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9201 | of his scheme" and "the ridiculousness of the proposal" create a situation in which the reader has "to consider just what perverted values and assumptions would allow such a diligent, thoughtful, and conventional man to propose so perverse a plan". Scholars have speculated about which earlier works Swift may have had in mind when he wrote "A Modest Proposal". James Johnson argued that "A Modest Proposal" was largely influenced and inspired by Tertullian's "Apology": a satirical attack against early Roman persecution of Christianity. James William Johnson believes that Swift saw major similarities between the two situations. Johnson notes Swift's obvious | "A Modest Proposal" | [
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9202 | affinity for Tertullian and the bold stylistic and structural similarities between the works "A Modest Proposal" and "Apology". In structure, Johnson points out the same central theme, that of cannibalism and the eating of babies as well as the same final argument, that "human depravity is such that men will attempt to justify their own cruelty by accusing their victims of being lower than human". Stylistically, Swift and Tertullian share the same command of sarcasm and language. In agreement with Johnson, Donald C. Baker points out the similarity between both authors' tones and use of irony. Baker notes the uncanny | "A Modest Proposal" | [
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9203 | way that both authors imply an ironic "justification by ownership" over the subject of sacrificing children—Tertullian while attacking pagan parents, and Swift while attacking the English mistreatment of the Irish poor. It has also been argued that "A Modest Proposal" was, at least in part, a response to the 1728 essay "The Generous Projector or, A Friendly Proposal to Prevent Murder and Other Enormous Abuses, By Erecting an Hospital for Foundlings and Bastard Children" by Swift's rival Daniel Defoe. Bernard Mandeville's "Modest Defence of Publick Stews" asked to introduce public and state controlled bordellos. The 1726 paper acknowledges women's interests | "A Modest Proposal" | [
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9204 | andwhile not being a complete satirical texthas also been discussed as an inspiration for Jonathan Swift's title. Mandeville had by 1705 already become famous for the Fable of The Bees and deliberations on private vices and public benefits. Locke commented: "Be it then as Sir Robert says, that Anciently, it was usual for Men to sell and Castrate their Children. Let it be, that they exposed them; Add to it, if you please, for this is still greater Power, "that they begat them for their Tables to fat and eat them": If this proves a right to do so, we | "A Modest Proposal" | [
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9205 | may, by the same Argument, justifie Adultery, Incest and Sodomy, for there are examples of these too, both Ancient and Modern; Sins, which I suppose, have the Principle Aggravation from this, that they cross the main intention of Nature, which willeth the increase of Mankind, and the continuation of the Species in the highest perfection, and the distinction of Families, with the Security of the Marriage Bed, as necessary thereunto". (First Treatise, sec. 59). Robert Phiddian's article "Have you eaten yet? The Reader in A Modest Proposal" focuses on two aspects of "A Modest Proposal": the voice of Swift and | "A Modest Proposal" | [
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9206 | the voice of the Proposer. Phiddian stresses that a reader of the pamphlet must learn to distinguish between the satirical voice of Jonathan Swift and the apparent economic projections of the Proposer. He reminds readers that "there is a gap between the narrator's meaning and the text's, and that a moral-political argument is being carried out by means of parody". While Swift's proposal is obviously not a serious economic proposal, George Wittkowsky, author of "Swift's Modest Proposal: The Biography of an Early Georgian Pamphlet", argues that to understand the piece fully it is important to understand the economics of Swift's | "A Modest Proposal" | [
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9207 | time. Wittowsky argues that not enough critics have taken the time to focus directly on the mercantilism and theories of labour in 18th century England. "[I]f one regards the "Modest Proposal" simply as a criticism of condition, about all one can say is that conditions were bad and that Swift's irony brilliantly underscored this fact". At the start of a new industrial age in the 18th century, it was believed that "people are the riches of the nation", and there was a general faith in an economy that paid its workers low wages because high wages meant workers would work | "A Modest Proposal" | [
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9208 | less. Furthermore, "in the mercantilist view no child was too young to go into industry". In those times, the "somewhat more humane attitudes of an earlier day had all but disappeared and the laborer had come to be regarded as a commodity". Landa composed a conducive analysis when he noted that it would have been healthier for the Irish economy to more appropriately utilize their human assets by giving the people an opportunity to “become a source of wealth to the nation” or else they “must turn to begging and thievery”. This opportunity may have included giving the farmers more | "A Modest Proposal" | [
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9209 | coin to work for, diversifying their professions, or even consider enslaving their people to lower coin usage and build up financial stock in Ireland. Landa wrote that, "Swift is maintaining that the maxim—people are the riches of a nation—applies to Ireland only if Ireland is permitted slavery or cannibalism" Louis A. Landa presents Swift's "A Modest Proposal" as a critique of the popular and unjustified maxim of mercantilism in the 18th century that "people are the riches of a nation". Swift presents the dire state of Ireland and shows that mere population itself, in Ireland's case, did not always mean | "A Modest Proposal" | [
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9210 | greater wealth and economy. The uncontrolled maxim fails to take into account that a person who does not produce in an economic or political way makes a country poorer, not richer. Swift also recognises the implications of this fact in making mercantilist philosophy a paradox: the wealth of a country is based on the poverty of the majority of its citizens. Swift however, Landa argues, is not merely criticising economic maxims but also addressing the fact that England was denying Irish citizens their natural rights and dehumanising them by viewing them as a mere commodity. Swift's writings created a backlash | "A Modest Proposal" | [
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9211 | within the community after its publication. The work was aimed at the aristocracy, and they responded in turn. Several members of society wrote to Swift regarding the work. Lord Bathurst's letter intimated that he certainly understood the message, and interpreted it as a work of comedy: February 12, 1729-30:"I did immediately propose it to Lady Bathurst, as your advice, particularly for her last boy, which was born the plumpest, finest thing, that could be seen; but she fell in a passion, and bid me send you word, that she would not follow your direction, but that she would breed him | "A Modest Proposal" | [
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9212 | up to be a parson, and he should live upon the fat of the land; or a lawyer, and then, instead of being eat himself, he should devour others. You know women in passion never mind what they say; but, as she is a very reasonable woman, I have almost brought her over now to your opinion; and having convinced her, that as matters stood, we could not possibly maintain all the nine, she does begin to think it reasonable the youngest should raise fortunes for the eldest: and upon that foot a man may perforin family duty with more | "A Modest Proposal" | [
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9213 | courage and zeal; for, if he should happen to get twins, the selling of one might provide for the other. Or if, by any accident, while his wife lies in with one child, he should get a second upon the body of another woman, he might dispose of the fattest of the two, and that would help to breed up the other.The more I think upon this scheme, the more reasonable it appears to me; and it ought by no means to be confined to Ireland; for, in all probability, we shall, in a very little time, be altogether as | "A Modest Proposal" | [
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9214 | poor here as you are there. I believe, indeed, we shall carry it farther, and not confine our luxury only to the eating of children; for I happened to peep the other day into a large assembly [Parliament] not far from Westminster-hall, and I found them roasting a great fat fellow, [Walpole again] For my own part, I had not the least inclination to a slice of him; but, if I guessed right, four or five of the company had a devilish mind to be at him. Well, adieu, you begin now to wish I had ended, when I might | "A Modest Proposal" | [
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9215 | have done it so conveniently". "A Modest Proposal" is included in many literature courses as an example of early modern western satire. It also serves as an exceptional introduction to the concept and use of argumentative language, lending itself well to secondary and post-secondary essay courses. Outside of the realm of English studies, "A Modest Proposal" is included in many comparative and global literature and history courses, as well as those of numerous other disciplines in the arts, humanities, and even the social sciences. The essay's approach has been copied many times. In his book "A Modest Proposal" (1984), the | "A Modest Proposal" | [
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9216 | evangelical author Frank Schaeffer emulated Swift's work in a social conservative polemic against abortion and euthanasia, imagining a future dystopia that advocates recycling of aborted embryos, fetuses, and some disabled infants with compound intellectual, physical and physiological difficulties. (Such Baby Doe Rules cases were then a major concern of the US pro-life movement of the early 1980s, which viewed selective treatment of those infants as disability discrimination). In his book "A Modest Proposal for America" (2013), statistician Howard Friedman opens with a satirical reflection of the extreme drive to fiscal stability by ultra-conservatives. In the 1998 edition of A Handmaid's | "A Modest Proposal" | [
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9217 | Tale by Margaret Atwood there is a quote from "A Modest Proposal" before the introduction. "A Modest Video Game Proposal" is the title of an open letter sent by activist/former attorney Jack Thompson on 10 October 2005. He proposed that someone should "create, manufacture, distribute, and sell a video game" that would allow players to act out a scenario in which the game character kills video game developers. Hunter S. Thompson's "Fear and Loathing in America: The Brutal Odyssey of an Outlaw Journalist includes" a letter in which he uses Swift's approach in connection with the Vietnam War. Thompson writes | "A Modest Proposal" | [
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9218 | a letter to a local Aspen newspaper informing them that, on Christmas Eve, he was going to use napalm to burn a number of dogs and hopefully any humans they find. The letter protests against the burning of Vietnamese people occurring overseas. The 2012 film "Butcher Boys," written by Kim Henkel, is said to be loosely based on Jonathan Swift's "A Modest Proposal." The film's opening scene takes place in a restaurant named "J. Swift's". On November 30, 2017, Jonathan Swift's 350th birthday, The Washington Post published a column entitled 'Why Alabamians should consider eating Democrats' babies", by the humorous | "A Modest Proposal" | [
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9219 | columnist Alexandra Petri. A Modest Proposal A Modest Proposal For preventing the Children of Poor People From being a Burthen to Their Parents or Country, and For making them Beneficial to the Publick, commonly referred to as A Modest Proposal, is a Juvenalian satirical essay written and published anonymously by Jonathan Swift in 1729. The essay suggests that the impoverished Irish might ease their economic troubles by selling their children as food for rich gentlemen and ladies. This satirical hyperbole mocked heartless attitudes towards the poor, as well as British policy toward the Irish in general. The primary target of | "A Modest Proposal" | [
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9220 | Alkali metal The alkali metals are a group (column) in the periodic table consisting of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). This group lies in the s-block of the periodic table of elements as all alkali metals have their outermost electron in an s-orbital: this shared electron configuration results in their having very similar characteristic properties. Indeed, the alkali metals provide the best example of group trends in properties in the periodic table, with elements exhibiting well-characterised homologous behaviour. The alkali metals are all shiny, soft, highly reactive metals at | "Alkali metal" | [
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9221 | standard temperature and pressure and readily lose their outermost electron to form cations with charge +1. They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to oxidation by atmospheric moisture and oxygen (and in the case of lithium, nitrogen). Because of their high reactivity, they must be stored under oil to prevent reaction with air, and are found naturally only in salts and never as the free elements. Caesium, the fifth alkali metal, is the most reactive of all the metals. In the modern IUPAC nomenclature, | "Alkali metal" | [
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9222 | the alkali metals comprise the group 1 elements, excluding hydrogen (H), which is nominally a group 1 element but not normally considered to be an alkali metal as it rarely exhibits behaviour comparable to that of the alkali metals. All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones. All of the discovered alkali metals occur in nature as their compounds: in order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and finally francium, which is very rare due to its extremely high radioactivity; francium occurs | "Alkali metal" | [
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9223 | only in the minutest traces in nature as an intermediate step in some obscure side branches of the natural decay chains. Experiments have been conducted to attempt the synthesis of ununennium (Uue), which is likely to be the next member of the group, but they have all met with failure. However, ununennium may not be an alkali metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements; even if it does turn out to be an alkali metal, it is predicted to have some differences in physical and chemical properties | "Alkali metal" | [
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9224 | from its lighter homologues. Most alkali metals have many different applications. One of the best-known applications of the pure elements is the use of rubidium and caesium in atomic clocks, of which caesium atomic clocks are the most accurate and precise representation of time. A common application of the compounds of sodium is the sodium-vapour lamp, which emits light very efficiently. Table salt, or sodium chloride, has been used since antiquity. Sodium and potassium are also essential elements, having major biological roles as electrolytes, and although the other alkali metals are not essential, they also have various effects on the | "Alkali metal" | [
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9225 | body, both beneficial and harmful. Sodium compounds have been known since ancient times; salt (sodium chloride) has been an important commodity in human activities, as testified by the English word "salary", referring to "salarium", money paid to Roman soldiers for the purchase of salt. While potash has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence which led him to suggest the fundamental difference of sodium and potassium salts in 1702, and Henri-Louis Duhamel du Monceau was able to | "Alkali metal" | [
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9226 | prove this difference in 1736. The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus Antoine Lavoisier did not include either alkali in his list of chemical elements in 1789. Pure potassium was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented voltaic pile. Previous attempts at electrolysis of the aqueous salt were unsuccessful due to potassium's extreme reactivity. Potassium was the | "Alkali metal" | [
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... |
9227 | first metal that was isolated by electrolysis. Later that same year, Davy reported extraction of sodium from the similar substance caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different. Petalite (Li Al SiO) was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada in a mine on the island of Utö, Sweden. However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jacob Berzelius, detected the presence of a new element while analysing petalite ore. This new element was noted | "Alkali metal" | [
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9228 | by him to form compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline than the other alkali metals. Berzelius gave the unknown material the name ""lithion"/"lithina"", from the Greek word "λιθoς" (transliterated as "lithos", meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium, which was known partly for its high abundance in animal blood. He named the metal inside the material ""lithium"". Lithium, sodium, and potassium were part of the discovery of periodicity, as | "Alkali metal" | [
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0.57071697711... |
9229 | they are among a series of triads of elements in the same group that were noted by Johann Wolfgang Döbereiner in 1850 as having similar properties. Rubidium and caesium were the first elements to be discovered using the spectroscope, invented in 1859 by Robert Bunsen and Gustav Kirchhoff. The next year, they discovered caesium in the mineral water from Bad Dürkheim, Germany. Their discovery of rubidium came the following year in Heidelberg, Germany, finding it in the mineral lepidolite. The names of rubidium and caesium come from the most prominent lines in their emission spectra: a bright red line for | "Alkali metal" | [
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9230 | rubidium (from the Latin word "rubidus", meaning dark red or bright red), and a sky-blue line for caesium (derived from the Latin word "caesius", meaning sky-blue). Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music, where notes an octave apart have similar musical functions. His version put all the alkali metals then known (lithium to caesium), as well as copper, silver, and thallium (which show the +1 oxidation | "Alkali metal" | [
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0.89290440082550... |
9231 | state characteristic of the alkali metals), together into a group. His table placed hydrogen with the halogens. After 1869, Dmitri Mendeleev proposed his periodic table placing lithium at the top of a group with sodium, potassium, rubidium, caesium, and thallium. Two years later, Mendeleev revised his table, placing hydrogen in group 1 above lithium, and also moving thallium to the boron group. In this 1871 version, copper, silver, and gold were placed twice, once as part of group IB, and once as part of a "group VIII" encompassing today's groups 8 to 11. After the introduction of the 18-column table, | "Alkali metal" | [
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1.01379072666168... |
9232 | the group IB elements were moved to their current position in the d-block, while alkali metals were left in "group IA". Later the group's name was changed to "group 1" in 1988. The trivial name "alkali metals" comes from the fact that the hydroxides of the group 1 elements are all strong alkalis when dissolved in water. There were at least four erroneous and incomplete discoveries before Marguerite Perey of the Curie Institute in Paris, France discovered francium in 1939 by purifying a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey | "Alkali metal" | [
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0.843820214... |
9233 | noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227. Perey then attempted to determine the proportion of beta decay to alpha | "Alkali metal" | [
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0.676839113235473... |
9234 | decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%. The next element below francium (eka-francium) in the periodic table would be ununennium (Uue), element 119. The synthesis of ununennium was first attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb. It is highly unlikely that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making | "Alkali metal" | [
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0.19420850276947021,
0.10371982306241989,
0.12396398186683655,
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0.671963810920... |
9235 | sufficient amounts of einsteinium-254, which is favoured for production of ultraheavy elements because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms, to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and has only been produced in laboratories, and in quantities smaller than those needed for effective synthesis of superheavy elements. However, given that ununennium is only the first period 8 element on the extended periodic table, it may well be discovered in the near future | "Alkali metal" | [
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0.53528249263763... |
9236 | through other reactions, and indeed an attempt to synthesise it is currently ongoing in Japan. Currently, none of the period 8 elements have been discovered yet, and it is also possible, due to drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible. No attempts at synthesis have been made for any heavier alkali metals: due to their extremely high atomic number, they would require new, more powerful methods and technology to make. The Oddo–Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with | "Alkali metal" | [
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0.5377228856... |
9237 | the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability. All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones | "Alkali metal" | [
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... |
9238 | as the alkali metals from rubidium onward can only be synthesised in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesised in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements. The Earth formed from the same cloud of | "Alkali metal" | [
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0.005382177885621786,
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0.56372147798... |
9239 | matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the solar system. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98 kg. It is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to planetary differentiation, the core region is believed to be primarily composed of iron (88.8%), with smaller | "Alkali metal" | [
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0.4626000225543... |
9240 | amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements. The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are lithophiles and therefore remain close to the Earth's surface because they combine readily with oxygen and so associate strongly with silica, forming relatively low-density minerals that do not sink down into the Earth's core. Potassium, rubidium and caesium are also incompatible elements due to their large ionic radii. Sodium and potassium are very abundant in earth, both being among the ten most common elements in Earth's crust; sodium makes | "Alkali metal" | [
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9241 | up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element. Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite. Many of these solid deposits occur as a result of ancient seas evaporating, which still occurs now in places such as Utah's Great Salt Lake | "Alkali metal" | [
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0.7598980665... |
9242 | and the Dead Sea. Despite their near-equal abundance in Earth's crust, sodium is far more common than potassium in the ocean, both because potassium's larger size makes its salts less soluble, and because potassium is bound by silicates in soil and what potassium leaches is absorbed far more readily by plant life than sodium. Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size. Due to its relatively low reactivity, it can be found in seawater in large amounts; it is estimated that seawater is approximately 0.14 to 0.25 parts per | "Alkali metal" | [
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0.12335514277219772,
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0.5191577076... |
9243 | million (ppm) or 25 micromolar. Its diagonal relationship with magnesium often allows it to replace magnesium in ferromagnesium minerals, where its crustal concentration is about 18 ppm, comparable to that of gallium and niobium. Commercially, the most important lithium mineral is spodumene, which occurs in large deposits worldwide. Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite, although none of these contain only rubidium and no other alkali metals. Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, | "Alkali metal" | [
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0.55419832468032... |
9244 | but is much less abundant than rubidium. Francium-223, the only naturally occurring isotope of francium, is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium minerals. In a given sample of uranium, there is estimated to be only one francium atom for every 10 uranium atoms. It has been calculated that there is at most 30 g of francium in the earth's crust at any time, due to its extremely short half-life of 22 minutes. The physical and chemical properties of the alkali metals can be readily explained by their having an | "Alkali metal" | [
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0.5371089577674... |
9245 | ns valence electron configuration, which results in weak metallic bonding. Hence, all the alkali metals are soft and have low densities, melting and boiling points, as well as heats of sublimation, vaporisation, and dissociation. They all crystallise in the body-centered cubic crystal structure, and have distinctive flame colours because their outer s electron is very easily excited. The ns configuration also results in the alkali metals having very large atomic and ionic radii, as well as very high thermal and electrical conductivity. Their chemistry is dominated by the loss of their lone valence electron in the outermost s-orbital to form | "Alkali metal" | [
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9246 | the +1 oxidation state, due to the ease of ionising this electron and the very high second ionisation energy. Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme radioactivity; thus, the presentation of its properties here is limited. What little is known about francium shows that it is very close in behaviour to caesium, as expected. The physical properties of francium are even sketchier because the bulk element has never been observed; hence any data that may be found in the | "Alkali metal" | [
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0.542576968669... |
9247 | literature are certainly speculative extrapolations. The alkali metals are more similar to each other than the elements in any other group are to each other. Indeed, the similarity is so great that it is quite difficult to separate potassium, rubidium, and caesium, due to their similar ionic radii; lithium and sodium are more distinct. For instance, when moving down the table, all known alkali metals show increasing atomic radius, decreasing electronegativity, increasing reactivity, and decreasing melting and boiling points as well as heats of fusion and vaporisation. In general, their densities increase when moving down the table, with the exception | "Alkali metal" | [
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9248 | that potassium is less dense than sodium. One of the very few properties of the alkali metals that does not display a very smooth trend is their reduction potentials: lithium's value is anomalous, being more negative than the others. This is because the Li ion has a very high hydration energy in the gas phase: though the lithium ion disrupts the structure of water significantly, causing a higher change in entropy, this high hydration energy is enough to make the reduction potentials indicate it as being the most electropositive alkali metal, despite the difficulty of ionising it in the gas | "Alkali metal" | [
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0.06599143892526627,
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0.360826939344... |
9249 | phase. The stable alkali metals are all silver-coloured metals except for caesium, which has a pale golden tint: it is one of only three metals that are clearly coloured (the other two being copper and gold). Additionally, the heavy alkaline earth metals calcium, strontium, and barium, as well as the divalent lanthanides europium and ytterbium, are pale yellow, though the colour is much less prominent than it is for caesium. Their lustre tarnishes rapidly in air due to oxidation. They all crystallise in the body-centered cubic crystal structure, and have distinctive flame colours because their outer s electron is very | "Alkali metal" | [
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0.058061420917510986,
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0.734114527702... |
9250 | easily excited. Indeed, these flame test colours are the most common way of identifying them since all their salts with common ions are soluble. All the alkali metals are highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored in mineral oil or kerosene (paraffin oil). They react aggressively with the halogens to form the alkali metal halides, which are white ionic crystalline compounds that are all soluble in water except lithium fluoride (Li F). The alkali metals also react with water to form strongly alkaline hydroxides and thus should be handled | "Alkali metal" | [
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0.35076454281806946,
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0.5798051357269... |
9251 | with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium if the same number of moles of each metal is used. The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table because of their low effective nuclear charge and the ability to attain a noble gas configuration by losing just one electron. Not only do the alkali metals react with water, but also with proton donors like alcohols and phenols, gaseous ammonia, and alkynes, the last | "Alkali metal" | [
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0.4921069145202... |
9252 | demonstrating the phenomenal degree of their reactivity. Their great power as reducing agents makes them very useful in liberating other metals from their oxides or halides. The second ionisation energy of all of the alkali metals is very high as it is in a full shell that is also closer to the nucleus; thus, they almost always lose a single electron, forming cations. The alkalides are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be | "Alkali metal" | [
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9253 | able to form anions and were thought to be able to appear in salts only as cations. The alkalide anions have filled s-subshells, which gives them enough stability to exist. All the stable alkali metals except lithium are known to be able to form alkalides, and the alkalides have much theoretical interest due to their unusual stoichiometry and low ionisation potentials. Alkalides are chemically similar to the electrides, which are salts with trapped electrons acting as anions. A particularly striking example of an alkalide is "inverse sodium hydride", HNa (both ions being complexed), as opposed to the usual sodium hydride, | "Alkali metal" | [
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9254 | NaH: it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be metastable or stable. In aqueous solution, the alkali metal ions form aqua ions of the formula [M(HO)], where "n" is the solvation number. Their coordination numbers and shapes agree well with those expected from their ionic radii. In aqueous solution the water molecules directly attached to the metal ion are said to belong to the first coordination sphere, also known as the first, or primary, solvation shell. The bond between a | "Alkali metal" | [
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9255 | water molecule and the metal ion is a dative covalent bond, with the oxygen atom donating both electrons to the bond. Each coordinated water molecule may be attached by hydrogen bonds to other water molecules. The latter are said to reside in the second coordination sphere. However, for the alkali metal cations, the second coordination sphere is not well-defined as the +1 charge on the cation is not high enough to polarise the water molecules in the primary solvation shell enough for them to form strong hydrogen bonds with those in the second coordination sphere, producing a more stable entity. | "Alkali metal" | [
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9256 | The solvation number for Li has been experimentally determined to be 4, forming the tetrahedral [Li(HO)]: while solvation numbers of 3 to 6 have been found for lithium aqua ions, solvation numbers less than 4 may be the result of the formation of contact ion pairs, and the higher solvation numbers may be interpreted in terms of water molecules that approach [Li(HO)] through a face of the tetrahedron, though molecular dynamic simulations may indicate the existence of an octahedral hexaaqua ion. There are also probably six water molecules in the primary solvation sphere of the sodium ion, forming the octahedral | "Alkali metal" | [
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9257 | [Na(HO)] ion. While it was previously thought that the heavier alkali metals also formed octahedral hexaaqua ions, it has since been found that potassium and rubidium probably form the [K(HO)] and [Rb(HO)] ions, which have the square antiprismatic structure, and that caesium forms the 12-coordinate [Cs(HO)] ion. The chemistry of lithium shows several differences from that of the rest of the group as the small Li cation polarises anions and gives its compounds a more covalent character. Lithium and magnesium have a diagonal relationship due to their similar atomic radii, so that they show some similarities. For example, lithium forms | "Alkali metal" | [
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9258 | a stable nitride, a property common among all the alkaline earth metals (magnesium's group) but unique among the alkali metals. In addition, among their respective groups, only lithium and magnesium form organometallic compounds with significant covalent character (e.g. LiMe and MgMe). Lithium fluoride is the only alkali metal halide that is poorly soluble in water, and lithium hydroxide is the only alkali metal hydroxide that is not deliquescent. Conversely, lithium perchlorate and other lithium salts with large anions that cannot be polarised are much more stable than the analogous compounds of the other alkali metals, probably because Li has a | "Alkali metal" | [
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9259 | high solvation energy. This effect also means that most simple lithium salts are commonly encountered in hydrated form, because the anhydrous forms are extremely hygroscopic: this allows salts like lithium chloride and lithium bromide to be used in dehumidifiers and air-conditioners. Francium is also predicted to show some differences due to its high atomic weight, causing its electrons to travel at considerable fractions of the speed of light and thus making relativistic effects more prominent. In contrast to the trend of decreasing electronegativities and ionisation energies of the alkali metals, francium's electronegativity and ionisation energy are predicted to be higher | "Alkali metal" | [
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9260 | than caesium's due to the relativistic stabilisation of the 7s electrons; also, its atomic radius is expected to be abnormally low. Thus, contrary to expectation, caesium is the most reactive of the alkali metals, not francium. All known physical properties of francium also deviate from the clear trends going from lithium to caesium, such as the first ionisation energy, electron affinity, and anion polarisability, though due to the paucity of known data about francium many sources give extrapolated values, ignoring that relativistic effects make the trend from lithium to caesium become inapplicable at francium. Some of the few properties of | "Alkali metal" | [
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9261 | francium that have been predicted taking relativity into account are the electron affinity (47.2 kJ/mol) and the enthalpy of dissociation of the Fr molecule (42.1 kJ/mol). The CsFr molecule is polarised as CsFr, showing that the 7s subshell of francium is much more strongly affected by relativistic effects than the 6s subshell of caesium. Additionally, francium superoxide (FrO) is expected to have significant covalent character, unlike the other alkali metal superoxides, because of bonding contributions from the 6p electrons of francium. All the alkali metals have odd atomic numbers; hence, their isotopes must be either odd–odd (both proton and neutron | "Alkali metal" | [
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9262 | number are odd) or odd–even (proton number is odd, but neutron number is even). Odd–odd nuclei have even mass numbers, whereas odd–even nuclei have odd mass numbers. Odd–odd primordial nuclides are rare because most odd–odd nuclei are highly unstable with respect to beta decay, because the decay products are even–even, and are therefore more strongly bound, due to nuclear pairing effects. Due to the great rarity of odd–odd nuclei, almost all the primordial isotopes of the alkali metals are odd–even (the exceptions being the light stable isotope lithium-6 and the long-lived radioisotope potassium-40). For a given odd mass number, there | "Alkali metal" | [
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0.72609168291091... |
9263 | can be only a single beta-stable nuclide, since there is not a difference in binding energy between even–odd and odd–even comparable to that between even–even and odd–odd, leaving other nuclides of the same mass number (isobars) free to beta decay toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 monoisotopic elements that have only a single stable isotope, all but one have an odd atomic number and all but one | "Alkali metal" | [
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9264 | also have an even number of neutrons. Beryllium is the single exception to both rules, due to its low atomic number. All of the alkali metals except lithium and caesium have at least one naturally occurring radioisotope: sodium-22 and sodium-24 are trace radioisotopes produced cosmogenically, potassium-40 and rubidium-87 have very long half-lives and thus occur naturally, and all isotopes of francium are radioactive. Caesium was also thought to be radioactive in the early 20th century, although it has no naturally occurring radioisotopes. (Francium had not been discovered yet at that time.) The natural long-lived radioisotope of potassium, potassium-40, makes up | "Alkali metal" | [
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9265 | about 0.012% of natural potassium, and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925. Natural rubidium is similarly slightly radioactive, with 27.83% being the long-lived radioisotope rubidium-87. Caesium-137, with a half-life of 30.17 years, is one of the two principal medium-lived fission products, along with strontium-90, which are responsible for most of the radioactivity of spent nuclear fuel after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left | "Alkali metal" | [
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0.806897163391... |
9266 | from the Chernobyl accident. Caesium-137 undergoes high-energy beta decay and eventually becomes stable barium-137. It is a strong emitter of gamma radiation. Caesium-137 has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay. Caesium-137 has been used as a tracer in hydrologic studies, analogous to the use of tritium. Small amounts of caesium-134 and caesium-137 were released into the environment during nearly all nuclear weapon tests and some nuclear accidents, most notably the Goiânia accident and the Chernobyl disaster. As of 2005, caesium-137 is the principal source | "Alkali metal" | [
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0.86158430576... |
9267 | of radiation in the zone of alienation around the Chernobyl nuclear power plant. Its chemical properties as one of the alkali metals make it one of most problematic of the short-to-medium-lifetime fission products because it easily moves and spreads in nature due to the high water solubility of its salts, and is taken up by the body, which mistakes it for its essential congeners sodium and potassium. The alkali metals are more similar to each other than the elements in any other group are to each other. For instance, when moving down the table, all known alkali metals show increasing | "Alkali metal" | [
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9268 | atomic radius, decreasing electronegativity, increasing reactivity, and decreasing melting and boiling points as well as heats of fusion and vaporisation. In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium. The atomic radii of the alkali metals increase going down the group. Because of the shielding effect, when an atom has more than one electron shell, each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus. In the alkali metals, the outermost electron only feels a net charge of +1, as some of | "Alkali metal" | [
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9269 | the nuclear charge (which is equal to the atomic number) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group. The ionic radii of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different electron shell than the | "Alkali metal" | [
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9270 | inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the effective nuclear charge has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases. The first ionisation energy of an element or molecule is the energy required to move the most loosely held electron from one mole of gaseous atoms of the element or molecules to form one mole of gaseous ions with electric charge +1. The factors affecting the first ionisation energy are the nuclear charge, the amount of shielding by | "Alkali metal" | [
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0.31627443432807... |
9271 | the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in main group elements. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feels the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases. (This trend is | "Alkali metal" | [
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9272 | broken in francium due to the relativistic stabilisation and contraction of the 7s orbital, bringing francium's valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium's outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.) The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled electron shell and is thus difficult to remove. The reactivities of the alkali metals increase going down the group. This is the result of | "Alkali metal" | [
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9273 | a combination of two factors: the first ionisation energies and atomisation energies of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in chemical reactions, thus increasing reactivity down the group. The atomisation energy measures the strength of the metallic bond of an element, which falls down the group as the atoms increase in radius and thus the metallic bond must increase in length, making the delocalised electrons further away from the attraction of the nuclei of the | "Alkali metal" | [
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9274 | heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the activation energy of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group. Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself. If the bond between sodium and chlorine in sodium chloride were covalent, the pair of shared electrons | "Alkali metal" | [
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9275 | would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an ionic bond). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms | "Alkali metal" | [
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9276 | (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them. As mentioned previously, francium is expected to be an exception. Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, lithium iodide (Li I) will dissolve in organic solvents, a property of most covalent compounds. Lithium fluoride (LiF) is the only alkali halide that is not soluble in water, and lithium hydroxide (LiOH) is the only alkali metal hydroxide that is not deliquescent. The melting point of a substance is the point where it | "Alkali metal" | [
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9277 | changes state from solid to liquid while the boiling point of a substance (in liquid state) is the point where the vapour pressure of the liquid equals the environmental pressure surrounding the liquid and all the liquid changes state to gas. As a metal is heated to its melting point, the metallic bonds keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point. Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali | "Alkali metal" | [
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0.3512519598007202,
-0.4854882061481476,
0.016841216012835503,
-0.2085903137922287,
-0.6531665921211243,
0.78755056858062... |
9278 | metals decreases down the group. This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons. As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points. (The increased nuclear charge is not a relevant factor due to the shielding effect.) The alkali metals all have the same crystal structure (body-centred | "Alkali metal" | [
-0.07741975784301758,
0.653624415397644,
0.21432779729366302,
-0.21031729876995087,
-0.5546543598175049,
-0.14949551224708557,
0.5063337087631226,
-0.6996979117393494,
0.298168420791626,
-0.32897958159446716,
-0.07753106951713562,
-0.30381616950035095,
-0.7808839082717896,
0.59908545017242... |
9279 | cubic) and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic | "Alkali metal" | [
-0.032003629952669144,
0.5423445105552673,
0.2004152536392212,
-0.17877830564975739,
-0.28406763076782227,
0.21269869804382324,
0.2806636691093445,
-0.6958383321762085,
0.10132051259279251,
-0.7022037506103516,
-0.14678116142749786,
-0.08805486559867859,
-0.6073347926139832,
0.565681695938... |
9280 | radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight and the largest atomic radius of all the elements in their periods, the alkali metals are the least dense metals in the periodic table. Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water: in fact, lithium is the least dense known | "Alkali metal" | [
-0.22815832495689392,
0.5875567197799683,
0.03150104358792305,
0.06880854070186615,
-0.17094534635543823,
0.005777287762612104,
0.3357499837875366,
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0.13886770606040955,
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-0.5492605566978455,
0.711886703968... |
9281 | solid at room temperature. The alkali metals form complete series of compounds with all usually encountered anions, which well illustrate group trends. These compounds can be described as involving the alkali metals losing electrons to acceptor species and forming monopositive ions. This description is most accurate for alkali halides and becomes less and less accurate as cationic and anionic charge increase, and as the anion becomes larger and more polarisable. For instance, ionic bonding gives way to metallic bonding along the series NaCl, NaO, NaS, NaP, NaAs, NaSb, NaBi, Na. All the alkali metals react vigorously or explosively with cold | "Alkali metal" | [
-0.007534942589700222,
0.7713237404823303,
0.2746686637401581,
-0.16559390723705292,
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0.5939639210700989,
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0.27596229314804077,
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-0.2587323784828186,
-0.2822684049606323,
-0.7704592943191528,
0.774379253387451... |
9282 | water, producing an aqueous solution of a strongly basic alkali metal hydroxide and releasing hydrogen gas. This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers. When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes | "Alkali metal" | [
0.0968199223279953,
0.6160027384757996,
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0.4106062650680542,
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-0.2667028307914734,
0.091841921210289,
-0.6512142419815063,
0.5913556218147278... |
9283 | place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water). The alkali metal hydroxides are the most basic known hydroxides. Recent research has suggested that the explosive behavior of alkali metals in water is driven by a Coulomb explosion rather | "Alkali metal" | [
0.17486830055713654,
0.7311539649963379,
-0.11780455708503723,
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0.47220614552497864,
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0.2817593812942505,
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-0.309775173664093,
0.07164627313613892,
-0.7074428796768188,
0.537759184837341... |
9284 | than solely by rapid generation of hydrogen itself. All alkali metals melt as a part of the reaction with water. Water molecules ionise the bare metallic surface of the liquid metal, leaving a positively charged metal surface and negatively charged water ions. The attraction between the charged metal and water ions will rapidly increase the surface area, causing an exponential increase of ionisation. When the repulsive forces within the liquid metal surface exceeds the forces of the surface tension, it vigorously explodes. The hydroxides themselves are the most basic hydroxides known, reacting with acids to give salts and with alcohols | "Alkali metal" | [
-0.04857484623789787,
0.8123049139976501,
0.1710134893655777,
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-0.4053851366043091,
-0.19446510076522827,
0.5207422971725464,
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0.3813965916633606,
-0.3576120138168335,
-0.13944023847579956,
-0.1734907031059265,
-0.908506453037262,
0.674996435642242... |
9285 | to give oligomeric alkoxides. They easily react with carbon dioxide to form carbonates or bicarbonates, or with hydrogen sulfide to form sulfides or bisulfides, and may be used to separate thiols from petroleum. They react with amphoteric oxides: for example, the oxides of aluminium, zinc, tin, and lead react with the alkali metal hydroxides to give aluminates, zincates, stannates, and plumbates. Silicon dioxide is acidic, and thus the alkali metal hydroxides can also attack silicate glass. The alkali metals form many intermetallic compounds with each other and the elements from groups 2 to 13 in the periodic table of varying | "Alkali metal" | [
0.014407675713300705,
0.81290602684021,
-0.0574369840323925,
0.14829838275909424,
-0.49599775671958923,
0.049599550664424896,
0.3137425184249878,
-0.3432522118091583,
0.6039220094680786,
-0.5204129219055176,
-0.17480401694774628,
0.2261100560426712,
-0.5435112118721008,
0.625213086605072,
... |
9286 | stoichiometries, such as the sodium amalgams with mercury, including NaHg and NaHg. Some of these have ionic characteristics: taking the alloys with gold, the most electronegative of metals, as an example, NaAu and KAu are metallic, but RbAu and CsAu are semiconductors. NaK is an alloy of sodium and potassium that is very useful because it is liquid at room temperature, although precautions must be taken due to its extreme reactivity towards water and air. The eutectic mixture melts at −12.6 °C. An alloy of 41% caesium, 47% sodium, and 12% potassium has the lowest known melting point of any | "Alkali metal" | [
-0.03901271894574165,
0.6181856393814087,
0.06520013511180878,
0.14240124821662903,
-0.5164956450462341,
0.11937087029218674,
0.43607157468795776,
-0.45551419258117676,
0.42411693930625916,
-0.27525442838668823,
0.09204324334859848,
-0.12098130583763123,
-0.31958362460136414,
0.52745801210... |
9287 | metal or alloy, −78 °C. The intermetallic compounds of the alkali metals with the heavier group 13 elements (aluminium, gallium, indium, and thallium), such as NaTl, are poor conductors or semiconductors, unlike the normal alloys with the preceding elements, implying that the alkali metal involved has lost an electron to the Zintl anions involved. Nevertheless, while the elements in group 14 and beyond tend to form discrete anionic clusters, group 13 elements tend to form polymeric ions with the alkali metal cations located between the giant ionic lattice. For example, NaTl consists of a polymeric anion (—Tl—) with a covalent | "Alkali metal" | [
0.28452321887016296,
0.205082967877388,
0.29965972900390625,
0.04487994313240051,
-0.5693366527557373,
-0.02450261451303959,
0.34277036786079407,
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0.19604144990444183,
-0.11490015685558319,
-0.15698179602622986,
-0.0921069085597992,
-0.6426495313644409,
0.53902053833007... |
9288 | diamond cubic structure with Na ions located between the anionic lattice. The larger alkali metals cannot fit similarly into an anionic lattice and tend to force the heavier group 13 elements to form anionic clusters. Boron is a special case, being the only nonmetal in group 13. The alkali metal borides tend to be boron-rich, involving appreciable boron–boron bonding involving deltahedral structures, and are thermally unstable due to the alkali metals having a very high vapour pressure at elevated temperatures. This makes direct synthesis problematic because the alkali metals do not react with boron below 700 °C, and thus this | "Alkali metal" | [
0.014493979513645172,
0.6564711332321167,
0.1315443515777588,
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-0.5853905081748962,
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0.2749078869819641,
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0.3189776539802551,
-0.2649298906326294,
-0.31228965520858765,
-0.07943294197320938,
-0.8747591376304626,
0.63442492485046... |
9289 | must be accomplished in sealed containers with the alkali metal in excess. Furthermore, exceptionally in this group, reactivity with boron decreases down the group: lithium reacts completely at 700 °C, but sodium at 900 °C and potassium not until 1200 °C, and the reaction is instantaneous for lithium but takes hours for potassium. Rubidium and caesium borides have not even been characterised. Various phases are known, such as LiB, NaB, NaB, and KB. Under high pressure the boron–boron bonding in the lithium borides changes from following Wade's rules to forming Zintl anions like the rest of group 13. Lithium and | "Alkali metal" | [
-0.12642808258533478,
0.5441271662712097,
0.25221583247184753,
-0.34216970205307007,
-0.41202831268310547,
-0.1679648756980896,
0.2992158830165863,
-0.7350479960441589,
0.2832816243171692,
0.06916449218988419,
-0.24450021982192993,
0.16180890798568726,
-0.47291237115859985,
0.5213533639907... |
9290 | sodium react with carbon to form acetylides, LiC and NaC, which can also be obtained by reaction of the metal with acetylene. Potassium, rubidium, and caesium react with graphite; their atoms are intercalated between the hexagonal graphite layers, forming graphite intercalation compounds of formulae MC (dark grey, almost black), MC (dark grey, almost black), MC (blue), MC (steel blue), and MC (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. ). Upon heating of | "Alkali metal" | [
0.17192915081977844,
0.6887418627738953,
0.3339855670928955,
0.1219400018453598,
-0.3766138255596161,
0.09702747315168381,
0.11971964687108994,
-0.5159651041030884,
0.2696550190448761,
-0.16922977566719055,
-0.18022961914539337,
-0.07900930941104889,
-0.6059457063674927,
0.8400565385818481... |
9291 | KC, the elimination of potassium atoms results in the conversion in sequence to KC, KC, KC and finally KC. KC is a very strong reducing agent and is pyrophoric and explodes on contact with water. While the larger alkali metals (K, Rb, and Cs) initially form MC, the smaller ones initially form MC, and indeed they require reaction of the metals with graphite at high temperatures around 500 °C to form. Apart from this, the alkali metals are such strong reducing agents that they can even reduce buckminsterfullerene to produce solid fullerides MC; sodium, potassium, rubidium, and caesium can form | "Alkali metal" | [
-0.10590221732854843,
0.7051925659179688,
0.35999271273612976,
0.18909581005573273,
-0.46072492003440857,
-0.047455091029405594,
0.33663201332092285,
-0.6388019919395447,
0.20790204405784607,
-0.3202517330646515,
-0.2890113890171051,
-0.03993876650929451,
-0.6538894772529602,
0.85261851549... |
9292 | fullerides where "n" = 2, 3, 4, or 6, and rubidium and caesium additionally can achieve "n" = 1. When the alkali metals react with the heavier elements in the carbon group (silicon, germanium, tin, and lead), ionic substances with cage-like structures are formed, such as the silicides MSi (M = K, Rb, or Cs), which contains M and tetrahedral ions. The chemistry of alkali metal germanides, involving the germanide ion Ge and other cluster (Zintl) ions such as , , , and [(Ge)], is largely analogous to that of the corresponding silicides. Alkali metal stannides are mostly ionic, sometimes | "Alkali metal" | [
-0.12875133752822876,
0.6566858291625977,
0.28846898674964905,
0.04653157666325569,
-0.45619305968284607,
0.17239700257778168,
0.34550854563713074,
-0.6388708353042603,
0.3548136353492737,
-0.34525033831596375,
-0.3417024612426758,
0.08780788630247116,
-0.4028809368610382,
0.76736485958099... |
9293 | with the stannide ion (Sn), and sometimes with more complex Zintl ions such as , which appears in tetrapotassium nonastannide (KSn). The monatomic plumbide ion (Pb) is unknown, and indeed its formation is predicted to be energetically unfavourable; alkali metal plumbides have complex Zintl ions, such as . These alkali metal germanides, stannides, and plumbides may be produced by reducing germanium, tin, and lead with sodium metal in liquid ammonia. Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with nitrogen at standard conditions, and its nitride is the only stable alkali metal nitride. Nitrogen | "Alkali metal" | [
-0.11763593554496765,
0.382945716381073,
0.0969931110739708,
0.24432288110256195,
-0.42647939920425415,
0.10571818798780441,
0.11210359632968903,
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0.4085191488265991,
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-0.2879790961742401,
0.15513718128204346,
-0.4319736361503601,
0.6048944592475891... |
9294 | is an unreactive gas because breaking the strong triple bond in the dinitrogen molecule (N) requires a lot of energy. The formation of an alkali metal nitride would consume the ionisation energy of the alkali metal (forming M ions), the energy required to break the triple bond in N and the formation of N ions, and all the energy released from the formation of an alkali metal nitride is from the lattice energy of the alkali metal nitride. The lattice energy is maximised with small, highly charged ions; the alkali metals do not form highly charged ions, only forming ions | "Alkali metal" | [
-0.04197796806693077,
0.48106661438941956,
0.15272659063339233,
-0.1650170236825943,
-0.41963350772857666,
0.07439824938774109,
0.23549053072929382,
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0.5246641039848328,
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-0.16546601057052612,
0.1174231469631195,
-0.7211953401565552,
0.7124626040458... |
9295 | with a charge of +1, so only lithium, the smallest alkali metal, can release enough lattice energy to make the reaction with nitrogen exothermic, forming lithium nitride. The reactions of the other alkali metals with nitrogen would not release enough lattice energy and would thus be endothermic, so they do not form nitrides at standard conditions. Sodium nitride (NaN) and potassium nitride (KN), while existing, are extremely unstable, being prone to decomposing back into their constituent elements, and cannot be produced by reacting the elements with each other at standard conditions. Steric hindrance forbids the existence of rubidium or caesium | "Alkali metal" | [
0.02328193187713623,
0.3310304880142212,
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0.03767324239015579,
0.260580837726593,
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0.5100975036621094,
-0.13939397037029266,
0.013194317929446697,
0.04275858774781227,
-0.5407944321632385,
0.71796292066574... |
9296 | nitride. However, sodium and potassium form colourless azide salts involving the linear anion; due to the large size of the alkali metal cations, they are thermally stable enough to be able to melt before decomposing. All the alkali metals react readily with phosphorus and arsenic to form phosphides and arsenides with the formula MPn (where M represents an alkali metal and Pn represents a pnictogen – phosphorus, arsenic, antimony, or bismuth). This is due to the greater size of the P and As ions, so that less lattice energy needs to be released for the salts to form. These are | "Alkali metal" | [
0.029147034510970116,
0.4563349485397339,
0.18070368468761444,
0.010866322554647923,
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0.2568001449108124,
0.35369551181793213,
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0.12082107365131378,
-0.3511580228805542,
-0.22227659821510315,
0.07310545444488525,
-0.6967618465423584,
0.72924160957336... |
9297 | not the only phosphides and arsenides of the alkali metals: for example, potassium has nine different known phosphides, with formulae KP, KP, KP, KP, KP, KP, KP, KP, and KP. While most metals form arsenides, only the alkali and alkaline earth metals form mostly ionic arsenides. The structure of NaAs is complex with unusually short Na–Na distances of 328–330 pm which are shorter than in sodium metal, and this indicates that even with these electropositive metals the bonding cannot be straightforwardly ionic. Other alkali metal arsenides not conforming to the formula MAs are known, such as LiAs, which has a | "Alkali metal" | [
-0.16946786642074585,
0.63249272108078,
0.02753760851919651,
-0.007301622070372105,
-0.6264808177947998,
0.44638004899024963,
0.3913274109363556,
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0.14523321390151978,
-0.23706457018852234,
-0.06271537393331528,
-0.23483145236968994,
-0.33859503269195557,
0.742834627628... |
9298 | metallic lustre and electrical conductivity indicating the presence of some metallic bonding. The antimonides are unstable and reactive as the Sb ion is a strong reducing agent; reaction of them with acids form the toxic and unstable gas stibine (SbH). Indeed, they have some metallic properties, and the alkali metal antimonides of stoichiometry MSb involve antimony atoms bonded in a spiral Zintl structure. Bismuthides are not even wholly ionic; they are intermetallic compounds containing partially metallic and partially ionic bonds. All the alkali metals react vigorously with oxygen at standard conditions. They form various types of oxides, such as simple | "Alkali metal" | [
0.17757989466190338,
0.683880090713501,
0.013736831955611706,
-0.0449078232049942,
-0.7066196799278259,
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0.6144564747810364,
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0.506278395652771,
-0.4200480282306671,
-0.31220731139183044,
0.08732477575540543,
-0.6543669104576111,
0.6422457098960876,... |
9299 | oxides (containing the O ion), peroxides (containing the ion, where there is a single bond between the two oxygen atoms), superoxides (containing the ion), and many others. Lithium burns in air to form lithium oxide, but sodium reacts with oxygen to form a mixture of sodium oxide and sodium peroxide. Potassium forms a mixture of potassium peroxide and potassium superoxide, while rubidium and caesium form the superoxide exclusively. Their reactivity increases going down the group: while lithium, sodium and potassium merely burn in air, rubidium and caesium are pyrophoric (spontaneously catch fire in air). The smaller alkali metals tend to | "Alkali metal" | [
0.1063280701637268,
0.7086466550827026,
0.026145242154598236,
-0.008827343583106995,
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0.22699962556362152,
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0.5108057856559753,
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-0.20245040953159332,
0.05089084431529045,
-0.6456126570701599,
0.49393224716186... |
9300 | polarise the larger anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, | "Alkali metal" | [
-0.02183729223906994,
0.6597999334335327,
0.15308550000190735,
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-0.2253090739250183,
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0.18989405035972595,
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0.49926528334617615,
-0.25740575790405273,
-0.01580711081624031,
-0.10649934411048889,
-0.6797134280204773,
0.4595291316... |
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