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pant in semiconductor devices.
Characteristics
Properties
Antimony is a member of group 15 of the periodic table, one of the elements called pnictogens, and has an electronegativity of 2.05. In accordance with periodic trends, it is more electronegative than tin or bismuth, and less electronegative than tellurium or arsenic. Antimony is stable in air at room temperature, but reacts with oxygen if heated to produce antimony trioxide, Sb2O3.
Antimony is a silvery, lustrous gray metalloid with a Mohs scale hardness of 3, which is too soft to make hard objects. Coins of antimony were issued in China's Guizhou province in 1931; durability was poor, and minting was soon discontinued. Antimony is resistant to attack by acids.
Four allotropes of antimony are known a stable metallic form, and three metastable forms explosive, black, and yellow. Elemental antimony is a brittle, silverwhite, shiny metalloid. When slowly cooled, molten antimony crystallizes into a trigonal cell, isomorphic with the gray allotrope of
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arsenic. A rare explosive form of antimony can be formed from the electrolysis of antimony trichloride. When scratched with a sharp implement, an exothermic reaction occurs and white fumes are given off as metallic antimony forms; when rubbed with a pestle in a mortar, a strong detonation occurs. Black antimony is formed upon rapid cooling of antimony vapor. It has the same crystal structure as red phosphorus and black arsenic; it oxidizes in air and may ignite spontaneously. At 100 C, it gradually transforms into the stable form. The yellow allotrope of antimony is the most unstable; it has been generated only by oxidation of stibine SbH3 at 90 C. Above this temperature and in ambient light, this metastable allotrope transforms into the more stable black allotrope.
Elemental antimony adopts a layered structure space group Rm No. 166 whose layers consist of fused, ruffled, sixmembered rings. The nearest and nextnearest neighbors form an irregular octahedral complex, with the three atoms in each double layer
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slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 6.697 gcm3, but the weak bonding between the layers leads to the low hardness and brittleness of antimony.
Isotopes
Antimony has two stable isotopes 121Sb with a natural abundance of 57.36 and 123Sb with a natural abundance of 42.64. It also has 35 radioisotopes, of which the longestlived is 125Sb with a halflife of 2.75 years. In addition, 29 metastable states have been characterized. The most stable of these is 120m1Sb with a halflife of 5.76 days. Isotopes that are lighter than the stable 123Sb tend to decay by decay, and those that are heavier tend to decay by decay, with some exceptions.
Occurrence
The abundance of antimony in the Earth's crust is estimated to be 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in more than 100 mineral species. Antimony is sometimes found natively e.g. o
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n Antimony Peak, but more frequently it is found in the sulfide stibnite Sb2S3 which is the predominant ore mineral.
Compounds
Antimony compounds are often classified according to their oxidation state SbIII and SbV. The 5 oxidation state is more stable.
Oxides and hydroxides
Antimony trioxide is formed when antimony is burnt in air. In the gas phase, the molecule of the compound is , but it polymerizes upon condensing. Antimony pentoxide can be formed only by oxidation with concentrated nitric acid. Antimony also forms a mixedvalence oxide, antimony tetroxide , which features both SbIII and SbV. Unlike oxides of phosphorus and arsenic, these oxides are amphoteric, do not form welldefined oxoacids, and react with acids to form antimony salts.
Antimonous acid is unknown, but the conjugate base sodium antimonite forms upon fusing sodium oxide and . Transition metal antimonites are also known. Antimonic acid exists only as the hydrate , forming salts as the antimonate anion . When a solution containing th
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is anion is dehydrated, the precipitate contains mixed oxides.
Many antimony ores are sulfides, including stibnite , pyrargyrite , zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is nonstoichiometric and features antimony in the 3 oxidation state and SS bonds. Several thioantimonides are known, such as and .
Halides
Antimony forms two series of halides and . The trihalides , , , and are all molecular compounds having trigonal pyramidal molecular geometry.
The trifluoride is prepared by the reaction of with HF
6 HF 2 3
It is Lewis acidic and readily accepts fluoride ions to form the complex anions and . Molten is a weak electrical conductor. The trichloride is prepared by dissolving in hydrochloric acid
6 HCl 2 3
The pentahalides and have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, is polymeric, whereas is monomeric. is a powerful Lewis acid used to make the superacid fluoroantimonic acid "H2SbF7".
Oxyhalides are more common
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for antimony than for arsenic and phosphorus. Antimony trioxide dissolves in concentrated acid to form oxoantimonyl compounds such as SbOCl and .
Antimonides, hydrides, and organoantimony compounds
Compounds in this class generally are described as derivatives of Sb3. Antimony forms antimonides with metals, such as indium antimonide InSb and silver antimonide . The alkali metal and zinc antimonides, such as Na3Sb and Zn3Sb2, are more reactive. Treating these antimonides with acid produces the highly unstable gas stibine,
3
Stibine can also be produced by treating salts with hydride reagents such as sodium borohydride. Stibine decomposes spontaneously at room temperature. Because stibine has a positive heat of formation, it is thermodynamically unstable and thus antimony does not react with hydrogen directly.
Organoantimony compounds are typically prepared by alkylation of antimony halides with Grignard reagents. A large variety of compounds are known with both SbIII and SbV centers, including mixed c
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hloroorganic derivatives, anions, and cations. Examples include SbC6H53 triphenylstibine, Sb2C6H54 with an SbSb bond, and cyclic SbC6H5n. Pentacoordinated organoantimony compounds are common, examples being SbC6H55 and several related halides.
History
AntimonyIII sulfide, Sb2S3, was recognized in predynastic Egypt as an eye cosmetic kohl as early as about 3100 BC, when the cosmetic palette was invented.
An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea part of presentday Iraq, and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892, commented that "we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable 'find' artifact mentioned above must represent the lost art of rendering antimony malleable."
The British archaeologist Roger Moorey was unco
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nvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object published in 1975, "attempted to relate the metal to Transcaucasian natural antimony" i.e. native metal and that "the antimony objects from Transcaucasia are all small personal ornaments." This weakens the evidence for a lost art "of rendering antimony malleable."
The Roman scholar Pliny the Elder described several ways of preparing antimony sulfide for medical purposes in his treatise Natural History, around 77 AD. Pliny the Elder also made a distinction between "male" and "female" forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony.
The Greek naturalist Pedanius Dioscorides mentioned that antimony sulfide could be roasted by heating by a current of air. It is thought that this produced metallic antimony.
The intentional isolation of antimony is described by Jabir ibn Hayyan b
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efore 815 AD. A description of a procedure for isolating antimony is later given in the 1540 book De la pirotechnia by Vannoccio Biringuccio, predating the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii The Triumphal Chariot of Antimony, describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to be written by a Benedictine monk, writing under the name Basilius Valentinus in the 15th century; if it were authentic, which it is not, it would predate Biringuccio.
The metal antimony was known to German chemist Andreas Libavius in 1615 who obtained it by adding iron to a molten mixture of antimony sulfide, salt and potassium tartrate. This procedure produced antimony with a crystalline or starred surface.
With the advent of challenges to phlogiston theory, it was recognized that antimony is an element forming sulfides, oxides,
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and other compounds, as do other metals.
The first discovery of naturally occurring pure antimony in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the typesample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Vstmanland, Sweden.
Etymology
The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from antimonachos or French antimoine, still has adherents; this would mean "monkkiller", and is explained by many early alchemists being monks, and antimony being poisonous. However, the low toxicity of antimony see below makes this unlikely.
Another popular etymology is the hypothetical Greek word antimonos, "against aloneness", explained as "not found as metal", or "not found unalloyed". Lippmann conjectured a hypo
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thetical Greek word anthemonion, which would mean "floret", and cites several examples of related Greek words but not that one which describe chemical or biological efflorescence.
The early uses of antimonium include the translations, in 10501100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical asstimmi, derived from or parallel to the Greek.
The standard chemical symbol for antimony Sb is credited to Jns Jakob Berzelius, who derived the abbreviation from stibium.
The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony.
The Egyptians called antimony mdmt; in hieroglyphs, the vowels are uncertain, but the Coptic form of the word is stm.
Egyptian stm O34D46G17F21D4
The Greek word, stimmi is used by Attic tragic poets of the 5th century BC, and is pos
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sibly a loan word from Arabic or from Egyptian stm. Later Greeks also used stibi, as did Celsus and Pliny, writing in Latin, in the first century AD. Pliny also gives the names stimi, larbaris, alabaster, and the "very common" platyophthalmos, "wideeye" from the effect of the cosmetic. Later Latin authors adapted the word to Latin as stibium.
The Arabic word for the substance, as opposed to the cosmetic, can appear as ithmid, athmoud, othmod, or uthmod. Littr suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi.
Production
Process
The extraction of antimony from ores depends on the quality and composition of the ore. Most antimony is mined as the sulfide; lowergrade ores are concentrated by froth flotation, while highergrade ores are heated to 500600 C, the temperature at which stibnite melts and separates from the gangue minerals. Antimony can be isolated from the crude antimony sulfide by reduction with scrap iron
3 Fe 2 Sb 3 FeS
The sulfide is converted
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to an oxide; the product is then roasted, sometimes for the purpose of vaporizing the volatile antimonyIII oxide, which is recovered. This material is often used directly for the main applications, impurities being arsenic and sulfide. Antimony is isolated from the oxide by a carbothermal reduction
2 3 C 4 Sb 3
The lowergrade ores are reduced in blast furnaces while the highergrade ores are reduced in reverberatory furnaces.
Top producers and production volumes
The British Geological Survey BGS reported that in 2005 China was the top producer of antimony with approximately 84 of the world share, followed at a distance by South Africa, Bolivia and Tajikistan. Xikuangshan Mine in Hunan province has the largest deposits in China with an estimated deposit of 2.1 million metric tons.
In 2016, according to the US Geological Survey, China accounted for 76.9 of total antimony production, followed in second place by Russia with 6.9 and Tajikistan with 6.2.
Chinese production of antimony is expected to decli
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ne in the future as mines and smelters are closed down by the government as part of pollution control. Especially due to an environmental protection law having gone into effect in January 2015 and revised "Emission Standards of Pollutants for Stanum, Antimony, and Mercury" having gone into effect, hurdles for economic production are higher. According to the National Bureau of Statistics in China, by September 2015 50 of antimony production capacity in the Hunan province the province with biggest antimony reserves in China had not been used.
Reported production of antimony in China has fallen and is unlikely to increase in the coming years, according to the Roskill report. No significant antimony deposits in China have been developed for about ten years, and the remaining economic reserves are being rapidly depleted.
The world's largest antimony producers, according to Roskill, are listed below
Reserves
Supply risk
For antimonyimporting regions such as Europe and the U.S., antimony is considered to be a cr
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itical mineral for industrial manufacturing that is at risk of supply chain disruption. With global production coming mainly from China 74, Tajikistan8, and Russia4, these sources are critical to supply.
European Union Antimony is considered a critical raw material for defense, automotive, construction and textiles. The E.U. sources are 100 imported, coming mainly from Turkey 62, Bolivia 20 and Guatemala 7.
United Kingdom The British Geological Survey's 2015 risk list ranks antimony second highest after rare earth elements on the relative supply risk index.
United States Antimony is a mineral commodity considered critical to the economic and national security. In 2021, no antimony was mined in the U.S.
Applications
About 60 of antimony is consumed in flame retardants, and 20 is used in alloys for batteries, plain bearings, and solders.
Flame retardants
Antimony is mainly used as the trioxide for flameproofing compounds, always in combination with halogenated flame retardants except in halogencontaining
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polymers. The flame retarding effect of antimony trioxide is produced by the formation of halogenated antimony compounds, which react with hydrogen atoms, and probably also with oxygen atoms and OH radicals, thus inhibiting fire. Markets for these flameretardants include children's clothing, toys, aircraft, and automobile seat covers. They are also added to polyester resins in fiberglass composites for such items as light aircraft engine covers. The resin will burn in the presence of an externally generated flame, but will extinguish when the external flame is removed.
Alloys
Antimony forms a highly useful alloy with lead, increasing its hardness and mechanical strength. For most applications involving lead, varying amounts of antimony are used as alloying metal. In leadacid batteries, this addition improves plate strength and charging characteristics. For sailboats, lead keels are used to provide righting moment, ranging from 600 lbs to over 200 tons for the largest sailing superyachts; to improve hardnes
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s and tensile strength of the lead keel, antimony is mixed with lead between 2 and 5 by volume. Antimony is used in antifriction alloys such as Babbitt metal, in bullets and lead shot, electrical cable sheathing, type metal for example, for linotype printing machines, solder some "leadfree" solders contain 5 Sb, in pewter, and in hardening alloys with low tin content in the manufacturing of organ pipes.
Other applications
Three other applications consume nearly all the rest of the world's supply. One application is as a stabilizer and catalyst for the production of polyethylene terephthalate. Another is as a fining agent to remove microscopic bubbles in glass, mostly for TV screens antimony ions interact with oxygen, suppressing the tendency of the latter to form bubbles. The third application is pigments.
In 1990s antimony was increasingly being used in semiconductors as a dopant in ntype silicon wafers for diodes, infrared detectors, and Halleffect devices. In the 1950s, the emitters and collectors of
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npn alloy junction transistors were doped with tiny beads of a leadantimony alloy. Indium antimonide is used as a material for midinfrared detectors.
Biology and medicine have few uses for antimony. Treatments containing antimony, known as antimonials, are used as emetics. Antimony compounds are used as antiprotozoan drugs. Potassium antimonyl tartrate, or tartar emetic, was once used as an antischistosomal drug from 1919 on. It was subsequently replaced by praziquantel. Antimony and its compounds are used in several veterinary preparations, such as anthiomaline and lithium antimony thiomalate, as a skin conditioner in ruminants. Antimony has a nourishing or conditioning effect on keratinized tissues in animals.
Antimonybased drugs, such as meglumine antimoniate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Besides having low therapeutic indices, the drugs have minimal penetration of the bone marrow, where some of the Leishmania amastigotes reside, and curing t
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he disease especially the visceral form is very difficult. Elemental antimony as an antimony pill was once used as a medicine. It could be reused by others after ingestion and elimination.
AntimonyIII sulfide is used in the heads of some safety matches. Antimony sulfides help to stabilize the friction coefficient in automotive brake pad materials. Antimony is used in bullets, bullet tracers, paint, glass art, and as an opacifier in enamel. Antimony124 is used together with beryllium in neutron sources; the gamma rays emitted by antimony124 initiate the photodisintegration of beryllium. The emitted neutrons have an average energy of 24 keV. Natural antimony is used in startup neutron sources.
Historically, the powder derived from crushed antimony kohl has been applied to the eyes with a metal rod and with one's spittle, thought by the ancients to aid in curing eye infections. The practice is still seen in Yemen and in other Muslim countries.
Precautions
The effects of antimony and its compounds on human
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and environmental health differ widely. Elemental antimony metal does not affect human and environmental health. Inhalation of antimony trioxide and similar poorly soluble SbIII dust particles such as antimony dust is considered harmful and suspected of causing cancer. However, these effects are only observed with female rats and after longterm exposure to high dust concentrations. The effects are hypothesized to be attributed to inhalation of poorly soluble Sb particles leading to impaired lung clearance, lung overload, inflammation and ultimately tumour formation, not to exposure to antimony ions OECD, 2008. Antimony chlorides are corrosive to skin. The effects of antimony are not comparable to those of arsenic; this might be caused by the significant differences of uptake, metabolism, and excretion between arsenic and antimony.
For oral absorption, ICRP 1994 has recommended values of 10 for tartar emetic and 1 for all other antimony compounds. Dermal absorption for metals is estimated to be at most 1 HER
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AG, 2007. Inhalation absorption of antimony trioxide and other poorly soluble SbIII substances such as antimony dust is estimated at 6.8 OECD, 2008, whereas a value 1 is derived for SbV substances. AntimonyV is not quantitatively reduced to antimonyIII in the cell, and both species exist simultaneously.
Antimony is mainly excreted from the human body via urine. Antimony and its compounds do not cause acute human health effects, with the exception of antimony potassium tartrate "tartar emetic", a prodrug that is intentionally used to treat leishmaniasis patients.
Prolonged skin contact with antimony dust may cause dermatitis. However, it was agreed at the European Union level that the skin rashes observed are not substancespecific, but most probably due to a physical blocking of sweat ducts ECHAPR0909, Helsinki, 6 July 2009. Antimony dust may also be explosive when dispersed in the air; when in a bulk solid it is not combustible.
Antimony is incompatible with strong acids, halogenated acids, and oxidizers;
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when exposed to newly formed hydrogen it may form stibine SbH3.
The 8hour timeweighted average TWA is set at 0.5 mgm3 by the American Conference of Governmental Industrial Hygienists and by the Occupational Safety and Health Administration OSHA as a legal permissible exposure limit PEL in the workplace. The National Institute for Occupational Safety and Health NIOSH has set a recommended exposure limit REL of 0.5 mgm3 as an 8hour TWA.
Antimony compounds are used as catalysts for polyethylene terephthalate PET production. Some studies report minor antimony leaching from PET bottles into liquids, but levels are below drinking water guidelines. Antimony concentrations in fruit juice concentrates were somewhat higher up to 44.7 gL of antimony, but juices do not fall under the drinking water regulations. The drinking water guidelines are
World Health Organization 20 gL
Japan 15 gL
United States Environmental Protection Agency, Health Canada and the Ontario Ministry of Environment 6 gL
EU and German Federal M
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inistry of Environment 5 gL
The tolerable daily intake TDI proposed by WHO is 6 g antimony per kilogram of body weight. The immediately dangerous to life or health IDLH value for antimony is 50 mgm3.
Toxicity
Certain compounds of antimony appear to be toxic, particularly antimony trioxide and antimony potassium tartrate. Effects may be similar to arsenic poisoning. Occupational exposure may cause respiratory irritation, pneumoconiosis, antimony spots on the skin, gastrointestinal symptoms, and cardiac arrhythmias. In addition, antimony trioxide is potentially carcinogenic to humans.
Adverse health effects have been observed in humans and animals following inhalation, oral, or dermal exposure to antimony and antimony compounds. Antimony toxicity typically occurs either due to occupational exposure, during therapy or from accidental ingestion. It is unclear if antimony can enter the body through the skin. The presence of low levels of antimony in saliva may also be associated with dental decay.
See also
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Phase change memory
Notes
References
Bibliography
Edmund Oscar von Lippmann 1919 Entstehung und Ausbreitung der Alchemie, teil 1. Berlin Julius Springer in German.
Public Health Statement for Antimony
External links
International Antimony Association vzw i2a
Chemistry in its element podcast MP3 from the Royal Society of Chemistry's Chemistry World Antimony
Antimony at The Periodic Table of Videos University of Nottingham
CDC NIOSH Pocket Guide to Chemical Hazards Antimony
Antimony Mineral data and specimen images
Chemical elements
Metalloids
Native element minerals
Nuclear materials
Pnictogens
Trigonal minerals
Minerals in space group 166
Materials that expand upon freezing
Chemical elements with rhombohedral structure
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Actinium is a chemical element with the symbol Ac and atomic number 89. It was first isolated by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance AndrLouis Debierne found in 1899 and called actinium. Actinium gave the name to the actinide series, a group of 15 similar elements between actinium and lawrencium in the periodic table. Together with polonium, radium, and radon, actinium was one of the first nonprimordial radioactive elements to be isolated.
A soft, silverywhite radioactive metal, actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxide that prevents further oxidation. As with most lanthanides and many actinides, actinium assumes oxidation state 3 in nearly all its chemical compounds. Actinium is found only in traces in uranium and thorium ores as the isotope 227Ac, which decays with a halflife of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 22
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8Ac, which is beta active with a halflife of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium227, and one tonne of thorium contains about 5 nanograms of actinium228. The close similarity of physical and chemical properties of actinium and lanthanum makes separation of actinium from the ore impractical. Instead, the element is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor. Owing to its scarcity, high price and radioactivity, actinium has no significant industrial use. Its current applications include a neutron source and an agent for radiation therapy.
History
AndrLouis Debierne, a French chemist, announced the discovery of a new element in 1899. He separated it from pitchblende residues left by Marie and Pierre Curie after they had extracted radium. In 1899, Debierne described the substance as similar to titanium and in 1900 as similar to thorium. Friedrich Oskar Giesel found in 1902 a substance similar to lanthanum and called it
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"emanium" in 1904. After a comparison of the substances' halflives determined by Debierne, Harriet Brooks in 1904, and Otto Hahn and Otto Sackur in 1905, Debierne's chosen name for the new element was retained because it had seniority, despite the contradicting chemical properties he claimed for the element at different times.
Articles published in the 1970s and later suggest that Debierne's results published in 1904 conflict with those reported in 1899 and 1900. Furthermore, the nowknown chemistry of actinium precludes its presence as anything other than a minor constituent of Debierne's 1899 and 1900 results; in fact, the chemical properties he reported make it likely that he had, instead, accidentally identified protactinium, which would not be discovered for another fourteen years, only to have it disappear due to its hydrolysis and adsorption onto his laboratory equipment. This has led some authors to advocate that Giesel alone should be credited with the discovery. A less confrontational vision of scie
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ntific discovery is proposed by Adloff. He suggests that hindsight criticism of the early publications should be mitigated by the then nascent state of radiochemistry highlighting the prudence of Debierne's claims in the original papers, he notes that nobody can contend that Debierne's substance did not contain actinium. Debierne, who is now considered by the vast majority of historians as the discoverer, lost interest in the element and left the topic. Giesel, on the other hand, can rightfully be credited with the first preparation of radiochemically pure actinium and with the identification of its atomic number 89.
The name actinium originates from the Ancient Greek aktis, aktinos , , meaning beam or ray. Its symbol Ac is also used in abbreviations of other compounds that have nothing to do with actinium, such as acetyl, acetate and sometimes acetaldehyde.
Properties
Actinium is a soft, silverywhite, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong
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radioactivity, actinium glows in the dark with a pale blue light, which originates from the surrounding air ionized by the emitted energetic particles. Actinium has similar chemical properties to lanthanum and other lanthanides, and therefore these elements are difficult to separate when extracting from uranium ores. Solvent extraction and ion chromatography are commonly used for the separation.
The first element of the actinides, actinium gave the group its name, much as lanthanum had done for the lanthanides. The group of elements is more diverse than the lanthanides and therefore it was not until 1945 that the most significant change to Dmitri Mendeleev's periodic table since the recognition of the lanthanides, the introduction of the actinides, was generally accepted after Glenn T. Seaborg's research on the transuranium elements although it had been proposed as early as 1892 by British chemist Henry Bassett.
Actinium reacts rapidly with oxygen and moisture in air forming a white coating of actinium oxi
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de that impedes further oxidation. As with most lanthanides and actinides, actinium exists in the oxidation state 3, and the Ac3 ions are colorless in solutions. The oxidation state 3 originates from the Rn6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closedshell structure of the noble gas radon. The rare oxidation state 2 is only known for actinium dihydride AcH2; even this may in reality be an electride compound like its lighter congener LaH2 and thus have actiniumIII. Ac3 is the largest of all known tripositive ions and its first coordination sphere contains approximately 10.9 0.5 water molecules.
Chemical compounds
Due to actinium's intense radioactivity, only a limited number of actinium compounds are known. These include AcF3, AcCl3, AcBr3, AcOF, AcOCl, AcOBr, Ac2S3, Ac2O3, AcPO4 and AcNO33. Except for AcPO4, they are all similar to the corresponding lanthanum compounds. They all contain actinium in the oxidation state 3. In partic
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ular, the lattice constants of the analogous lanthanum and actinium compounds differ by only a few percent.
Here a, b and c are lattice constants, No is space group number and Z is the number of formula units per unit cell. Density was not measured directly but calculated from the lattice parameters.
Oxides
Actinium oxide Ac2O3 can be obtained by heating the hydroxide at 500 C or the oxalate at 1100 C, in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rareearth metals.
Halides
Actinium trifluoride can be produced either in solution or in solid reaction. The former reaction is carried out at room temperature, by adding hydrofluoric acid to a solution containing actinium ions. In the latter method, actinium metal is treated with hydrogen fluoride vapors at 700 C in an allplatinum setup. Treating actinium trifluoride with ammonium hydroxide at 9001000 C yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at 800 C for
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an hour, similar treatment of actinium trifluoride yields no AcOF and only results in melting of the initial product.
AcF3 2 NH3 H2O AcOF 2 NH4F
Actinium trichloride is obtained by reacting actinium hydroxide or oxalate with carbon tetrachloride vapors at temperatures above 960 C. Similar to oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium trichloride with ammonium hydroxide at 1000 C. However, in contrast to the oxyfluoride, the oxychloride could well be synthesized by igniting a solution of actinium trichloride in hydrochloric acid with ammonia.
Reaction of aluminium bromide and actinium oxide yields actinium tribromide
Ac2O3 2 AlBr3 2 AcBr3 Al2O3
and treating it with ammonium hydroxide at 500 C results in the oxybromide AcOBr.
Other compounds
Actinium hydride was obtained by reduction of actinium trichloride with potassium at 300 C, and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain.
Mixin
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g monosodium phosphate NaH2PO4 with a solution of actinium in hydrochloric acid yields whitecolored actinium phosphate hemihydrate AcPO40.5H2O, and heating actinium oxalate with hydrogen sulfide vapors at 1400 C for a few minutes results in a black actinium sulfide Ac2S3. It may possibly be produced by acting with a mixture of hydrogen sulfide and carbon disulfide on actinium oxide at 1000 C.
Isotopes
Naturally occurring actinium is composed of two radioactive isotopes; from the radioactive family of and a granddaughter of . decays mainly as a beta emitter with a very small energy, but in 1.38 of cases it emits an alpha particle, so it can readily be identified through alpha spectrometry. Thirtysix radioisotopes have been identified, the most stable being with a halflife of 21.772 years, with a halflife of 10.0 days and with a halflife of 29.37 hours. All remaining radioactive isotopes have halflives that are less than 10 hours and the majority of them have halflives shorter than one minute. The shor
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testlived known isotope of actinium is halflife of 69 nanoseconds which decays through alpha decay. Actinium also has two known meta states. The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac.
Purified comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772year halflife emitting mostly beta 98.62 and some alpha particles 1.38; the successive decay products are part of the actinium series. Owing to the low available amounts, low energy of its beta particles maximum 44.8 keV and low intensity of alpha radiation, is difficult to detect directly by its emission and it is therefore traced via its decay products. The isotopes of actinium range in atomic weight from 205 u to 236 u .
Occurrence and synthesis
Actinium is found only in traces in uranium ores one tonne of uranium in ore contains about 0.2 milligrams of 227Ac and in thorium ores, which contain about 5 nanograms of 228Ac per one tonne of thorium. The actinium isotope 227Ac
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is a transient member of the uraniumactinium series decay chain, which begins with the parent isotope 235U or 239Pu and ends with the stable lead isotope 207Pb. The isotope 228Ac is a transient member of the thorium series decay chain, which begins with the parent isotope 232Th and ends with the stable lead isotope 208Pb. Another actinium isotope 225Ac is transiently present in the neptunium series decay chain, beginning with 237Np or 233U and ending with thallium 205Tl and nearstable bismuth 209Bi; even though all primordial 237Np has decayed away, it is continuously produced by neutron knockout reactions on natural 238U.
The low natural concentration, and the close similarity of physical and chemical properties to those of lanthanum and other lanthanides, which are always abundant in actiniumbearing ores, render separation of actinium from the ore impractical, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of in a nuclear reactor
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.
22688Ra 10n 22788Ra beta42.2 cemin 22789Ac
The reaction yield is about 2 of the radium weight. 227Ac can further capture neutrons resulting in small amounts of 228Ac. After the synthesis, actinium is separated from radium and from the products of decay and nuclear fusion, such as thorium, polonium, lead and bismuth. The extraction can be performed with thenoyltrifluoroacetonebenzene solution from an aqueous solution of the radiation products, and the selectivity to a certain element is achieved by adjusting the pH to about 6.0 for actinium. An alternative procedure is anion exchange with an appropriate resin in nitric acid, which can result in a separation factor of 1,000,000 for radium and actinium vs. thorium in a twostage process. Actinium can then be separated from radium, with a ratio of about 100, using a low crosslinking cation exchange resin and nitric acid as eluant.
225Ac was first produced artificially at the Institute for Transuranium Elements ITU in Germany using a cyclotron and at St Georg
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e Hospital in Sydney using a linac in 2000. This rare isotope has potential applications in radiation therapy and is most efficiently produced by bombarding a radium226 target with 2030 MeV deuterium ions. This reaction also yields 226Ac which however decays with a halflife of 29 hours and thus does not contaminate 225Ac.
Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor in vacuum at a temperature between 1100 and 1300 C. Higher temperatures resulted in evaporation of the product and lower ones lead to an incomplete transformation. Lithium was chosen among other alkali metals because its fluoride is most volatile.
Applications
Owing to its scarcity, high price and radioactivity, 227Ac currently has no significant industrial use, but 225Ac is currently being studied for use in cancer treatments such as targeted alpha therapies.
227Ac is highly radioactive and was therefore studied for use as an active element of radioisotope thermoelectric generators, for example in sp
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acecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americiumberyllium and radiumberyllium pairs. In all those applications, 227Ac a beta source is merely a progenitor which generates alphaemitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large crosssection for the ,n nuclear reaction
94Be 42He 126C 10n gamma
The 227AcBe neutron sources can be applied in a neutron probe a standard device for measuring the quantity of water present in soil, as well as moisturedensity for quality control in highway construction. Such probes are also used in well logging applications, in neutron radiography, tomography and other radiochemical investigations.
225Ac is applied in medicine to produce in a reusable generator or can be used alone as an agent for radiation therapy, in particular targeted alpha therapy TAT. This isotope has a halflife of 10 days, making it much more suitable
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for radiation therapy than 213Bi halflife 46 minutes. Additionally, 225Ac decays to nontoxic 209Bi rather than stable but toxic lead, which is the final product in the decay chains of several other candidate isotopes, namely 227Th, 228Th, and 230U. Not only 225Ac itself, but also its daughters, emit alpha particles which kill cancer cells in the body. The major difficulty with application of 225Ac was that intravenous injection of simple actinium complexes resulted in their accumulation in the bones and liver for a period of tens of years. As a result, after the cancer cells were quickly killed by alpha particles from 225Ac, the radiation from the actinium and its daughters might induce new mutations. To solve this problem, 225Ac was bound to a chelating agent, such as citrate, ethylenediaminetetraacetic acid EDTA or diethylene triamine pentaacetic acid DTPA. This reduced actinium accumulation in the bones, but the excretion from the body remained slow. Much better results were obtained with such chelating ag
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ents as HEHA or DOTA coupled to trastuzumab, a monoclonal antibody that interferes with the HER2neu receptor. The latter delivery combination was tested on mice and proved to be effective against leukemia, lymphoma, breast, ovarian, neuroblastoma and prostate cancers.
The medium halflife of 227Ac 21.77 years makes it very convenient radioactive isotope in modeling the slow vertical mixing of oceanic waters. The associated processes cannot be studied with the required accuracy by direct measurements of current velocities of the order 50 meters per year. However, evaluation of the concentration depthprofiles for different isotopes allows estimating the mixing rates. The physics behind this method is as follows oceanic waters contain homogeneously dispersed 235U. Its decay product, 231Pa, gradually precipitates to the bottom, so that its concentration first increases with depth and then stays nearly constant. 231Pa decays to 227Ac; however, the concentration of the latter isotope does not follow the 231Pa dep
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th profile, but instead increases toward the sea bottom. This occurs because of the mixing processes which raise some additional 227Ac from the sea bottom. Thus analysis of both 231Pa and 227Ac depth profiles allows researchers to model the mixing behavior.
There are theoretical predictions that AcHx hydrides in this case with very high pressure are a candidate for a near roomtemperature superconductor as they have Tc significantly higher than H3S, possibly near 250 K.
Precautions
227Ac is highly radioactive and experiments with it are carried out in a specially designed laboratory equipped with a tight glove box. When actinium trichloride is administered intravenously to rats, about 33 of actinium is deposited into the bones and 50 into the liver. Its toxicity is comparable to, but slightly lower than that of americium and plutonium. For trace quantities, fume hoods with good aeration suffice; for gram amounts, hot cells with shielding from the intense gamma radiation emitted by 227Ac are necessary.
See a
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lso
Actinium series
Notes
References
Bibliography
Meyer, Gerd and Morss, Lester R. 1991 Synthesis of lanthanide and actinide compounds, Springer.
External links
Actinium at The Periodic Table of Videos University of Nottingham
NLM Hazardous Substances Databank Actinium, Radioactive
Actinium in
Chemical elements
Actinides
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Americium is a synthetic radioactive chemical element with the symbol Am and atomic number 95. It is a transuranic member of the actinide series, in the periodic table located under the lanthanide element europium, and thus by analogy was named after the Americas.
Americium was first produced in 1944 by the group of Glenn T. Seaborg from Berkeley, California, at the Metallurgical Laboratory of the University of Chicago, as part of the Manhattan Project. Although it is the third element in the transuranic series, it was discovered fourth, after the heavier curium. The discovery was kept secret and only released to the public in November 1945. Most americium is produced by uranium or plutonium being bombarded with neutrons in nuclear reactors one tonne of spent nuclear fuel contains about 100 grams of americium. It is widely used in commercial ionization chamber smoke detectors, as well as in neutron sources and industrial gauges. Several unusual applications, such as nuclear batteries or fuel for space ships
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with nuclear propulsion, have been proposed for the isotope 242mAm, but they are as yet hindered by the scarcity and high price of this nuclear isomer.
Americium is a relatively soft radioactive metal with silvery appearance. Its most common isotopes are 241Am and 243Am. In chemical compounds, americium usually assumes the oxidation state 3, especially in solutions. Several other oxidation states are known, ranging from 2 to 7, and can be identified by their characteristic optical absorption spectra. The crystal lattice of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by selfirradiation with alpha particles, which accumulates with time; this can cause a drift of some material properties over time, more noticeable in older samples.
History
Although americium was likely produced in previous nuclear experiments, it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn
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T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60inch cyclotron at the University of California, Berkeley. The element was chemically identified at the Metallurgical Laboratory now Argonne National Laboratory of the University of Chicago. Following the lighter neptunium, plutonium, and heavier curium, americium was the fourth transuranium element to be discovered. At the time, the periodic table had been restructured by Seaborg to its present layout, containing the actinide row below the lanthanide one. This led to americium being located right below its twin lanthanide element europium; it was thus by analogy named after the Americas "The name americium after the Americas and the symbol Am are suggested for the element on the basis of its position as the sixth member of the actinide rareearth series, analogous to europium, Eu, of the lanthanide series."
The new element was isolated from its oxides in a complex, multistep process. First plutonium239 nitrate 239PuNO3 solution was
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coated on a platinum foil of about 0.5 cm2 area, the solution was evaporated and the residue was converted into plutonium dioxide PuO2 by calcining. After cyclotron irradiation, the coating was dissolved with nitric acid, and then precipitated as the hydroxide using concentrated aqueous ammonia solution. The residue was dissolved in perchloric acid. Further separation was carried out by ion exchange, yielding a certain isotope of curium. The separation of curium and americium was so painstaking that those elements were initially called by the Berkeley group as pandemonium from Greek for all demons or hell and delirium from Latin for madness.
Initial experiments yielded four americium isotopes 241Am, 242Am, 239Am and 238Am. Americium241 was directly obtained from plutonium upon absorption of two neutrons. It decays by emission of a particle to 237Np; the halflife of this decay was first determined as years but then corrected to 432.2 years.
The times are halflives
The second isotope 242Am was produced u
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pon neutron bombardment of the alreadycreated 241Am. Upon rapid decay, 242Am converts into the isotope of curium 242Cm which had been discovered previously. The halflife of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h.
The discovery of americium and curium in 1944 was closely related to the Manhattan Project; the results were confidential and declassified only in 1945. Seaborg leaked the synthesis of the elements 95 and 96 on the U.S. radio show for children Quiz Kids five days before the official presentation at an American Chemical Society meeting on 11 November 1945, when one of the listeners asked whether any new transuranium element besides plutonium and neptunium had been discovered during the war. After the discovery of americium isotopes 241Am and 242Am, their production and compounds were patented listing only Seaborg as the inventor. The initial americium samples weighed a few micrograms; they were barely visible and were identified
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by their radioactivity. The first substantial amounts of metallic americium weighing 40200 micrograms were not prepared until 1951 by reduction of americiumIII fluoride with barium metal in high vacuum at 1100 C.
Occurrence
The longestlived and most common isotopes of americium, 241Am and 243Am, have halflives of 432.2 and 7,370 years, respectively. Therefore, any primordial americium americium that was present on Earth during its formation should have decayed by now. Trace amounts of americium probably occur naturally in uranium minerals as a result of nuclear reactions, though this has not been confirmed.
Existing americium is concentrated in the areas used for the atmospheric nuclear weapons tests conducted between 1945 and 1980, as well as at the sites of nuclear incidents, such as the Chernobyl disaster. For example, the analysis of the debris at the testing site of the first U.S. hydrogen bomb, Ivy Mike, 1 November 1952, Enewetak Atoll, revealed high concentrations of various actinides including amer
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icium; but due to military secrecy, this result was not published until later, in 1956. Trinitite, the glassy residue left on the desert floor near Alamogordo, New Mexico, after the plutoniumbased Trinity nuclear bomb test on 16 July 1945, contains traces of americium241. Elevated levels of americium were also detected at the crash site of a US Boeing B52 bomber aircraft, which carried four hydrogen bombs, in 1968 in Greenland.
In other regions, the average radioactivity of surface soil due to residual americium is only about 0.01 picocuriesg 0.37 mBqg. Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900 times higher concentration of americium inside sandy soil particles than in the water present in the soil pores; an even higher ratio was measured in loam soils.
Americium is produced mostly artificially in small quantities, for research purposes. A tonne of spent nuclear fuel contains about 100 grams of various americ
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ium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other longlived actinides, must be neutralized. The associated procedure may involve several steps, where americium is first separated and then converted by neutron bombardment in special reactors to shortlived nuclides. This procedure is well known as nuclear transmutation, but it is still being developed for americium. The transuranic elements from americium to fermium occurred naturally in the natural nuclear fission reactor at Oklo, but no longer do so.
Americium is also one of the elements that have been detected in Przybylski's Star.
Synthesis and extraction
Isotope nucleosynthesis
Americium has been produced in small quantities in nuclear reactors for decades, and kilograms of its 241Am and 243Am isotopes have been accumulated by now. Nevertheless, since it was first offered for sale in 1962, its price, about US1,500 per gram of 241Am, remains almost unchanged
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owing to the very complex separation procedure. The heavier isotope 243Am is produced in much smaller amounts; it is thus more difficult to separate, resulting in a higher cost of the order 100,000160,000 USDg.
Americium is not synthesized directly from uranium the most common reactor material but from the plutonium isotope 239Pu. The latter needs to be produced first, according to the following nuclear process
23892U cen,gamma 23992U beta23.5 cemin 23993Np beta2.3565 ced 23994Pu
The capture of two neutrons by 239Pu a socalled n, reaction, followed by a decay, results in 241Am
23994Pu ce2n,gamma 24194Pu beta14.35 ceyr 24195Am
The plutonium present in spent nuclear fuel contains about 12 of 241Pu. Because it spontaneously converts to 241Am, 241Pu can be extracted and may be used to generate further 241Am. However, this process is rather slow half of the original amount of 241Pu decays to 241Am after about 15 years, and the 241Am amount reaches a maximum after 70 years.
The obtained 241Am can be us
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ed for generating heavier americium isotopes by further neutron capture inside a nuclear reactor. In a light water reactor LWR, 79 of 241Am converts to 242Am and 10 to its nuclear isomer 242mAm
Americium242 has a halflife of only 16 hours, which makes its further conversion to 243Am extremely inefficient. The latter isotope is produced instead in a process where 239Pu captures four neutrons under high neutron flux
23994Pu ce4n,gamma 24394Pu beta4.956 ceh 24395Am
Metal generation
Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of americium can be separated. In a typical procedure, the spent reactor fuel e.g. MOX fuel is dissolved in nitric acid, and the bulk of uranium and plutonium is removed using a PUREXtype extraction PlutoniumURanium EXtraction with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue raffinate by a diamidebased extraction, to give, after stripping, a mixtu
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re of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multistep chromatographic and centrifugation techniques with an appropriate reagent. A large amount of work has been done on the solvent extraction of americium. For example, a 2003 EUfunded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bistriazinyl bipyridine complex was proposed in 2009 as such a reagent is highly selective to americium and curium. Separation of americium from the highly similar curium can be achieved by treating a slurry of their hydroxides in aqueous sodium bicarbonate with ozone, at elevated temperatures. Both Am and Cm are mostly present in solutions in the 3 valence state; whereas curium remains unchanged, americium oxidizes to soluble AmIV complexes which can be washed away.
Metallic americium is obtained by reduction from its compounds. AmericiumIII fluoride was first used for this purpose. The reaction was conducted using element
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al barium as reducing agent in a water and oxygenfree environment inside an apparatus made of tantalum and tungsten.
An alternative is the reduction of americium dioxide by metallic lanthanum or thorium
Physical properties
In the periodic table, americium is located to the right of plutonium, to the left of curium, and below the lanthanide europium, with which it shares many physical and chemical properties. Americium is a highly radioactive element. When freshly prepared, it has a silverywhite metallic lustre, but then slowly tarnishes in air. With a density of 12 gcm3, americium is less dense than both curium 13.52 gcm3 and plutonium 19.8 gcm3; but has a higher density than europium 5.264 gcm3mostly because of its higher atomic mass. Americium is relatively soft and easily deformable and has a significantly lower bulk modulus than the actinides before it Th, Pa, U, Np and Pu. Its melting point of 1173 C is significantly higher than that of plutonium 639 C and europium 826 C, but lower than for curium
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1340 C.
At ambient conditions, americium is present in its most stable form which has a hexagonal crystal symmetry, and a space group P63mmc with cell parameters a 346.8 pm and c 1124 pm, and four atoms per unit cell. The crystal consists of a doublehexagonal close packing with the layer sequence ABAC and so is isotypic with lanthanum and several actinides such as curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, Am transforms to the modification, which has a facecentered cubic fcc symmetry, space group Fmm and lattice constant a 489 pm. This fcc structure is equivalent to the closest packing with the sequence ABC. Upon further compression to 23 GPa, americium transforms to an orthorhombic Am structure similar to that of uranium. There are no further transitions observed up to 52 GPa, except for an appearance of a monoclinic phase at pressures between 10 and 15 GPa. There is no consistency on the status of this phase in the li
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terature, which also sometimes lists the , and phases as I, II and III. The transition is accompanied by a 6 decrease in the crystal volume; although theory also predicts a significant volume change for the transition, it is not observed experimentally. The pressure of the transition decreases with increasing temperature, and when americium is heated at ambient pressure, at 770 C it changes into an fcc phase which is different from Am, and at 1075 C it converts to a bodycentered cubic structure. The pressuretemperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium.
As with many other actinides, selfdamage of the crystal structure due to alphaparticle irradiation is intrinsic to americium. It is especially noticeable at low temperatures, where the mobility of the produced structure defects is relatively low, by broadening of Xray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as el
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ectrical resistivity. So for americium241, the resistivity at 4.2 K increases with time from about 2 Ohmcm to 10 Ohmcm after 40 hours, and saturates at about 16 Ohmcm after 140 hours. This effect is less pronounced at room temperature, due to annihilation of radiation defects; also heating to room temperature the sample which was kept for hours at low temperatures restores its resistivity. In fresh samples, the resistivity gradually increases with temperature from about 2 Ohmcm at liquid helium to 69 Ohmcm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which show a rapid rise up to 60 K followed by saturation. The room temperature value for americium is lower than that of neptunium, plutonium and curium, but higher than for uranium, thorium and protactinium.
Americium is paramagnetic in a wide temperature range, from that of liquid helium, to room temperature and above. This behavior is markedly different from
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that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to along the shorter a axis and for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is , from which the standard enthalpy change of formation fH of aqueous Am3 ion is . The standard potential Am3Am0 is .
Chemical properties
Americium metal readily reacts with oxygen and dissolves in aqueous acids. The most stable oxidation state for americium is 3,. The chemistry of americiumIII has many similarities to the chemistry of lanthanideIII compounds. For example, trivalent americium forms insoluble fluoride, oxalate, iodate, hydroxide, phosphate and other salts. Compounds of americium in oxidation states 2, 4, 5, 6 and 7 have also been studied. This is the widest range that has been observed with actinide elements. The color of americium compounds in aqueous solution is as follows
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Am3 yellowreddish, Am4 yellowreddish, AmV; yellow, AmVI brown and AmVII dark green. The absorption spectra have sharp peaks, due to ff transitions' in the visible and nearinfrared regions. Typically, AmIII has absorption maxima at ca. 504 and 811 nm, AmV at ca. 514 and 715 nm, and AmVI at ca. 666 and 992 nm.
Americium compounds with oxidation state 4 and higher are strong oxidizing agents, comparable in strength to the permanganate ion in acidic solutions. Whereas the Am4 ions are unstable in solutions and readily convert to Am3, compounds such as americium dioxide AmO2 and americiumIV fluoride AmF4 are stable in the solid state.
The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the ion is unstable with respect to disproportionation. The reaction
3AmO2 4H 2AmO22 Am3 2H2O
is typical. The chemistry of AmV and AmVI is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like Li3AmO4 and Li6AmO6 are comparable to u
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ranates and the ion AmO22 is comparable to the uranyl ion, UO22. Such compounds can be prepared by oxidation of AmIII in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silverI oxide, ozone and sodium persulfate.
Chemical compounds
Oxygen compounds
Three americium oxides are known, with the oxidation states 2 AmO, 3 Am2O3 and 4 AmO2. AmericiumII oxide was prepared in minute amounts and has not been characterized in detail. AmericiumIII oxide is a redbrown solid with a melting point of 2205 C. AmericiumIV oxide is the main form of solid americium which is used in nearly all its applications. As most other actinide dioxides, it is a black solid with a cubic fluorite crystal structure.
The oxalate of americiumIII, vacuum dried at room temperature, has the chemical formula Am2C2O437H2O. Upon heating in vacuum, it loses water at 240 C and starts decomposing into AmO2 at 300 C, the decomposition completes at about 470 C. The initial oxalate dissolves in nitric acid
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with the maximum solubility of 0.25 gL.
Halides
Halides of americium are known for the oxidation states 2, 3 and 4, where the 3 is most stable, especially in solutions.
Reduction of AmIII compounds with sodium amalgam yields AmII salts the black halides AmCl2, AmBr2 and AmI2. They are very sensitive to oxygen and oxidize in water, releasing hydrogen and converting back to the AmIII state. Specific lattice constants are
Orthorhombic AmCl2 a , b and c
Tetragonal AmBr2 a and c . They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X Cl, Br or I
Am undersetmercury halideHgX2 atop 400 500 circ ce C AmX2 Hg
AmericiumIII fluoride AmF3 is poorly soluble and precipitates upon reaction of Am3 and fluoride ions in weak acidic solutions
Am3 3F AmF3v
The tetravalent americiumIV fluoride AmF4 is obtained by reacting solid americiumIII fluoride with molecular fluorine
2AmF3 F2 2AmF4
Another known form of solid tetravalent americium fluoride is
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KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black AmOH4 was dissolved in 15M NH4F with the americium concentration of 0.01 M. The resulting reddish solution had a characteristic optical absorption spectrum which is similar to that of AmF4 but differed from other oxidation states of americium. Heating the AmIV solution to 90 C did not result in its disproportionation or reduction, however a slow reduction was observed to AmIII and assigned to selfirradiation of americium by alpha particles.
Most americiumIII halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride AmCl3 is reddish and has a structure isotypic to uraniumIII chloride space group P63m and the melting point of 715 C. The fluoride is isotypic to LaF3 space group P63mmc and the iodide to BiI3 space group R. The bromide is an exception with the orthorhombic PuBr3type structure and space group Cmcm. Crystals of americium hexahydrate AmCl3
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6H2O can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellowreddish color and a monoclinic crystal structure.
Oxyhalides of americium in the form AmVIO2X2, AmVO2X, AmIVOX2 and AmIIIOX can be obtained by reacting the corresponding americium halide with oxygen or Sb2O3, and AmOCl can also be produced by vapor phase hydrolysis
AmCl3 H2O AmOCl 2HCl
Chalcogenides and pnictides
The known chalcogenides of americium include the sulfide AmS2, selenides AmSe2 and Am3Se4, and tellurides Am2Te3 and AmTe2. The pnictides of americium 243Am of the AmX type are known for the elements phosphorus, arsenic, antimony and bismuth. They crystallize in the rocksalt lattice.
Silicides and borides
Americium monosilicide AmSi and "disilicide" nominally AmSix with 1.87 x 2.0 were obtained by reduction of americiumIII fluoride with elementary silicon in vacuum at 1050 C AmSi and 11501200 C AmSix. AmSi is a black solid isomorphic with LaSi, i
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t has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice space group I41amd, it is isomorphic with PuSi2 and ThSi2. Borides of americium include AmB4 and AmB6. The tetraboride can be obtained by heating an oxide or halide of americium with magnesium diboride in vacuum or inert atmosphere.
Organoamericium compounds
Analogous to uranocene, americium forms the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula 8C8H82Am. A cyclopentadienyl complex is also known that is likely to be stoichiometrically AmCp3.
Formation of the complexes of the type AmnC3H7BTP3, where BTP stands for 2,6di1,2,4triazin3ylpyridine, in solutions containing nC3H7BTP and Am3 ions has been confirmed by EXAFS. Some of these BTPtype complexes selectively interact with americium and therefore are useful in its selective separation from lanthanides and another actinides.
Biological aspects
Americium is an artificial element of recent origin, and t
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hus does not have a biological requirement. It is harmful to life. It has been proposed to use bacteria for removal of americium and other heavy metals from rivers and streams. Thus, Enterobacteriaceae of the genus Citrobacter precipitate americium ions from aqueous solutions, binding them into a metalphosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi.
Fission
The isotope 242mAm halflife 141 years has the largest cross sections for absorption of thermal neutrons 5,700 barns, that results in a small critical mass for a sustained nuclear chain reaction. The critical mass for a bare 242mAm sphere is about 914 kg the uncertainty results from insufficient knowledge of its material properties. It can be lowered to 35 kg with a metal reflector and should become even smaller with a water reflector. Such small critical mass is favorable for portable nuclear weapons, but those based on 242mAm are not known yet, probably be
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cause of its scarcity and high price. The critical masses of two other readily available isotopes, 241Am and 243Am, are relatively high 57.6 to 75.6 kg for 241Am and 209 kg for 243Am. Scarcity and high price yet hinder application of americium as a nuclear fuel in nuclear reactors.
There are proposals of very compact 10kW highflux reactors using as little as 20 grams of 242mAm. Such lowpower reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals.
Isotopes
About 19 isotopes and 8 nuclear isomers are known for americium. There are two longlived alphaemitters; 243Am has a halflife of 7,370 years and is the most stable isotope, and 241Am has a halflife of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long halflife of 141 years. The halflives of other isotopes and isomers range from 0.64 microseconds for 245m1Am to 50.8 hours for 240Am. As with most other actinides, the isotopes of americium with odd number of neutrons have relatively high rate of nucl
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ear fission and low critical mass.
Americium241 decays to 237Np emitting alpha particles of 5 different energies, mostly at 5.486 MeV 85.2 and 5.443 MeV 12.8. Because many of the resulting states are metastable, they also emit gamma rays with the discrete energies between 26.3 and 158.5 keV.
Americium242 is a shortlived isotope with a halflife of 16.02 h. It mostly 82.7 converts by decay to 242Cm, but also by electron capture to 242Pu 17.3. Both 242Cm and 242Pu transform via nearly the same decay chain through 238Pu down to 234U.
Nearly all 99.541 of 242m1Am decays by internal conversion to 242Am and the remaining 0.459 by decay to 238Np. The latter subsequently decays to 238Pu and then to 234U.
Americium243 transforms by emission into 239Np, which converts by decay to 239Pu, and the 239Pu changes into 235U by emitting an particle.
Applications
Ionizationtype smoke detector
Americium is used in the most common type of household smoke detector, which uses 241Am in the form of americium dioxide as its so
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urce of ionizing radiation. This isotope is preferred over 226Ra because it emits 5 times more alpha particles and relatively little harmful gamma radiation.
The amount of americium in a typical new smoke detector is 1 microcurie 37 kBq or 0.29 microgram. This amount declines slowly as the americium decays into neptunium237, a different transuranic element with a much longer halflife about 2.14 million years. With its halflife of 432.2 years, the americium in a smoke detector includes about 3 neptunium after 19 years, and about 5 after 32 years. The radiation passes through an ionization chamber, an airfilled space between two electrodes, and permits a small, constant current between the electrodes. Any smoke that enters the chamber absorbs the alpha particles, which reduces the ionization and affects this current, triggering the alarm. Compared to the alternative optical smoke detector, the ionization smoke detector is cheaper and can detect particles which are too small to produce significant light scatter
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ing; however, it is more prone to false alarms.
Radionuclide
As 241Am has a roughly similar halflife to 238Pu 432.2 years vs. 87 years, it has been proposed as an active element of radioisotope thermoelectric generators, for example in spacecraft. Although americium produces less heat and electricity the power yield is 114.7 mWg for 241Am and 6.31 mWg for 243Am cf. 390 mWg for 238Pu and its radiation poses more threat to humans owing to neutron emission, the European Space Agency is considering using americium for its space probes.
Another proposed spacerelated application of americium is a fuel for space ships with nuclear propulsion. It relies on the very high rate of nuclear fission of 242mAm, which can be maintained even in a micrometerthick foil. Small thickness avoids the problem of selfabsorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alphaparticles. The fission products of 242mAm can either directly propel the spaceship or
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they can heat a thrusting gas. They can also transfer their energy to a fluid and generate electricity through a magnetohydrodynamic generator.
One more proposal which utilizes the high nuclear fission rate of 242mAm is a nuclear battery. Its design relies not on the energy of the emitted by americium alpha particles, but on their charge, that is the americium acts as the selfsustaining "cathode". A single 3.2 kg 242mAm charge of such battery could provide about 140 kW of power over a period of 80 days. Even with all the potential benefits, the current applications of 242mAm are as yet hindered by the scarcity and high price of this particular nuclear isomer.
In 2019, researchers at the UK National Nuclear Laboratory and the University of Leicester demonstrated the use of heat generated by americium to illuminate a small light bulb. This technology could lead to systems to power missions with durations up to 400 years into interstellar space, where solar panels do not function.
Neutron source
The oxide of
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241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large crosssection for the ,n nuclear reaction
24195Am 23793Np 42He gamma
94Be 42He 126C 10n gamma
The most widespread use of 241AmBe neutron sources is a neutron probe a device used to measure the quantity of water present in soil, as well as moisturedensity for quality control in highway construction. 241Am neutron sources are also used in well logging applications, as well as in neutron radiography, tomography and other radiochemical investigations.
Production of other elements
Americium is a starting material for the production of other transuranic elements and transactinides for example, 82.7 of 242Am decays to 242Cm and 17.3 to 242Pu. In the nuclear reactor, 242Am is also upconverted by neutron capture to 243Am and 244Am, which transforms by decay to 244Cm
24395Am cen,gamma 24495Am beta10.1 ceh 24496Cm
Irradiation of 241Am by 12C or 22Ne i
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ons yields the isotopes 247Es einsteinium or 260Db dubnium, respectively. Furthermore, the element berkelium 243Bk isotope had been first intentionally produced and identified by bombarding 241Am with alpha particles, in 1949, by the same Berkeley group, using the same 60inch cyclotron. Similarly, nobelium was produced at the Joint Institute for Nuclear Research, Dubna, Russia, in 1965 in several reactions, one of which included irradiation of 243Am with 15N ions. Besides, one of the synthesis reactions for lawrencium, discovered by scientists at Berkeley and Dubna, included bombardment of 243Am with 18O.
Spectrometer
Americium241 has been used as a portable source of both gamma rays and alpha particles for a number of medical and industrial uses. The 59.5409 keV gamma ray emissions from 241Am in such sources can be used for indirect analysis of materials in radiography and Xray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers. For example, the
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element has been employed to gauge glass thickness to help create flat glass. Americium241 is also suitable for calibration of gammaray spectrometers in the lowenergy range, since its spectrum consists of nearly a single peak and negligible Compton continuum at least three orders of magnitude lower intensity. Americium241 gamma rays were also used to provide passive diagnosis of thyroid function. This medical application is however obsolete.
Health concerns
As a highly radioactive element, americium and its compounds must be handled only in an appropriate laboratory under special arrangements. Although most americium isotopes predominantly emit alpha particles which can be blocked by thin layers of common materials, many of the daughter products emit gammarays and neutrons which have a long penetration depth.
If consumed, most of the americium is excreted within a few days, with only 0.05 absorbed in the blood, of which roughly 45 goes to the liver and 45 to the bones, and the remaining 10 is excreted. The
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uptake to the liver depends on the individual and increases with age. In the bones, americium is first deposited over cortical and trabecular surfaces and slowly redistributes over the bone with time. The biological halflife of 241Am is 50 years in the bones and 20 years in the liver, whereas in the gonads testicles and ovaries it remains permanently; in all these organs, americium promotes formation of cancer cells as a result of its radioactivity.
Americium often enters landfills from discarded smoke detectors. The rules associated with the disposal of smoke detectors are relaxed in most jurisdictions. In 1994, 17yearold David Hahn extracted the americium from about 100 smoke detectors in an attempt to build a breeder nuclear reactor. There have been a few cases of exposure to americium, the worst case being that of chemical operations technician Harold McCluskey, who at the age of 64 was exposed to 500 times the occupational standard for americium241 as a result of an explosion in his lab. McCluskey died
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at the age of 75 of unrelated preexisting disease.
See also
Actinides in the environment
CategoryAmericium compounds
Notes
References
Bibliography
Penneman, R. A. and Keenan T. K. The radiochemistry of americium and curium, University of California, Los Alamos, California, 1960
Further reading
Nuclides and Isotopes 14th Edition, GE Nuclear Energy, 1989.
External links
Americium at The Periodic Table of Videos University of Nottingham
ATSDR Public Health Statement Americium
World Nuclear Association Smoke Detectors and Americium
Chemical elements
Actinides
Carcinogens
Synthetic elements
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Astatine is a chemical element with the symbol At and atomic number 85. It is the rarest naturally occurring element in the Earth's crust, occurring only as the decay product of various heavier elements. All of astatine's isotopes are shortlived; the most stable is astatine210, with a halflife of 8.1 hours. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity.
The bulk properties of astatine are not known with certainty. Many of them have been estimated based on the element's position on the periodic table as a heavier analog of iodine, and a member of the halogens the group of elements including fluorine, chlorine, bromine, and iodine. However, astatine also falls roughly along the dividing line between metals and nonmetals, and some metallic behavior has also been observed and predicted for it. Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal. Chemica
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lly, several anionic species of astatine are known and most of its compounds resemble those of iodine, but it also sometimes displays metallic characteristics and shows some similarities to silver.
The first synthesis of the element was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segr at the University of California, Berkeley, who named it from the Ancient Greek 'unstable'. Four isotopes of astatine were subsequently found to be naturally occurring, although much less than one gram is present at any given time in the Earth's crust. Neither the most stable isotope astatine210, nor the medically useful astatine211, occur naturally; they can only be produced synthetically, usually by bombarding bismuth209 with alpha particles.
Characteristics
Astatine is an extremely radioactive element; all its isotopes have halflives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium, or radon. Most of its isotopes are very unstable, with halflives of one second or less. Of
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the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than francium are in any case synthetic and do not occur in nature.
The bulk properties of astatine are not known with any certainty. Research is limited by its short halflife, which prevents the creation of weighable quantities. A visible piece of astatine would immediately vaporize itself because of the heat generated by its intense radioactivity. It remains to be seen if, with sufficient cooling, a macroscopic quantity of astatine could be deposited as a thin film. Astatine is usually classified as either a nonmetal or a metalloid; metal formation has also been predicted.
Physical
Most of the physical properties of astatine have been estimated by interpolation or extrapolation, using theoretically or empirically derived methods. For example, halogens get darker with increasing atomic weight fluorine is nearly colorless, chlorine is yellow green, bromine is red brown, and iodine is dark
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grayviolet. Astatine is sometimes described as probably being a black solid assuming it follows this trend, or as having a metallic appearance if it is a metalloid or a metal.
Astatine sublimes less readily than does iodine, having a lower vapor pressure. Even so, half of a given quantity of astatine will vaporize in approximately an hour if put on a clean glass surface at room temperature. The absorption spectrum of astatine in the middle ultraviolet region has lines at 224.401 and 216.225 nm, suggestive of 6p to 7s transitions.
The structure of solid astatine is unknown. As an analogue of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor with a band gap of 0.7 eV. Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic facecentered cubic structure; in this structure it may well be a superconductor, like the similar highpressure phase of iodine. Evidence for or against the existence o
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f diatomic astatine At2 is sparse and inconclusive. Some sources state that it does not exist, or at least has never been observed, while other sources assert or imply its existence. Despite this controversy, many properties of diatomic astatine have been predicted; for example, its bond length would be , dissociation energy , and heat of vaporization Hvap 54.39 kJmol. Many values have been predicted for the melting and boiling points of astatine, but only for At2.
Chemical
The chemistry of astatine is "clouded by the extremely low concentrations at which astatine experiments have been conducted, and the possibility of reactions with impurities, walls and filters, or radioactivity byproducts, and other unwanted nanoscale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 1010 molL1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, su
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ch as plating onto a cathode, and coprecipitating with metal sulfides in hydrochloric acid. It forms complexes with EDTA, a metal chelating agent, and is capable of acting as a metal in antibody radiolabeling; in some respects astatine in the 1 state is akin to silver in the same state. Most of the organic chemistry of astatine is, however, analogous to that of iodine. It has been suggested that astatine can form a stable monatomic cation in aqueous solution, but electromigration evidence suggests that the cationic AtI species is protonated hypoastatous acid H2OAt, showing analogy to iodine.
Astatine has an electronegativity of 2.2 on the revised Pauling scale lower than that of iodine 2.66 and the same as hydrogen. In hydrogen astatide HAt, the negative charge is predicted to be on the hydrogen atom, implying that this compound could be referred to as astatine hydride according to certain nomenclatures. That would be consistent with the electronegativity of astatine on the AllredRochow scale 1.9 being less
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than that of hydrogen 2.2. However, official IUPAC stoichiometric nomenclature is based on an idealized convention of determining the relative electronegativities of the elements by the mere virtue of their position within the periodic table. According to this convention, astatine is handled as though it is more electronegative than hydrogen, irrespective of its true electronegativity. The electron affinity of astatine, at 233 kJ mol1, is 21 less than that of iodine. In comparison, the value of Cl 349 is 6.4 higher than F 328; Br 325 is 6.9 less than Cl; and I 295 is 9.2 less than Br. The marked reduction for At was predicted as being due to spinorbit interactions. The first ionisation energy of astatine is about 899 kJ mol1, which continues the trend of decreasing first ionisation energies down the halogen group fluorine, 1681; chlorine, 1251; bromine, 1140; iodine, 1008.
Compounds
Less reactive than iodine, astatine is the least reactive of the halogens. Its compounds have been synthesized in microscopic
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amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques such as filtration and precipitation to work. Like iodine, astatine has been shown to adopt oddnumbered oxidation states ranging from 1 to 7.
Only a few compounds with metals have been reported, in the form of astatides of sodium, palladium, silver, thallium, and lead. Some characteristic properties of silver and sodium astatide, and the other hypothetical alkali and alkaline earth astatides, have been estimated by extrapolation from other metal halides.
The formation of an astatine compound with hydrogen usually referred to as hydrogen astatide was noted by the pioneers of astatine chemistry. As mentioned, there are grounds for instead referring to this compound as astatine hydride. It is easily oxidized;
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acidification by dilute nitric acid gives the At0 or At forms, and the subsequent addition of silverI may only partially, at best, precipitate astatine as silverI astatide AgAt. Iodine, in contrast, is not oxidized, and precipitates readily as silverI iodide.
Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with AtB bonds, these being more stable than AtC bonds. Astatine can replace a hydrogen atom in benzene to form astatobenzene C6H5At; this may be oxidized to C6H5AtCl2 by chlorine. By treating this compound with an alkaline solution of hypochlorite, C6H5AtO2 can be produced. The dipyridineastatineI cation, AtC5H5N2, forms ionic compounds with perchlorate a noncoordinating anion and with nitrate, AtC5H5N2NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatineI centre with each of the pyridine rings via their nitrogen atoms.
With oxygen, there is evidence of the species AtO and AtO in aqueous
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solution, formed by the reaction of astatine with an oxidant such as elemental bromine or in the last case by sodium persulfate in a solution of perchloric acid the latter species might also be protonated astatous acid, . The species previously thought to be has since been determined to be , a hydrolysis product of AtO another such hydrolysis product being AtOOH. The well characterized anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate LaAtO33, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of , such as by xenon difluoride in a hot alkaline solution or periodate in a neutral or alkaline solution, yields the perastatate ion ; this is only stable in neutral or alkaline solutions. Astatine is also thought to be capable of forming cations in salts with oxyanions such as iodate or dichromate; this is based on the observation that, in acidic sol
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utions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silverI iodate or thalliumI dichromate.
Astatine may form bonds to the other chalcogens; these include S7At and with sulfur, a coordination selenourea compound with selenium, and an astatinetellurium colloid with tellurium.
Astatine is known to react with its lighter homologs iodine, bromine, and chlorine in the vapor state; these reactions produce diatomic interhalogen compounds with formulas AtI, AtBr, and AtCl. The first two compounds may also be produced in water astatine reacts with iodineiodide solution to form AtI, whereas AtBr requires aside from astatine an iodineiodine monobromidebromide solution. The excess of iodides or bromides may lead to and ions, or in a chloride solution, they may produce species like or via equilibrium reactions with the chlorides. Oxidation of the element with dichromate in nitric acid solution showed that adding chloride turned the astatine
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into a molecule likely to be either AtCl or AtOCl. Similarly, or may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed to have been precipitated. In a plasma ion source mass spectrometer, the ions AtI, AtBr, and AtCl have been formed by introducing lighter halogen vapors into a heliumfilled cell containing astatine, supporting the existence of stable neutral molecules in the plasma ion state. No astatine fluorides have been discovered yet. Their absence has been speculatively attributed to the extreme reactivity of such compounds, including the reaction of an initially formed fluoride with the walls of the glass container to form a nonvolatile product. Thus, although the synthesis of an astatine fluoride is thought to be possible, it may require a liquid halogen fluoride solvent, as has already been used for the characterization of radon fluoride.
History
In 1869, when Dmitri Mendeleev published his periodic table, the space under iodine was empty; after Niels Bohr est
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ablished the physical basis of the classification of chemical elements, it was suggested that the fifth halogen belonged there. Before its officially recognized discovery, it was called "ekaiodine" from Sanskrit eka "one" to imply it was one space under iodine in the same manner as ekasilicon, ekaboron, and others. Scientists tried to find it in nature; given its extreme rarity, these attempts resulted in several false discoveries.
The first claimed discovery of ekaiodine was made by Fred Allison and his associates at the Alabama Polytechnic Institute now Auburn University in 1931. The discoverers named element 85 "alabamine", and assigned it the symbol Ab, designations that were used for a few years. In 1934, H. G. MacPherson of University of California, Berkeley disproved Allison's method and the validity of his discovery. There was another claim in 1937, by the chemist Rajendralal De. Working in Dacca in British India now Dhaka in Bangladesh, he chose the name "dakin" for element 85, which he claimed to
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have isolated as the thorium series equivalent of radium F polonium210 in the radium series. The properties he reported for dakin do not correspond to those of astatine; moreover, astatine is not found in the thorium series, and the true identity of dakin is not known.
In 1936, the team of Romanian physicist Horia Hulubei and French physicist Yvette Cauchois claimed to have discovered element 85 via Xray analysis. In 1939, they published another paper which supported and extended previous data. In 1944, Hulubei published a summary of data he had obtained up to that time, claiming it was supported by the work of other researchers. He chose the name "dor", presumably from the Romanian for "longing" for peace, as World War II had started five years earlier. As Hulubei was writing in French, a language which does not accommodate the "ine" suffix, dor would likely have been rendered in English as "dorine", had it been adopted. In 1947, Hulubei's claim was effectively rejected by the Austrian chemist Friedrich Pan
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eth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine, his means to detect it were too weak, by current standards, to enable correct identification. He had also been involved in an earlier false claim as to the discovery of element 87 francium and this is thought to have caused other researchers to downplay his work.
In 1940, the Swiss chemist Walter Minder announced the discovery of element 85 as the beta decay product of radium A polonium218, choosing the name "helvetium" from , the Latin name of Switzerland. Berta Karlik and Traude Bernert were unsuccessful in reproducing his experiments, and subsequently attributed Minder's results to contamination of his radon stream radon222 is the parent isotope of polonium218. In 1942, Minder, in collaboration with the English scientist Alice LeighSmith, announced the discovery of another isotope of element 85, presumed to be the product of thorium A polonium216 beta decay. They
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named this substance "anglohelvetium", but Karlik and Bernert were again unable to reproduce these results.
Later in 1940, Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segr isolated the element at the University of California, Berkeley. Instead of searching for the element in nature, the scientists created it by bombarding bismuth209 with alpha particles in a cyclotron particle accelerator to produce, after emission of two neutrons, astatine211. The discoverers, however, did not immediately suggest a name for the element. The reason for this was that at the time, an element created synthetically in "invisible quantities" that had not yet been discovered in nature was not seen as a completely valid one; in addition, chemists were reluctant to recognize radioactive isotopes as legitimately as stable ones. In 1943, astatine was found as a product of two naturally occurring decay chains by Berta Karlik and Traude Bernert, first in the socalled uranium series, and then in the actinium series. Since then, as
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tatine was also found in a third decay chain, the neptunium series. Friedrich Paneth in 1946 called to finally recognize synthetic elements, quoting, among other reasons, recent confirmation of their natural occurrence, and proposed that the discoverers of the newly discovered unnamed elements name these elements. In early 1947, Nature published the discoverers' suggestions; a letter from Corson, MacKenzie, and Segr suggested the name "astatine" coming from the Greek astatos meaning "unstable", because of its propensity for radioactive decay, with the ending "ine", found in the names of the four previously discovered halogens. The name was also chosen to continue the tradition of the four stable halogens, where the name referred to a property of the element.
Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry. Subsequent investigators reported iodinelike, cationic, or amphoteric behavior. In a 2003 retrospective, Corson wrote that "some of the properties of asta
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tine are similar to iodine it also exhibits metallic properties, more like its metallic neighbors Po and Bi."
Isotopes
There are 39 known isotopes of astatine, with atomic masses mass numbers of 191229. Theoretical modeling suggests that 37 more isotopes could exist. No stable or longlived astatine isotope has been observed, nor is one expected to exist.
Astatine's alpha decay energies follow the same trend as for other heavy elements. Lighter astatine isotopes have quite high energies of alpha decay, which become lower as the nuclei become heavier. Astatine211 has a significantly higher energy than the previous isotope, because it has a nucleus with 126 neutrons, and 126 is a magic number corresponding to a filled neutron shell. Despite having a similar halflife to the previous isotope 8.1 hours for astatine210 and 7.2 hours for astatine211, the alpha decay probability is much higher for the latter 41.81 against only 0.18. The two following isotopes release even more energy, with astatine213 releasing t
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he most energy. For this reason, it is the shortestlived astatine isotope. Even though heavier astatine isotopes release less energy, no longlived astatine isotope exists, because of the increasing role of beta decay electron emission. This decay mode is especially important for astatine; as early as 1950 it was postulated that all isotopes of the element undergo beta decay, though nuclear mass measurements indicate that 215At is in fact betastable, as it has the lowest mass of all isobars with A 215. A beta decay mode has been found for all other astatine isotopes except for astatine213, astatine214, and astatine216m. Astatine210 and lighter isotopes exhibit beta plus decay positron emission, astatine216 and heavier isotopes exhibit beta minus decay, and astatine212 decays via both modes, while astatine211 undergoes electron capture.
The most stable isotope is astatine210, which has a halflife of 8.1 hours. The primary decay mode is beta plus, to the relatively longlived in comparison to astatine isotopes
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alpha emitter polonium210. In total, only five isotopes have halflives exceeding one hour astatine207 to 211. The least stable ground state isotope is astatine213, with a halflife of 125 nanoseconds. It undergoes alpha decay to the extremely longlived bismuth209.
Astatine has 24 known nuclear isomers, which are nuclei with one or more nucleons protons or neutrons in an excited state. A nuclear isomer may also be called a "metastate", meaning the system has more internal energy than the "ground state" the state with the lowest possible internal energy, making the former likely to decay into the latter. There may be more than one isomer for each isotope. The most stable of these nuclear isomers is astatine202m1, which has a halflife of about 3 minutes, longer than those of all the ground states bar those of isotopes 203211 and 220. The least stable is astatine214m1; its halflife of 265 nanoseconds is shorter than those of all ground states except that of astatine213.
Natural occurrence
Astatine is the rares
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t naturally occurring element. The total amount of astatine in the Earth's crust quoted mass 2.36 1025 grams is estimated by some to be less than one gram at any given time. Other sources estimate the amount of ephemeral astatine, present on earth at any given moment, to be up to one ounce about 28 grams.
Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes astatine215, 217, 218 and 219 are instead continuously produced as a result of the decay of radioactive thorium and uranium ores, and trace quantities of neptunium237. The landmass of North and South America combined, to a depth of 16 kilometers 10 miles, contains only about one trillion astatine215 atoms at any given time around 3.5 1010 grams. Astatine217 is produced via the radioactive decay of neptunium237. Primordial remnants of the latter isotopedue to its relatively short halflife of 2.14 million yearsare no longer present on Earth. However, trace amounts occur naturally as a product
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of transmutation reactions in uranium ores. Astatine218 was the first astatine isotope discovered in nature. Astatine219, with a halflife of 56 seconds, is the longest lived of the naturally occurring isotopes.
Isotopes of astatine are sometimes not listed as naturally occurring because of misconceptions that there are no such isotopes, or discrepancies in the literature. Astatine216 has been counted as a naturally occurring isotope but reports of its observation which were described as doubtful have not been confirmed.
Synthesis
Formation
Astatine was first produced by bombarding bismuth209 with energetic alpha particles, and this is still the major route used to create the relatively longlived isotopes astatine209 through astatine211. Astatine is only produced in minuscule quantities, with modern techniques allowing production runs of up to 6.6 giga becquerels about 86 nanograms or 2.47 1014 atoms. Synthesis of greater quantities of astatine using this method is constrained by the limited availabilit
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y of suitable cyclotrons and the prospect of melting the target. Solvent radiolysis due to the cumulative effect of astatine decay is a related problem. With cryogenic technology, microgram quantities of astatine might be able to be generated via proton irradiation of thorium or uranium to yield radon211, in turn decaying to astatine211. Contamination with astatine210 is expected to be a drawback of this method.
The most important isotope is astatine211, the only one in commercial use. To produce the bismuth target, the metal is sputtered onto a gold, copper, or aluminium surface at 50 to 100 milligrams per square centimeter. Bismuth oxide can be used instead; this is forcibly fused with a copper plate. The target is kept under a chemically neutral nitrogen atmosphere, and is cooled with water to prevent premature astatine vaporization. In a particle accelerator, such as a cyclotron, alpha particles are collided with the bismuth. Even though only one bismuth isotope is used bismuth209, the reaction may occur
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in three possible ways, producing astatine209, astatine210, or astatine211. In order to eliminate undesired nuclides, the maximum energy of the particle accelerator is set to a value optimally 29.17 MeV above that for the reaction producing astatine211 to produce the desired isotope and below the one producing astatine210 to avoid producing other astatine isotopes.
Separation methods
Since astatine is the main product of the synthesis, after its formation it must only be separated from the target and any significant contaminants. Several methods are available, "but they generally follow one of two approachesdry distillation or wet acid treatment of the target followed by solvent extraction." The methods summarized below are modern adaptations of older procedures, as reviewed by Kugler and Keller. Pre1985 techniques more often addressed the elimination of coproduced toxic polonium; this requirement is now mitigated by capping the energy of the cyclotron irradiation beam.
Dry
The astatinecontaining cyclotr
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on target is heated to a temperature of around 650 C. The astatine volatilizes and is condensed in typically a cold trap. Higher temperatures of up to around 850 C may increase the yield, at the risk of bismuth contamination from concurrent volatilization. Redistilling the condensate may be required to minimize the presence of bismuth as bismuth can interfere with astatine labeling reactions. The astatine is recovered from the trap using one or more low concentration solvents such as sodium hydroxide, methanol or chloroform. Astatine yields of up to around 80 may be achieved. Dry separation is the method most commonly used to produce a chemically useful form of astatine.
Wet
The irradiated bismuth or sometimes bismuth trioxide target is first dissolved in, for example, concentrated nitric or perchloric acid. Following this first step, the acid can be distilled away to leave behind a white residue that contains both bismuth and the desired astatine product. This residue is then dissolved in a concentrated ac
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