in an aqueous solution of an oxidizing metal salt such as copper(II) or iron(III) chloride: [ citation needed ]
Two less well-known laboratory methods of producing sulfuric acid, albeit in dilute form and requiring some extra effort in purification, rely on electrolysis . A solution of copper(II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper at cathode and oxygen gas at the anode. The solution of dilute sulfuric acid indicates completion of the reaction when it turns from blue to clear (production of hydrogen at cathode is another sign): [ citation needed ]
More costly, dangerous, and troublesome is the electrobromine method, which employs a mixture of sulfur , water, and hydrobromic acid as the electrolyte. The sulfur is pushed to bottom of container under the acid solution. Then the copper cathode and platinum/graphite anode are used with the cathode near the surface and the anode is positioned at the bottom of the electrolyte to apply the current. This may take longer and emits toxic bromine /sulfur-bromide vapors, but the reactant acid is recyclable. Overall, only the sulfur and water are converted to sulfuric acid and hydrogen (omitting losses of acid as vapors): [ citation needed ]
Sulfuric acid is a very important commodity chemical, and a nation's sulfuric acid production was as recently as 2002 believed to be a good indicator of its industrial strength. [ 33 ] World production in the year 2004 was about 180 million tonnes , with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%. [ 34 ] World production in 2022 was estimated at about 260 million tonnes. [ 35 ]
As of the late 20th century, most of the produced amount (≈60%) was consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze , as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, and water treatment. About 6% of uses are related to pigments and include paints, enamels , printing inks, coated fabrics and paper, while the rest is dispersed into a multitude of applications such as production of explosives, cellophane , acetate and viscose textiles, lubricants, non-ferrous metals , and batteries. [ 36 ]
The dominant use for sulfuric acid is in the "wet method" for the production of phosphoric acid , used for manufacture of phosphate fertilizers . In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite , though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate , hydrogen fluoride (HF) and phosphoric acid . The HF is removed as hydrofluoric acid . The overall process can be represented as:
Ammonium sulfate , an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.
Sulfuric acid is also important in the manufacture of dyestuffs solutions.
Sulfuric acid is used in steelmaking and other metallurgical industries as a pickling agent for removal of rust and fouling . [ 37 ] Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid [ clarification needed ] with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide ( SO 2 ) and sulfur trioxide ( SO 3 ) which are then used to manufacture "new" sulfuric acid.
Hydrogen peroxide ( H 2 O 2 ) can be added to sulfuric acid to produce piranha solution , a powerful but potentially hazardous cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used in the microelectronics industry, and also in laboratory settings to clean glassware.
Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam , used for making nylon . It is used for making hydrochloric acid from salt via the Mannheim process . Much H 2 SO 4 is used in petroleum refining , for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane , a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidizing agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.
Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator):
At anode :
At cathode :
Overall:
Sulfuric acid at high concentrations is frequently the major ingredient in domestic acidic drain cleaners [ 12 ] which are used to remove lipids , hair , tissue paper , etc. Similar to their alkaline versions , such drain openers can dissolve fats and proteins via hydrolysis . Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.
The study of vitriols (hydrated sulfates of various metals forming glassy minerals from which sulfuric acid can be derived) began in ancient times . Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis , in the treatise Phisica et Mystica , and the Leyden papyrus X . [ 38 ] Medieval Islamic alchemists like the authors writing under the name of Jabir ibn Hayyan (died c. 806 – c. 816 AD, known in Latin as Geber), Abu Bakr al-Razi (865 – 925 AD, known in Latin as Rhazes), Ibn Sina (980 – 1037 AD, known in Latin as Avicenna), and Muhammad ibn Ibrahim al-Watwat (1234 – 1318 AD) included vitriol in their mineral classification lists. [ 39 ]
The Jabirian authors and al-Razi experimented extensively with the distillation of various substances, including vitriols. [ 40 ] In one recipe recorded in his Kitāb al-Asrār ( 'Book of Secrets' ), al-Razi may have created sulfuric acid without being aware of it: [ 41 ]
Take white (Yemeni) alum , dissolve it and purify it by filtration. Then distil (green?) vitriol with copper-green (the acetate), and mix (the distillate) with the filtered solution of the purified alum, afterwards let it solidify (or crystallise) in the glass beaker. You will get the best qalqadis (white alum) that may be had. [ 42 ]
In an anonymous Latin work variously attributed to Aristotle (under the title Liber Aristotilis , 'Book of Aristotle'), [ 43 ] to al-Razi (under the title Lumen luminum magnum , 'Great Light of Lights'), or to Ibn Sina, [ 44 ] the author speaks of an 'oil' ( oleum ) obtained through the distillation of iron(II) sulfate (green vitriol), which was likely 'oil of vitriol' or sulfuric acid. [ 45 ] The work refers multiple times to Jabir ibn Hayyan's Seventy Books ( Liber de septuaginta ), one of the few Arabic Jabir works that were translated into Latin. [ 46 ] The author of the version attributed to al-Razi also refers to the Liber de septuaginta as his own work, showing that he erroneously believed the Liber de septuaginta to be a work by al-Razi. [ 47 ] There are several indications that the anonymous work was an original composition in Latin, [ 48 ] although according to one manuscript it was translated by a certain Raymond of Marseilles, meaning that it may also have been a translation from the Arabic. [ 49 ]
According to Ahmad Y. al-Hassan , three recipes for sulfuric acid occur in an anonymous Garshuni manuscript containing a compilation taken from several authors and dating from before c. 1100 AD . [ 50 ] One of them runs as follows:
The water of vitriol and sulphur which is used to irrigate the drugs: yellow vitriol three parts, yellow sulphur one part, grind them and distil them in the manner of rose-water. [ 51 ]
A recipe for the preparation of sulfuric acid is mentioned in Risālat Jaʿfar al-Sādiq fī ʿilm al-ṣanʿa , an Arabic treatise falsely attributed to the Shi'i Imam Ja'far al-Sadiq (died 765). Julius Ruska dated this treatise to the 13th century, but according to Ahmad Y. al-Hassan it likely dates from an earlier period: [ 52 ]
Then distil green vitriol in a cucurbit and alembic, using medium fire; take what you obtain from the distillate, and you will find it clear with a greenish tint. [ 51 ]
Sulfuric acid was called 'oil of vitriol' by medieval European alchemists because it was prepared by roasting iron(II) sulfate or green vitriol in an iron retort . The first allusions to it in works that are European in origin appear in the thirteenth century AD, as for example in the works of Vincent of Beauvais , in the Compositum de Compositis ascribed to Albertus Magnus , and in pseudo-Geber 's Summa perfectionis . [ 53 ]
A method of producing oleum sulphuris per campanam, or "oil of sulfur by the bell", was known by the 16th century: it involved burning sulfur under a glass bell in moist weather (or, later, under a moistened bell). However, it was very inefficient (according to Gesner , 5 pounds (2.3 kg) of sulfur converted into less than 1 ounce (0.03 kg) of acid), and the resulting product was contaminated by sulfurous acid (or rather, solution of sulfur dioxide ) so most alchemists (including, for example, Isaac Newton) didn't consider it equivalent with the "oil of vitriol".
In the 17th century, Johann Rudolf Glauber discovered that adding saltpeter ( potassium nitrate , KNO 3 ) significantly improves the output, also replacing moisture with steam. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid. In 1736, Joshua Ward , a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead -lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries with a purity of 62% and a conversion of 75%. [ 4 ]
Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS 2 ) was heated in air to yield iron(II) sulfate, FeSO 4 , which was oxidized by further heating in air to form iron(III) sulfate , Fe 2 (SO 4 ) 3 , which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid. [ 4 ]
In 1831, British vinegar merchant Peregrine Phillips patented the contact process , which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method. [ 33 ]
In the early to mid 19th century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim , Northern Ireland , where it was used as a bleach for linen. Early bleaching of linen was done using lactic acid from sour milk but this was a slow process and the use of vitriol sped up the bleaching process. [ 54 ]
Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations . In common with other corrosive acids and alkali , it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues , such as skin and flesh . In addition, it exhibits a strong dehydrating property on carbohydrates , liberating extra heat and causing secondary thermal burns . [ 8 ] [ 9 ] Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes . If ingested, it damages internal organs irreversibly and may even be fatal. [ 7 ] Personal protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials. [ 8 ] Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids , such as hydrochloric acid and nitric acid .
Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time.
The standard first aid treatment for acid spills on the skin is, as for other corrosive agents , irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.
Preparation of diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid is usually added to water and not the other way around. A saying used to remember this is "Do like you oughta, add the acid to the water". [ 55 ] [ better source needed ] [ 56 ] Water has a higher heat capacity than the acid, and so a vessel of cold water will absorb heat as acid is added.
Also, because the acid is denser than water, it sinks to the bottom. Heat is generated at the interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection , which cools the interface, and prevents boiling of either acid or water.
In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol , or worse, an explosion .
Preparation of solutions greater than 6 M (35%) in concentration is dangerous, unless the acid is added slowly enough to allow the mixture sufficient time to cool. Otherwise, the heat produced may be sufficient to boil the mixture. Efficient mechanical stirring and external cooling (such as an ice bath) are essential.
Reaction rates double for about every 10-degree Celsius increase in temperature . [ 57 ] Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction.
On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.
Sulfuric acid is non-flammable.
The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent. [ 58 ] In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m 3 : limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination , loss of axons and gliosis .
International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988 , which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances. [ 59 ] | https://en.wikipedia.org/wiki/H2S04 |
Hydrogen disulfide is the inorganic compound with the formula H 2 S 2 . This hydrogen chalcogenide is a pale yellow volatile liquid with a camphor-like odor. It decomposes readily to hydrogen sulfide ( H 2 S ) and elemental sulfur . [ 1 ]
The connection of atoms in the hydrogen disulfide molecule is H−S−S−H . The structure of hydrogen disulfide is similar to that of hydrogen peroxide , with C 2 point group symmetry . Both molecules are distinctly nonplanar. The dihedral angle between the H a −S−S and S−S−H b planes is 90.6°, compared with 111.5° in H 2 O 2 . The H−S−S bond angle is 92°, close to 90° for unhybridized divalent sulfur. [ 1 ]
Hydrogen disulfide can be synthesised by cracking polysulfanes ( H 2 S n ) according to this idealized equation:
The main impurity is trisulfane ( H 2 S 3 ). [ 1 ] The precursor polysulfane is produced by the reaction of hydrochloric acid with aqueous sodium polysulfide . The polysulfane precipitates as an oil. [ 1 ] [ 2 ]
Upon contact with water or alcohols , hydrogen disulfide readily decomposes under ambient conditions to hydrogen sulfide and sulfur.
It is more acidic than hydrogen sulfide , but the p K a has not been reported. [ 1 ]
In organosulfur chemistry , hydrogen disulfide adds to alkenes to give disulfides and thiols . [ 3 ]
The deuterated form of hydrogen disulfide, deuterium disulfide D−S−S−D (dideuterodisulfane), has a similar geometry to H−S−S−H , but its tunneling time is slower, making it a convenient test case for the quantum Zeno effect , in which frequent observation of a quantum system suppresses its normal evolution. Trost and Hornberger [ 4 ] have calculated that while an isolated D−S−S−D molecule would spontaneously oscillate between left and right chiral forms with a period of 5.6 milliseconds, the presence of a small amount of inert helium gas should stabilize the chiral states, the collisions of the helium atoms in effect "observing" the molecule's momentary chirality and so suppressing spontaneous evolution to the other chiral state. [ 5 ]
In high concentrations, it can cause dizziness, disorientation and ultimately unconsciousness. [ 6 ] | https://en.wikipedia.org/wiki/H2S2 |
Thiosulfurous acid is a hypothetical chemical compound with the formula HS−S(=O)−OH or HO−S(=S)−OH. Attempted synthesis leads to polymers. [ 3 ] It is a low oxidation state (+1) sulfur acid. [ 4 ] It is the Arrhenius acid for disulfur monoxide . Salts derived from thiosulfurous acid, which are also unknown, are named "thiosulfites", "thionosulfites" or "sulfurothioites". The ion is S=SO 2− 2 .
Other possible isomers are dihydroxydisulfane or hypodithionous acid HOSSOH, a linear chain, and thiothionyl hydroxide (S=S(OH) 2 ) a tautomer where the hydrogen has moved from a sulfur to an oxygen. [ 5 ] HOSSOH can have two different rotamers with symmetry C 1 and C 2 . The isomer with one hydrogen on sulfur and one on oxygen is the most stable according to calculations. [ 6 ]
The sulfur analog of thiosulfuric acid ( HS−S(=O) 2 −OH ), in which two sulfur atoms branch off of the central sulfur atom of a linear dihydrogen trisulfide structure (tetrathiosulfuric acid, (S=) 2 S(−SH) 2 or ( − S−) 2 S 2+ (−SH) 2 ) has also been computationally studied. [ 7 ]
It apparently decomposes to polysulfane oxide or polythionic acids in water, which is termed Wakenroder's liquid . [ 5 ]
In alkaline conditions thiosulfurous acid rapidly deteriorates forming a mixture of sulfide, sulfur, sulfite, and thiosulfate. In acidic conditions it will form hydrogen sulfide and sulfur dioxide as well. Some of these react to form pentathionate and other polythionates . Thiosulfurous acid reacts with sulfurous acid to give tetrathionate, and with thiosulfuric acid to make hexathionate. [ 8 ]
Four isomers are possible for R 2 S 2 O 2 , at least restricting sulfur to di- and tetravalency: (RO) 2 S=S, ROSSOR, RS(O) 2 SR, and RS(O)SOR. For the first two, the R groups are equivalent, and in the latter two they are nonequivalent. A simple example is diethylthiosulfite, (EtO) 2 S=S. It is also known as diethylthionosulfite. It is a stereochemically rigid on the NMR timescale to about 140 °C, somewhat similar to diethylsulfoxide. Many derivatives have been prepared from glycols. From meso-hydrobenzoin (PhCH(OH)−CH(OH)Ph), one obtains two isomers; a third isomer results from d , l -PhCH(OH)−CH(OH)Ph. [ 9 ] [ 10 ]
The reaction with simple alkoxide sources with disulfur dichloride gives the unbranched ROSSOR. They are immiscible in water, but dissolve in benzene or carbon tetrachloride. [ 8 ] These species are less rigid than the thiosulfite esters. | https://en.wikipedia.org/wiki/H2S2O2 |
Thiosulfuric acid is the inorganic compound with the formula H 2 S 2 O 3 . It has attracted academic interest as a simple, easily accessed compound that is labile . It has few practical uses.
The acid cannot be made by acidifying aqueous thiosulfate salt solutions as the acid readily decomposes in water. The decomposition products can include sulfur , sulfur dioxide , hydrogen sulfide , polysulfanes , sulfuric acid and polythionates , depending on the reaction conditions. [ 6 ] Anhydrous methods of producing the acid were developed by Max Schmidt: [ 6 ] [ 7 ]
The anhydrous acid also decomposes above −5 °C: [ 6 ]
The isomer (O=) 2 S(−OH)(−SH) is more stable than the isomer (O=)(S=)S(−OH) 2 as established by Hartree–Fock / ab initio calculations with a 6-311 G** basis set and MP2 to MP4 refinements. [ 8 ] [ clarification needed ] The theoretically predicted structure conforms with the double bond rule .
An isomer of thiosulfuric acid is the adduct of hydrogen sulfide and sulfur trioxide , H 2 S·SO 3 , which can also be prepared at low temperature. It is a white crystalline solid. [ 6 ] | https://en.wikipedia.org/wiki/H2S2O3 |
Dithionous acid is a sulfur oxoacid with the chemical formula H 2 S 2 O 4 . It has not been observed experimentally. [ 2 ] It is the conjugate acid of the dithionite dianion, S 2 O 4 2- . Dithionite is a well-known reducing agent.
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2S2O4 |
Disulfurous acid , metabisulfurous acid or pyrosulfurous acid is an oxoacid of sulfur with the formula H 2 S 2 O 5 . Its structure is HO−S(=O) 2 −S(=O)−OH . The salts of disulfurous acid are called disulfites or metabisulfites. Disulfurous acid is, like sulfurous acid ( H 2 SO 3 ), a phantom acid, which does not exist in the free state. [ 2 ] In contrast to disulfate ( S 2 O 2− 7 ), disulfite has two directly connected sulfur atoms. The oxidation state of the sulfur atom bonded to three oxygen atoms is +5 and its valence is 6, while that of the other sulfur is +3 and 4 respectively.
This article about theoretical chemistry is a stub . You can help Wikipedia by expanding it .
This article about an acid is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2S2O5 |
Dithionic acid , H 2 S 2 O 6 , is the inorganic compound with the formula H 2 S 2 O 6 . It is the doubly protonated derivative of dithionate , a well-characterized dianion. Dithionic acid is mainly observed and characterized as an aqueous solution . [ 3 ]
Dithionates can be made by oxidizing a sulfite (from the +4 to the +5 oxidation state ), but on a larger scale they are made by oxidizing a cooled aqueous solution of sulfur dioxide with manganese dioxide :
The manganese dithionate solution formed can then be converted to dithionate salts of other metals by metathesis reactions :
Concentrated solutions of dithionic acid can subsequently be obtained treating a barium dithionate solution with sulfuric acid: | https://en.wikipedia.org/wiki/H2S2O6 |
Disulfuric acid (alternative spelling disulphuric acid ) or pyrosulfuric acid (alternative spelling pyrosulphuric acid ), also named oleum , is a sulfur oxoacid . [ 3 ] It is a major constituent of fuming sulfuric acid, oleum , and this is how most chemists encounter it. As confirmed by X-ray crystallography , the molecule consists of a pair of SO 2 (OH) groups joined by an oxide. [ 4 ]
It is also a minor constituent of liquid anhydrous sulfuric acid due to the equilibria:
The acid is prepared by reacting excess sulfur trioxide (SO 3 ) with sulfuric acid:
Disulfuric acid can be seen as the sulfuric acid analog of an acid anhydride . The mutual electron-withdrawing effects of each sulfuric acid unit on its neighbour causes a marked increase in acidity. Disulfuric acid is strong enough to protonate "normal" sulfuric acid in the (anhydrous) sulfuric acid solvent system. There are salts of disulfuric acid, commonly called pyrosulfates , e.g. potassium pyrosulfate .
There are other related acids with the general formula H 2 O·(SO 3 ) x though none can be isolated. | https://en.wikipedia.org/wiki/H2S2O7 |
Peroxydisulfuric acid is an inorganic compound with a chemical formula (HO 3 SO) 2 . It is also called Marshall's acid after Professor Hugh Marshall , who discovered it in 1891. [ 1 ]
This oxoacid features sulfur in its +6 oxidation state and a peroxide group. Sulfur adopts the usual tetrahedral geometry. [ 2 ]
The acid is prepared by the reaction of chlorosulfuric acid with hydrogen peroxide : [ 3 ]
Another method is the electrolysis of moderately concentrated sulfuric acid (60-70%) with platinum electrodes at high current density and voltage:
Peroxydisulfuric acid is a precursor to several salts including sodium peroxydisulfate, potassium peroxydisulfate, and ammonium peroxydisulfate. These salts are used to initiate the polymerization of acrylonitrile , styrene , and related monomers. This application exploits the tendency of the peroxydisulfate anion to undergo homolysis to produce radicals. They are also used for cleaning circuit boards. [ 3 ] | https://en.wikipedia.org/wiki/H2S2O8 |
Trisulfane is the inorganic compound with the formula H 2 S 3 . It is a pale yellow volatile liquid with a camphor-like odor. It decomposes readily to hydrogen sulfide ( H 2 S ) and elemental sulfur . It is produced by distillation of the polysulfane oil obtained by acidification of polysulfide salts. [ 3 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2S3 |
Trithionic acid is a polythionic acid with three sulfur atoms. It can be viewed as two bisulfite radicals bridged by a sulfur atom.
This article about an acid is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2S3O6 |
Sulfuric(IV) acid ( United Kingdom spelling: sulphuric(IV) acid ), also known as sulfurous (UK: sulphurous ) acid and thionic acid , [ citation needed ] is the chemical compound with the formula H 2 SO 3 .
Raman spectra of solutions of sulfur dioxide in water show only signals due to the SO 2 molecule and the bisulfite ion, HSO − 3 . [ 2 ] The intensities of the signals are consistent with the following equilibrium :
17 O NMR spectroscopy provided evidence that solutions of sulfurous acid and protonated sulfites contain a mixture of isomers, which is in equilibrium: [ 3 ]
Attempts to concentrate the solutions of sulfurous acid simply reverse the equilibrium, producing sulfur dioxide and water vapor. A clathrate with the formula 4SO 2 ·23H 2 O has been crystallised. It decomposes above 7 °C.
Sulfurous acid is commonly known not to exist in its free state, and owing to this, it is stated in textbooks that it cannot be isolated in the water-free form. [ 4 ] However, the molecule has been detected in the gas phase in 1988 by the dissociative ionization of diethyl sulfite . [ 5 ] The conjugate bases of this elusive acid are, however, common anions, bisulfite (or hydrogen sulfite) and sulfite . Sulfurous acid is an intermediate species in the formation of acid rain from sulfur dioxide. [ 6 ]
Aqueous solutions of sulfur dioxide, which sometimes are referred to as sulfurous acid, are used as reducing agents and as disinfectants, as are solutions of bisulfite and sulfite salts. They are oxidised to sulfuric acid or sulfate by accepting another oxygen atom. [ 7 ] | https://en.wikipedia.org/wiki/H2SO3 |
Sulfuric acid ( American spelling and the preferred IUPAC name ) or sulphuric acid ( Commonwealth spelling ), known in antiquity as oil of vitriol , is a mineral acid composed of the elements sulfur , oxygen , and hydrogen , with the molecular formula H 2 SO 4 . It is a colorless, odorless, and viscous liquid that is miscible with water. [ 7 ]
Pure sulfuric acid does not occur naturally due to its strong affinity to water vapor ; it is hygroscopic and readily absorbs water vapor from the air . [ 7 ] Concentrated sulfuric acid is a strong oxidant with powerful dehydrating properties, making it highly corrosive towards other materials, from rocks to metals. Phosphorus pentoxide is a notable exception in that it is not dehydrated by sulfuric acid but, to the contrary, dehydrates sulfuric acid to sulfur trioxide . Upon addition of sulfuric acid to water, a considerable amount of heat is released; thus, the reverse procedure of adding water to the acid is generally avoided since the heat released may boil the solution, spraying droplets of hot acid during the process. Upon contact with body tissue, sulfuric acid can cause severe acidic chemical burns and secondary thermal burns due to dehydration. [ 8 ] [ 9 ] Dilute sulfuric acid is substantially less hazardous without the oxidative and dehydrating properties; though, it is handled with care for its acidity.
Many methods for its production are known, including the contact process , the wet sulfuric acid process , and the lead chamber process . [ 10 ] Sulfuric acid is also a key substance in the chemical industry . It is most commonly used in fertilizer manufacture [ 11 ] but is also important in mineral processing , oil refining , wastewater treating , and chemical synthesis . It has a wide range of end applications, including in domestic acidic drain cleaners , [ 12 ] as an electrolyte in lead-acid batteries , as a dehydrating compound, and in various cleaning agents .
Sulfuric acid can be obtained by dissolving sulfur trioxide in water.
Although nearly 100% sulfuric acid solutions can be made, the subsequent loss of SO 3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade, which is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are: [ 13 ] [ 14 ]
"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process , chamber acid being the acid produced in the lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid ) and tower acid being the acid recovered from the bottom of the Glover tower. [ 13 ] [ 14 ] They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10 M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations ), is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. [ 14 ]
Sulfuric acid contains not only H 2 SO 4 molecules, but is actually an equilibrium of many other chemical species, as it is shown in the table below.
Sulfuric acid is a colorless oily liquid, and has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, [ 16 ] and 98% sulfuric acid has a vapor pressure of <1 mmHg at 40 °C. [ 17 ]
In the solid state, sulfuric acid is a molecular solid that forms monoclinic crystals with nearly trigonal lattice parameters. The structure consists of layers parallel to the (010) plane, in which each molecule is connected by hydrogen bonds to two others. [ 3 ] Hydrates H 2 SO 4 · n H 2 O are known for n = 1, 2, 3, 4, 6.5, and 8, although most intermediate hydrates are stable against disproportionation . [ 18 ]
Anhydrous H 2 SO 4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity , a consequence of autoprotolysis , i.e. self- protonation : [ 15 ]
The equilibrium constant for autoprotolysis (25 °C) is: [ 15 ]
The corresponding equilibrium constant for water , K w is 10 −14 , a factor of 10 10 (10 billion) smaller.
In spite of the viscosity of the acid, the effective conductivities of the H 3 SO + 4 and HSO − 4 ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.
The hydration reaction of sulfuric acid is highly exothermic . [ 19 ]
As indicated by its acid dissociation constant , sulfuric acid is a strong acid:
The product of this ionization is HSO − 4 , the bisulfate anion. Bisulfate is a far weaker acid:
The product of this second dissociation is SO 2− 4 , the sulfate anion.
Concentrated sulfuric acid has a powerful dehydrating property, removing water ( H 2 O ) from other chemical compounds such as table sugar ( sucrose ) and other carbohydrates , to produce carbon , steam , and heat. Dehydration of table sugar (sucrose) is a common laboratory demonstration. [ 21 ] The sugar darkens as carbon is formed, and a rigid column of black, porous carbon called a carbon snake may emerge. [ 22 ]
Similarly, mixing starch into concentrated sulfuric acid gives elemental carbon and water. The effect of this can also be seen when concentrated sulfuric acid is spilled on paper. Paper is composed of cellulose , a polysaccharide related to starch. The cellulose reacts to give a burnt appearance in which the carbon appears much like soot that results from fire.
Although less dramatic, the action of the acid on cotton , even in diluted form, destroys the fabric.
The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystals change into white powder as water is removed.
Sulfuric acid reacts with most bases to give the corresponding sulfate or bisulfate.
Aluminium sulfate , also known as paper maker's alum, is made by treating bauxite with sulfuric acid:
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate , for example, displaces acetic acid , CH 3 COOH , and forms sodium bisulfate :
Similarly, treating potassium nitrate with sulfuric acid produces nitric acid . Sulfuric acid reacts with sodium chloride , and gives hydrogen chloride gas and sodium bisulfate :
When combined with nitric acid , sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO + 2 , which is important in nitration reactions involving electrophilic aromatic substitution . This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols.
When allowed to react with superacids , sulfuric acid can act as a base and can be protonated, forming the [H 3 SO 4 ] + ion. Salts of [H 3 SO 4 ] + have been prepared (e.g. trihydroxyoxosulfonium hexafluoroantimonate(V) [H 3 SO 4 ] + [SbF 6 ] − ) using the following reaction in liquid HF :
The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply fluoroantimonic acid , however, has met with failure, as pure sulfuric acid undergoes self-ionization to give [H 3 O] + ions:
which prevents the conversion of H 2 SO 4 to [H 3 SO 4 ] + by the HF/ SbF 5 system. [ 23 ]
Even diluted sulfuric acid reacts with many metals via a single displacement reaction, like other typical acids , producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series ) such as iron , aluminium , zinc , manganese , magnesium , and nickel .
Concentrated sulfuric acid can serve as an oxidizing agent , releasing sulfur dioxide: [ 8 ]
Lead and tungsten , however, are resistant to sulfuric acid.
Hot concentrated sulfuric acid oxidizes carbon [ 24 ] (as bituminous coal ) and sulfur :
Benzene and many derivatives undergo electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids : [ 25 ]
Sulfuric acid can be used to produce hydrogen from water :
The compounds of sulfur and iodine are recovered and reused, hence the process is called the sulfur–iodine cycle . This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur–iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy . It is an alternative to electrolysis , and does not require hydrocarbons like current methods of steam reforming . But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it. [ 26 ] [ 27 ]
Sulfuric acid is rarely encountered naturally on Earth in anhydrous form, due to its great affinity for water . Dilute sulfuric acid is a constituent of acid rain , which is formed by atmospheric oxidation of sulfur dioxide in the presence of water —i.e. oxidation of sulfurous acid . When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water).
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as pyrite :
The resulting highly acidic water is called acid mine drainage (AMD) or acid rock drainage (ARD).
The Fe 2+ can be further oxidized to Fe 3+ :
The Fe 3+ produced can be precipitated as the hydroxide or hydrous iron oxide :
The iron(III) ion (" ferric iron ") can also oxidize pyrite:
FeS 2 (s) + 14 Fe 3+ + 8 H 2 O → 15 Fe 2+ + 2 SO 2− 4 + 16 H +
When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process.
ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.
Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales ) concentrates sulfuric acid in cell vacuoles. [ 28 ]
In the stratosphere , the atmosphere's second layer that is generally between 10–50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical : [ 29 ]
Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer. [ 29 ]
The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. [ 30 ] Sulfuric acid ice has been detected on Jupiter 's moon Europa , where it forms when sulfur ions from Jupiter's magnetosphere implant into the icy surface. [ 31 ]
Sulfuric acid is produced from sulfur , oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).
In the first step, sulfur is burned to produce sulfur dioxide.
The sulfur dioxide is oxidized to sulfur trioxide by oxygen in the presence of a vanadium(V) oxide catalyst . This reaction is reversible and the formation of the sulfur trioxide is exothermic.
The sulfur trioxide is absorbed into 97–98% H 2 SO 4 to form oleum ( H 2 S 2 O 7 ), also known as fuming sulfuric acid or pyrosulphuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.
Directly dissolving SO 3 in water, called the " wet sulfuric acid process ", is rarely practiced because the reaction is extremely exothermic, resulting in a hot aerosol of sulfuric acid that requires condensation and separation.
In the first step, sulfur is burned to produce sulfur dioxide:
or, alternatively, hydrogen sulfide ( H 2 S ) gas is incinerated to SO 2 gas:
The sulfur dioxide then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst .
The sulfur trioxide is hydrated into sulfuric acid H 2 SO 4 :
The last step is the condensation of the sulfuric acid to liquid 97–98% H 2 SO 4 :
Burning sulfur together with saltpeter ( potassium nitrate , KNO 3 ), in the presence of steam, has been used historically. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid.
Prior to 1900, most sulfuric acid was manufactured by the lead chamber process . [ 32 ] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.
A wide variety of laboratory syntheses are known, and typically begin from sulfur dioxide or an equivalent salt . In the metabisulfite method, hydrochloric acid reacts with metabisulfite to produce sulfur dioxide vapors. The gas is bubbled through nitric acid , which will release brown/red vapors of nitrogen dioxide as the reaction proceeds. The completion of the reaction is indicated by the ceasing of the fumes. This method conveniently does not produce an inseparable mist. [ citation needed ]
Alternatively, dissolving sulfur dioxide in an aqueous solution of an oxidizing metal salt such as copper(II) or iron(III) chloride: [ citation needed ]
Two less well-known laboratory methods of producing sulfuric acid, albeit in dilute form and requiring some extra effort in purification, rely on electrolysis . A solution of copper(II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper at cathode and oxygen gas at the anode. The solution of dilute sulfuric acid indicates completion of the reaction when it turns from blue to clear (production of hydrogen at cathode is another sign): [ citation needed ]
More costly, dangerous, and troublesome is the electrobromine method, which employs a mixture of sulfur , water, and hydrobromic acid as the electrolyte. The sulfur is pushed to bottom of container under the acid solution. Then the copper cathode and platinum/graphite anode are used with the cathode near the surface and the anode is positioned at the bottom of the electrolyte to apply the current. This may take longer and emits toxic bromine /sulfur-bromide vapors, but the reactant acid is recyclable. Overall, only the sulfur and water are converted to sulfuric acid and hydrogen (omitting losses of acid as vapors): [ citation needed ]
Sulfuric acid is a very important commodity chemical, and a nation's sulfuric acid production was as recently as 2002 believed to be a good indicator of its industrial strength. [ 33 ] World production in the year 2004 was about 180 million tonnes , with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%. [ 34 ] World production in 2022 was estimated at about 260 million tonnes. [ 35 ]
As of the late 20th century, most of the produced amount (≈60%) was consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze , as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, and water treatment. About 6% of uses are related to pigments and include paints, enamels , printing inks, coated fabrics and paper, while the rest is dispersed into a multitude of applications such as production of explosives, cellophane , acetate and viscose textiles, lubricants, non-ferrous metals , and batteries. [ 36 ]
The dominant use for sulfuric acid is in the "wet method" for the production of phosphoric acid , used for manufacture of phosphate fertilizers . In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite , though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate , hydrogen fluoride (HF) and phosphoric acid . The HF is removed as hydrofluoric acid . The overall process can be represented as:
Ammonium sulfate , an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.
Sulfuric acid is also important in the manufacture of dyestuffs solutions.
Sulfuric acid is used in steelmaking and other metallurgical industries as a pickling agent for removal of rust and fouling . [ 37 ] Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid [ clarification needed ] with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide ( SO 2 ) and sulfur trioxide ( SO 3 ) which are then used to manufacture "new" sulfuric acid.
Hydrogen peroxide ( H 2 O 2 ) can be added to sulfuric acid to produce piranha solution , a powerful but potentially hazardous cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used in the microelectronics industry, and also in laboratory settings to clean glassware.
Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam , used for making nylon . It is used for making hydrochloric acid from salt via the Mannheim process . Much H 2 SO 4 is used in petroleum refining , for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane , a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidizing agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.
Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator):
At anode :
At cathode :
Overall:
Sulfuric acid at high concentrations is frequently the major ingredient in domestic acidic drain cleaners [ 12 ] which are used to remove lipids , hair , tissue paper , etc. Similar to their alkaline versions , such drain openers can dissolve fats and proteins via hydrolysis . Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.
The study of vitriols (hydrated sulfates of various metals forming glassy minerals from which sulfuric acid can be derived) began in ancient times . Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis , in the treatise Phisica et Mystica , and the Leyden papyrus X . [ 38 ] Medieval Islamic alchemists like the authors writing under the name of Jabir ibn Hayyan (died c. 806 – c. 816 AD, known in Latin as Geber), Abu Bakr al-Razi (865 – 925 AD, known in Latin as Rhazes), Ibn Sina (980 – 1037 AD, known in Latin as Avicenna), and Muhammad ibn Ibrahim al-Watwat (1234 – 1318 AD) included vitriol in their mineral classification lists. [ 39 ]
The Jabirian authors and al-Razi experimented extensively with the distillation of various substances, including vitriols. [ 40 ] In one recipe recorded in his Kitāb al-Asrār ( 'Book of Secrets' ), al-Razi may have created sulfuric acid without being aware of it: [ 41 ]
Take white (Yemeni) alum , dissolve it and purify it by filtration. Then distil (green?) vitriol with copper-green (the acetate), and mix (the distillate) with the filtered solution of the purified alum, afterwards let it solidify (or crystallise) in the glass beaker. You will get the best qalqadis (white alum) that may be had. [ 42 ]
In an anonymous Latin work variously attributed to Aristotle (under the title Liber Aristotilis , 'Book of Aristotle'), [ 43 ] to al-Razi (under the title Lumen luminum magnum , 'Great Light of Lights'), or to Ibn Sina, [ 44 ] the author speaks of an 'oil' ( oleum ) obtained through the distillation of iron(II) sulfate (green vitriol), which was likely 'oil of vitriol' or sulfuric acid. [ 45 ] The work refers multiple times to Jabir ibn Hayyan's Seventy Books ( Liber de septuaginta ), one of the few Arabic Jabir works that were translated into Latin. [ 46 ] The author of the version attributed to al-Razi also refers to the Liber de septuaginta as his own work, showing that he erroneously believed the Liber de septuaginta to be a work by al-Razi. [ 47 ] There are several indications that the anonymous work was an original composition in Latin, [ 48 ] although according to one manuscript it was translated by a certain Raymond of Marseilles, meaning that it may also have been a translation from the Arabic. [ 49 ]
According to Ahmad Y. al-Hassan , three recipes for sulfuric acid occur in an anonymous Garshuni manuscript containing a compilation taken from several authors and dating from before c. 1100 AD . [ 50 ] One of them runs as follows:
The water of vitriol and sulphur which is used to irrigate the drugs: yellow vitriol three parts, yellow sulphur one part, grind them and distil them in the manner of rose-water. [ 51 ]
A recipe for the preparation of sulfuric acid is mentioned in Risālat Jaʿfar al-Sādiq fī ʿilm al-ṣanʿa , an Arabic treatise falsely attributed to the Shi'i Imam Ja'far al-Sadiq (died 765). Julius Ruska dated this treatise to the 13th century, but according to Ahmad Y. al-Hassan it likely dates from an earlier period: [ 52 ]
Then distil green vitriol in a cucurbit and alembic, using medium fire; take what you obtain from the distillate, and you will find it clear with a greenish tint. [ 51 ]
Sulfuric acid was called 'oil of vitriol' by medieval European alchemists because it was prepared by roasting iron(II) sulfate or green vitriol in an iron retort . The first allusions to it in works that are European in origin appear in the thirteenth century AD, as for example in the works of Vincent of Beauvais , in the Compositum de Compositis ascribed to Albertus Magnus , and in pseudo-Geber 's Summa perfectionis . [ 53 ]
A method of producing oleum sulphuris per campanam, or "oil of sulfur by the bell", was known by the 16th century: it involved burning sulfur under a glass bell in moist weather (or, later, under a moistened bell). However, it was very inefficient (according to Gesner , 5 pounds (2.3 kg) of sulfur converted into less than 1 ounce (0.03 kg) of acid), and the resulting product was contaminated by sulfurous acid (or rather, solution of sulfur dioxide ) so most alchemists (including, for example, Isaac Newton) didn't consider it equivalent with the "oil of vitriol".
In the 17th century, Johann Rudolf Glauber discovered that adding saltpeter ( potassium nitrate , KNO 3 ) significantly improves the output, also replacing moisture with steam. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid. In 1736, Joshua Ward , a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead -lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries with a purity of 62% and a conversion of 75%. [ 4 ]
Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS 2 ) was heated in air to yield iron(II) sulfate, FeSO 4 , which was oxidized by further heating in air to form iron(III) sulfate , Fe 2 (SO 4 ) 3 , which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid. [ 4 ]
In 1831, British vinegar merchant Peregrine Phillips patented the contact process , which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method. [ 33 ]
In the early to mid 19th century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim , Northern Ireland , where it was used as a bleach for linen. Early bleaching of linen was done using lactic acid from sour milk but this was a slow process and the use of vitriol sped up the bleaching process. [ 54 ]
Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations . In common with other corrosive acids and alkali , it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues , such as skin and flesh . In addition, it exhibits a strong dehydrating property on carbohydrates , liberating extra heat and causing secondary thermal burns . [ 8 ] [ 9 ] Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes . If ingested, it damages internal organs irreversibly and may even be fatal. [ 7 ] Personal protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials. [ 8 ] Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids , such as hydrochloric acid and nitric acid .
Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time.
The standard first aid treatment for acid spills on the skin is, as for other corrosive agents , irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.
Preparation of diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid is usually added to water and not the other way around. A saying used to remember this is "Do like you oughta, add the acid to the water". [ 55 ] [ better source needed ] [ 56 ] Water has a higher heat capacity than the acid, and so a vessel of cold water will absorb heat as acid is added.
Also, because the acid is denser than water, it sinks to the bottom. Heat is generated at the interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection , which cools the interface, and prevents boiling of either acid or water.
In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol , or worse, an explosion .
Preparation of solutions greater than 6 M (35%) in concentration is dangerous, unless the acid is added slowly enough to allow the mixture sufficient time to cool. Otherwise, the heat produced may be sufficient to boil the mixture. Efficient mechanical stirring and external cooling (such as an ice bath) are essential.
Reaction rates double for about every 10-degree Celsius increase in temperature . [ 57 ] Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction.
On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.
Sulfuric acid is non-flammable.
The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent. [ 58 ] In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m 3 : limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination , loss of axons and gliosis .
International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988 , which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances. [ 59 ] | https://en.wikipedia.org/wiki/H2SO4 |
Peroxymonosulfuric acid , also known as persulfuric acid , peroxysulfuric acid is the inorganic compound with the formula H 2 SO 5 . It is a white solid. It is a component of Caro's acid , which is a solution of peroxymonosulfuric acid in sulfuric acid containing small amounts of water. [ 4 ] Peroxymonosulfuric acid is a very strong oxidant ( E 0 = +2.51 V).
In peroxymonosulfuric acid, the S(VI) center adopts its characteristic tetrahedral geometry; the connectivity is indicated by the formula HO–O–S(O) 2 –OH. The S-O- H proton is more acidic. [ 4 ]
The German chemist Heinrich Caro first reported investigations of mixtures of hydrogen peroxide and sulfuric acid. [ 5 ]
One laboratory scale preparation of Caro's acid involves the combination of chlorosulfuric acid and hydrogen peroxide : [ 6 ]
Patents include more than one reaction for preparation of Caro's acid, usually as an intermediate for the production of potassium monopersulfate (PMPS) , a bleaching and oxidizing agent. One route employs the following reaction: [ 7 ]
This reaction is related to " piranha solution ".
H 2 SO 5 and Caro's acid have been used for a variety of disinfectant and cleaning applications, e.g., swimming pool treatment and denture cleaning. It is used in gold mining to destroy the cyanide in the waste stream (" Tailings ").
Alkali metal salts of H 2 SO 5 , especially oxone , are widely investigated.
These peroxy acids can be explosive. Explosions have been reported at Brown University [ 8 ] and Sun Oil . As with all strong oxidizing agents, peroxysulfuric acid is incompatible with organic compounds . | https://en.wikipedia.org/wiki/H2SO5 |
Hydrogen selenide is an inorganic compound with the formula H 2 Se. This hydrogen chalcogenide is the simplest and most commonly encountered hydride of selenium . H 2 Se is a colorless, flammable gas under standard conditions. It is the most toxic selenium compound [ 3 ] with an exposure limit of 0.05 ppm over an 8-hour period. [ 4 ] [ 5 ] Even at extremely low concentrations, this compound has a very irritating smell resembling that of decayed horseradish or "leaking gas", but smells of rotten eggs at higher concentrations.
H 2 Se adopts a bent structure with a H−Se−H bond angle of 91° [ citation needed ] . Consistent with this structure, three IR -active vibrational bands are observed: 2358, 2345, and 1034 cm −1 . [ 6 ]
The properties of H 2 S and H 2 Se are similar, although the selenide is more acidic with p K a = 3.89 and the second p K a = 11, [ 6 ] or 15.05 ± 0.02 at 25 °C. [ 7 ]
Industrially, it is produced by treating elemental selenium at T > 300 °C with hydrogen gas. [ 8 ] A number of routes to H 2 Se have been reported, which are suitable for both large and small scale preparations. In the laboratory, H 2 Se is usually prepared by the action of water on Al 2 Se 3 , concomitant with formation of hydrated alumina . A related reaction involves the acid hydrolysis of FeSe. [ 9 ]
H 2 Se can also be prepared by means of different methods based on the in situ generation in aqueous solution using boron hydride , Marsh test and Devarda's alloy . According to the Sonoda method, H 2 Se is generated from the reaction of H 2 O and CO on Se in the presence of Et 3 N . [ 10 ] H 2 Se can be purchased in cylinders.
Elemental selenium can be recovered from H 2 Se through a reaction with aqueous sulfur dioxide (SO 2 ).
Its decomposition is used to prepare the highly pure element.
H 2 Se is commonly used in the synthesis of Se-containing compounds. It adds across alkenes. Illustrative is the synthesis of selenoureas from cyanamides : [ 11 ]
H 2 Se gas is used to dope semiconductors with selenium.
Hydrogen selenide is hazardous, being the most toxic selenium compound [ 3 ] and far more toxic than its congener hydrogen sulfide . The threshold limit value is 0.05 ppm. The gas acts as an irritant at concentrations higher than 0.3 ppm, which is the main warning sign of exposure; below 1 ppm, this is "insufficient to prevent exposure", while at 1.5 ppm the irritation is "intolerable". [ 5 ] Exposure at high concentrations, even for less than a minute, causes the gas to attack the eyes and mucous membranes; this causes cold-like symptoms for at least a few days afterwards. In Germany, the limit in drinking water is 0.008 mg/L, and the US EPA recommends a maximum contamination of 0.01 mg/L. [ 8 ] [ 12 ]
Despite being extremely toxic, no human fatalities have yet been reported. It is suspected that this is due to the gas' tendency to oxidise to form red selenium in mucous membranes; elemental selenium is less toxic than selenides are. [ 4 ] | https://en.wikipedia.org/wiki/H2Se |
Hydrogen diselenide is an inorganic selenium compound with a chemical formula H 2 Se 2 or (SeH) 2 . [ 1 ] [ 2 ] At room temperature , hydrogen diselenide dissociates easily to hydrogen selenide ( H 2 Se ) and elemental selenium , and is therefore not stable. However, hydrogen diselenide can be stable in some solutions. [ 3 ] | https://en.wikipedia.org/wiki/H2Se2 |
Selenous acid (or selenious acid ) is the chemical compound with the formula H 2 SeO 3 . Structurally, it is more accurately described by O=Se(OH) 2 . It is the principal oxoacid of selenium ; the other being selenic acid .
Selenous acid is analogous to sulfurous acid , but it is more readily isolated. Selenous acid is easily formed upon the addition of selenium dioxide to water. As a crystalline solid, the compound can be seen as pyramidal molecules that are interconnected with hydrogen bonds. In solution it is a diprotic acid: [ 3 ]
It is moderately oxidizing in nature, but kinetically slow. In 1 M H + :
In 1 M OH − :
Selenous acid is hygroscopic . [ 4 ] [ 5 ]
The major use is in protecting and changing the color of steel, especially steel parts on firearms. [ 6 ] The so-called cold-bluing process uses selenous acid, copper(II) nitrate , and nitric acid to change the color of the steel from silver-grey to blue-grey or black. Alternative procedures use copper sulfate and phosphoric acid instead. This process deposits a coating of copper selenide and is fundamentally different from other bluing processes which generate black iron oxide . Some older razor blades were also made of blued steel. [ 6 ]
Another use for selenious acid is the chemical darkening and patination of copper, brass and bronze, producing a rich dark brown color that can be further enhanced with mechanical abrasion. [ citation needed ]
It is used in organic synthesis as an oxidizing agent for the synthesis of 1,2-dicarbonyl compounds, e.g. in laboratory preparation of glyoxal (oxaldehyde) from acetaldehyde . [ 7 ]
Selenious acid is a key component of the Mecke reagent used for drug checking. [ 8 ] [ 9 ]
Selenous acid can supply the trace element indicated in people as a source of selenium. [ 10 ] [ 11 ]
Like many selenium compounds, selenous acid is highly toxic in excessive quantities, and ingestion of any significant quantity of selenous acid is usually fatal, however it is an approved dietary source in proper amounts. Symptoms of selenium poisoning can occur several hours after exposure, and may include stupor , nausea , severe hypotension and death. | https://en.wikipedia.org/wiki/H2SeO3 |
Selenic acid is the inorganic compound with the formula H 2 SeO 4 . It is an oxoacid of selenium , and its structure is more accurately described as O 2 Se(OH) 2 . It is a colorless compound. Although it has few uses, one of its salts, sodium selenate is used in the production of glass and animal feeds. [ 3 ]
The molecule is tetrahedral, as predicted by VSEPR theory . The Se–O bond length is 161 pm . [ 4 ] In the solid state, it crystallizes in an orthorhombic structure. [ 5 ]
It is prepared by oxidising selenium compounds in lower oxidation states. One method involves the oxidation of selenium dioxide with hydrogen peroxide :
Unlike the production sulfuric acid by hydration of sulfur trioxide , the hydration of selenium trioxide is an impractical method. [ 4 ] Instead, selenic acid may also be prepared by the oxidation of selenous acid ( H 2 SeO 3 ) with halogens, such as chlorine or bromine , or with potassium permanganate . [ 6 ] Using chlorine or bromine as the oxidising agents also produces hydrochloric or hydrobromic acid as a side-product, which needs to be removed from the solution since they can reduce the selenic acid to selenous acid. [ 7 ]
To obtain the anhydrous acid as a crystalline solid, the resulting solution is evaporated at temperatures below 140 °C (413 K; 284 °F) in a vacuum. [ 8 ]
Like sulfuric acid , selenic acid is a strong acid that is hygroscopic and extremely soluble in water. Concentrated solutions are viscous. Crystalline mono- and di- hydrates are known. [ 6 ] The monohydrate melts at 26 °C, and the dihydrate melts at −51.7 °C. [ 4 ]
Selenic acid is a stronger oxidizer than sulfuric acid , [ 9 ] capable of liberating chlorine from chloride ions , being reduced to selenous acid in the process:
It decomposes above 200 °C, liberating oxygen gas and being reduced to selenous acid: [ 6 ]
Selenic acid reacts with barium salts to precipitate solid BaSeO 4 , analogous to the sulfate. In general, selenate salts resemble sulfate salts, but are more soluble. Many selenate salts have the same crystal structure as the corresponding sulfate salts. [ 4 ]
Treatment with fluorosulfuric acid gives selenoyl fluoride : [ 8 ]
Hot, concentrated selenic acid reacts with gold , forming a reddish-yellow solution of gold(III) selenate: [ 10 ]
Selenic acid is used as a specialized oxidizing agent. | https://en.wikipedia.org/wiki/H2SeO4 |
Disilyne is a low valent silicon compound with the chemical formula Si 2 R 2 where oxidation state of Si is +1. Several isomers are possible, but none are sufficiently stable to be of practical value. Substituted disilynes contain a formal silicon–silicon triple bond and as such are sometimes written R 2 Si 2 (where R is a substituent group). They are the silicon analogues of alkynes .
The term silyne has two diverse meanings. Some chemists use it to refer to compounds containing a silicon–silicon triple bond, [ 1 ] by analogy to the carbon–carbon triple bond in alkynes , whereas others use the term to refer to compounds containing a silicon–carbon triple bond [ 2 ] by analogy to silene, which often refers to compounds containing silicon–carbon double bonds. [ 3 ] The term polysilyne can refer to the layer polymer (SiH) n or substituted derivatives. [ 1 ]
The first substituted disilyne to be isolated and characterised by X-ray crystallography is one with an additional trisubstituted silicon group on each silicon of the disilyne core. The structure R′ 2 R′′Si−Si≡Si−SiR′′R′ 2 , where R′ = HC(SiMe 3 ) 2 and R′′ = HCMe 2 , is an emerald green crystalline compound reported in 2004. [ 4 ]
It was prepared by the reduction of the related tetrabrominated precursor by potassium graphite (KC 8 ). It is air- and moisture-sensitive but is a stable solid up to 128 °C.
The geometry of disilynes is unlike that of analogous carbon structures. Whereas substituted alkynes , such as 2-butyne , are linear, having a 180° bond angle at each end of the carbon–carbon triple bond, the Si−Si≡Si−Si chain is bent to 137° at each end. The four silicon atoms in the chain are however perfectly coplanar, with the first and fourth silicon atoms trans to one another. The central triple bond length is 206 pm, which is around 4% shorter than the typical bond-length of Si–Si double bonds (214 pm)) and the Si–Si single bonds are 237 pm. The color is attributed to a weak π–π * transition.
Calculations show a bond order of 2.6. An alternative calculation of the bond order by a different group describes the bonding as essentially due to only two electron pairs, with the third pair in a non-bonding orbital . [ 5 ] [ 6 ] [ 7 ] Reaction of this compound with phenylacetylene produced a 1,2-di silabenzene . [ 8 ] Other workers [ 9 ] have also reported another related compound which contains a hexasila-3-yne chain:
In this compound, the Si–Si triple bond length was calculated as 207 pm.
Triple bonded compounds of the heavier members of group 14 have also been prepared; lead , [ 10 ] and tin [ 11 ] and germanium ( digermyne ) [ 12 ] The cores of the disilyne, digermyne, distannyne, and diplumbyne have similarly bent geometries. These findings are generally consistent with the absence of conventional triple bonds. | https://en.wikipedia.org/wiki/H2Si2 |
Sodium fluorosilicate
Hexafluorosilicic acid is an inorganic compound with the chemical formula H 2 SiF 6 . Aqueous solutions of hexafluorosilicic acid consist of salts of the cation and hexafluorosilicate anion. These salts and their aqueous solutions are colorless.
Hexafluorosilicic acid is produced naturally on a large scale in volcanoes. [ 2 ] [ 3 ] It is manufactured as a coproduct in the production of phosphate fertilizers . The resulting hexafluorosilicic acid is almost exclusively consumed as a precursor to aluminum trifluoride and synthetic cryolite , which are used in aluminium processing. Salts derived from hexafluorosilicic acid are called hexafluorosilicates .
Hexafluorosilicic acid has been crystallized as various hydrates. These include ( H 5 O 2 ) 2 SiF 6 , the more complicated (H 5 O 2 ) 2 SiF 6 ·2H 2 O, and (H 5 O 2 )(H 7 O 3 )SiF 6 ·4.5H 2 O. In all of these salts, the octahedral hexafluorosilicate anion is hydrogen bonded to the cations. [ 4 ]
Aqueous solutions of hexafluorosilicic acid are often described as H 2 SiF 6 .
Hexafluorosilicic acid is produced commercially from fluoride-containing minerals that also contain silicates. Specifically, apatite and fluorapatite are treated with sulfuric acid to give phosphoric acid , a precursor to several water-soluble fertilizers. This is called the wet phosphoric acid process . [ 5 ] As a by-product, approximately 50 kg of hexafluorosilicic acid is produced per tonne of HF owing to reactions involving silica-containing mineral impurities. [ 6 ] : 3
Some of the hydrogen fluoride (HF) produced during this process in turn reacts with silicon dioxide (SiO 2 ) impurities, which are unavoidable constituents of the mineral feedstock, to give silicon tetrafluoride . Thus formed, the silicon tetrafluoride reacts further with HF. [ citation needed ] The net process can be described as: [ 7 ] [ page needed ]
Hexafluorosilicic acid can also be produced by treating silicon tetrafluoride with hydrofluoric acid. [ 7 ]
Hexafluorosilic acid is only stable in hydrogen fluoride or acidic aqueous solutions. In any other circumstance, it acts as a source of hydrofluoric acid . Thus, for example, hexafluorosilicic acid pure or in oleum solution evolves silicon tetrafluoride until the residual hydrogen fluoride re-establishes equilibrium: [ 7 ]
In alkaline-to-neutral aqueous solutions, hexafluorosilicic acid readily hydrolyzes to fluoride anions and amorphous, hydrated silica ("SiO 2 "). Strong bases give fluorosilicate salts at first, but any stoichiometric excess begins hydrolysis. [ 7 ] At the concentrations usually used for water fluoridation , 99% hydrolysis occurs: [ 6 ] [ 8 ]
Neutralization of solutions of hexafluorosilicic acid with alkali metal bases produces the corresponding alkali metal fluorosilicate salts:
The resulting salt Na 2 SiF 6 is mainly used in water fluoridation . Related ammonium and barium salts are produced similarly for other applications. At room temperature 15–30% concentrated hexafluorosilicic acid undergoes similar reactions with chlorides , hydroxides , and carbonates of alkali and alkaline earth metals . [ 9 ]
Sodium hexafluorosilicate for instance may be produced by treating sodium chloride ( NaCl ) by hexafluorosilicic acid: [ 6 ] : 3 [ 10 ] : 7
Heating sodium hexafluorosilicate gives silicon tetrafluoride : [ 10 ] : 8
The majority of the hexafluorosilicic acid is converted to aluminium fluoride and synthetic cryolite . These materials are central to the conversion of aluminium ore into aluminium metal. The conversion to aluminium trifluoride is described as: [ 7 ]
Hexafluorosilicic acid is also converted to a variety of useful hexafluorosilicate salts. The potassium salt, Potassium fluorosilicate , is used in the production of porcelains, the magnesium salt for hardened concretes and as an insecticide, and the barium salts for phosphors.
Hexafluorosilicic acid and the salts are used as wood preservation agents. [ 11 ]
Hexafluorosilicic acid is also used as an electrolyte in the Betts electrolytic process for refining lead.
Hexafluorosilicic acid (identified as hydrofluorosilicic acid on the label) along with oxalic acid are the active ingredients used in Iron Out rust-removing cleaning products, which are essentially varieties of laundry sour .
H 2 SiF 6 is a specialized reagent in organic synthesis for cleaving Si–O bonds of silyl ethers . It is more reactive for this purpose than HF. It reacts faster with t - butyldimethysilyl ( TBDMS ) ethers than triisopropylsilyl ( TIPS ) ethers. [ 12 ]
The application of hexafluorosilicic acid to a calcium-rich surface such as concrete will give that surface some resistance to acid attack. [ 13 ]
Calcium fluoride (CaF 2 ) is an insoluble solid that is acid resistant.
Some rare minerals, encountered either within volcanic or coal-fire fumaroles, are salts of the hexafluorosilicic acid. Examples include ammonium hexafluorosilicate that naturally occurs as two polymorphs: cryptohalite and bararite . [ 14 ] [ 15 ] [ 16 ]
Hexafluorosilicic acid can release hydrogen fluoride (HF) when evaporated, so it has similar risks. Inhalation of the vapors may cause lung edema . Like hydrogen fluoride, it attacks glass and stoneware . [ 17 ] The LD 50 value of hexafluorosilicic acid is 430 mg/kg. [ 6 ] | https://en.wikipedia.org/wiki/H2SiF6 |
Hydrogen telluride is the inorganic compound with the formula H 2 Te . A hydrogen chalcogenide and the simplest hydride of tellurium , it is a colorless gas. Although unstable in ambient air, the gas can exist long enough to be readily detected by the odour of rotting garlic at extremely low concentrations; or by the revolting odour of rotting leeks at somewhat higher concentrations. Most compounds with Te–H bonds ( tellurols ) are unstable with respect to loss of H 2 . H 2 Te is chemically and structurally similar to hydrogen selenide , both are acidic. The H–Te–H angle is about 90°. Volatile tellurium compounds often have unpleasant odours, reminiscent of decayed leeks or garlic. [ 2 ]
Electrolytic methods have been developed. [ 3 ]
H 2 Te can also be prepared by hydrolysis of the telluride derivatives of electropositive metals. [ 4 ] The typical hydrolysis is that of aluminium telluride :
Other salts of Te 2− such as MgTe and sodium telluride can also be used. Na 2 Te can be made by the reaction of Na and Te in anhydrous ammonia . [ 5 ] The intermediate in the hydrolysis, HTe − , can be isolated as salts as well. NaHTe can be made by reducing tellurium with NaBH 4 . [ 5 ]
Hydrogen telluride cannot be efficiently prepared from its constituent elements, in contrast to H 2 Se. [ 3 ]
H 2 Te is an endothermic compound, degrading to the elements at room temperature:
Light accelerates the decomposition. It is unstable in air, being oxidized to water and elemental tellurium: [ 6 ]
It is almost as acidic as phosphoric acid ( K a = 8.1×10 −3 ), having a K a value of about 2.3×10 −3 . [ 6 ] It reacts with many metals to form tellurides. [ 7 ] | https://en.wikipedia.org/wiki/H2Te |
Hydrogen ditelluride or ditellane is an unstable hydrogen dichalcogenide containing two tellurium atoms per molecule, with structure H−Te−Te−H or (TeH) 2 . Hydrogen ditelluride is interesting to theorists because its molecule is simple yet asymmetric (with no centre of symmetry ) and is predicted to be one of the easiest to detect parity violation , in which the left handed molecule has differing properties to the right handed one due to the effects of the weak force .
Hydrogen ditelluride can possibly be formed at the tellurium cathode in electrolysis in acid. [ 2 ] When electrolysed in alkaline solutions, a tellurium cathode produces ditelluride Te 2− 2 ions, as well as Te 2− and a red polytelluride. The greatest amount of ditelluride is made when pH is over 12. [ 3 ]
Apart from its speculative detection in electrolysis , ditellane has been detected in the gas phase produced from di- sec -butylditellane. [ 1 ] [ 4 ]
Hydrogen ditelluride has been investigated theoretically, with various properties predicted. The molecule is twisted with a C 2 symmetry . There are two enantiomers . Hydrogen ditelluride is one of the simplest possible unsymmetrical molecules; any simpler molecule will not have the required low symmetry. The equilibrium geometry (not counting zero point energy or vibrational energy) has bond lengths of 2.879 Å between the tellurium atoms and 1.678 Å between hydrogen and tellurium. The H−Te−Te angle is 94.93°. The angle of lowest energy between the two H−Te bonds (the dihedral angle between the H a −Te−Te and Te−Te−H b planes) is 89.32°. The trans configuration is higher in energy (3.71 kcal/mol), and the cis would be even higher (4.69 kcal/mol). [ 5 ]
Being chiral , the molecule is predicted to show evidence of parity violation , though this may get interference from stereomutation tunneling , where the P enantiomer and M enantiomer spontaneously convert into each other by quantum tunneling . The parity violation effect on energy comes about from virtual Z boson exchanges between the nucleus and electrons. [ 6 ] It is proportional to the cube of the atomic number, so is stronger in tellurium molecules than others higher up in the periodic table (O, S, Se). Because of parity violation, the energy of the two enantiomers differs, and is likely to be higher in this molecule than most molecules, so an effort is underway to observe this so-far undetected effect. The tunneling effect is reduced by higher masses, so that the deuterium form, D 2 Te 2 will show less tunneling. In a torsional vibrational mode, the molecule can twist back and forward storing energy. Seven different quantum vibration levels are predicted below the energy to jump to the other enantiomer. The levels are numbered v t = 0 up to 6. The sixth level is predicted to be split into two energy levels because of quantum tunneling. [ 7 ] The parity violation energy is calculated as 3 × 10 −9 cm −1 or 90 Hz. [ 7 ]
The different vibrational modes for H 2 Te are symmetrical stretch of H−Te , symmetrical bend of H−Te−Te , torsion, stretch Te−Te , asymmetrical stretch H−Te , asymmetrical bend of H−Te−Te . [ 7 ] The time to tunnel between enantiomers is only 0.6 ms for 1 H 2 Te 2 , but is 66 000 seconds (18 h 20 min) for the tritium isotopomer T 2 Te 2 . [ 7 ]
There are organic derivatives, in which the hydrogen is replaced by organic groups. One example is bis(2,4,6-tributylphenyl)ditellane. [ 8 ] Others are diphenyl ditelluride and 1,2-bis(cyclohexylmethyl)ditellane. A ligand -TeTeH is known in some transition metal complexes. IUPAC nomenclature calls this "ditellanido". | https://en.wikipedia.org/wiki/H2Te2 |
Tellurous acid is an inorganic compound with the formula H 2 TeO 3 . It is the oxoacid of tellurium(IV). [ 2 ] This compound is not well characterized. An alternative way of writing its formula is (HO) 2 TeO. In principle, tellurous acid would form by treatment of tellurium dioxide with water, that is by hydrolysis. The related conjugate base is well known in the form of several salts such as potassium hydrogen tellurite, KHTeO 3 .
In contrast to the analogous compound selenous acid , tellurous acid is only metastable. Most tellurite salts contain the TeO 2− 3 ion. Oxidation of its aqueous solution with hydrogen peroxide gives the tellurate ion. It is usually prepared as an aqueous solution where it acts as a weak acid. [ 1 ] [ 3 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2TeO3 |
Telluric acid , or more accurately orthotelluric acid , is a chemical compound with the formula Te(OH) 6 , often written as H 6 TeO 6 . It is a white crystalline solid made up of octahedral Te(OH) 6 molecules which persist in aqueous solution. [ 3 ] In the solid state, there are two forms, rhombohedral and monoclinic, and both contain octahedral Te(OH) 6 molecules, [ 4 ] containing one hexavalent tellurium (Te) atom in the +6 oxidation state, attached to six hydroxyl (–OH) groups, thus, it can be called tellurium(VI) hydroxide.
Telluric acid is a weak acid which is dibasic , forming tellurate salts with strong bases and hydrogen tellurate salts with weaker bases or upon hydrolysis of tellurates in water. [ 4 ] [ 5 ] It is used as tellurium-source in the synthesis of oxidation catalysts.
Telluric acid is formed by the oxidation of tellurium or tellurium dioxide with a powerful oxidising agent such as hydrogen peroxide , chromium trioxide or sodium peroxide . [ 4 ]
Crystallization of telluric acid solutions below 10 °C gives telluric acid tetrahydrate Te(OH) 6 ·4H 2 O . [ 3 ] It is an oxidising agent, as shown by the electrode potential for the reaction below, although it is kinetically slow in its oxidations. [ 4 ]
Chlorine , by comparison, is +1.36 V and selenous acid is +0.74 V in oxidizing conditions.
The anhydrous acid is stable in air at 100 °C but above this it dehydrates to form polymetatelluric acid, a white hygroscopic powder (approximate composition (H 2 TeO 4 ) 10 ), and allotelluric acid, an acid syrup of unknown structure (approximate composition 3·H 2 TeO 4 ·4H 2 O ). [ 6 ] [ 3 ]
Typical salts of the acid contains the anions [Te(O)(OH) 5 ] − and [Te(O) 2 (OH) 4 ] 2− . The presence of the tellurate ion TeO 2− 4 has been confirmed in the solid state structure of Rb 6 [TeO 5 ][TeO 4 ] . [ 7 ] Strong heating at over 300 °C produces the α crystalline modification of tellurium trioxide , α- TeO 3 . [ 5 ] Reaction with diazomethane gives the hexamethyl ester, Te(OCH 3 ) 6 . [ 3 ]
Telluric acid and its salts mostly contain hexacoordinate tellurium . [ 4 ] This is true even for salts such as magnesium tellurate, MgTeO 4 , which is isostructural with magnesium molybdate and contains TeO 6 octahedra. [ 4 ]
Metatelluric acid, H 2 TeO 4 , the tellurium analogue of sulfuric acid , H 2 SO 4 , is unknown. Allotelluric acid of approximate composition 3·H 2 TeO 4 ·4H 2 O , is not well characterised and may be a mixture of Te(OH) 6 and (H 2 TeO 4 ) n . [ 3 ]
Tellurous acid H 2 TeO 3 , containing tellurium in its +4 oxidation state, is known but not well characterised. Hydrogen telluride is an unstable gas that forms hydrotelluric acid upon addition to water. | https://en.wikipedia.org/wiki/H2TeO4 |
In chemistry , a tungstate is a compound that contains an oxyanion of tungsten or is a mixed oxide containing tungsten. The simplest tungstate ion is WO 2− 4 , "orthotungstate". [ 1 ] Many other tungstates belong to a large group of polyatomic ions that are termed polyoxometalates , ("POMs"), and specifically termed isopolyoxometalates as they contain, along with oxygen and maybe hydrogen, only one other element. Almost all useful tungsten ores are tungstates. [ 2 ]
Orthotungstates feature tetrahedral W(VI) centres with short W–O distances of 1.79 Å . Structurally, they resemble sulfates. Six-coordinate, octahedral tungsten dominates in the polyoxotungstates. In these compounds, the W–O distances are elongated. [ 1 ]
Some examples of tungstate ions: [ 3 ]
See the tungstates category for a list of tungstates.
Tungstates occur naturally with molybdates . Scheelite , the mineral calcium tungstate, often contains a small amount of molybdate. Wolframite is manganese and iron tungstate, and all these are valuable sources of tungsten. Powellite is a mineral form of calcium molybdate containing a small amount of tungstate.
Solutions of tungstates, like those of molybdates , give intensely blue solutions of complex tungstate(V,VI) analogous to the molybdenum blues when reduced by most organic materials. [ 1 ]
Unlike chromate , tungstate is not a good oxidizer , but like chromate, solutions of tungstate condense to give the isopolytungstates upon acidification. | https://en.wikipedia.org/wiki/H2W12O40 |
In chemistry , paratungstate refers to the anion with the formula [W 12 O 42 ] 12- and salts derived from this anion. The term also refers to protonated derivatives of this anion, including [H 2 W 12 O 42 ] 10- . Ammonium paratungstate (or APT), (NH 4 ) 10 [H 2 W 12 O 42 ] is a key intermediate in the purification of tungsten from its ores . [ 1 ]
The salt (NH 4 ) 10 (W 12 O 42 )·4H 2 O has been characterized by X-ray crystallography . [ 2 ]
The unprotonated anion [W 12 O 42 ] 12- has C 2h symmetry . | https://en.wikipedia.org/wiki/H2W12O42 |
Tungstic acid refers to hydrated forms of tungsten trioxide , WO 3 . Both a monohydrate (WO 3 ·H 2 O) and hemihydrate (WO 3 · 1 / 2 H 2 O) [ 1 ] are known. Molecular species akin to sulfuric acid , i.e. (HO) 2 WO 2 are not observed.
The solid-state structure of WO 3 ·H 2 O consists of layers of octahedrally coordinated WO 5 (H 2 O) units where 4 vertices are shared. [ 2 ] The dihydrate has the same layer structure with the extra H 2 O molecule intercalated . [ 2 ] The monohydrate is a yellow solid and insoluble in water. The classical name for this acid is 'acid of wolfram'. Salts of tungstic acid are tungstates .
The acid was discovered by Carl Wilhelm Scheele in 1781. [ 3 ]
Tungstic acid is obtained by the action of strong acids on solutions of alkali metallic tungstates. It may also be prepared from the reaction between hydrogen carbonate and sodium tungstate . It can also be obtained from pure tungsten by reaction with hydrogen peroxide . [ 4 ]
It is used as a mordant and a dye in textiles. | https://en.wikipedia.org/wiki/H2WO4 |
Xenic acid is a proposed noble gas compound with the chemical formula H 2 XeO 4 or XeO 2 (OH) 2 . It has not been isolated, and the published characterization data are ambiguous. [ 2 ]
Salts of xenic acid are called xenates , containing the HXeO − 4 anion, such as monosodium xenate . They tend to disproportionate into xenon gas and perxenates : [ 3 ]
The energy given off is sufficient to form ozone from diatomic oxygen:
Salts containing the deprotonated anion XeO 2− 4 are presently unknown. [ 3 ]
This inorganic compound –related article is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2XeO4 |
Sulfuric acid ( American spelling and the preferred IUPAC name ) or sulphuric acid ( Commonwealth spelling ), known in antiquity as oil of vitriol , is a mineral acid composed of the elements sulfur , oxygen , and hydrogen , with the molecular formula H 2 SO 4 . It is a colorless, odorless, and viscous liquid that is miscible with water. [ 7 ]
Pure sulfuric acid does not occur naturally due to its strong affinity to water vapor ; it is hygroscopic and readily absorbs water vapor from the air . [ 7 ] Concentrated sulfuric acid is a strong oxidant with powerful dehydrating properties, making it highly corrosive towards other materials, from rocks to metals. Phosphorus pentoxide is a notable exception in that it is not dehydrated by sulfuric acid but, to the contrary, dehydrates sulfuric acid to sulfur trioxide . Upon addition of sulfuric acid to water, a considerable amount of heat is released; thus, the reverse procedure of adding water to the acid is generally avoided since the heat released may boil the solution, spraying droplets of hot acid during the process. Upon contact with body tissue, sulfuric acid can cause severe acidic chemical burns and secondary thermal burns due to dehydration. [ 8 ] [ 9 ] Dilute sulfuric acid is substantially less hazardous without the oxidative and dehydrating properties; though, it is handled with care for its acidity.
Many methods for its production are known, including the contact process , the wet sulfuric acid process , and the lead chamber process . [ 10 ] Sulfuric acid is also a key substance in the chemical industry . It is most commonly used in fertilizer manufacture [ 11 ] but is also important in mineral processing , oil refining , wastewater treating , and chemical synthesis . It has a wide range of end applications, including in domestic acidic drain cleaners , [ 12 ] as an electrolyte in lead-acid batteries , as a dehydrating compound, and in various cleaning agents .
Sulfuric acid can be obtained by dissolving sulfur trioxide in water.
Although nearly 100% sulfuric acid solutions can be made, the subsequent loss of SO 3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade, which is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are: [ 13 ] [ 14 ]
"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process , chamber acid being the acid produced in the lead chamber itself (<70% to avoid contamination with nitrosylsulfuric acid ) and tower acid being the acid recovered from the bottom of the Glover tower. [ 13 ] [ 14 ] They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10 M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations ), is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher. [ 14 ]
Sulfuric acid contains not only H 2 SO 4 molecules, but is actually an equilibrium of many other chemical species, as it is shown in the table below.
Sulfuric acid is a colorless oily liquid, and has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, [ 16 ] and 98% sulfuric acid has a vapor pressure of <1 mmHg at 40 °C. [ 17 ]
In the solid state, sulfuric acid is a molecular solid that forms monoclinic crystals with nearly trigonal lattice parameters. The structure consists of layers parallel to the (010) plane, in which each molecule is connected by hydrogen bonds to two others. [ 3 ] Hydrates H 2 SO 4 · n H 2 O are known for n = 1, 2, 3, 4, 6.5, and 8, although most intermediate hydrates are stable against disproportionation . [ 18 ]
Anhydrous H 2 SO 4 is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity , a consequence of autoprotolysis , i.e. self- protonation : [ 15 ]
The equilibrium constant for autoprotolysis (25 °C) is: [ 15 ]
The corresponding equilibrium constant for water , K w is 10 −14 , a factor of 10 10 (10 billion) smaller.
In spite of the viscosity of the acid, the effective conductivities of the H 3 SO + 4 and HSO − 4 ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.
The hydration reaction of sulfuric acid is highly exothermic . [ 19 ]
As indicated by its acid dissociation constant , sulfuric acid is a strong acid:
The product of this ionization is HSO − 4 , the bisulfate anion. Bisulfate is a far weaker acid:
The product of this second dissociation is SO 2− 4 , the sulfate anion.
Concentrated sulfuric acid has a powerful dehydrating property, removing water ( H 2 O ) from other chemical compounds such as table sugar ( sucrose ) and other carbohydrates , to produce carbon , steam , and heat. Dehydration of table sugar (sucrose) is a common laboratory demonstration. [ 21 ] The sugar darkens as carbon is formed, and a rigid column of black, porous carbon called a carbon snake may emerge. [ 22 ]
Similarly, mixing starch into concentrated sulfuric acid gives elemental carbon and water. The effect of this can also be seen when concentrated sulfuric acid is spilled on paper. Paper is composed of cellulose , a polysaccharide related to starch. The cellulose reacts to give a burnt appearance in which the carbon appears much like soot that results from fire.
Although less dramatic, the action of the acid on cotton , even in diluted form, destroys the fabric.
The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystals change into white powder as water is removed.
Sulfuric acid reacts with most bases to give the corresponding sulfate or bisulfate.
Aluminium sulfate , also known as paper maker's alum, is made by treating bauxite with sulfuric acid:
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate , for example, displaces acetic acid , CH 3 COOH , and forms sodium bisulfate :
Similarly, treating potassium nitrate with sulfuric acid produces nitric acid . Sulfuric acid reacts with sodium chloride , and gives hydrogen chloride gas and sodium bisulfate :
When combined with nitric acid , sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO + 2 , which is important in nitration reactions involving electrophilic aromatic substitution . This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols.
When allowed to react with superacids , sulfuric acid can act as a base and can be protonated, forming the [H 3 SO 4 ] + ion. Salts of [H 3 SO 4 ] + have been prepared (e.g. trihydroxyoxosulfonium hexafluoroantimonate(V) [H 3 SO 4 ] + [SbF 6 ] − ) using the following reaction in liquid HF :
The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply fluoroantimonic acid , however, has met with failure, as pure sulfuric acid undergoes self-ionization to give [H 3 O] + ions:
which prevents the conversion of H 2 SO 4 to [H 3 SO 4 ] + by the HF/ SbF 5 system. [ 23 ]
Even diluted sulfuric acid reacts with many metals via a single displacement reaction, like other typical acids , producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series ) such as iron , aluminium , zinc , manganese , magnesium , and nickel .
Concentrated sulfuric acid can serve as an oxidizing agent , releasing sulfur dioxide: [ 8 ]
Lead and tungsten , however, are resistant to sulfuric acid.
Hot concentrated sulfuric acid oxidizes carbon [ 24 ] (as bituminous coal ) and sulfur :
Benzene and many derivatives undergo electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids : [ 25 ]
Sulfuric acid can be used to produce hydrogen from water :
The compounds of sulfur and iodine are recovered and reused, hence the process is called the sulfur–iodine cycle . This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. The sulfur–iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy . It is an alternative to electrolysis , and does not require hydrocarbons like current methods of steam reforming . But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it. [ 26 ] [ 27 ]
Sulfuric acid is rarely encountered naturally on Earth in anhydrous form, due to its great affinity for water . Dilute sulfuric acid is a constituent of acid rain , which is formed by atmospheric oxidation of sulfur dioxide in the presence of water —i.e. oxidation of sulfurous acid . When sulfur-containing fuels such as coal or oil are burned, sulfur dioxide is the main byproduct (besides the chief products carbon oxides and water).
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as pyrite :
The resulting highly acidic water is called acid mine drainage (AMD) or acid rock drainage (ARD).
The Fe 2+ can be further oxidized to Fe 3+ :
The Fe 3+ produced can be precipitated as the hydroxide or hydrous iron oxide :
The iron(III) ion (" ferric iron ") can also oxidize pyrite:
FeS 2 (s) + 14 Fe 3+ + 8 H 2 O → 15 Fe 2+ + 2 SO 2− 4 + 16 H +
When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process.
ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals.
Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales ) concentrates sulfuric acid in cell vacuoles. [ 28 ]
In the stratosphere , the atmosphere's second layer that is generally between 10–50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical : [ 29 ]
Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer. [ 29 ]
The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. [ 30 ] Sulfuric acid ice has been detected on Jupiter 's moon Europa , where it forms when sulfur ions from Jupiter's magnetosphere implant into the icy surface. [ 31 ]
Sulfuric acid is produced from sulfur , oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).
In the first step, sulfur is burned to produce sulfur dioxide.
The sulfur dioxide is oxidized to sulfur trioxide by oxygen in the presence of a vanadium(V) oxide catalyst . This reaction is reversible and the formation of the sulfur trioxide is exothermic.
The sulfur trioxide is absorbed into 97–98% H 2 SO 4 to form oleum ( H 2 S 2 O 7 ), also known as fuming sulfuric acid or pyrosulphuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.
Directly dissolving SO 3 in water, called the " wet sulfuric acid process ", is rarely practiced because the reaction is extremely exothermic, resulting in a hot aerosol of sulfuric acid that requires condensation and separation.
In the first step, sulfur is burned to produce sulfur dioxide:
or, alternatively, hydrogen sulfide ( H 2 S ) gas is incinerated to SO 2 gas:
The sulfur dioxide then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst .
The sulfur trioxide is hydrated into sulfuric acid H 2 SO 4 :
The last step is the condensation of the sulfuric acid to liquid 97–98% H 2 SO 4 :
Burning sulfur together with saltpeter ( potassium nitrate , KNO 3 ), in the presence of steam, has been used historically. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid.
Prior to 1900, most sulfuric acid was manufactured by the lead chamber process . [ 32 ] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.
A wide variety of laboratory syntheses are known, and typically begin from sulfur dioxide or an equivalent salt . In the metabisulfite method, hydrochloric acid reacts with metabisulfite to produce sulfur dioxide vapors. The gas is bubbled through nitric acid , which will release brown/red vapors of nitrogen dioxide as the reaction proceeds. The completion of the reaction is indicated by the ceasing of the fumes. This method conveniently does not produce an inseparable mist. [ citation needed ]
Alternatively, dissolving sulfur dioxide in an aqueous solution of an oxidizing metal salt such as copper(II) or iron(III) chloride: [ citation needed ]
Two less well-known laboratory methods of producing sulfuric acid, albeit in dilute form and requiring some extra effort in purification, rely on electrolysis . A solution of copper(II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper at cathode and oxygen gas at the anode. The solution of dilute sulfuric acid indicates completion of the reaction when it turns from blue to clear (production of hydrogen at cathode is another sign): [ citation needed ]
More costly, dangerous, and troublesome is the electrobromine method, which employs a mixture of sulfur , water, and hydrobromic acid as the electrolyte. The sulfur is pushed to bottom of container under the acid solution. Then the copper cathode and platinum/graphite anode are used with the cathode near the surface and the anode is positioned at the bottom of the electrolyte to apply the current. This may take longer and emits toxic bromine /sulfur-bromide vapors, but the reactant acid is recyclable. Overall, only the sulfur and water are converted to sulfuric acid and hydrogen (omitting losses of acid as vapors): [ citation needed ]
Sulfuric acid is a very important commodity chemical, and a nation's sulfuric acid production was as recently as 2002 believed to be a good indicator of its industrial strength. [ 33 ] World production in the year 2004 was about 180 million tonnes , with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%. [ 34 ] World production in 2022 was estimated at about 260 million tonnes. [ 35 ]
As of the late 20th century, most of the produced amount (≈60%) was consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze , as well as in various processes such as oil well acidicizing, aluminium reduction, paper sizing, and water treatment. About 6% of uses are related to pigments and include paints, enamels , printing inks, coated fabrics and paper, while the rest is dispersed into a multitude of applications such as production of explosives, cellophane , acetate and viscose textiles, lubricants, non-ferrous metals , and batteries. [ 36 ]
The dominant use for sulfuric acid is in the "wet method" for the production of phosphoric acid , used for manufacture of phosphate fertilizers . In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite , though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate , hydrogen fluoride (HF) and phosphoric acid . The HF is removed as hydrofluoric acid . The overall process can be represented as:
Ammonium sulfate , an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.
Sulfuric acid is also important in the manufacture of dyestuffs solutions.
Sulfuric acid is used in steelmaking and other metallurgical industries as a pickling agent for removal of rust and fouling . [ 37 ] Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid [ clarification needed ] with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide ( SO 2 ) and sulfur trioxide ( SO 3 ) which are then used to manufacture "new" sulfuric acid.
Hydrogen peroxide ( H 2 O 2 ) can be added to sulfuric acid to produce piranha solution , a powerful but potentially hazardous cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used in the microelectronics industry, and also in laboratory settings to clean glassware.
Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam , used for making nylon . It is used for making hydrochloric acid from salt via the Mannheim process . Much H 2 SO 4 is used in petroleum refining , for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane , a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidizing agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.
Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator):
At anode :
At cathode :
Overall:
Sulfuric acid at high concentrations is frequently the major ingredient in domestic acidic drain cleaners [ 12 ] which are used to remove lipids , hair , tissue paper , etc. Similar to their alkaline versions , such drain openers can dissolve fats and proteins via hydrolysis . Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.
The study of vitriols (hydrated sulfates of various metals forming glassy minerals from which sulfuric acid can be derived) began in ancient times . Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis , in the treatise Phisica et Mystica , and the Leyden papyrus X . [ 38 ] Medieval Islamic alchemists like the authors writing under the name of Jabir ibn Hayyan (died c. 806 – c. 816 AD, known in Latin as Geber), Abu Bakr al-Razi (865 – 925 AD, known in Latin as Rhazes), Ibn Sina (980 – 1037 AD, known in Latin as Avicenna), and Muhammad ibn Ibrahim al-Watwat (1234 – 1318 AD) included vitriol in their mineral classification lists. [ 39 ]
The Jabirian authors and al-Razi experimented extensively with the distillation of various substances, including vitriols. [ 40 ] In one recipe recorded in his Kitāb al-Asrār ( 'Book of Secrets' ), al-Razi may have created sulfuric acid without being aware of it: [ 41 ]
Take white (Yemeni) alum , dissolve it and purify it by filtration. Then distil (green?) vitriol with copper-green (the acetate), and mix (the distillate) with the filtered solution of the purified alum, afterwards let it solidify (or crystallise) in the glass beaker. You will get the best qalqadis (white alum) that may be had. [ 42 ]
In an anonymous Latin work variously attributed to Aristotle (under the title Liber Aristotilis , 'Book of Aristotle'), [ 43 ] to al-Razi (under the title Lumen luminum magnum , 'Great Light of Lights'), or to Ibn Sina, [ 44 ] the author speaks of an 'oil' ( oleum ) obtained through the distillation of iron(II) sulfate (green vitriol), which was likely 'oil of vitriol' or sulfuric acid. [ 45 ] The work refers multiple times to Jabir ibn Hayyan's Seventy Books ( Liber de septuaginta ), one of the few Arabic Jabir works that were translated into Latin. [ 46 ] The author of the version attributed to al-Razi also refers to the Liber de septuaginta as his own work, showing that he erroneously believed the Liber de septuaginta to be a work by al-Razi. [ 47 ] There are several indications that the anonymous work was an original composition in Latin, [ 48 ] although according to one manuscript it was translated by a certain Raymond of Marseilles, meaning that it may also have been a translation from the Arabic. [ 49 ]
According to Ahmad Y. al-Hassan , three recipes for sulfuric acid occur in an anonymous Garshuni manuscript containing a compilation taken from several authors and dating from before c. 1100 AD . [ 50 ] One of them runs as follows:
The water of vitriol and sulphur which is used to irrigate the drugs: yellow vitriol three parts, yellow sulphur one part, grind them and distil them in the manner of rose-water. [ 51 ]
A recipe for the preparation of sulfuric acid is mentioned in Risālat Jaʿfar al-Sādiq fī ʿilm al-ṣanʿa , an Arabic treatise falsely attributed to the Shi'i Imam Ja'far al-Sadiq (died 765). Julius Ruska dated this treatise to the 13th century, but according to Ahmad Y. al-Hassan it likely dates from an earlier period: [ 52 ]
Then distil green vitriol in a cucurbit and alembic, using medium fire; take what you obtain from the distillate, and you will find it clear with a greenish tint. [ 51 ]
Sulfuric acid was called 'oil of vitriol' by medieval European alchemists because it was prepared by roasting iron(II) sulfate or green vitriol in an iron retort . The first allusions to it in works that are European in origin appear in the thirteenth century AD, as for example in the works of Vincent of Beauvais , in the Compositum de Compositis ascribed to Albertus Magnus , and in pseudo-Geber 's Summa perfectionis . [ 53 ]
A method of producing oleum sulphuris per campanam, or "oil of sulfur by the bell", was known by the 16th century: it involved burning sulfur under a glass bell in moist weather (or, later, under a moistened bell). However, it was very inefficient (according to Gesner , 5 pounds (2.3 kg) of sulfur converted into less than 1 ounce (0.03 kg) of acid), and the resulting product was contaminated by sulfurous acid (or rather, solution of sulfur dioxide ) so most alchemists (including, for example, Isaac Newton) didn't consider it equivalent with the "oil of vitriol".
In the 17th century, Johann Rudolf Glauber discovered that adding saltpeter ( potassium nitrate , KNO 3 ) significantly improves the output, also replacing moisture with steam. As saltpeter decomposes, it oxidizes the sulfur to SO 3 , which combines with water to produce sulfuric acid. In 1736, Joshua Ward , a London pharmacist, used this method to begin the first large-scale production of sulfuric acid.
In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead -lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries with a purity of 62% and a conversion of 75%. [ 4 ]
Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS 2 ) was heated in air to yield iron(II) sulfate, FeSO 4 , which was oxidized by further heating in air to form iron(III) sulfate , Fe 2 (SO 4 ) 3 , which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid. [ 4 ]
In 1831, British vinegar merchant Peregrine Phillips patented the contact process , which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method. [ 33 ]
In the early to mid 19th century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim , Northern Ireland , where it was used as a bleach for linen. Early bleaching of linen was done using lactic acid from sour milk but this was a slow process and the use of vitriol sped up the bleaching process. [ 54 ]
Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations . In common with other corrosive acids and alkali , it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues , such as skin and flesh . In addition, it exhibits a strong dehydrating property on carbohydrates , liberating extra heat and causing secondary thermal burns . [ 8 ] [ 9 ] Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes . If ingested, it damages internal organs irreversibly and may even be fatal. [ 7 ] Personal protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials. [ 8 ] Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids , such as hydrochloric acid and nitric acid .
Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time.
The standard first aid treatment for acid spills on the skin is, as for other corrosive agents , irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.
Preparation of diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid is usually added to water and not the other way around. A saying used to remember this is "Do like you oughta, add the acid to the water". [ 55 ] [ better source needed ] [ 56 ] Water has a higher heat capacity than the acid, and so a vessel of cold water will absorb heat as acid is added.
Also, because the acid is denser than water, it sinks to the bottom. Heat is generated at the interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection , which cools the interface, and prevents boiling of either acid or water.
In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol , or worse, an explosion .
Preparation of solutions greater than 6 M (35%) in concentration is dangerous, unless the acid is added slowly enough to allow the mixture sufficient time to cool. Otherwise, the heat produced may be sufficient to boil the mixture. Efficient mechanical stirring and external cooling (such as an ice bath) are essential.
Reaction rates double for about every 10-degree Celsius increase in temperature . [ 57 ] Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction.
On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.
Sulfuric acid is non-flammable.
The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent. [ 58 ] In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m 3 : limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination , loss of axons and gliosis .
International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988 , which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances. [ 59 ] | https://en.wikipedia.org/wiki/H2o4s |
H 2 PTM , or Hypertext Hypermedia Products Tools and Methods , is an international conference on hypermedia . The first congress was held in Paris in 1989 and was organized by the Paragraphe Lab at the University of Paris VIII . [ 1 ]
In 2009, H 2 PTM celebrated its 20th anniversary with a conference held in Paris in November.
This article about a computer conference is a stub . You can help Wikipedia by expanding it . | https://en.wikipedia.org/wiki/H2ptm |
Arsine ( IUPAC name: arsane ) is an inorganic compound with the formula As H 3 . This flammable, pyrophoric , and highly toxic pnictogen hydride gas is one of the simplest compounds of arsenic . [ 4 ] Despite its lethality, it finds some applications in the semiconductor industry and for the synthesis of organoarsenic compounds . The term arsine is commonly used to describe a class of organoarsenic compounds of the formula AsH 3− x R x , where R = aryl or alkyl . For example, As(C 6 H 5 ) 3 , called triphenylarsine , is referred to as "an arsine".
In its standard state arsine is a colorless, denser-than-air gas that is slightly soluble in water (2% at 20 °C) [ 1 ] and in many organic solvents as well. [ citation needed ] Arsine itself is odorless, [ 5 ] but it oxidizes in air and this creates a slight garlic or fish-like scent when the compound is present above 0.5 ppm . [ 6 ] This compound is kinetically stable: at room temperature it decomposes only slowly. At temperatures of ca. 230 °C, decomposition to arsenic and hydrogen is sufficiently rapid to be the basis of the Marsh test for arsenic presence. Similar to stibine , the decomposition of arsine is autocatalytic, as the arsenic freed during the reaction acts as a catalyst for the same reaction. [ 7 ] Several other factors, such as humidity , presence of light and certain catalysts (namely alumina ) facilitate the rate of decomposition. [ 8 ]
AsH 3 is a trigonal pyramidal molecule with H–As–H angles of 91.8° and three equivalent As–H bonds, each of 1.519 Å length. [ 9 ]
AsH 3 is generally prepared by the reaction of As 3+ sources with H − equivalents. [ 10 ]
As reported in 1775, Carl Scheele reduced arsenic(III) oxide with zinc in the presence of acid. [ 11 ] This reaction is a prelude to the Marsh test .
Alternatively, sources of As 3− react with protonic reagents to also produce this gas. Zinc arsenide and sodium arsenide are suitable precursors: [ 12 ]
The understanding of the chemical properties of AsH 3 is well developed and can be anticipated based on an average of the behavior of pnictogen counterparts, such as PH 3 and SbH 3 .
Typical for a heavy hydride (e.g., SbH 3 , H 2 Te , SnH 4 ), AsH 3 is unstable with respect to its elements. In other words, it is stable kinetically but not thermodynamically.
This decomposition reaction is the basis of the Marsh test, which detects elemental As.
Continuing the analogy to SbH 3 , AsH 3 is readily oxidized by concentrated O 2 or the dilute O 2 concentration in air:
Arsine will react violently in presence of strong oxidizing agents, such as potassium permanganate , sodium hypochlorite , or nitric acid . [ 8 ]
AsH 3 is used as a precursor to metal complexes of "naked" (or "nearly naked") arsenic. An example is the dimanganese species [(C 5 H 5 )Mn(CO) 2 ] 2 AsH, wherein the Mn 2 AsH core is planar. [ 13 ]
A characteristic test for arsenic involves the reaction of AsH 3 with Ag + , called the Gutzeit test for arsenic. [ 14 ] Although this test has become obsolete in analytical chemistry , the underlying reactions further illustrate the affinity of AsH 3 for "soft" metal cations. In the Gutzeit test, AsH 3 is generated by reduction of aqueous arsenic compounds, typically arsenites , with Zn in the presence of H 2 SO 4 . The evolved gaseous AsH 3 is then exposed to AgNO 3 either as powder or as a solution. With solid AgNO 3 , AsH 3 reacts to produce yellow Ag 4 AsNO 3 , whereas AsH 3 reacts with a solution of AgNO 3 to give black Ag 3 As.
The acidic properties of the As–H bond are often exploited. Thus, AsH 3 can be deprotonated:
Upon reaction with the aluminium trialkyls, AsH 3 gives the trimeric [R 2 AlAsH 2 ] 3 , where R = (CH 3 ) 3 C. [ 15 ] This reaction is relevant to the mechanism by which GaAs forms from AsH 3 (see below).
AsH 3 is generally considered non-basic, but it can be protonated by superacids to give isolable salts of the tetrahedral species [AsH 4 ] + . [ 16 ]
Reactions of arsine with the halogens ( fluorine and chlorine ) or some of their compounds, such as nitrogen trichloride , are extremely dangerous and can result in explosions. [ 8 ]
In contrast to the behavior of PH 3 , AsH 3 does not form stable chains, although diarsine (or diarsane) H 2 As–AsH 2 , and even triarsane H 2 As–As(H)–AsH 2 have been detected. The diarsine is unstable above −100 °C.
AsH 3 is used in the synthesis of semiconducting materials related to microelectronics and solid-state lasers . Related to phosphorus , arsenic is an n-dopant for silicon and germanium. [ 8 ] More importantly, AsH 3 is used to make the semiconductor GaAs by chemical vapor deposition (CVD) at 700–900 °C:
For microelectronic applications, arsine can be provided by a sub-atmospheric gas source (a source that supplies less than atmospheric pressure). In this type of gas package, the arsine is adsorbed on a solid microporous adsorbent inside a gas cylinder. This method allows the gas to be stored without pressure, significantly reducing the risk of an arsine gas leak from the cylinder. With this apparatus, arsine is obtained by applying vacuum to the gas cylinder valve outlet. For semiconductor manufacturing , this method is feasible, as processes such as ion implantation operate under high vacuum.
Since before WWII AsH 3 was proposed as a possible chemical warfare weapon. The gas is colorless, almost odorless, and 2.5 times denser than air, as required for a blanketing effect sought in chemical warfare. It is also lethal in concentrations far lower than those required to smell its garlic -like scent. In spite of these characteristics, arsine was never officially used as a weapon, because of its high flammability and its lower efficacy when compared to the non-flammable alternative phosgene . On the other hand, several organic compounds based on arsine, such as lewisite (β-chlorovinyldichloroarsine), adamsite (diphenylaminechloroarsine), Clark 1 ( diphenylchloroarsine ) and Clark 2 ( diphenylcyanoarsine ) have been effectively developed for use in chemical warfare. [ 17 ]
AsH 3 is well known in forensic science because it is a chemical intermediate in the detection of arsenic poisoning. The old (but extremely sensitive) Marsh test generates AsH 3 in the presence of arsenic. [ 4 ] This procedure, published in 1836 by James Marsh , [ 18 ] is based upon treating an As-containing sample of a victim's body (typically the stomach contents) with As-free zinc and dilute sulfuric acid : if the sample contains arsenic, gaseous arsine will form. The gas is swept into a glass tube and decomposed by means of heating around 250–300 °C. The presence of As is indicated by formation of a deposit in the heated part of the equipment. On the other hand, the appearance of a black mirror deposit in the cool part of the equipment indicates the presence of antimony (the highly unstable SbH 3 decomposes even at low temperatures).
The Marsh test was widely used by the end of the 19th century and the start of the 20th; nowadays more sophisticated techniques such as atomic spectroscopy , inductively coupled plasma , and x-ray fluorescence analysis are employed in the forensic field. Though neutron activation analysis was used to detect trace levels of arsenic in the mid 20th century, it has since fallen out of use in modern forensics.
The toxicity of arsine is distinct from that of other arsenic compounds. The main route of exposure is by inhalation, although poisoning after skin contact has also been described. Arsine attacks hemoglobin in the red blood cells , causing them to be destroyed by the body. [ 19 ] [ 20 ]
The first signs of exposure, which can take several hours to become apparent, are headaches , vertigo , and nausea , followed by the symptoms of haemolytic anaemia (high levels of unconjugated bilirubin ), haemoglobinuria and nephropathy . In severe cases, the damage to the kidneys can be long-lasting. [ 1 ]
Exposure to arsine concentrations of 250 ppm is rapidly fatal: concentrations of 25–30 ppm are fatal for 30 min exposure, and concentrations of 10 ppm can be fatal at longer exposure times. [ 3 ] Symptoms of poisoning appear after exposure to concentrations of 0.5 ppm. There is little information on the chronic toxicity of arsine, although it is reasonable to assume that, in common with other arsenic compounds, a long-term exposure could lead to arsenicosis . [ citation needed ]
Arsine can cause pneumonia in two different ways either the "extensive edema of the acute stage may become diffusely infiltrated with polymorphonuclear leucocytes, and the edema may change to ringed with leucocytes, their epithelium degenerated, their walls infiltrated, and each bronchiole the center of a small focus or nodule of pneumonic consolidation", and In the second Case "the areas involved are practically always the anterior tips of the middle and upper lobes, while the posterior portions of these lobes and the whole of the lower lobes present an air-containing and emphysematous condition, sometimes with slight congestion, sometimes with none." which can result in death. [ 21 ]
It is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities. [ 22 ] | https://en.wikipedia.org/wiki/H3As |
Arsenous acid (or arsorous acid ) is the inorganic compound with the formula H 3 AsO 3 . It is known to occur in aqueous solutions , but it has not been isolated as a pure material, although this fact does not detract from the significance of As(OH) 3 . [ 2 ]
As(OH) 3 is a pyramidal molecule consisting of three hydroxyl groups bonded to arsenic. The 1 H NMR spectrum of arsenous acid solutions consists of a single signal consistent with the molecule's high symmetry. [ 3 ] In contrast, the nominally related phosphorous acid H 3 PO 3 adopts the structure HPO(OH) 2 . The structural analogue of arsenous acid (P(OH) 3 ) is a very minor equilibrium component of such solutions. The differing behaviors of the As and P compounds reflect a trend whereby high oxidation states are more stable for lighter members of main group elements than their heavier congeners. [ 4 ]
One tautomer of arsenous acid is HAsO(OH) 2 , which is called arsonic acid . It has not been isolated or well-characterized.
The preparation of As(OH) 3 involves a slow hydrolysis of arsenic trioxide in water. Addition of base converts arsenous acid to the arsenite ions [AsO(OH) 2 ] − , [AsO 2 (OH)] 2− , and [AsO 3 ] 3− .
With its first p K a being 9.2, As(OH) 3 is a weak acid. [ 4 ] Reactions attributed to aqueous arsenic trioxide are due to arsenous acid and its conjugate bases.
Like arsenic trioxide, arsenous acid is sometimes amphoteric . For example, it reacts with hydrochloric, hydrobromic, and hydroiodic acids to produce arsenic trichloride, tribromide, and triiodide.
Reaction of arsenous acid with methyl iodide gives methylarsonic acid . This historically significant conversion is the Meyer reaction : [ 5 ]
Alkylation occurs at arsenic, and the oxidation state of arsenic increases from +3 to +5.
Arsenic-containing compounds are highly toxic and carcinogenic . The anhydride form of arsenous acid, arsenic trioxide , is used as a herbicide , pesticide , and rodenticide . | https://en.wikipedia.org/wiki/H3AsO3 |
6 mg/kg (rabbit, oral)
Arsenic acid or arsoric acid is the chemical compound with the formula H 3 AsO 4 . More descriptively written as AsO(OH) 3 , this colorless acid is the arsenic analogue of phosphoric acid . Arsenate and phosphate salts behave very similarly. Arsenic acid as such has not been isolated, but is only found in solution, where it is largely ionized. Its hemihydrate form ( 2H 3 AsO 4 ·H 2 O ) does form stable crystals. Crystalline samples dehydrate with condensation at 100 °C. [ 3 ]
It is a tetrahedral species of idealized symmetry C 3v with As–O bond lengths ranging from 1.66 to 1.71 Å. [ 4 ]
Being a triprotic acid, its acidity is described by three equilibria:
These p K a values are close to those for phosphoric acid . The highly basic arsenate ion ( AsO 3− 4 ) is the product of the third ionization. Unlike phosphoric acid, arsenic acid is an oxidizer, as illustrated by its ability to convert iodide to iodine .
Arsenic acid is prepared by treating arsenic trioxide with concentrated nitric acid . Dinitrogen trioxide is produced as a by-product. [ 5 ]
The resulting solution is cooled to give colourless crystals of the hemihydrate H 3 AsO 4 ·0.5H 2 O (or 2H 3 AsO 4 ·H 2 O ), although the dihydrate H 3 AsO 4 ·2H 2 O is produced when crystallisation occurs at lower temperatures. [ 5 ]
Arsenic acid is slowly formed when arsenic pentoxide is dissolved in water, and when meta - or pyroarsenic acid ( H 4 As 2 O 7 ) is treated with cold water. Arsenic acid can also be prepared directly from elemental arsenic by moistening it and treating with ozone .
Commercial applications of arsenic acid are limited by its toxicity. It is a precursor to a variety of pesticides. It has found occasional use as a wood preservative , a broad-spectrum biocide , a finishing agent for glass and metal, and a reagent in the synthesis of some dyestuffs and organic arsenic compounds. [ 6 ]
Arsenic acid is extremely toxic and carcinogenic, like all arsenic compounds. It is also corrosive . The LD 50 in rabbits is 6 mg/kg (0.006 g/kg). [ 7 ] | https://en.wikipedia.org/wiki/H3AsO4 |
Borane carbonyl is the inorganic compound with the formula H 3 B C O . This colorless gas is the adduct of borane and carbon monoxide . It is usually prepared by combining borane-ether complexes and CO. The compound is mainly of theoretical and pedagogical interest. [ 2 ]
The structure of the molecule of borane carbonyl is H 3 B − −C≡O + . The B−C≡O linkage is linear . The coordination geometry around the boron atom is tetrahedral . The bond distances are 114.0 pm for the C≡O bond, 152.9 pm for the C−B bond, and 119.4 pm for the B−H bonds. The H−B−H bond angle is 113.7°. The C≡O vibrational band is at 2164.7 cm −1 , around 22 cm −1 higher than that of free CO . [ 3 ]
Borane carbonyl has an enthalpy of vaporization of 19.7 kJ / mol (4750 cal /mol). [ 4 ] It has electronic state 1 A 1 and point group symmetry C 3v . [ 5 ]
Borane carbonyl was discovered in 1937 by reacting diborane with excess carbon monoxide, with the equation:
The reaction quickly reaches equilibrium at 100°C, but at room temperature, the reverse reaction is slow enough to isolate borane carbonyl. This reaction is performed at high pressures, typically with a maximum pressure observed of 1000 to 1600 psi (68.95 to 110.32 bar ). [ 6 ] It can also be performed at atmospheric pressure, with ethers as a catalyst . [ 7 ] [ 8 ]
A more recent synthesis of borane carbonyl involves slowly bubbling carbon monoxide through a 1 M H 3 B− THF solution. The resulting gas stream can be condensed and subsequently bubbled through ethanolic potassium hydroxide to produce the boranocarbonate anion ( [H 3 BCO 2 ] 2− or H 3 B − −CO − 2 ). [ 8 ] | https://en.wikipedia.org/wiki/H3BCO |
Ammonia borane (also systematically named ammoniotrihydroborate [ citation needed ] ), also called borazane , is the chemical compound with the formula H 3 NBH 3 . The colourless or white solid is the simplest molecular boron - nitrogen - hydride compound. It has attracted attention as a source for hydrogen fuel, but is otherwise primarily of academic interest.
Reaction of diborane with ammonia mainly gives the diammoniate salt [H 2 B(NH 3 ) 2 ] + [BH 4 ] − (diammoniodihydroboronium tetrahydroborate). Ammonia borane is the main product when an adduct of borane is employed in place of diborane: [ 5 ]
It can also be synthesized from sodium borohydride . [ 6 ] [ 7 ] [ 8 ]
The molecule adopts a structure similar to that of ethane , with which it is isoelectronic . The B−N distance is 1.58(2) Å. The B−H and N−H distances are 1.15 and 0.96 Å, respectively. Its similarity to ethane is tenuous since ammonia borane is a solid and ethane is a gas: their melting points differing by 284 °C. This difference is consistent with the highly polar nature of ammonia borane. The H atoms attached to boron are hydridic (negatively charged) and those attached to nitrogen are acidic (positively charged). [ 9 ]
The structure of the solid indicates a close association of the N H and the B H centers. The closest H−H distance is 1.990 Å, which can be compared with the H−H bonding distance of 0.74 Å. This interaction is called a dihydrogen bond . [ 10 ] [ 11 ] The original crystallographic analysis of this compound reversed the assignments of B and N. The updated structure was arrived at with improved data using the technique of neutron diffraction that allowed the hydrogen atoms to be located with greater precision.
Ammonia borane has been suggested as a storage medium for hydrogen , e.g. for when the gas is used to fuel motor vehicles. It can be made to release hydrogen on heating, being polymerized first to (NH 2 BH 2 ) n , then to (NHBH) n , [ 15 ] which ultimately decomposes to boron nitride (BN) at temperatures above 1000 °C. [ 16 ] It is more hydrogen-dense than liquid hydrogen and also able to exist at normal temperatures and pressures. [ 17 ]
Ammonia borane finds some use in organic synthesis as an air-stable derivative of diborane. [ 18 ] It can be used as a reducing agent in transfer hydrogenation reactions, often in the presence of a transition metal catalyst. [ 19 ]
Many analogues have been prepared from primary, secondary, and even tertiary amines :
The first amine adduct of borane was derived from trimethylamine . Borane tert-butylamine complex is prepared by the reaction of sodium borohydride with t-butylammonium chloride. Generally adduct are more robust with more basic amines. Variations are also possible for the boron component, although primary and secondary boranes are less common. [ 8 ] | https://en.wikipedia.org/wiki/H3BNH3 |
Boric acid , more specifically orthoboric acid , is a compound of boron , oxygen , and hydrogen with formula B(OH) 3 . It may also be called hydrogen orthoborate , trihydroxidoboron or boracic acid . [ 3 ] It is usually encountered as colorless crystals or a white powder, that dissolves in water , and occurs in nature as the mineral sassolite . It is a weak acid that yields various borate anions and salts , and can react with alcohols to form borate esters .
Boric acid is often used as an antiseptic , insecticide , flame retardant , neutron absorber , or precursor to other boron compounds.
The term "boric acid" is also used generically for any oxyacid of boron, such as metaboric acid HBO 2 and tetraboric acid H 2 B 4 O 7 .
Orthoboric acid was first prepared by Wilhelm Homberg (1652–1715) from borax , by the action of mineral acids, and was given the name sal sedativum Hombergi ("sedative salt of Homberg"). However, boric acid and borates have been used since the time of the ancient Greeks for cleaning, preserving food, and other uses. [ 4 ]
The three oxygen atoms form a trigonal planar geometry around the boron. The B-O bond length is 136 pm, and the O-H is 97 pm. The molecular point group is C 3h . [ 5 ]
Two crystalline forms of orthoboric acid are known: triclinic with space group P 1 , and trigonal with space group P3 2 . The former is the most common; the second, which is a bit more stable thermodynamically, can be obtained with a special preparation method. [ 6 ]
The triclinic form of boric acid consists of layers of B(OH) 3 molecules held together by hydrogen bonds with an O...O separation of 272 pm. The distance between two adjacent layers is 318 pm. [ 7 ] While the layers of the triclinic phase are nearly trigonal with γ = 119.76° , a = 701.87 pm , and b = 703.5 pm (compared to a = 704.53(4) pm for the trigonal form), the stacking of the layers is somewhat offset in the triclinic phase, with α = 92.49° and β = 101.46° . The triclinic phase has c = 634.72 pm and the trigonal one has a = 956.08(7) pm . [ 8 ] [ 9 ]
Boric acid may be prepared by reacting borax (sodium tetraborate decahydrate) with a mineral acid , such as hydrochloric acid :
It is also formed as a byproduct of hydrolysis of boron trihalides and diborane : [ 10 ]
When heated, orthoboric acid undergoes a three-step dehydration. The reported transition temperatures vary substantially from source to source. [ citation needed ]
When heated above 140 °C, orthoboric acid yields metaboric acid ( HBO 2 ) with loss of one water molecule: [ 11 ] [ 12 ]
Heating metaboric acid above about 180 °C eliminates another water molecule forming tetraboric acid , also called pyroboric acid ( H 2 B 4 O 7 ): [ 11 ] [ 12 ]
Further heating (to about 530 °C) leads to boron trioxide : [ 13 ] [ 11 ] [ 12 ]
When orthoboric acid is dissolved in water, it partially dissociates to give metaboric acid :
The solution is mildly acidic due to the ionization of the acids:
However, Raman spectroscopy of strongly alkaline solutions has shown the presence of [B(OH) 4 ] − ions , [ 14 ] leading some to conclude that the acidity is exclusively due to the abstraction of OH − from water: [ 14 ]
Equivalently,
Or, more properly,
This reaction occurs in two steps, with the neutral complex aquatrihydroxyboron B(OH) 3 (OH 2 ) as an intermediate: [ 15 ]
This reaction may be characterized as Lewis acidity of boron toward HO − , rather than as Brønsted acidity . [ 16 ] [ 17 ] [ 18 ] However, some of its behaviour towards some chemical reactions suggest it to be a tribasic acid in the Brønsted-Lowry sense as well.
Boric acid, mixed with borax Na 2 B 4 O 7 ·10H 2 O (more properly Na 2 B 4 O 5 (OH) 4 ·8H 2 O ) in the weight ratio of 4:5, is highly soluble in water, though they are not so soluble separately. [ 19 ]
Boric acid also dissolves in anhydrous sulfuric acid according to the equation: [ 7 ]
The product is an extremely strong acid, even stronger than the original sulfuric acid. [ 7 ]
Boric acid reacts with alcohols to form borate esters , B(OR) 3 where R is alkyl or aryl . The reaction is typically driven by a dehydrating agent, such as concentrated sulfuric acid : [ 20 ]
The acidity of boric acid solutions is considerably increased in the presence of cis - vicinal diols ( organic compounds containing similarly oriented hydroxyl groups in adjacent carbon atoms, (R 1 ,R 2 )=C(OH)−C(OH)=(R 3 ,R 4 ) ) such as glycerol and mannitol . [ 21 ] [ 7 ] [ 22 ] [ 23 ]
The tetrahydroxyborate anion formed in the dissolution spontaneously reacts with these diols to form relatively stable anion esters containing one or two five-member −B−O−C−C−O− rings. For example, the reaction with mannitol H(HCOH) 6 H , whose two middle hydroxyls are in cis orientation, can be written as:
Giving the overall reaction:
The stability of these mannitoborate ester anions shifts the equilibrium to the right, thereby increasing the solution's acidity by five orders of magnitude compared to that of pure boric oxide. This lowers the p K a from 9 to below 4 for a sufficient concentration of mannitol. [ 21 ] [ 7 ] [ 22 ] [ 23 ] The resulting solution is referred to as mannitoboric acid.
The addition of mannitol to an initially neutral solution containing boric acid or simple borates lowers its pH enough for it to be titrated by a strong base such as NaOH, including with an automated potentiometric titrator . This property is used in analytical chemistry to determine the borate content of aqueous solutions, for example to monitor the depletion of boric acid by neutrons in the water of the primary circuit of light-water reactor when the compound is added as a neutron poison during refueling operations. [ 7 ]
Based on mammalian median lethal dose (LD 50 ) rating of 2,660 mg/kg body mass, boric acid is only poisonous if taken internally or inhaled in large quantities. The Fourteenth Edition of the Merck Index indicates that the LD 50 of boric acid is 5.14 g/kg for oral dosages given to rats, and that 5 to 20 g/kg has produced death in adult humans. For a 70 kg adult, at the lower 5 g/kg limit, 350 g could produce death in humans. For comparison's sake, the LD 50 of salt is reported to be 3.75 g/kg in rats according to the Merck Index . According to the Agency for Toxic Substances and Disease Registry , "The minimal lethal dose of ingested boron (as boric acid) was reported to be 2–3 g in infants, 5–6 g in children, and 15–20 g in adults. [...] However, a review of 784 human poisonings with boric acid (10–88 g) reported no fatalities, with 88% of cases being asymptomatic." [ 24 ] Human studies in three borate exposure-rich comparison groups (U.S. Borax mine and production facility workers, Chinese boron workers, Turkish residents living near boron-rich regions) produced no indicators of developmental toxicity in blood and semen tests. The highest estimated exposure was 5 mg B/kg/day, likely due to eating in contaminated workplaces, more than 100 times the average daily exposure. [ 25 ]
Long-term exposure to boric acid may be of more concern, causing kidney damage and eventually kidney failure (see links below). Although it does not appear to be carcinogenic , studies in dogs have reported testicular atrophy after exposure to 32 mg/(kg⋅day) for 90 days. This level, were it applicable to humans at like dose, would equate to a cumulative dose of 202 g over 90 days for a 70 kg adult, not far lower than the above LD 50 . [ 26 ]
According to the CLH report for boric acid published by the Bureau for Chemical Substances Lodz, Poland, boric acid in high doses shows significant developmental toxicity and teratogenicity in rabbit, rat, and mouse fetuses, as well as cardiovascular defects, skeletal variations, and mild kidney lesions. [ 25 ] As a consequence in the 30th ATP to EU directive 67/548/EEC of August 2008, the European Commission decided to amend its classification as reprotoxic category 2 and to apply the risk phrases R60 (may impair fertility) and R61 (may cause harm to the unborn child). [ 27 ] [ 28 ] [ 29 ] [ 30 ] [ 31 ]
At a 2010 European Diagnostics Manufacturing Association (EDMA) Meeting, several new additions to the substance of very high concern (SVHC) candidate list in relation to the Registration, Evaluation, Authorisation and Restriction of Chemicals Regulations 2007 (REACH) were discussed. Following the registration and review completed as part of REACH, the classification of boric acid CAS 10043-35-3 / 11113-50-1 is listed from 1 December 2010 is H360FD (May damage fertility. May damage the unborn child) . [ 32 ] [ 33 ]
Sound absorption in oceans not only depends on water molecules but also on dissolved salts present in low concentration in seawater . [ 34 ] [ 35 ] Boric acid and borate ([B] ≈ 4 × 10 -4 M in seawater) [ 36 ] relaxation contributes to absorbing sounds in the low‐ frequency region (0.2–10 kHz ). [ 37 ] At higher frequencies, between 10 and 1000 kHz magnesium sulfate (formed by the second most abundant cation and anion species in seawater) is the main contributor to the absorption of acoustic waves in seawater. [ 38 ]
The primary industrial use of boric acid is in the manufacture of monofilament fiberglass , which is usually referred to as textile fiberglass. Textile fiberglass is used to reinforce plastics in applications that range from boats to industrial piping to computer circuit boards. [ 39 ]
In the jewelry industry, boric acid is often used in combination with denatured alcohol to reduce surface oxidation and formation of firescale on metals during annealing and soldering operations. [ 40 ] [ 41 ]
Boric acid is used in the production of glass in LCD flat panel displays . [ 42 ] [ 43 ]
In electroplating , boric acid is used as part of some proprietary formulas. One known formula uses about a 1 to 10 ratio of H 3 BO 3 to NiSO 4 , a very small portion of sodium lauryl sulfate and a small portion of H 2 SO 4 .
The solution of orthoboric acid and borax in 4:5 ratio is used as a fire retarding agent of wood by impregnation. [ 44 ] Also, it is used in combination with other chemicals for the fire retardancy of wood-based materials. [ 45 ]
It is also used in the manufacturing of ramming mass , a fine silica -containing powder used for producing induction furnace linings and ceramics .
Boric acid is added to borax for use as welding flux by blacksmiths . [ 46 ]
Boric acid, in combination with polyvinyl alcohol (PVA) or silicone oil , is used to manufacture Silly Putty . [ 47 ]
Boric acid is also present in the list of chemical additives used for hydraulic fracturing (fracking) in the Marcellus Shale in Pennsylvania. [ 48 ] It is often used in conjunction with guar gum as cross-linking and gelling agent for controlling the viscosity and the rheology of the fracking fluid injected at high pressure in the well. It is important to control the fluid viscosity for keeping in suspension on long transport distances the grains of the propping agents aimed at maintaining the cracks in the shales sufficiently open to facilitate the gas extraction after the hydraulic pressure is relieved. [ 49 ] [ 50 ] [ 51 ] The rheological properties of borate cross-linked guar gum hydrogel mainly depend on the pH value. [ 52 ]
Boric acid is used in some expulsion-type electrical fuses as a de-ionization/extinguishing agent. [ 53 ] During an electrical fault in an expulsion-type fuse, a plasma arc is generated by the disintegration and rapid spring -loaded separation of the fusible element, which is typically a specialized metal rod that passes through a compressed mass of boric acid within the fuse assembly. The high-temperature plasma causes the boric acid to rapidly decompose into water vapor and boric anhydride , and in-turn, the vaporization products de-ionize the plasma, helping to interrupt the electrical fault. [ 54 ]
Boric acid can be used as an antiseptic for minor burns or cuts and is sometimes used in salves and dressings , such as boracic lint . Boric acid is applied in a very dilute solution as an eye wash. Boric acid vaginal suppositories can be used for recurrent candidiasis due to non-albicans candida as a second line treatment when conventional treatment has failed. [ 55 ] [ 56 ] It is less effective than conventional treatment overall. [ 55 ] Boric acid largely spares lactobacilli within the vagina. [ 57 ] As TOL-463 , it is under development as an intravaginal medication for the treatment for vulvovaginal candidiasis . [ 58 ] [ 59 ] [ 60 ]
As an antibacterial compound, boric acid can also be used as an acne treatment. It is also used as a prevention of athlete's foot , by inserting powder in the socks or stockings. Various preparations can be used to treat some kinds of otitis externa (ear infection) in both humans and animals. [ 61 ] The preservative in urine sample bottles in the UK is boric acid. [ 62 ]
Boric acid solutions used as an eye wash or on abraded skin are known to be toxic, particularly to infants, especially after repeated use; this is because of its slow elimination rate. [ 63 ]
Boric acid is one of the most commonly used substances that can counteract the harmful effects of reactive hydrofluoric acid (HF) after accidental contact with the skin. It works by forcing the free F − anions into the inert tetrafluoroborate anion. This process defeats the extreme toxicity of hydrofluoric acid, particularly its ability to sequester ionic calcium from blood serum which can lead to cardiac arrest and bone decomposition; such an event can occur from just minor skin contact with HF. [ 64 ] [ failed verification ]
Boric acid was first registered in the US as an insecticide in 1948 for control of cockroaches , termites , fire ants , fleas , silverfish , and many other insects . The product is generally considered safe in household kitchens to control cockroaches and ants. It acts as a stomach poison affecting the insects' metabolism . The dry powder is abrasive to the insects' exoskeletons . [ 65 ] [ 66 ] [ 67 ] It is in non-specific IRAC group 8D. Boric acid is also known as "the gift that keeps on killing" because cockroaches cross over lightly dusted areas and do not die immediately. Still, the effect is like shards of glass cutting them apart. This often allows a cockroach to return to the nest, where it soon dies. Cockroaches, being cannibalistic , eat others killed by contact or consumption of boric acid, consuming the powder trapped in the dead roach and killing them, too. [ citation needed ]
Boric acid is also widely used in wood treatment to protect against termites. The full complexity of its mechanism is not fully understood. Still, aside from causing dose-dependent mortality, boric acid causes dysbiosis in the Eastern Subterranean termite , leading to the opportunistic rise of insect pathogens that could be contributing to mortality. [ 68 ] In Japan the practice of laying newspapers treated with o-boric acid and borax under buildings has been effective in controlling Coptotermes formosanus and Reticulitermes speratus populations. Decaying wood treated with 0.25 to 0.5 percent disodium octaborate ( Na 2 B 8 O 13 ·4H 2 O , commonly abbreviated DOT) is also effective for baiting Heterotermes aureus populations. A 1997 paper concluded: "Borate baits would undoubtedly be helpful in the long-term, but do not appear sufficient as a sole method of structural protection." [ 69 ]
In combination with its use as an insecticide, boric acid also prevents and destroys existing wet and dry rot in timbers. It can be used in combination with an ethylene glycol carrier to treat external wood against fungal and insect attack. It is possible to buy borate-impregnated rods for insertion into wood via drill holes where dampness and moisture are known to collect and sit. It is available in a gel form and injectable paste form for treating rot-affected wood without replacing the timber. Concentrates of borate-based treatments can be used to prevent slime, mycelium, and algae growth, even in marine environments. [ citation needed ]
Boric acid is added to salt in the curing of cattle hides, calfskins , and sheepskins . This helps to control bacterial development and insects. [ citation needed ]
Boric acid in equilibrium with its conjugate base the borate ion is widely used (in the concentration range 50–100 ppm boron equivalents) as a primary or adjunct pH buffer system in swimming pools . Boric acid is a weak acid, with p K a (the pH at which buffering is strongest because the free acid and borate ion are in equal concentrations) of 9.24 in pure water at 25 °C. But apparent p K a is substantially lower in swimming pool or ocean waters because of interactions with various other molecules in solution. It will be around 9.0 in a saltwater pool. No matter which form of soluble boron is added, within the acceptable range of pH and boron concentration for swimming pools, boric acid is the predominant form in aqueous solution, as shown in the accompanying figure. The boric acid – borate system can be useful as a primary buffer system (substituting for the bicarbonate system with p K a 1 = 6.0 and p K a 2 = 9.4 under typical salt-water pool conditions) in pools with salt-water chlorine generators that tend to show upward drift in pH from a working range of pH 7.5–8.2. Buffer capacity is greater against rising pH (towards the pKa around 9.0), as illustrated in the accompanying graph. The use of boric acid in this concentration range does not allow any reduction in free HOCl concentration needed for pool sanitation. Still, it may add marginally to the photo-protective effects of cyanuric acid and confer other benefits through anti-corrosive activity or perceived water softness, depending on overall pool solute composition. [ 70 ]
Colloidal suspensions of nanoparticles of boric acid dissolved in petroleum or vegetable oil can form a remarkable lubricant on ceramic or metal surfaces [ 71 ] with a coefficient of sliding friction that decreases with increasing pressure to a value ranging from 0.10 to 0.02. Self-lubricating B(OH) 3 films result from a spontaneous chemical reaction between water molecules and B 2 O 3 coatings in a humid environment. On a bulk scale, an inverse relationship exists between the friction coefficient and the Hertzian contact pressure induced by the applied load. [ citation needed ]
Boric acid is used to lubricate carrom and novuss boards, allowing for faster play. [ 72 ]
Boric acid is used in some nuclear power plants as a neutron poison . The boron in boric acid reduces the probability of thermal fission by absorbing some thermal neutrons . Fission chain reactions are generally driven by the probability that free neutrons will result in fission and is determined by the material and geometric properties of the reactor. Natural boron consists of approximately 20% boron-10 and 80% boron-11 isotopes. Boron-10 has a high cross-section for absorption of low-energy (thermal) neutrons. By increasing boric acid concentration in the reactor coolant, the probability that a neutron will cause fission is reduced. Changes in boric acid concentration can effectively regulate the rate of fission taking place in the reactor. During normal at power operation, boric acid is used only in pressurized water reactors (PWRs), whereas boiling water reactors (BWRs) employ control rod pattern and coolant flow for power control. However, BWRs can use an aqueous solution of boric acid and borax or sodium pentaborate for an emergency shutdown system if the control rods fail to insert. Boric acid may be dissolved in spent fuel pools used to store spent fuel elements. The concentration is high enough to keep neutron multiplication at a minimum. Boric acid was dumped over Reactor 4 of the Chernobyl nuclear power plant after its meltdown to prevent another reaction from occurring. [ citation needed ]
Boron is used in pyrotechnics to prevent the amide -forming reaction between aluminium and nitrates . A small amount of boric acid is added to the composition to neutralize alkaline amides that can react with the aluminium.
Boric acid can be used as a colorant to make fire green. For example, when dissolved in methanol , it is popularly used by fire jugglers and fire spinners to create a deep green flame much stronger than copper sulfate. [ 73 ]
Boric acid is used to treat or prevent boron deficiencies in plants. It is also used in the preservation of grains such as rice and wheat. [ 74 ] | https://en.wikipedia.org/wiki/H3BO3 |
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