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converters when magnetic fields are also involved. Studies conducted in wind tunnels involve most of the time low atmospheric pressure similar to an altitude of 20–50 km, typical of hypersonic flight, where the electrical conductivity of air is higher, hence non-thermal weakly ionized plasmas can be easily produced with a fewer energy expense. === Catalysis === Atmospheric pressure non-thermal plasma can be used to promote chemical reactions. Collisions between hot temperature electrons and cold gas molecules can lead to dissociation reactions and the subsequent formation of radicals. This kind of discharge exhibits reacting properties that are usually seen in high temperature discharge systems. Non-thermal plasma is also used in conjunction with a catalyst to further enhance the chemical conversion of reactants or to alter the products chemical composition. Among the different application fields, there are ozone production at a commercial level; pollution abatement, both solid (PM, VOC) and gaseous (SOx, NOx); CO2 conversion in fuels (methanol, syngas) or value added chemicals; nitrogen fixation; methanol synthesis; liquid fuels synthesis from lighter hydrocarbons (e.g. methane), hydrogen production via hydrocarbons reforming ==== Configurations ==== The coupling between the two different mechanisms can be done in two different ways: two-stage configuration, also called post-plasma catalysis (PPC) and one-stage configuration, also called in-plasma catalysis (IPC) or plasma enhanced catalysis (PEC). In the first case the catalytic reactor is placed after the plasma chamber. This means that only the long-lived species can reach the catalyst surface and react, while short-lived radicals, ions and excited species decay in the first part of the reactor. As an example, the oxygen ground state atom O(3P) has a lifetime of about 14 μs in a dry air atmospheric pressure plasma. This means that only a small region of the catalyst is in contact with active radicals. In a such two-stage
|
{
"page_id": 6161283,
"source": null,
"title": "Nonthermal plasma"
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set-up, the main role of the plasma is to alter the gas composition fed to the catalytic reactor. In a PEC system, synergistic effects are greater since short-lived excited species are formed near the catalyst surface. The way the catalyst is inserted in the PEC reactor influence the overall performance. It can be placed inside the reactor in different ways: in powder form (packed bed), deposited on foams, deposited on structured material (honeycomb), and coating of the reactor walls Packed bed plasma-catalytic reactor are commonly used for fundamental studies and a scale-up to industrial applications is difficult since the pressure drop increase with the flow rate. ==== Plasma-catalysis interactions ==== In a PEC system, the way the catalyst is positioned in relation to the plasma can affect the process in different ways. The catalyst can positively influence the plasma and vice versa resulting in an output that cannot be obtained using each process individually. The synergy that is established is ascribed to different cross effects. Plasma effects on catalyst: Change in the physiochemical properties. Plasma change the adsorption/desorption equilibrium on the catalyst surface leading to higher adsorption capabilities. An interpretation to this phenomenon is not yet clear. Higher catalyst surface area. A catalyst exposed to a discharge can give rise to the formation of nanoparticles. The higher surface/volume ratio leads to better catalyst performances. Higher adsorption probability. Change in the catalyst oxidation state. Some metallic catalyst (e.g. Ni, Fe) are more active in their metallic form. The presence of a plasma discharge can induce a reduction of the catalyst metal oxides, improving the catalytic activity. Reduced coke formation. When dealing with hydrocarbons, coke formation leads to a progressive deactivation of the catalyst. The reduced coke formation in presence of plasma reduces the poisoning/deactivation rate and thus extending the life of
|
{
"page_id": 6161283,
"source": null,
"title": "Nonthermal plasma"
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a catalyst. Presence of new gas phase species. In a plasma discharge a wide range of new species is produced allowing the catalyst to be exposed to them. Ions, vibrationally and rotationally excited species do not affect the catalyst since they lose charge and the additional energy they possess when they reach a solid surface. Radicals, instead, show high sticking coefficients for chemisorption, increasing the catalytic activity. Catalyst effects on plasma: Local electric field enhancement. This aspect is mainly related to a packed-bed PEC configuration. The presence of a packing material inside an electric field generates local field enhancements due to the presence of asperities, solid material surface inhomogeneities, presence of pores and other physical aspects. This phenomenon is related to surface charge accumulation on the packing material surface and it is present even if a packed-bed is used without a catalyst. Despite this is a physical aspect, it also affects the chemistry since it alters the electron energy distribution in proximity of the asperities. Discharges formation inside pores. This aspect is strictly related to the previous one. Small void spaces inside a packing material affect the electric field strength. The enhancement can also lead to a change in the discharge characteristics, which can be different from the discharge condition of the bulk region (i.e. far from the solid material). The high intensity of the electric field can also lead to the production of different species that are not observed in the bulk. Change in the discharge type. Inserting a dielectric material in a discharge region leads to a shifting in the discharge type. From a filamentary regime a mixed filamentary/surface discharge is established. Ions, excited species and radicals are formed in a wider region if a surface discharge regime is present. Catalyst effects on plasma are mostly related to
|
{
"page_id": 6161283,
"source": null,
"title": "Nonthermal plasma"
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|
the presence of a dielectric material inside the discharge region and do not necessarily require the presence of a catalyst. == See also == Anisothermal plasma Gliding Arc Plasmatron == References ==
|
{
"page_id": 6161283,
"source": null,
"title": "Nonthermal plasma"
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Actinium is a chemical element; it has symbol Ac and atomic number 89. It was discovered by Friedrich Oskar Giesel in 1902, who gave it the name emanium; the element got its name by being wrongly identified with a substance André-Louis Debierne found in 1899 and called actinium. The actinide series, a set of 15 elements between actinium and lawrencium in the periodic table, are named for actinium. Together with polonium, radium, and radon, actinium was one of the first non-primordial radioactive elements to be discovered. A soft, silvery-white 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 half-life of 21.772 years, predominantly emitting beta and sometimes alpha particles, and 228Ac, which is beta active with a half-life of 6.15 hours. One tonne of natural uranium in ore contains about 0.2 milligrams of actinium-227, and one tonne of thorium contains about 5 nanograms of actinium-228. 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 226Ra 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 == André-Louis 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
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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|
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 "emanium" in 1904. After a comparison of the substances' half-lives 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 now-known 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 scientific 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 (ακτίς, ακτίνος), meaning beam or ray. Its symbol Ac is also
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
}
|
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, silvery-white, radioactive, metallic element. Its estimated shear modulus is similar to that of lead. Owing to its strong 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 set its name, much as lanthanum had done for the lanthanides. The actinides are much 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 oxide 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 [Rn] 6d17s2 electronic configuration of actinium, with three valence electrons that are easily donated to give the stable closed-shell structure of the noble gas radon. Although the 5f orbitals are unoccupied in an actinium atom, it can be used as a valence orbital in actinium complexes and hence it is generally considered the first 5f element by authors working on it. Ac3+ is the largest of all
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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|
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 Ac(NO3)3. They all contain actinium in the oxidation state +3. In particular, 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 (932 °F) or the oxalate at 1,100 °C (2,010 °F), in vacuum. Its crystal lattice is isotypic with the oxides of most trivalent rare-earth 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 (1,292 °F) in an all-platinum setup. Treating actinium trifluoride with ammonium hydroxide at 900–1,000 °C (1,650–1,830 °F) yields oxyfluoride AcOF. Whereas lanthanum oxyfluoride can be easily obtained by burning lanthanum trifluoride in air at 800 °C (1,470 °F) for 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 (1,760 °F). Similarly to the oxyfluoride, actinium oxychloride can be prepared by hydrolyzing actinium
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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trichloride with ammonium hydroxide at 1,000 °C (1,830 °F). 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 (932 °F) results in the oxybromide AcOBr. === Other compounds === Actinium hydride was obtained by reduction of actinium trichloride with potassium at 300 °C (572 °F), and its structure was deduced by analogy with the corresponding LaH2 hydride. The source of hydrogen in the reaction was uncertain. Mixing monosodium phosphate (NaH2PO4) with a solution of actinium in hydrochloric acid yields white-colored actinium phosphate hemihydrate (AcPO4·0.5H2O), and heating actinium oxalate with hydrogen sulfide vapors at 1,400 °C (2,550 °F) 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 1,000 °C (1,830 °F). == Isotopes == Naturally occurring actinium is principally composed of two radioactive isotopes; 227Ac (from the radioactive family of 235U) and 228Ac (a granddaughter of 232Th). 227Ac 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. Thirty-three radioisotopes have been identified, the most stable being 227Ac with a half-life of 21.772 years, 225Ac with a half-life of 10.0 days and 226Ac with a half-life of 29.37 hours. All remaining radioactive isotopes have half-lives that are less than 10 hours and the majority of them have half-lives shorter than one minute. The shortest-lived known isotope of actinium is 217Ac (half-life of 69 nanoseconds) which decays
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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|
through alpha decay. Actinium also has two known meta states. The most significant isotopes for chemistry are 225Ac, 227Ac, and 228Ac. Purified 227Ac comes into equilibrium with its decay products after about a half of year. It decays according to its 21.772-year half-life 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, 227Ac 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 203 u (203Ac) to 236 u (236Ac). == 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 is a transient member of the uranium-actinium 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 near-stable bismuth (209Bi); even though all primordial 237Np has decayed away, it is continuously produced by neutron knock-out 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 actinium-bearing ores, render separation of actinium from the ore
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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impractical. The most concentrated actinium sample prepared from raw material consisted of 7 micrograms of 227Ac in less than 0.1 milligrams of La2O3, and complete separation was never achieved. Instead, actinium is prepared, in milligram amounts, by the neutron irradiation of 226Ra in a nuclear reactor. Ra 88 226 + n 0 1 ⟶ Ra 88 227 → 42.2 min β − Ac 89 227 {\displaystyle {\ce {^{226}_{88}Ra + ^{1}_{0}n -> ^{227}_{88}Ra ->[\beta^-][42.2 \ {\ce {min}}] ^{227}_{89}Ac}}} 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 thenoyltrifluoroacetone-benzene 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 two-stage process. Actinium can then be separated from radium, with a ratio of about 100, using a low cross-linking 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 George 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 radium-226 target with 20–30 MeV deuterium ions. This reaction also yields 226Ac which however decays with a half-life of 29 hours and thus does not contaminate 225Ac. Actinium metal has been prepared by the reduction of actinium fluoride with lithium vapor
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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in vacuum at a temperature between 1,100 and 1,300 °C (2,010 and 2,370 °F). 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 spacecraft. The oxide of 227Ac pressed with beryllium is also an efficient neutron source with the activity exceeding that of the standard americium-beryllium and radium-beryllium pairs. In all those applications, 227Ac (a beta source) is merely a progenitor which generates alpha-emitting isotopes upon its decay. Beryllium captures alpha particles and emits neutrons owing to its large cross-section for the (α,n) nuclear reaction: Be 4 9 + He 2 4 ⟶ C 6 12 + n 0 1 + γ {\displaystyle {\ce {^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \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 moisture/density 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 213Bi 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 half-life of 10 days, making it much more suitable for radiation therapy than 213Bi (half-life 46 minutes). Additionally, 225Ac decays to nontoxic 209Bi rather than toxic lead, which
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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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 agents as HEHA (1,4,7,10,13,16-hexaazacyclohexadecane-N,N′,N″,N‴,N‴′,N‴″-hexaacetic acid) or DOTA (1,4,7,10-tetraazacyclododecane-1,4,7,10-tetraacetic acid) coupled to trastuzumab, a monoclonal antibody that interferes with the HER2/neu 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 half-life of 227Ac (21.77 years) makes it a 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 depth-profiles 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 depth profile, but instead increases
|
{
"page_id": 899,
"source": null,
"title": "Actinium"
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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 room-temperature 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 also == Actinium series == Notes == == References == == Bibliography == Meyer, Gerd; Morss, Lester R. (1991). Synthesis of lanthanide and actinide compounds. Springer. ISBN 0-7923-1018-7. == External links == Actinium at The Periodic Table of Videos (University of Nottingham) NLM Hazardous Substances Databank – Actinium, Radioactive Actinium in Kirby, H. W.; Morss, L. R. (2006). Morss; Edelstein, Norman M.; Fuger, Jean (eds.). The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer. ISBN 978-1-4020-3555-5.
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"page_id": 899,
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Americium is a synthetic chemical element; it has symbol Am and atomic number 95. It is radioactive and a transuranic member of the actinide series in the periodic table, located under the lanthanide element europium and was thus named after the Americas by analogy. 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 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 a 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 lattices of solid americium and its compounds contain small intrinsic radiogenic defects, due to metamictization induced by self-irradiation 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,
|
{
"page_id": 900,
"source": null,
"title": "Americium"
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it was first intentionally synthesized, isolated and identified in late autumn 1944, at the University of California, Berkeley, by Glenn T. Seaborg, Leon O. Morgan, Ralph A. James, and Albert Ghiorso. They used a 60-inch 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 rare-earth series, analogous to europium, Eu, of the lanthanide series." The new element was isolated from its oxides in a complex, multi-step process. First plutonium-239 nitrate (239PuNO3) solution was 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. Americium-241 was directly obtained from plutonium upon absorption of
|
{
"page_id": 900,
"source": null,
"title": "Americium"
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two neutrons. It decays by emission of a α-particle to 237Np; the half-life of this decay was first determined as 510±20 years but then corrected to 432.2 years. Pu 94 239 → ( n , γ ) Pu 94 240 → ( n , γ ) Pu 94 241 → 14.35 yr β − Am 95 241 ( → 432.2 yr α Np 93 237 ) {\displaystyle {\ce {^{239}_{94}Pu ->[{\ce {(n,\gamma)}}] ^{240}_{94}Pu ->[{\ce {(n,\gamma)}}] ^{241}_{94}Pu ->[\beta^-][14.35\ {\ce {yr}}] ^{241}_{95}Am}}\ \left({\ce {->[\alpha][432.2\ {\ce {yr}}] ^{237}_{93}Np}}\right)} The times are half-lives The second isotope 242Am was produced upon neutron bombardment of the already-created 241Am. Upon rapid β-decay, 242Am converts into the isotope of curium 242Cm (which had been discovered previously). The half-life of this decay was initially determined at 17 hours, which was close to the presently accepted value of 16.02 h. Am 95 241 → ( n , γ ) Am 95 242 ( → 16.02 h β − Cm 96 242 ) {\displaystyle {\ce {^{241}_{95}Am ->[{\ce {(n,\gamma)}}] ^{242}_{95}Am}}\ \left({\ce {->[\beta^-][16.02\ {\ce {h}}] ^{242}_{96}Cm}}\right)} 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 by their radioactivity. The first substantial amounts of metallic americium weighing 40–200
|
{
"page_id": 900,
"source": null,
"title": "Americium"
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micrograms were not prepared until 1951 by reduction of americium(III) fluoride with barium metal in high vacuum at 1100 °C. == Occurrence == The longest-lived and most common isotopes of americium, 241Am and 243Am, have half-lives 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 neutron capture and beta decay (238U → 239Pu → 240Pu → 241Am), though the quantities would be tiny and this has not been confirmed. Extraterrestrial long-lived 247Cm is probably also deposited on Earth and has 243Am as one of its intermediate decay products, but again 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 americium; 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 plutonium-based Trinity nuclear bomb test on 16 July 1945, contains traces of americium-241. Elevated levels of americium were also detected at the crash site of a US Boeing B-52 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 picocuries per gram (0.37 mBq/g). Atmospheric americium compounds are poorly soluble in common solvents and mostly adhere to soil particles. Soil analysis revealed about 1,900
|
{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
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 americium isotopes, mostly 241Am and 243Am. Their prolonged radioactivity is undesirable for the disposal, and therefore americium, together with other long-lived 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 short-lived 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 theoretically 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 US$1,500 per gram (US$43,000/oz) of 241Am, remains almost unchanged 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 US$100,000–US$160,000 per gram (US$2,800,000–US$4,500,000/oz). 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: U 92 238 → ( n , γ ) U 92 239 → 23.5 min β − Np 93 239 → 2.3565 d
|
{
"page_id": 900,
"source": null,
"title": "Americium"
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β − Pu 94 239 {\displaystyle {\ce {^{238}_{92}U ->[{\ce {(n,\gamma)}}] ^{239}_{92}U ->[\beta^-][23.5 \ {\ce {min}}] ^{239}_{93}Np ->[\beta^-][2.3565 \ {\ce {d}}] ^{239}_{94}Pu}}} The capture of two neutrons by 239Pu (a so-called (n,γ) reaction), followed by a β-decay, results in 241Am: Pu 94 239 → 2 ( n , γ ) Pu 94 241 → 14.35 yr β − Am 95 241 {\displaystyle {\ce {^{239}_{94}Pu ->[{\ce {2(n,\gamma)}}] ^{241}_{94}Pu ->[\beta^-][14.35 \ {\ce {yr}}] ^{241}_{95}Am}}} The plutonium present in spent nuclear fuel contains about 12% of 241Pu. Because it beta-decays 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 used 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: { 79 % : Am 95 241 → ( n , γ ) Am 95 242 10 % : Am 95 241 → ( n , γ ) Am 95 242 m {\displaystyle {\begin{cases}79\%:&{\ce {^{241}_{95}Am ->[{\ce {(n,\gamma)}}] ^{242}_{95}Am}}\\10\%:&{\ce {^{241}_{95}Am ->[{\ce {(n,\gamma)}}] ^{242 m}_{95}Am}}\end{cases}}} Americium-242 has a half-life 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: Pu 94 239 → 4 ( n , γ ) 94 243 Pu → 4.956 h β − Am 95 243 {\displaystyle {\ce {^{239}_{94}Pu ->[{\ce {4(n,\gamma)}}] \ ^{243}_{94}Pu ->[\beta^-][4.956 \ {\ce {h}}] ^{243}_{95}Am}}} === Metal generation === Most synthesis routines yield a mixture of different actinide isotopes in oxide forms, from which isotopes of
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{
"page_id": 900,
"source": null,
"title": "Americium"
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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 PUREX-type extraction (Plutonium–URanium EXtraction) with tributyl phosphate in a hydrocarbon. The lanthanides and remaining actinides are then separated from the aqueous residue (raffinate) by a diamide-based extraction, to give, after stripping, a mixture of trivalent actinides and lanthanides. Americium compounds are then selectively extracted using multi-step 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 EU-funded project codenamed "EUROPART" studied triazines and other compounds as potential extraction agents. A bis-triazinyl 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 Am(IV) complexes which can be washed away. Metallic americium is obtained by reduction from its compounds. Americium(III) fluoride was first used for this purpose. The reaction was conducted using elemental barium as reducing agent in a water- and oxygen-free environment inside an apparatus made of tantalum and tungsten. 2 A m F 3 + 3 B a ⟶ 2 A m + 3 B a F 2 {\displaystyle \mathrm {2\ AmF_{3}\ +\ 3\ Ba\ \longrightarrow \ 2\ Am\ +\ 3\ BaF_{2}} } An alternative is the reduction of americium dioxide by metallic lanthanum or thorium: 3 A m O 2 + 4 L a ⟶ 3 A m + 2 L a 2 O 3 {\displaystyle
|
{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
\mathrm {3\ AmO_{2}\ +\ 4\ La\ \longrightarrow \ 3\ Am\ +\ 2\ La_{2}O_{3}} } == 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 silvery-white metallic lustre, but then slowly tarnishes in air. With a density of 12 g/cm3, americium is less dense than both curium (13.52 g/cm3) and plutonium (19.8 g/cm3); but has a higher density than europium (5.264 g/cm3)—mostly 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 (1340 °C). At ambient conditions, americium is present in its most stable α form which has a hexagonal crystal symmetry, and a space group P63/mmc with cell parameters a = 346.8 pm and c = 1124 pm, and four atoms per unit cell. The crystal consists of a double-hexagonal close packing with the layer sequence ABAC and so is isotypic with α-lanthanum and several actinides such as α-curium. The crystal structure of americium changes with pressure and temperature. When compressed at room temperature to 5 GPa, α-Am transforms to the β modification, which has a face-centered cubic (fcc) symmetry, space group Fm3m 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 γ-Am structure similar to that of α-uranium. There are no further transitions observed up
|
{
"page_id": 900,
"source": null,
"title": "Americium"
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|
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 literature, which also sometimes lists the α, β and γ phases as I, II and III. The β-γ transition is accompanied by a 6% decrease in the crystal volume; although theory also predicts a significant volume change for the α-β transition, it is not observed experimentally. The pressure of the α-β transition decreases with increasing temperature, and when α-americium is heated at ambient pressure, at 770 °C it changes into an fcc phase which is different from β-Am, and at 1075 °C it converts to a body-centered cubic structure. The pressure-temperature phase diagram of americium is thus rather similar to those of lanthanum, praseodymium and neodymium. As with many other actinides, self-damage of the crystal structure due to alpha-particle 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 X-ray diffraction peaks. This effect makes somewhat uncertain the temperature of americium and some of its properties, such as electrical resistivity. So for americium-241, the resistivity at 4.2 K increases with time from about 2 μOhm·cm to 10 μOhm·cm after 40 hours, and saturates at about 16 μOhm·cm 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 μOhm·cm at liquid helium to 69 μOhm·cm at room temperature; this behavior is similar to that of neptunium, uranium, thorium and protactinium, but is different from plutonium and curium which
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{
"page_id": 900,
"source": null,
"title": "Americium"
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|
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 that of its neighbor curium which exhibits antiferromagnetic transition at 52 K. The thermal expansion coefficient of americium is slightly anisotropic and amounts to (7.5±0.2)×10−6 /°C along the shorter a axis and (6.2±0.4)×10−6 /°C for the longer c hexagonal axis. The enthalpy of dissolution of americium metal in hydrochloric acid at standard conditions is −620.6±1.3 kJ/mol, from which the standard enthalpy change of formation (ΔfH°) of aqueous Am3+ ion is −621.2±2.0 kJ/mol. The standard potential Am3+/Am0 is −2.08±0.01 V. == 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 americium(III) has many similarities to the chemistry of lanthanide(III) 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: Am3+ (yellow-reddish), Am4+ (yellow-reddish), AmVO+2; (yellow), AmVIO2+2 (brown) and AmVIIO5−6 (dark green). The absorption spectra have sharp peaks, due to f-f transitions' in the visible and near-infrared regions. Typically, Am(III) has absorption maxima at ca. 504 and 811 nm, Am(V) at ca. 514 and 715 nm, and Am(VI) 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
|
{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
(MnO−4) in acidic solutions. Whereas the Am4+ ions are unstable in solutions and readily convert to Am3+, compounds such as americium dioxide (AmO2) and americium(IV) fluoride (AmF4) are stable in the solid state. The pentavalent oxidation state of americium was first observed in 1951. In acidic aqueous solution the AmO+2 ion is unstable with respect to disproportionation. The reaction 3[AmO2]+ + 4H+ → 2[AmO2]2+ + Am3+ + 2H2O is typical. The chemistry of Am(V) and Am(VI) is comparable to the chemistry of uranium in those oxidation states. In particular, compounds like Li3AmO4 and Li6AmO6 are comparable to uranates and the ion AmO2+2 is comparable to the uranyl ion, UO2+2. Such compounds can be prepared by oxidation of Am(III) in dilute nitric acid with ammonium persulfate. Other oxidising agents that have been used include silver(I) 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). Americium(II) oxide was prepared in minute amounts and has not been characterized in detail. Americium(III) oxide is a red-brown solid with a melting point of 2205 °C. Americium(IV) 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 americium(III), vacuum dried at room temperature, has the chemical formula Am2(C2O4)3·7H2O. 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 with the maximum solubility of 0.25 g/L. === Halides === Halides of americium are known for the oxidation states +2, +3 and +4, where the +3 is most stable, especially in
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{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
solutions. Reduction of Am(III) compounds with sodium amalgam yields Am(II) 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 Am(III) state. Specific lattice constants are: Orthorhombic AmCl2: a = 896.3±0.8 pm, b = 757.3±0.8 pm and c = 453.2±0.6 pm Tetragonal AmBr2: a = 1159.2±0.4 pm and c = 712.1±0.3 pm. They can also be prepared by reacting metallic americium with an appropriate mercury halide HgX2, where X = Cl, Br or I: Am + HgX 2 mercury halide → 400 − 500 ∘ C AmX 2 + Hg {\displaystyle {\ce {{Am}+{\underset {mercury\ halide}{HgX2}}->[{} \atop 400-500^{\circ }{\ce {C}}]{AmX2}+{Hg}}}} Americium(III) fluoride (AmF3) is poorly soluble and precipitates upon reaction of Am3+ and fluoride ions in weak acidic solutions: Am 3 + + 3 F − ⟶ AmF 3 ↓ {\displaystyle {\ce {Am^3+ + 3F^- -> AmF3(v)}}} The tetravalent americium(IV) fluoride (AmF4) is obtained by reacting solid americium(III) fluoride with molecular fluorine: 2 AmF 3 + F 2 ⟶ 2 AmF 4 {\displaystyle {\ce {2AmF3 + F2 -> 2AmF4}}} Another known form of solid tetravalent americium fluoride is KAmF5. Tetravalent americium has also been observed in the aqueous phase. For this purpose, black Am(OH)4 was dissolved in 15-M 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 Am(IV) solution to 90 °C did not result in its disproportionation or reduction, however a slow reduction was observed to Am(III) and assigned to self-irradiation of americium by alpha particles. Most americium(III) halides form hexagonal crystals with slight variation of the color and exact structure between the halogens. So, chloride (AmCl3) is
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{
"page_id": 900,
"source": null,
"title": "Americium"
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|
reddish and has a structure isotypic to uranium(III) chloride (space group P63/m) and the melting point of 715 °C. The fluoride is isotypic to LaF3 (space group P63/mmc) and the iodide to BiI3 (space group R3). The bromide is an exception with the orthorhombic PuBr3-type structure and space group Cmcm. Crystals of americium(III) chloride hexahydrate (AmCl3·6H2O) can be prepared by dissolving americium dioxide in hydrochloric acid and evaporating the liquid. Those crystals are hygroscopic and have yellow-reddish 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 rock-salt lattice. === Silicides and borides === Americium monosilicide (AmSi) and "disilicide" (nominally AmSix with: 1.87 < x < 2.0) were obtained by reduction of americium(III) fluoride with elementary silicon in vacuum at 1050 °C (AmSi) and 1150−1200 °C (AmSix). AmSi is a black solid isomorphic with LaSi, it has an orthorhombic crystal symmetry. AmSix has a bright silvery lustre and a tetragonal crystal lattice (space group I41/amd), 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 is predicted to form the organometallic compound amerocene with two cyclooctatetraene ligands, with the chemical formula (η8-C8H8)2Am. A cyclopentadienyl
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{
"page_id": 900,
"source": null,
"title": "Americium"
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|
complex is also known that is likely to be stoichiometrically AmCp3. Formation of the complexes of the type Am(n-C3H7-BTP)3, where BTP stands for 2,6-di(1,2,4-triazin-3-yl)pyridine, in solutions containing n-C3H7-BTP and Am3+ ions has been confirmed by EXAFS. Some of these BTP-type 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 thus 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 metal-phosphate complex at their cell walls. Several studies have been reported on the biosorption and bioaccumulation of americium by bacteria and fungi. In the laboratory, both americium and curium were found to support the growth of methylotrophs. == Fission == The isotope 242mAm (half-life 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 9–14 kg (the uncertainty results from insufficient knowledge of its material properties). It can be lowered to 3–5 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 because of its scarcity and high price. The critical masses of the two 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
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{
"page_id": 900,
"source": null,
"title": "Americium"
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reactors. There are proposals of very compact 10-kW high-flux reactors using as little as 20 grams of 242mAm. Such low-power reactors would be relatively safe to use as neutron sources for radiation therapy in hospitals. == Isotopes == About 18 isotopes and 11 nuclear isomers are known for americium, having mass numbers 229, 230, and 232 through 247. There are two long-lived alpha-emitters; 243Am has a half-life of 7,370 years and is the most stable isotope, and 241Am has a half-life of 432.2 years. The most stable nuclear isomer is 242m1Am; it has a long half-life of 141 years. The half-lives 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 nuclear fission and low critical mass. Americium-241 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. Americium-242 is a short-lived isotope with a half-life of 16.02 h. It mostly (82.7%) converts by β-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 α-decay to 238Np. The latter subsequently decays to 238Pu and then to 234U. Americium-243 transforms by α-emission into 239Np, which converts by β-decay to 239Pu, and the 239Pu changes into 235U by emitting an α-particle. == Applications == === Ionization-type smoke detector === Americium is used in the most common type of household smoke detector, which uses
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{
"page_id": 900,
"source": null,
"title": "Americium"
}
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241Am in the form of americium dioxide as its source 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 neptunium-237, a different transuranic element with a much longer half-life (about 2.14 million years). With its half-life 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 air-filled 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 scattering; however, it is more prone to false alarms. === Radionuclide === As 241Am has a roughly similar half-life 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 mW/g for 241Am and 6.31 mW/g for 243Am (cf. 390 mW/g 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 space-related 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 micrometer-thick foil. Small thickness avoids the
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{
"page_id": 900,
"source": null,
"title": "Americium"
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|
problem of self-absorption of emitted radiation. This problem is pertinent to uranium or plutonium rods, in which only surface layers provide alpha-particles. The fission products of 242mAm can either directly propel the spaceship or 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 self-sustaining "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 241Am pressed with beryllium is an efficient neutron source. Here americium acts as the alpha source, and beryllium produces neutrons owing to its large cross-section for the (α,n) nuclear reaction: Am 95 241 ⟶ Np 93 237 + He 2 4 + γ {\displaystyle {\ce {^{241}_{95}Am -> ^{237}_{93}Np + ^{4}_{2}He + \gamma}}} Be 4 9 + He 2 4 ⟶ C 6 12 + n 0 1 + γ {\displaystyle {\ce {^{9}_{4}Be + ^{4}_{2}He -> ^{12}_{6}C + ^{1}_{0}n + \gamma}}} The most widespread use of 241AmBe neutron sources is a neutron probe – a
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{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
device used to measure the quantity of water present in soil, as well as moisture/density 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 up-converted by neutron capture to 243Am and 244Am, which transforms by β-decay to 244Cm: Am 95 243 → ( n , γ ) Am 95 244 → 10.1 h β − Cm 96 244 {\displaystyle {\ce {^{243}_{95}Am ->[{\ce {(n,\gamma)}}] ^{244}_{95}Am ->[\beta^-][10.1 \ {\ce {h}}] ^{244}_{96}Cm}}} Irradiation of 241Am by 12C or 22Ne ions 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 60-inch 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 === Americium-241 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 X-ray fluorescence spectroscopy, as well as for quality control in fixed nuclear density gauges and nuclear densometers. For example, the element has been employed to gauge glass thickness
|
{
"page_id": 900,
"source": null,
"title": "Americium"
}
|
to help create flat glass. Americium-241 is also suitable for calibration of gamma-ray spectrometers in the low-energy range, since its spectrum consists of nearly a single peak and negligible Compton continuum (at least three orders of magnitude lower intensity). Americium-241 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 gamma-rays 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 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 half-life 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, 17-year-old 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
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{
"page_id": 900,
"source": null,
"title": "Americium"
}
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was exposed to 500 times the occupational standard for americium-241 as a result of an explosion in his lab. McCluskey died at the age of 75 of unrelated pre-existing disease. == See also == Actinides in the environment Category:Americium compounds == Notes == == References == == Bibliography == Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. Penneman, R. A. and Keenan T. K. The radiochemistry of americium and curium, University of California, Los Alamos, California, 1960 Wiberg, Nils (2007). Lehrbuch Der Anorganischen Chemie. De Gruyter. ISBN 978-3-11-017770-1. == Further reading == Nuclides and Isotopes – 14th Edition, GE Nuclear Energy, 1989. Fioni, Gabriele; Cribier, Michel & Marie, Frédéric. "Can the minor actinide, americium-241, be transmuted by thermal neutrons?". Commissariat à l'énergie atomique. Archived from the original on 11 November 2007. Stwertka, Albert (1999). A Guide to the Elements. Oxford University Press, USA. ISBN 978-0-19-508083-4. == External links == Americium at The Periodic Table of Videos (University of Nottingham) ATSDR – Public Health Statement: Americium World Nuclear Association – Smoke Detectors and Americium
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{
"page_id": 900,
"source": null,
"title": "Americium"
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|
In molecular biology mir-202 microRNA is a short RNA molecule. MicroRNAs function to regulate the expression levels of other genes by several mechanisms. The pre-miR-202 in the mouse genome is located fully within an exon, whereas in human it lies across a splice junction. This implies that human miR-202 is exposed to a negative regulation by splicing, whereas murine miR-202 is not. == See also == MicroRNA == References == == Further reading == == External links == Page for mir-202 microRNA precursor family at Rfam
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{
"page_id": 36373382,
"source": null,
"title": "Mir-202 microRNA precursor family"
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Astatine is a chemical element; it has 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 short-lived; the most stable is astatine-210, with a half-life of 8.1 hours. Consequently, a solid sample of the element has never been seen, because any macroscopic specimen would be immediately vaporized by the heat of its radioactivity. The bulk properties of astatine are not known with certainty. Many of them have been estimated from its position on the periodic table as a heavier analog of fluorine, chlorine, bromine, and iodine, the four stable halogens. 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. Chemically, 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 astatine was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley. They named it from the Ancient Greek astatos (ἄστατος) '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, astatine-210, nor the medically useful astatine-211 occur naturally; they are usually produced by bombarding bismuth-209 with alpha particles. == Characteristics == Astatine is an extremely radioactive element; all its isotopes have half-lives of 8.1 hours or less, decaying into other astatine isotopes, bismuth, polonium, or radon.
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Most of its isotopes are very unstable, with half-lives of seconds or less. Of the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than the longest-lived francium isotopes (205–211At) 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 half-life, 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 gray/violet. 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 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 analog of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap
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of 0.7 eV). Alternatively, if condensed astatine forms a metallic phase, as has been predicted, it may have a monatomic face-centered cubic structure; in this structure, it may well be a superconductor, like the similar high-pressure phase of iodine. Metallic astatine is expected to have a density of 8.91–8.95 g/cm3. Evidence for (or against) the existence of 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 300±10 pm, dissociation energy <50 kJ/mol, and heat of vaporization (∆Hvap) 54.39 kJ/mol. 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 by-products, and other unwanted nano-scale interactions". Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions, typically less than 10−10 mol·L−1. Some properties, such as anion formation, align with other halogens. Astatine has some metallic characteristics as well, such 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. Astatine has an electronegativity of 2.2 on the
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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 Allred–Rochow scale (1.9) being less 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 mol−1, 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 spin–orbit interactions. The first ionization energy of astatine is about 899 kJ mol−1, which continues the trend of decreasing first ionization 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; the chemical properties of tennessine, the next-heavier group 17 element, have not yet been investigated, however. Astatine compounds have been synthesized in nano-scale 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
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adopt odd-numbered 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; acidification by dilute nitric acid gives the At0 or At+ forms, and the subsequent addition of silver(I) may only partially, at best, precipitate astatine as silver(I) astatide (AgAt). Iodine, in contrast, is not oxidized, and precipitates readily as silver(I) iodide. Astatine is known to bind to boron, carbon, and nitrogen. Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C 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 dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate (a non-coordinating anion) and with nitrate, [At(C5H5N)2]NO3. This cation exists as a coordination complex in which two dative covalent bonds separately link the astatine(I) centre with each of the pyridine rings via their nitrogen atoms. With oxygen, there is evidence of the species AtO− and AtO+ in aqueous 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 species previously thought to be AtO−2 has since been determined to be AtO(OH)−2, a hydrolysis
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product of AtO+ (another such hydrolysis product being AtOOH). The well characterized AtO−3 anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide. Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported. Further oxidation of AtO−3, such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion AtO−4; 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 solutions, monovalent or intermediate positive states of astatine coprecipitate with the insoluble salts of metal cations such as silver(I) iodate or thallium(I) dichromate. Astatine may form bonds to the other chalcogens; these include S7At+ and At(CSN)−2 with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium 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 iodine/iodide solution to form AtI, whereas AtBr requires (aside from astatine) an iodine/iodine monobromide/bromide solution. The excess of iodides or bromides may lead to AtBr−2 and AtI−2 ions, or in a chloride solution, they may produce species like AtCl−2 or AtBrCl− via equilibrium reactions with the chlorides. Oxidation of the element with dichromate (in nitric acid solution) showed that adding chloride turned the astatine into a molecule likely to be either AtCl or AtOCl. Similarly, AtOCl−2 or AtCl−2 may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed
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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 helium-filled 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 non-volatile 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 established 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 "eka-iodine" (from Sanskrit eka 'one') to imply it was one space under iodine (in the same manner as eka-silicon, eka-boron, 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 eka-iodine 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 have
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isolated as the thorium series equivalent of radium F (polonium-210) in the radium series. The properties he reported for dakin do not correspond to those of astatine, and astatine's radioactivity would have prevented him from handling it in the quantities he claimed. 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 by observing its X-ray emission lines. 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 Paneth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine-218, his means to detect it were too weak, by current standards, to enable correct identification; moreover, he could not perform chemical tests on the element. 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 (polonium-218), choosing the name "helvetium" (from Helvetia, the Latin name of Switzerland). Berta Karlik and Traude
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Bernert were unsuccessful in reproducing his experiments, and subsequently attributed Minder's results to contamination of his radon stream (radon-222 is the parent isotope of polonium-218). In 1942, Minder, in collaboration with the English scientist Alice Leigh-Smith, announced the discovery of another isotope of element 85, presumed to be the product of thorium A (polonium-216) beta decay. They named this substance "anglo-helvetium", but Karlik and Bernert were again unable to reproduce these results. Later in 1940, Dale R. Corson, Kenneth Ross MacKenzie, and Emilio Segrè isolated the element at the University of California, Berkeley. Instead of searching for the element in nature, the scientists created it by bombarding bismuth-209 with alpha particles in a cyclotron (particle accelerator) to produce, after emission of two neutrons, astatine-211. 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 so-called uranium series, and then in the actinium series. (Since then, astatine 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 Segrè suggested the name "astatine" coming from the Ancient Greek astatos (ἄστατος) meaning 'unstable', because of its propensity for radioactive decay, with the ending
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"-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 iodine-like, cationic, or amphoteric behavior. In a 2003 retrospective, Corson wrote that "some of the properties [of astatine] are similar to iodine ... it also exhibits metallic properties, more like its metallic neighbors Po and Bi." == Isotopes == There are 41 known isotopes of astatine, with mass numbers of 188 and 190–229. Theoretical modeling suggests that about 37 more isotopes could exist. No stable or long-lived 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. Astatine-211 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 half-life to the previous isotope (8.1 hours for astatine-210 and 7.2 hours for astatine-211), 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 astatine-213 releasing the most energy. For this reason, it is the shortest-lived astatine isotope. Even though heavier astatine isotopes release less energy, no long-lived 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
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that 215At is in fact beta-stable, as it has the lowest mass of all isobars with A = 215. Astatine-210 and most of the lighter isotopes exhibit beta plus decay (positron emission), astatine-217 and heavier isotopes except astatine-218 exhibit beta minus decay, while astatine-211 undergoes electron capture. The most stable isotope is astatine-210, which has a half-life of 8.1 hours. The primary decay mode is beta plus, to the relatively long-lived (in comparison to astatine isotopes) alpha emitter polonium-210. In total, only five isotopes have half-lives exceeding one hour (astatine-207 to -211). The least stable ground state isotope is astatine-213, with a half-life of 125 nanoseconds. It undergoes alpha decay to the extremely long-lived bismuth-209. 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 "meta-state", 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 astatine-202m1, which has a half-life of about 3 minutes, longer than those of all the ground states bar those of isotopes 203–211 and 220. The least stable is astatine-213m1; its half-life of 110 nanoseconds is shorter than 125 nanoseconds for astatine-213, the shortest-lived ground state. == Natural occurrence == Astatine is the rarest 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).
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Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes (astatine-215, -217, -218 and -219) are instead continuously produced as a result of the decay of radioactive thorium and uranium ores, and trace quantities of neptunium-237. The landmass of North and South America combined, to a depth of 16 kilometers (10 miles), contains only about one trillion astatine-215 atoms at any given time (around 3.5 × 10−10 grams). Astatine-217 is produced via the radioactive decay of neptunium-237. Primordial remnants of the latter isotope—due to its relatively short half-life of 2.14 million years—are no longer present on Earth. However, trace amounts occur naturally as a product of transmutation reactions in uranium ores. Astatine-218 was the first astatine isotope discovered in nature. Astatine-219, with a half-life 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. Astatine-216 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 bismuth-209 with energetic alpha particles, and this is still the major route used to create the relatively long-lived isotopes astatine-209 through astatine-211. Astatine is only produced in minuscule quantities, with modern techniques allowing production runs of up to 6.6 gigabecquerels (about 86 nanograms or 2.47×1014 atoms). Synthesis of greater quantities of astatine using this method is constrained by the limited availability 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
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generated via proton irradiation of thorium or uranium to yield radon-211, in turn decaying to astatine-211. Contamination with astatine-210 is expected to be a drawback of this method. The most important isotope is astatine-211, 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 (bismuth-209), the reaction may occur in three possible ways, producing astatine-209, astatine-210, or astatine-211. Although higher energies can produce more astatine-211, it will produce unwanted astatine-210 that decays to toxic polonium-210 as well. Instead, the maximum energy of the particle accelerator is set to be below or slightly above the threshold of astatine-210 production, in order to maximize the production of astatine-211 while keeping the amount of astatine-210 at an acceptable level. === 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 approaches—dry 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. Pre-1985 techniques more often addressed the elimination of co-produced toxic polonium; this requirement is now mitigated by capping the energy of the cyclotron irradiation beam. ==== Dry ==== The astatine-containing cyclotron target is heated to a temperature of around
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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 acid, such as hydrochloric acid. Astatine is extracted from this acid using an organic solvent such as dibutyl ether, diisopropyl ether (DIPE), or thiosemicarbazide. Using liquid-liquid extraction, the astatine product can be repeatedly washed with an acid, such as HCl, and extracted into the organic solvent layer. A separation yield of 93% using nitric acid has been reported, falling to 72% by the time purification procedures were completed (distillation of nitric acid, purging residual nitrogen oxides, and redissolving bismuth nitrate to enable liquid–liquid extraction). Wet methods involve "multiple radioactivity handling steps" and have not been considered well suited for isolating larger quantities of astatine. However, wet extraction methods are being examined for use in production of larger quantities of astatine-211, as it is thought that wet extraction methods can provide more consistency. They can enable the production of astatine in a specific oxidation state
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and may have greater applicability in experimental radiochemistry. == Uses and precautions == Newly formed astatine-211 is the subject of ongoing research in nuclear medicine. It must be used quickly as it decays with a half-life of 7.2 hours; this is long enough to permit multistep labeling strategies. Astatine-211 has potential for targeted alpha-particle therapy, since it decays either via emission of an alpha particle (to bismuth-207), or via electron capture (to an extremely short-lived nuclide, polonium-211, which undergoes further alpha decay), very quickly reaching its stable granddaughter lead-207. Polonium X-rays emitted as a result of the electron capture branch, in the range of 77–92 keV, enable the tracking of astatine in animals and patients. Although astatine-210 has a slightly longer half-life, it is wholly unsuitable because it usually undergoes beta plus decay to the extremely toxic polonium-210. The principal medicinal difference between astatine-211 and iodine-131 (a radioactive iodine isotope also used in medicine) is that iodine-131 emits high-energy beta particles, and astatine does not. Beta particles have much greater penetrating power through tissues than do the much heavier alpha particles. An average alpha particle released by astatine-211 can travel up to 70 μm through surrounding tissues; an average-energy beta particle emitted by iodine-131 can travel nearly 30 times as far, to about 2 mm. The short half-life and limited penetrating power of alpha radiation through tissues offers advantages in situations where the "tumor burden is low and/or malignant cell populations are located in close proximity to essential normal tissues." Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine-211 atoms bound per cell. Several obstacles have been encountered in the development of astatine-based radiopharmaceuticals for cancer treatment. World War II delayed research for close to a decade. Results of early experiments
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indicated that a cancer-selective carrier would need to be developed and it was not until the 1970s that monoclonal antibodies became available for this purpose. Unlike iodine, astatine shows a tendency to dehalogenate from molecular carriers such as these, particularly at sp3 carbon sites (less so from sp2 sites). Given the toxicity of astatine accumulated and retained in the body, this emphasized the need to ensure it remained attached to its host molecule. While astatine carriers that are slowly metabolized can be assessed for their efficacy, more rapidly metabolized carriers remain a significant obstacle to the evaluation of astatine in nuclear medicine. Mitigating the effects of astatine-induced radiolysis of labeling chemistry and carrier molecules is another area requiring further development. A practical application for astatine as a cancer treatment would potentially be suitable for a "staggering" number of patients; production of astatine in the quantities that would be required remains an issue. Animal studies show that astatine, similarly to iodine—although to a lesser extent, perhaps because of its slightly more metallic nature—is preferentially (and dangerously) concentrated in the thyroid gland. Unlike iodine, astatine also shows a tendency to be taken up by the lungs and spleen, possibly because of in-body oxidation of At– to At+. If administered in the form of a radiocolloid it tends to concentrate in the liver. Experiments in rats and monkeys suggest that astatine-211 causes much greater damage to the thyroid gland than does iodine-131, with repetitive injection of the nuclide resulting in necrosis and cell dysplasia within the gland. Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue; this conclusion remained controversial for many years. General agreement was later reached that this was likely caused by the effect of breast tissue irradiation combined with hormonal changes due to
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irradiation of the ovaries. Trace amounts of astatine can be handled safely in fume hoods if they are well-aerated; biological uptake of the element must be avoided. == See also == Radiation protection == Notes == == References == == Bibliography == Corson, D. R.; MacKenzie, K. R.; Segrè, E. (1940). "Artificially Radioactive Element 85". Physical Review. 58 (8): 672–678. Bibcode:1940PhRv...58..672C. doi:10.1103/PhysRev.58.672. (subscription required) Greenwood, N. N.; Earnshaw, A. (2002). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-7506-3365-9. Kugler, H. K.; Keller, C. (1985). 'At, Astatine', System No. 8a. Gmelin Handbook of Inorganic and Organometallic Chemistry. Vol. 8 (8th ed.). Springer-Verlag. ISBN 978-3-540-93516-2. Lavrukhina, Avgusta Konstantinovna; Pozdnyakov, Aleksandr Aleksandrovich (1970). Analytical Chemistry of Technetium, Promethium, Astatine, and Francium. Translated by R. Kondor. Ann Arbor–Humphrey Science Publishers. ISBN 978-0-250-39923-9. Scerri, Eric (2013). A Tale of Seven Elements. Oxford University Press, ISBN 9780195391312. Vértes, A.; Nagy, S.; Klencsár, Z. (2003). Handbook of Nuclear Chemistry. Vol. 4. Springer. ISBN 978-1-4020-1316-4. Zuckerman, J. J.; Hagen, A. P. (1989). Inorganic Reactions and Methods, Volume 3, The Formation of Bonds to Halogens (Part 1). John Wiley & Sons. ISBN 978-0-471-18656-4. Zuckerman, J. J.; Hagen, A. P. (1990). Inorganic Reactions and Methods, Volume 4, The Formation of Bonds to Halogens (Part 2). John Wiley & Sons. ISBN 978-0-471-18657-1. == External links == Astatine at The Periodic Table of Videos (University of Nottingham) Astatine: Halogen or Metal?
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{
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"title": "Astatine"
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In biology, trimorphism is the existence in certain plants and animals of three distinct forms, especially in connection with the reproductive organs. In trimorphic plants there are three forms, differing in the lengths of their pistils and stamens, in size and color of their pollen grains, and in some other respects; and, as in each of the three forms there are two sets of stamens, the three forms possess altogether six sets of stamens and three kinds of pistils. These organs are so proportioned in length to each other that half the stamens in two of the forms stand on a level with the stigma of the third form. To obtain full fertility with these plants, it is necessary that the stigma of the one should be fertilized by pollen taken from the stamens of corresponding height in another form. Hence six unions are legitimate, that is, fully fertile, and 12 are illegitimate, or more or less unfertile. Wallace has shown that the females of certain butterflies from the Malay Archipelago appear in three conspicuously distinct forms without intermediate links. In crystallography, trimorphism refers to the occurrence of certain forms in minerals which have the same chemical composition, but are referable to three systems of crystallization. == See also == Sexual dimorphism == Notes == Text from Collier's New Encyclopedia (1921).
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{
"page_id": 1377160,
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"title": "Trimorphism"
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Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element. Atoms are extremely small, typically around 100 picometers across. A human hair is about a million carbon atoms wide. Atoms are smaller than the shortest wavelength of visible light, which means humans cannot see atoms with conventional microscopes. They are so small that accurately predicting their behavior using classical physics is not possible due to quantum effects. More than 99.9994% of an atom's mass is in the nucleus. Protons have a positive electric charge and neutrons have no charge, so the nucleus is positively charged. The electrons are negatively charged, and this opposing charge is what binds them to the nucleus. If the numbers of protons and electrons are equal, as they normally are, then the atom is electrically neutral as a whole. If an atom has more electrons than protons, then it has an overall negative charge and is called a negative ion (or anion). Conversely, if it has more protons than electrons, it has a positive charge and is called a positive ion (or cation). The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged
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protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus splits and leaves behind different elements. This is a form of nuclear decay. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to attach and detach from each other is responsible for most of the physical changes observed in nature. Chemistry is the science that studies these changes. == History of atomic theory == === In philosophy === The basic idea that matter is made up of tiny indivisible particles is an old idea that appeared in many ancient cultures. The word atom is derived from the ancient Greek word atomos, which means "uncuttable". But this ancient idea was based in philosophical reasoning rather than scientific reasoning. Modern atomic theory is not based on these old concepts. In the early 19th century, the scientist John Dalton found evidence that matter really is composed of discrete units, and so applied the word atom to those units. === Dalton's law of multiple proportions === In the early 1800s, John Dalton compiled experimental data gathered by him and other scientists and discovered a pattern now known as the "law of multiple proportions". He noticed that in any group of chemical compounds which all contain two particular chemical elements, the amount of Element A per measure of Element B will differ across these compounds by ratios of small whole numbers. This pattern suggested that each element combines with other elements in multiples of a basic unit of weight, with each element having a unit of unique weight. Dalton decided to call these units "atoms". For example, there are two types of tin oxide: one is a
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grey powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures, in the grey powder there is about 13.5 g of oxygen for every 100 g of tin, and in the white powder there is about 27 g of oxygen for every 100 g of tin. 13.5 and 27 form a ratio of 1:2. Dalton concluded that in the grey oxide there is one atom of oxygen for every atom of tin, and in the white oxide there are two atoms of oxygen for every atom of tin (SnO and SnO2). Dalton also analyzed iron oxides. There is one type of iron oxide that is a black powder which is 78.1% iron and 21.9% oxygen; and there is another iron oxide that is a red powder which is 70.4% iron and 29.6% oxygen. Adjusting these figures, in the black powder there is about 28 g of oxygen for every 100 g of iron, and in the red powder there is about 42 g of oxygen for every 100 g of iron. 28 and 42 form a ratio of 2:3. Dalton concluded that in these oxides, for every two atoms of iron, there are two or three atoms of oxygen respectively. These substances are known today as iron(II) oxide and iron(III) oxide, and their formulas are FeO and Fe2O3 respectively. Iron(II) oxide's formula is normally written as FeO, but since it is a crystalline substance we could alternately write it as Fe2O2, and when we contrast that with Fe2O3, the 2:3 ratio for the oxygen is plain to see. As a final example: nitrous oxide is 63.3% nitrogen and 36.7% oxygen, nitric oxide is 44.05% nitrogen and 55.95% oxygen, and nitrogen dioxide is 29.5% nitrogen and
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70.5% oxygen. Adjusting these figures, in nitrous oxide there is 80 g of oxygen for every 140 g of nitrogen, in nitric oxide there is about 160 g of oxygen for every 140 g of nitrogen, and in nitrogen dioxide there is 320 g of oxygen for every 140 g of nitrogen. 80, 160, and 320 form a ratio of 1:2:4. The respective formulas for these oxides are N2O, NO, and NO2. === Discovery of the electron === In 1897, J. J. Thomson discovered that cathode rays can be deflected by electric and magnetic fields, which meant that cathode rays are not a form of light but made of electrically charged particles, and their charge was negative given the direction the particles were deflected in. He measured these particles to be 1,700 times lighter than hydrogen (the lightest atom). He called these new particles corpuscles but they were later renamed electrons since these are the particles that carry electricity. Thomson also showed that electrons were identical to particles given off by photoelectric and radioactive materials. Thomson explained that an electric current is the passing of electrons from one atom to the next, and when there was no current the electrons embedded themselves in the atoms. This in turn meant that atoms were not indivisible as scientists thought. The atom was composed of electrons whose negative charge was balanced out by some source of positive charge to create an electrically neutral atom. Ions, Thomson explained, must be atoms which have an excess or shortage of electrons. === Discovery of the nucleus === The electrons in the atom logically had to be balanced out by a commensurate amount of positive charge, but Thomson had no idea where this positive charge came from, so he tentatively proposed that it was everywhere in the
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atom, the atom being in the shape of a sphere. This was the mathematically simplest hypothesis to fit the available evidence, or lack thereof. Following from this, Thomson imagined that the balance of electrostatic forces would distribute the electrons throughout the sphere in a more or less even manner. Thomson's model is popularly known as the plum pudding model, though neither Thomson nor his colleagues used this analogy. Thomson's model was incomplete, it was unable to predict any other properties of the elements such as emission spectra and valencies. It was soon rendered obsolete by the discovery of the atomic nucleus. Between 1908 and 1913, Ernest Rutherford and his colleagues Hans Geiger and Ernest Marsden performed a series of experiments in which they bombarded thin foils of metal with a beam of alpha particles. They did this to measure the scattering patterns of the alpha particles. They spotted a small number of alpha particles being deflected by angles greater than 90°. This shouldn't have been possible according to the Thomson model of the atom, whose charges were too diffuse to produce a sufficiently strong electric field. The deflections should have all been negligible. Rutherford proposed that the positive charge of the atom is concentrated in a tiny volume at the center of the atom and that the electrons surround this nucleus in a diffuse cloud. This nucleus carried almost all of the atom's mass. Only such an intense concentration of charge, anchored by its high mass, could produce an electric field that could deflect the alpha particles so strongly. === Bohr model === A problem in classical mechanics is that an accelerating charged particle radiates electromagnetic radiation, causing the particle to lose kinetic energy. Circular motion counts as acceleration, which means that an electron orbiting a central charge should spiral
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down into that nucleus as it loses speed. In 1913, the physicist Niels Bohr proposed a new model in which the electrons of an atom were assumed to orbit the nucleus but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy corresponding to absorption or radiation of a photon. This quantization was used to explain why the electrons' orbits are stable and why elements absorb and emit electromagnetic radiation in discrete spectra. Bohr's model could only predict the emission spectra of hydrogen, not atoms with more than one electron. === Discovery of protons and neutrons === Back in 1815, William Prout observed that the atomic weights of many elements were multiples of hydrogen's atomic weight, which is in fact true for all of them if one takes isotopes into account. In 1898, J. J. Thomson found that the positive charge of a hydrogen ion is equal to the negative charge of an electron, and these were then the smallest known charged particles. Thomson later found that the positive charge in an atom is a positive multiple of an electron's negative charge. In 1913, Henry Moseley discovered that the frequencies of X-ray emissions from an excited atom were a mathematical function of its atomic number and hydrogen's nuclear charge. In 1919, Rutherford bombarded nitrogen gas with alpha particles and detected hydrogen ions being emitted from the gas, and concluded that they were produced by alpha particles hitting and splitting the nuclei of the nitrogen atoms. These observations led Rutherford to conclude that the hydrogen nucleus is a singular particle with a positive charge equal to the electron's negative charge. He named this particle "proton" in 1920. The number of protons in an atom (which Rutherford called the "atomic
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{
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number") was found to be equal to the element's ordinal number on the periodic table and therefore provided a simple and clear-cut way of distinguishing the elements from each other. The atomic weight of each element is higher than its proton number, so Rutherford hypothesized that the surplus weight was carried by unknown particles with no electric charge and a mass equal to that of the proton. In 1928, Walter Bothe observed that beryllium emitted a highly penetrating, electrically neutral radiation when bombarded with alpha particles. It was later discovered that this radiation could knock hydrogen atoms out of paraffin wax. Initially it was thought to be high-energy gamma radiation, since gamma radiation had a similar effect on electrons in metals, but James Chadwick found that the ionization effect was too strong for it to be due to electromagnetic radiation, so long as energy and momentum were conserved in the interaction. In 1932, Chadwick exposed various elements, such as hydrogen and nitrogen, to the mysterious "beryllium radiation", and by measuring the energies of the recoiling charged particles, he deduced that the radiation was actually composed of electrically neutral particles which could not be massless like the gamma ray, but instead were required to have a mass similar to that of a proton. Chadwick now claimed these particles as Rutherford's neutrons. === The current consensus model === In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics (matrix mechanics). One year earlier, Louis de Broglie had proposed that all particles behave like waves to some extent, and in 1926 Erwin Schrödinger used this idea to develop the Schrödinger equation, which describes electrons as three-dimensional waveforms rather than points in space. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values
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{
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for both the position and momentum of a particle at a given point in time. This became known as the uncertainty principle, formulated by Werner Heisenberg in 1927. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be found. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. == Structure == === Subatomic particles === Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the proton, and the neutron. The electron is the least massive of these particles by four orders of magnitude at 9.11×10−31 kg, with a negative electrical charge and a size that is too small to be measured using available techniques. It was the lightest particle with a positive rest mass measured, until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details. Protons have a positive charge and a mass of 1.6726×10−27 kg. The number of protons in an atom is
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called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton. Neutrons have no electrical charge and have a mass of 1.6749×10−27 kg. Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons (collectively known as nucleons) have comparable dimensions—on the order of 2.5×10−15 m—although the 'surface' of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick. In the Standard Model of physics, electrons are truly elementary particles with no internal structure, whereas protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +2/3) and one down quark (with a charge of −1/3). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The quarks are held together by the strong interaction (or strong force), which is mediated by gluons. The protons and neutrons, in turn, are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces. === Nucleus === All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately
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equal to 1.07 A 3 {\displaystyle 1.07{\sqrt[{3}]{A}}} femtometres, where A {\displaystyle A} is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 105 fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other. Atoms of the same element have the same number of protons, called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. The proton, the electron, and the neutron are classified as fermions. Fermions obey the Pauli exclusion principle which prohibits identical fermions, such as multiple protons, from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus and to all electrons of the electron cloud. A nucleus that has a different number of protons than neutrons can potentially drop to a lower energy state through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus. The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high
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{
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energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei. For example, at the core of the Sun protons require energies of 3 to 10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei—usually through radioactive decay. The nucleus can also be modified through bombardment by high energy subatomic particles or photons. If this modifies the number of protons in a nucleus, the atom changes to a different chemical element. If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy (such as a gamma ray, or the kinetic energy of a beta particle), as described by Albert Einstein's mass–energy equivalence formula, E = mc2, where m is the mass loss and c is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon begins to decrease. That means that a fusion process producing a nucleus that has an atomic number higher than
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about 26, and a mass number higher than about 60, is an endothermic process. Thus, more massive nuclei cannot undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. === Electron cloud === The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force. Hence electrons bound near the center of the potential well require more energy to escape than those at greater separations. Electrons, like other particles, have properties of both a particle and a wave. The electron cloud is a region inside the potential well where each electron forms a type of three-dimensional standing wave—a wave form that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterises the probability that an electron appears to be at a particular location when its position is measured. Only a discrete (or quantized) set of these orbitals exist around the nucleus, as other possible wave patterns rapidly decay into a more stable form. Orbitals can have one or more ring or node structures, and differ from each other in size, shape and orientation. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic
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energy values, defined by the differences in the energies of the quantum states, are responsible for atomic spectral lines. The amount of energy needed to remove or add an electron—the electron binding energy—is far less than the binding energy of nucleons. For example, it requires only 13.6 eV to strip a ground-state electron from a hydrogen atom, compared to 2.23 million eV for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons. Atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. == Properties == === Nuclear properties === By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible. About 339 nuclides occur naturally on Earth, of which 251 (about 74%) have not been observed to decay, and are referred to as "stable isotopes". Only 90 nuclides are stable theoretically, while another 161 (bringing
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{
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the total to 251) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 35 radioactive nuclides have half-lives longer than 100 million years, and are long-lived enough to have been present since the birth of the Solar System. This collection of 286 nuclides are known as primordial nuclides. Finally, an additional 53 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or as products of natural energetic processes on Earth, such as cosmic ray bombardment (for example, carbon-14). For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.1 stable isotopes per element. Twenty-six "monoisotopic elements" have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 43, 61, and all elements numbered 83 or higher have no stable isotopes.: 1–12 Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain "magic numbers" of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus; filled shells, such as the filled shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 251 known stable nuclides, only four have both an odd number of protons and odd number of neutrons: hydrogen-2 (deuterium), lithium-6, boron-10, and nitrogen-14. (Tantalum-180m is odd-odd and observationally stable, but is predicted to decay with a very long half-life.) Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over
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a billion years: potassium-40, vanadium-50, lanthanum-138, and lutetium-176. Most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects. === Mass === The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles (called "nucleons") in a given atom is called the mass number. It is a positive integer and dimensionless (instead of having dimension of mass), because it expresses a count. An example of use of a mass number is "carbon-12," which has 12 nucleons (six protons and six neutrons). The actual mass of an atom at rest is often expressed in daltons (Da), also called the unified atomic mass unit (u). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately 1.66×10−27 kg. Hydrogen-1 (the lightest isotope of hydrogen which is also the nuclide with the lowest mass) has an atomic weight of 1.007825 Da. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal (within 1%) to its mass number times the atomic mass unit (for example the mass of a nitrogen-14 is roughly 14 Da), but this number will not be exactly an integer except (by definition) in the case of carbon-12. The heaviest stable atom is lead-208, with a mass of 207.9766521 Da. As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms (about 6.022×1023). This number was chosen so that if an element has an atomic mass of 1 u, a
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mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon-12 atom has an atomic mass of exactly 12 Da, and so a mole of carbon-12 atoms weighs exactly 0.012 kg. === Shape and size === Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. This assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right). Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm. When subjected to external forces, like electrical fields, the shape of an atom may deviate from spherical symmetry. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites. Significant ellipsoidal deformations have been shown to occur for sulfur ions and chalcogen ions in pyrite-type compounds. Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700 nm) so they cannot be viewed using
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an optical microscope, although individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. A single drop of water contains about 2 sextillion (2×1021) atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of 2×10−4 kg contains about 10 sextillion (1022) atoms of carbon. If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple. === Radioactive decay === Every element has one or more isotopes that have unstable nuclei that are subject to radioactive decay, causing the nucleus to emit particles or electromagnetic radiation. Radioactivity can occur when the radius of a nucleus is large compared with the radius of the strong force, which only acts over distances on the order of 1 fm. The most common forms of radioactive decay are: Alpha decay: this process is caused when the nucleus emits an alpha particle, which is a helium nucleus consisting of two protons and two neutrons. The result of the emission is a new element with a lower atomic number. Beta decay (and electron capture): these processes are regulated by the weak force, and result from a transformation of a neutron into a proton, or a proton into a neutron. The neutron to proton transition is accompanied by the emission of an electron and an antineutrino, while proton to neutron transition (except in electron capture) causes the emission of a positron and a neutrino. The electron or positron emissions are called beta particles. Beta decay either increases or decreases the atomic number of the nucleus by one. Electron capture is more common than positron
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|
emission, because it requires less energy. In this type of decay, an electron is absorbed by the nucleus, rather than a positron emitted from the nucleus. A neutrino is still emitted in this process, and a proton changes to a neutron. Gamma decay: this process results from a change in the energy level of the nucleus to a lower state, resulting in the emission of electromagnetic radiation. The excited state of a nucleus which results in gamma emission usually occurs following the emission of an alpha or a beta particle. Thus, gamma decay usually follows alpha or beta decay. Other more rare types of radioactive decay include ejection of neutrons or protons or clusters of nucleons from a nucleus, or more than one beta particle. An analog of gamma emission which allows excited nuclei to lose energy in a different way, is internal conversion—a process that produces high-speed electrons that are not beta rays, followed by production of high-energy photons that are not gamma rays. A few large nuclei explode into two or more charged fragments of varying masses plus several neutrons, in a decay called spontaneous nuclear fission. Each radioactive isotope has a characteristic decay time period—the half-life—that is determined by the amount of time needed for half of a sample to decay. This is an exponential decay process that steadily decreases the proportion of the remaining isotope by 50% every half-life. Hence after two half-lives have passed only 25% of the isotope is present, and so forth. === Magnetic moment === Elementary particles possess an intrinsic quantum mechanical property known as spin. This is analogous to the angular momentum of an object that is spinning around its center of mass, although strictly speaking these particles are believed to be point-like and cannot be said to be rotating. Spin
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{
"page_id": 902,
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|
is measured in units of the reduced Planck constant (ħ), with electrons, protons and neutrons all having spin 1⁄2 ħ, or "spin-1⁄2". In an atom, electrons in motion around the nucleus possess orbital angular momentum in addition to their spin, while the nucleus itself possesses angular momentum due to its nuclear spin. The magnetic field produced by an atom—its magnetic moment—is determined by these various forms of angular momentum, just as a rotating charged object classically produces a magnetic field, but the most dominant contribution comes from electron spin. Due to the nature of electrons to obey the Pauli exclusion principle, in which no two electrons may be found in the same quantum state, bound electrons pair up with each other, with one member of each pair in a spin up state and the other in the opposite, spin down state. Thus these spins cancel each other out, reducing the total magnetic dipole moment to zero in some atoms with even number of electrons. In ferromagnetic elements such as iron, cobalt and nickel, an odd number of electrons leads to an unpaired electron and a net overall magnetic moment. The orbitals of neighboring atoms overlap and a lower energy state is achieved when the spins of unpaired electrons are aligned with each other, a spontaneous process known as an exchange interaction. When the magnetic moments of ferromagnetic atoms are lined up, the material can produce a measurable macroscopic field. Paramagnetic materials have atoms with magnetic moments that line up in random directions when no magnetic field is present, but the magnetic moments of the individual atoms line up in the presence of a field. The nucleus of an atom will have no spin when it has even numbers of both neutrons and protons, but for other cases of odd numbers,
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{
"page_id": 902,
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the nucleus may have a spin. Normally nuclei with spin are aligned in random directions because of thermal equilibrium, but for certain elements (such as xenon-129) it is possible to polarize a significant proportion of the nuclear spin states so that they are aligned in the same direction—a condition called hyperpolarization. This has important applications in magnetic resonance imaging. === Energy levels === The potential energy of an electron in an atom is negative relative to when the distance from the nucleus goes to infinity; its dependence on the electron's position reaches the minimum inside the nucleus, roughly in inverse proportion to the distance. In the quantum-mechanical model, a bound electron can occupy only a set of states centered on the nucleus, and each state corresponds to a specific energy level; see time-independent Schrödinger equation for a theoretical explanation. An energy level can be measured by the amount of energy needed to unbind the electron from the atom, and is usually given in units of electronvolts (eV). The lowest energy state of a bound electron is called the ground state, i.e., stationary state, while an electron transition to a higher level results in an excited state. The electron's energy increases along with n because the (average) distance to the nucleus increases. Dependence of the energy on ℓ is caused not by the electrostatic potential of the nucleus, but by interaction between electrons. For an electron to transition between two different states, e.g. ground state to first excited state, it must absorb or emit a photon at an energy matching the difference in the potential energy of those levels, according to the Niels Bohr model, what can be precisely calculated by the Schrödinger equation. Electrons jump between orbitals in a particle-like fashion. For example, if a single photon strikes the electrons,
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{
"page_id": 902,
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only a single electron changes states in response to the photon; see Electron properties. The energy of an emitted photon is proportional to its frequency, so these specific energy levels appear as distinct bands in the electromagnetic spectrum. Each element has a characteristic spectrum that can depend on the nuclear charge, subshells filled by electrons, the electromagnetic interactions between the electrons and other factors. When a continuous spectrum of energy is passed through a gas or plasma, some of the photons are absorbed by atoms, causing electrons to change their energy level. Those excited electrons that remain bound to their atom spontaneously emit this energy as a photon, traveling in a random direction, and so drop back to lower energy levels. Thus the atoms behave like a filter that forms a series of dark absorption bands in the energy output. (An observer viewing the atoms from a view that does not include the continuous spectrum in the background, instead sees a series of emission lines from the photons emitted by the atoms.) Spectroscopic measurements of the strength and width of atomic spectral lines allow the composition and physical properties of a substance to be determined. Close examination of the spectral lines reveals that some display a fine structure splitting. This occurs because of spin–orbit coupling, which is an interaction between the spin and motion of the outermost electron. When an atom is in an external magnetic field, spectral lines become split into three or more components; a phenomenon called the Zeeman effect. This is caused by the interaction of the magnetic field with the magnetic moment of the atom and its electrons. Some atoms can have multiple electron configurations with the same energy level, which thus appear as a single spectral line. The interaction of the magnetic field with the
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{
"page_id": 902,
"source": null,
"title": "Atom"
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atom shifts these electron configurations to slightly different energy levels, resulting in multiple spectral lines. The presence of an external electric field can cause a comparable splitting and shifting of spectral lines by modifying the electron energy levels, a phenomenon called the Stark effect. If a bound electron is in an excited state, an interacting photon with the proper energy can cause stimulated emission of a photon with a matching energy level. For this to occur, the electron must drop to a lower energy state that has an energy difference matching the energy of the interacting photon. The emitted photon and the interacting photon then move off in parallel and with matching phases. That is, the wave patterns of the two photons are synchronized. This physical property is used to make lasers, which can emit a coherent beam of light energy in a narrow frequency band. === Valence and bonding behavior === Valency is the combining power of an element. It is determined by the number of bonds it can form to other atoms or groups. The outermost electron shell of an atom in its uncombined state is known as the valence shell, and the electrons in that shell are called valence electrons. The number of valence electrons determines the bonding behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills (or empties) their outer valence shells. For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with one-electron more than a filled shell, and others that are one-electron short of a full shell, such as occurs in the compound sodium chloride and other chemical ionic salts. Many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds.
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{
"page_id": 902,
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Thus, chemical bonding between these elements takes many forms of electron-sharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds. The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that is aligned in the same column of the table. (The horizontal rows correspond to the filling of a quantum shell of electrons.) The elements at the far right of the table have their outer shell completely filled with electrons, which results in chemically inert elements known as the noble gases. === States === Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the conditions, materials can transition between solids, liquids, gases, and plasmas. Within a state, a material can also exist in different allotropes. An example of this is solid carbon, which can exist as graphite or diamond. Gaseous allotropes exist as well, such as dioxygen and ozone. At temperatures close to absolute zero, atoms can form a Bose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale. This super-cooled collection of atoms then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior. == Identification == While atoms are too small to be seen, devices such as the scanning tunneling microscope (STM) enable their visualization at the surfaces of solids. The microscope uses the quantum tunneling phenomenon, which allows particles to pass through a barrier that would be insurmountable in the classical perspective. Electrons tunnel through the vacuum between two biased electrodes, providing a tunneling current that is
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{
"page_id": 902,
"source": null,
"title": "Atom"
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exponentially dependent on their separation. One electrode is a sharp tip ideally ending with a single atom. At each point of the scan of the surface the tip's height is adjusted so as to keep the tunneling current at a set value. How much the tip moves to and away from the surface is interpreted as the height profile. For low bias, the microscope images the averaged electron orbitals across closely packed energy levels—the local density of the electronic states near the Fermi level. Because of the distances involved, both electrodes need to be extremely stable; only then periodicities can be observed that correspond to individual atoms. The method alone is not chemically specific, and cannot identify the atomic species present at the surface. Atoms can be easily identified by their mass. If an atom is ionized by removing one of its electrons, its trajectory when it passes through a magnetic field will bend. The radius by which the trajectory of a moving ion is turned by the magnetic field is determined by the mass of the atom. The mass spectrometer uses this principle to measure the mass-to-charge ratio of ions. If a sample contains multiple isotopes, the mass spectrometer can determine the proportion of each isotope in the sample by measuring the intensity of the different beams of ions. Techniques to vaporize atoms include inductively coupled plasma atomic emission spectroscopy and inductively coupled plasma mass spectrometry, both of which use a plasma to vaporize samples for analysis. The atom-probe tomograph has sub-nanometer resolution in 3-D and can chemically identify individual atoms using time-of-flight mass spectrometry. Electron emission techniques such as X-ray photoelectron spectroscopy (XPS) and Auger electron spectroscopy (AES), which measure the binding energies of the core electrons, are used to identify the atomic species present in a sample
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{
"page_id": 902,
"source": null,
"title": "Atom"
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in a non-destructive way. With proper focusing both can be made area-specific. Another such method is electron energy loss spectroscopy (EELS), which measures the energy loss of an electron beam within a transmission electron microscope when it interacts with a portion of a sample. Spectra of excited states can be used to analyze the atomic composition of distant stars. Specific light wavelengths contained in the observed light from stars can be separated out and related to the quantized transitions in free gas atoms. These colors can be replicated using a gas-discharge lamp containing the same element. Helium was discovered in this way in the spectrum of the Sun 23 years before it was found on Earth. == Origin and current state == Baryonic matter forms about 4% of the total energy density of the observable universe, with an average density of about 0.25 particles/m3 (mostly protons and electrons). Within a galaxy such as the Milky Way, particles have a much higher concentration, with the density of matter in the interstellar medium (ISM) ranging from 105 to 109 atoms/m3. The Sun is believed to be inside the Local Bubble, so the density in the solar neighborhood is only about 103 atoms/m3. Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium. Up to 95% of the Milky Way's baryonic matter are concentrated inside stars, where conditions are unfavorable for atomic matter. The total baryonic mass is about 10% of the mass of the galaxy; the remainder of the mass is an unknown dark matter. High temperature inside stars makes most "atoms" fully ionized, that is, separates all electrons from the nuclei. In stellar remnants—with exception of their surface layers—an immense pressure
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{
"page_id": 902,
"source": null,
"title": "Atom"
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make electron shells impossible. === Formation === Electrons are thought to exist in the Universe since early stages of the Big Bang. Atomic nuclei forms in nucleosynthesis reactions. In about three minutes Big Bang nucleosynthesis produced most of the helium, lithium, and deuterium in the Universe, and perhaps some of the beryllium and boron. Ubiquitousness and stability of atoms relies on their binding energy, which means that an atom has a lower energy than an unbound system of the nucleus and electrons. Where the temperature is much higher than ionization potential, the matter exists in the form of plasma—a gas of positively charged ions (possibly, bare nuclei) and electrons. When the temperature drops below the ionization potential, atoms become statistically favorable. Atoms (complete with bound electrons) became to dominate over charged particles 380,000 years after the Big Bang—an epoch called recombination, when the expanding Universe cooled enough to allow electrons to become attached to nuclei. Since the Big Bang, which produced no carbon or heavier elements, atomic nuclei have been combined in stars through the process of nuclear fusion to produce more of the element helium, and (via the triple-alpha process) the sequence of elements from carbon up to iron; see stellar nucleosynthesis for details. Isotopes such as lithium-6, as well as some beryllium and boron are generated in space through cosmic ray spallation. This occurs when a high-energy proton strikes an atomic nucleus, causing large numbers of nucleons to be ejected. Elements heavier than iron were produced in supernovae and colliding neutron stars through the r-process, and in AGB stars through the s-process, both of which involve the capture of neutrons by atomic nuclei. Elements such as lead formed largely through the radioactive decay of heavier elements. === Earth === Most of the atoms that make up the Earth
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{
"page_id": 902,
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and its inhabitants were present in their current form in the nebula that collapsed out of a molecular cloud to form the Solar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine the age of the Earth through radiometric dating. Most of the helium in the crust of the Earth (about 99% of the helium from gas wells, as shown by its lower abundance of helium-3) is a product of alpha decay. There are a few trace atoms on Earth that were not present at the beginning (i.e., not "primordial"), nor are results of radioactive decay. Carbon-14 is continuously generated by cosmic rays in the atmosphere. Some atoms on Earth have been artificially generated either deliberately or as by-products of nuclear reactors or explosions. Of the transuranic elements—those with atomic numbers greater than 92—only plutonium and neptunium occur naturally on Earth. Transuranic elements have radioactive lifetimes shorter than the current age of the Earth and thus identifiable quantities of these elements have long since decayed, with the exception of traces of plutonium-244 possibly deposited by cosmic dust. Natural deposits of plutonium and neptunium are produced by neutron capture in uranium ore. The Earth contains approximately 1.33×1050 atoms. Although small numbers of independent atoms of noble gases exist, such as argon, neon, and helium, 99% of the atmosphere is bound in the form of molecules, including carbon dioxide and diatomic oxygen and nitrogen. At the surface of the Earth, an overwhelming majority of atoms combine to form various compounds, including water, salt, silicates, and oxides. Atoms can also combine to create materials that do not consist of discrete molecules, including crystals and liquid or solid metals. This atomic matter forms networked arrangements that lack the particular type of small-scale interrupted order associated
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{
"page_id": 902,
"source": null,
"title": "Atom"
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with molecular matter. === Rare and theoretical forms === ==== Superheavy elements ==== All nuclides with atomic numbers higher than 82 (lead) are known to be radioactive. No nuclide with an atomic number exceeding 92 (uranium) exists on Earth as a primordial nuclide, and heavier elements generally have shorter half-lives. Nevertheless, an "island of stability" encompassing relatively long-lived isotopes of superheavy elements with atomic numbers 110 to 114 might exist. Predictions for the half-life of the most stable nuclide on the island range from a few minutes to millions of years. In any case, superheavy elements (with Z > 104) would not exist due to increasing Coulomb repulsion (which results in spontaneous fission with increasingly short half-lives) in the absence of any stabilizing effects. ==== Exotic matter ==== Each particle of matter has a corresponding antimatter particle with the opposite electrical charge. Thus, the positron is a positively charged antielectron and the antiproton is a negatively charged equivalent of a proton. When a matter and corresponding antimatter particle meet, they annihilate each other. Because of this, along with an imbalance between the number of matter and antimatter particles, the latter are rare in the universe. The first causes of this imbalance are not yet fully understood, although theories of baryogenesis may offer an explanation. As a result, no antimatter atoms have been discovered in nature. In 1996, the antimatter counterpart of the hydrogen atom (antihydrogen) was synthesized at the CERN laboratory in Geneva. Other exotic atoms have been created by replacing one of the protons, neutrons or electrons with other particles that have the same charge. For example, an electron can be replaced by a more massive muon, forming a muonic atom. These types of atoms can be used to test fundamental predictions of physics. == See also == ==
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{
"page_id": 902,
"source": null,
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Notes == == References == == Bibliography == Oliver Manuel (2001). Origin of Elements in the Solar System: Implications of Post-1957 Observations. Springer. ISBN 978-0-306-46562-8. OCLC 228374906. Andrew G. van Melsen (2004) [1952]. From Atomos to Atom: The History of the Concept Atom. Translated by Henry J. Koren. Dover Publications. ISBN 0-486-49584-1. J.P. Millington (1906). John Dalton. J. M. Dent & Co. (London); E. P. Dutton & Co. (New York). Charles H. Holbrow; James N. Lloyd; Joseph C. Amato; Enrique Galvez; M. Elizabeth Parks (2010). Modern Introductory Physics. Springer Science & Business Media. ISBN 978-0-387-79079-4. John Dalton (1808). A New System of Chemical Philosophy vol. 1. John Dalton (1817). A New System of Chemical Philosophy vol. 2. John L. Heilbron (2003). Ernest Rutherford and the Explosion of Atoms. Oxford University Press. ISBN 0-19-512378-6. Jaume Navarro (2012). A History of the Electron: J. J. and G. P. Thomson. Cambridge University Press. ISBN 978-1-107-00522-8. Bernard Pullman (1998). The Atom in the History of Human Thought. Translated by Axel Reisinger. Oxford University Press. ISBN 0-19-511447-7. Jean Perrin (1910) [1909]. Brownian Movement and Molecular Reality. Translated by F. Soddy. Taylor and Francis. Eric R. Scerri (2020). The Periodic Table, Its Story and Its Significance (2nd ed.). New York: Oxford University Press. ISBN 978-0-190-91436-3. == Further reading == == External links == Atoms in Motion – The Feynman Lectures on Physics Sharp, Tim (8 August 2017). "What is an Atom?". Live Science.
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{
"page_id": 902,
"source": null,
"title": "Atom"
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A Bland–Altman plot (difference plot) in analytical chemistry or biomedicine is a method of data plotting used in analyzing the agreement between two different assays. It is identical to a Tukey mean-difference plot, the name by which it is known in other fields, but was popularised in medical statistics by J. Martin Bland and Douglas G. Altman. == Construction == Consider a sample consisting of n {\displaystyle n} observations (for example, objects of unknown volume). Both assays (for example, different methods of volume measurement) are performed on each sample, resulting in 2 n {\displaystyle 2n} data points. Each of the n {\displaystyle n} samples is then represented on the graph by assigning the mean of the two measurements as the x {\displaystyle x} -value, and the difference between the two values as the y {\displaystyle y} -value. The Cartesian coordinates of a given sample S {\displaystyle S} with values of S 1 {\displaystyle S_{1}} and S 2 {\displaystyle S_{2}} determined by the two assays is S ( x , y ) = ( S 1 + S 2 2 , S 1 − S 2 ) . {\displaystyle S(x,y)=\left({\frac {S_{1}+S_{2}}{2}},S_{1}-S_{2}\right).} For comparing the dissimilarities between the two sets of samples independently from their mean values, it is more appropriate to look at the ratio of the pairs of measurements. Log transformation (base 2) of the measurements before the analysis will enable the standard approach to be used; so the plot will be given by the following equation: S ( x , y ) = ( log 2 S 1 + log 2 S 2 2 , log 2 S 1 − log 2 S 2 ) . {\displaystyle S(x,y)=\left({\frac {\log _{2}S_{1}+\log _{2}S_{2}}{2}},\log _{2}S_{1}-\log _{2}S_{2}\right).} This version of the plot is used in MA plot. == Interpretation
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"page_id": 3146632,
"source": null,
"title": "Bland–Altman plot"
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|
== Interpretation of a Bland-Altman plot is contingent on the construction of the plot and data at hand. Variations to the default plot have introduced throughout the years and each should be interpreted accordingly. === Original Construction === The original plot displays a scatter plot of differences between individual data points. The differences should be of the new reference system minus a gold standard. An average of the differences is plotted horizontally with limits of agreement plotted parallel to this mean difference line. The limits of agreement represent a confidence interval for which most of the differences lie between systems. The mean difference represents a general bias between the two systems; a positive mean difference indicates the reference system generally produces larger values relative to the golden standard, and a negative mean difference indicating the reference system generally produces lower values than the verified system. A mean difference closet to 0 indicates agreement between two systems, though the limits of agreement illustrate more nuance. === Limits of Agreement === Since the limits of agreement are by-default contingent on the standard deviation of the data, the distribution of the differences must follow a normal distribution. In the event that the distribution of differences are not normal, limits of agreement not contingent on normal distribution may be used instead. Bland and Altman's follow up paper on the topic explains that percentile of differences are a suitable replacement in such cases. In any case, the limits of agreement more accurately illustrate the agreement between systems as opposed to just the mean difference. A novel reference system is said to be an appropriate substitute for a golden standard system if the limits of agreement are within a predetermined threshold. The threshold depends extensively on the magnitude of the data, the nature of the systems,
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{
"page_id": 3146632,
"source": null,
"title": "Bland–Altman plot"
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|
and the contexts in which they are to be used. The 95% limits of agreement can be unreliable estimates of the population parameters especially for small sample sizes so, when comparing methods or assessing repeatability, it is important to calculate confidence intervals for 95% limits of agreement. This can be done by Bland and Altman's approximate method or by more precise methods. === Visualization Variations === In the case that the differences grow proportionally to the magnitude of the data, then the data is said to have a 'proportional bias'. There are many methods for visualizing the plot and subsequent analysis to accommodate for it. Firstly, a linear regression could illustrate any relevant trends. If the distribution of differences are equal at all points around the regression the data is said to be homoscedastic and the trend is a simple proportional bias. Inversely, if the data has wider spread at different magnitudes of the data, then the differences are said to be heteroscedastic, which has further implications. Statistical tests such as the Breusch–Pagan test or the White test can provide statistical indicators of heteroscedasticity.One typical example of a plot with heteroscedastic data is one whose variation of differences grows proportional to the magnitude of the data, visualized as an expanding 'v' shape. In such cases, it may be suitable to visualize the proportion of data points between systems as opposed to the raw differences. Similarly, the plot of differences could be visualized logarithmically. In either case, the relationship between the two systems illustrates a multiplicative relationship as opposed to linear one. This also indicates that the magnitude of the data correlates with variations of accuracy for the systems. == Application == One primary application of the Bland-Altman plot is to compare two clinical measurements that produce continuous output. It can
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{
"page_id": 3146632,
"source": null,
"title": "Bland–Altman plot"
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|
be used to compare a new reference system, technique, or method with a verified gold standard, but a gold standard does not imply it to be without error. In order for the plot to be used to verify a reference system, a threshold is typically predetermined for which the limits of agreement must fall under. The value for the threshold is contingent on a myriad of contexts in which the systems and data exist within. The ability to verify a reference system lends the plot to a broad applicability and prominence across many fields. Over the years, it has gained prominence in Optometry, nutritional science, radiology, environmental sciences, surgery, medicine, veterinary medicine, engineering, and psychology, to name a few. Many recommendations and scholarly articles have also been published in efforts of polishing the technique, the underlying statistical construction, and validity of the plot. See Analyse-it, MedCalc, NCSS, GraphPad Prism, R, StatsDirect, or JASP for software providing Bland–Altman plots. == See also == MA plot Gardner–Altman plot == Notes == A similar method was proposed in 1981 by Eksborg. This method was based on Deming regression—a method introduced by Adcock in 1878. Bland and Altman's Lancet paper was number 29 in a list of the top 100 most-cited papers of all time with over 23,000 citations. == References ==
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{
"page_id": 3146632,
"source": null,
"title": "Bland–Altman plot"
}
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Aluminium (or aluminum in North American English) is a chemical element; it has symbol Al and atomic number 13. It has a density lower than that of other common metals, about one-third that of steel. Aluminium has a great affinity towards oxygen, forming a protective layer of oxide on the surface when exposed to air. It visually resembles silver, both in its color and in its great ability to reflect light. It is soft, nonmagnetic, and ductile. It has one stable isotope, 27Al, which is highly abundant, making aluminium the 12th-most abundant element in the universe. The radioactivity of 26Al leads to it being used in radiometric dating. Chemically, aluminium is a post-transition metal in the boron group; as is common for the group, aluminium forms compounds primarily in the +3 oxidation state. The aluminium cation Al3+ is small and highly charged; as such, it has more polarizing power, and bonds formed by aluminium have a more covalent character. The strong affinity of aluminium for oxygen leads to the common occurrence of its oxides in nature. Aluminium is found on Earth primarily in rocks in the crust, where it is the third-most abundant element, after oxygen and silicon, rather than in the mantle, and virtually never as the free metal. It is obtained industrially by mining bauxite, a sedimentary rock rich in aluminium minerals. The discovery of aluminium was announced in 1825 by Danish physicist Hans Christian Ørsted. The first industrial production of aluminium was initiated by French chemist Henri Étienne Sainte-Claire Deville in 1856. Aluminium became much more available to the public with the Hall–Héroult process developed independently by French engineer Paul Héroult and American engineer Charles Martin Hall in 1886, and the mass production of aluminium led to its extensive use in industry and everyday life. In the
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
First and Second World Wars, aluminium was a crucial strategic resource for aviation. In 1954, aluminium became the most produced non-ferrous metal, surpassing copper. In the 21st century, most aluminium was consumed in transportation, engineering, construction, and packaging in the United States, Western Europe, and Japan. Despite its prevalence in the environment, no living organism is known to metabolize aluminium salts, but this aluminium is well tolerated by plants and animals. Because of the abundance of these salts, the potential for a biological role for them is of interest, and studies are ongoing. == Physical characteristics == === Isotopes === Of aluminium isotopes, only 27Al is stable. This situation is common for elements with an odd atomic number. It is the only primordial aluminium isotope, i.e. the only one that has existed on Earth in its current form since the formation of the planet. It is therefore a mononuclidic element and its standard atomic weight is virtually the same as that of the isotope. This makes aluminium very useful in nuclear magnetic resonance (NMR), as its single stable isotope has a high NMR sensitivity. The standard atomic weight of aluminium is low in comparison with many other metals. All other isotopes of aluminium are radioactive. The most stable of these is 26Al: while it was present along with stable 27Al in the interstellar medium from which the Solar System formed, having been produced by stellar nucleosynthesis as well, its half-life is only 717,000 years and therefore a detectable amount has not survived since the formation of the planet. However, minute traces of 26Al are produced from argon in the atmosphere by spallation caused by cosmic ray protons. The ratio of 26Al to 10Be has been used for radiodating of geological processes over 105 to 106 year time scales, in particular
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
transport, deposition, sediment storage, burial times, and erosion. Most meteorite scientists believe that the energy released by the decay of 26Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago. The remaining isotopes of aluminium, with mass numbers ranging from 21 to 43, all have half-lives well under an hour. Three metastable states are known, all with half-lives under a minute. === Electron shell === An aluminium atom has 13 electrons, arranged in an electron configuration of [Ne] 3s2 3p1, with three electrons beyond a stable noble gas configuration. Accordingly, the combined first three ionization energies of aluminium are far lower than the fourth ionization energy alone. Such an electron configuration is shared with the other well-characterized members of its group, boron, gallium, indium, and thallium; it is also expected for nihonium. Aluminium can surrender its three outermost electrons in many chemical reactions (see below). The electronegativity of aluminium is 1.61 (Pauling scale). A free aluminium atom has a radius of 143 pm. With the three outermost electrons removed, the radius shrinks to 39 pm for a 4-coordinated atom or 53.5 pm for a 6-coordinated atom. At standard temperature and pressure, aluminium atoms (when not affected by atoms of other elements) form a face-centered cubic crystal system bound by metallic bonding provided by atoms' outermost electrons; hence aluminium (at these conditions) is a metal. This crystal system is shared by many other metals, such as lead and copper; the size of a unit cell of aluminium is comparable to that of those other metals. The system, however, is not shared by the other members of its group: boron has ionization energies too high to allow metallization, thallium has a hexagonal close-packed structure, and gallium and indium have unusual structures that are not
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
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close-packed like those of aluminium and thallium. The few electrons that are available for metallic bonding in aluminium are a probable cause for it being soft with a low melting point and low electrical resistivity. === Bulk === Aluminium metal has an appearance ranging from silvery white to dull gray depending on its surface roughness. Aluminium mirrors provides high reflectivity for light in the ultraviolet, visible (on par with silver), and the far infrared region. Aluminium is also good at reflecting solar radiation, although prolonged exposure to sunlight in air can deteriorate the reflectivity of the metal; this may be prevented if aluminium is anodized, which adds a protective layer of oxide on the surface. The density of aluminium is 2.70 g/cm3, about 1/3 that of steel, much lower than other commonly encountered metals, making aluminium parts easily identifiable through their lightness. Aluminium's low density compared to most other metals arises from the fact that its nuclei are much lighter, while difference in the unit cell size does not compensate for this difference. The only lighter metals are the metals of groups 1 and 2, which apart from beryllium and magnesium are too reactive for structural use (and beryllium is very toxic). Aluminium is not as strong or stiff as steel, but the low density makes up for this in the aerospace industry and for many other applications where light weight and relatively high strength are crucial. Pure aluminium is quite soft and lacking in strength. In most applications various aluminium alloys are used instead because of their higher strength and hardness. The yield strength of pure aluminium is 7–11 MPa, while aluminium alloys have yield strengths ranging from 200 MPa to 600 MPa. Aluminium is ductile, with a percent elongation of 50–70%, and malleable allowing it to be easily
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
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drawn and extruded. It is also easily machined and cast. Aluminium is an excellent thermal and electrical conductor, having around 60% the conductivity of copper, both thermal and electrical, while having only 30% of copper's density. Aluminium is capable of superconductivity, with a superconducting critical temperature of 1.2 kelvin and a critical magnetic field of about 100 gauss (10 milliteslas). It is paramagnetic and thus essentially unaffected by static magnetic fields. The high electrical conductivity, however, means that it is strongly affected by alternating magnetic fields through the induction of eddy currents. == Chemistry == Aluminium combines characteristics of pre- and post-transition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a post-transition metal, with longer-than-expected interatomic distances. Furthermore, as Al3+ is a small and highly charged cation, it is strongly polarizing and bonding in aluminium compounds tends towards covalency; this behavior is similar to that of beryllium (Be2+), and the two display an example of a diagonal relationship. The underlying core under aluminium's valence shell is that of the preceding noble gas, whereas those of its heavier congeners gallium, indium, thallium, and nihonium also include a filled d-subshell and in some cases a filled f-subshell. Hence, the inner electrons of aluminium shield the valence electrons almost completely, unlike those of aluminium's heavier congeners. As such, aluminium is the most electropositive metal in its group, and its hydroxide is in fact more basic than that of gallium. Aluminium also bears minor similarities to the metalloid boron in the same group: AlX3 compounds are valence isoelectronic to BX3 compounds (they have the same valence electronic structure), and both behave as Lewis acids and readily form adducts. Additionally, one of the main motifs of boron chemistry is regular
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class. Aluminium has a high chemical affinity to oxygen, which renders it suitable for use as a reducing agent in the thermite reaction. A fine powder of aluminium reacts explosively on contact with liquid oxygen; under normal conditions, however, aluminium forms a thin oxide layer (~5 nm at room temperature) that protects the metal from further corrosion by oxygen, water, or dilute acid, a process termed passivation. Aluminium is not attacked by oxidizing acids because of its passivation. This allows aluminium to be used to store reagents such as nitric acid, concentrated sulfuric acid, and some organic acids. In hot concentrated hydrochloric acid, aluminium reacts with water with evolution of hydrogen, and in aqueous sodium hydroxide or potassium hydroxide at room temperature to form aluminates—protective passivation under these conditions is negligible. Aqua regia also dissolves aluminium. Aluminium is corroded by dissolved chlorides, such as common sodium chloride. The oxide layer on aluminium is also destroyed by contact with mercury due to amalgamation or with salts of some electropositive metals. As such, the strongest aluminium alloys are less corrosion-resistant due to galvanic reactions with alloyed copper, and aluminium's corrosion resistance is greatly reduced by aqueous salts, particularly in the presence of dissimilar metals. Aluminium reacts with most nonmetals upon heating, forming compounds such as aluminium nitride (AlN), aluminium sulfide (Al2S3), and the aluminium halides (AlX3). It also forms a wide range of intermetallic compounds involving metals from every group on the periodic table. === Inorganic compounds === The vast majority of compounds, including all aluminium-containing minerals and all commercially significant aluminium compounds, feature aluminium in the oxidation state 3+. The coordination number of such compounds varies, but generally Al3+ is either six- or four-coordinate. Almost
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
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|
all compounds of aluminium(III) are colorless. In aqueous solution, Al3+ exists as the hexaaqua cation [Al(H2O)6]3+, which has an approximate Ka of 10−5. Such solutions are acidic as this cation can act as a proton donor and progressively hydrolyze until a precipitate of aluminium hydroxide, Al(OH)3, forms. This is useful for clarification of water, as the precipitate nucleates on suspended particles in the water, hence removing them. Increasing the pH even further leads to the hydroxide dissolving again as aluminate, [Al(H2O)2(OH)4]−, is formed. Aluminium hydroxide forms both salts and aluminates and dissolves in acid and alkali, as well as on fusion with acidic and basic oxides. This behavior of Al(OH)3 is termed amphoterism and is characteristic of weakly basic cations that form insoluble hydroxides and whose hydrated species can also donate their protons. One effect of this is that aluminium salts with weak acids are hydrolyzed in water to the aquated hydroxide and the corresponding nonmetal hydride: for example, aluminium sulfide yields hydrogen sulfide. However, some salts like aluminium carbonate exist in aqueous solution but are unstable as such; and only incomplete hydrolysis takes place for salts with strong acids, such as the halides, nitrate, and sulfate. For similar reasons, anhydrous aluminium salts cannot be made by heating their "hydrates": hydrated aluminium chloride is in fact not AlCl3·6H2O but [Al(H2O)6]Cl3, and the Al–O bonds are so strong that heating is not sufficient to break them and form Al–Cl bonds. This reaction is observed instead: 2[Al(H2O)6]Cl3 heat→ Al2O3 + 6 HCl + 9 H2O All four trihalides are well known. Unlike the structures of the three heavier trihalides, aluminium fluoride (AlF3) features six-coordinate aluminium, which explains its involatility and insolubility as well as high heat of formation. Each aluminium atom is surrounded by six fluorine atoms in a distorted octahedral arrangement,
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
with each fluorine atom being shared between the corners of two octahedra. Such {AlF6} units also exist in complex fluorides such as cryolite, Na3AlF6. AlF3 melts at 1,290 °C (2,354 °F) and is made by reaction of aluminium oxide with hydrogen fluoride gas at 700 °C (1,300 °F). With heavier halides, the coordination numbers are lower. The other trihalides are dimeric or polymeric with tetrahedral four-coordinate aluminium centers. Aluminium trichloride (AlCl3) has a layered polymeric structure below its melting point of 192.4 °C (378 °F) but transforms on melting to Al2Cl6 dimers. At higher temperatures those increasingly dissociate into trigonal planar AlCl3 monomers similar to the structure of BCl3. Aluminium tribromide and aluminium triiodide form Al2X6 dimers in all three phases and hence do not show such significant changes of properties upon phase change. These materials are prepared by treating aluminium with the halogen. The aluminium trihalides form many addition compounds or complexes; their Lewis acidic nature makes them useful as catalysts for the Friedel–Crafts reactions. Aluminium trichloride has major industrial uses involving this reaction, such as in the manufacture of anthraquinones and styrene; it is also often used as the precursor for many other aluminium compounds and as a reagent for converting nonmetal fluorides into the corresponding chlorides (a transhalogenation reaction). Aluminium forms one stable oxide with the chemical formula Al2O3, commonly called alumina. It can be found in nature in the mineral corundum, α-alumina; there is also a γ-alumina phase. Its crystalline form, corundum, is very hard (Mohs hardness 9), has a high melting point of 2,045 °C (3,713 °F), has very low volatility, is chemically inert, and a good electrical insulator, it is often used in abrasives (such as toothpaste), as a refractory material, and in ceramics, as well as being the starting material for the electrolytic
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
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|
production of aluminium. Sapphire and ruby are impure corundum contaminated with trace amounts of other metals. The two main oxide-hydroxides, AlO(OH), are boehmite and diaspore. There are three main trihydroxides: bayerite, gibbsite, and nordstrandite, which differ in their crystalline structure (polymorphs). Many other intermediate and related structures are also known. Most are produced from ores by a variety of wet processes using acid and base. Heating the hydroxides leads to formation of corundum. These materials are of central importance to the production of aluminium and are themselves extremely useful. Some mixed oxide phases are also very useful, such as spinel (MgAl2O4), Na-β-alumina (NaAl11O17), and tricalcium aluminate (Ca3Al2O6, an important mineral phase in Portland cement). The only stable chalcogenides under normal conditions are aluminium sulfide (Al2S3), selenide (Al2Se3), and telluride (Al2Te3). All three are prepared by direct reaction of their elements at about 1,000 °C (1,800 °F) and quickly hydrolyze completely in water to yield aluminium hydroxide and the respective hydrogen chalcogenide. As aluminium is a small atom relative to these chalcogens, these have four-coordinate tetrahedral aluminium with various polymorphs having structures related to wurtzite, with two-thirds of the possible metal sites occupied either in an orderly (α) or random (β) fashion; the sulfide also has a γ form related to γ-alumina, and an unusual high-temperature hexagonal form where half the aluminium atoms have tetrahedral four-coordination and the other half have trigonal bipyramidal five-coordination. Four pnictides – aluminium nitride (AlN), aluminium phosphide (AlP), aluminium arsenide (AlAs), and aluminium antimonide (AlSb) – are known. They are all III-V semiconductors isoelectronic to silicon and germanium, all of which but AlN have the zinc blende structure. All four can be made by high-temperature (and possibly high-pressure) direct reaction of their component elements. Aluminium alloys well with most other metals (with the exception of most
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
alkali metals and group 13 metals) and over 150 intermetallics with other metals are known. Preparation involves heating fixed metals together in certain proportion, followed by gradual cooling and annealing. Bonding in them is predominantly metallic and the crystal structure primarily depends on efficiency of packing. There are few compounds with lower oxidation states. A few aluminium(I) compounds exist: AlF, AlCl, AlBr, and AlI exist in the gaseous phase when the respective trihalide is heated with aluminium, and at cryogenic temperatures. A stable derivative of aluminium monoiodide is the cyclic adduct formed with triethylamine, Al4I4(NEt3)4. Al2O and Al2S also exist but are very unstable. Very simple aluminium(II) compounds are invoked or observed in the reactions of Al metal with oxidants. For example, aluminium monoxide, AlO, has been detected in the gas phase after explosion and in stellar absorption spectra. More thoroughly investigated are compounds of the formula R4Al2 which contain an Al–Al bond and where R is a large organic ligand. === Organoaluminium compounds and related hydrides === A variety of compounds of empirical formula AlR3 and AlR1.5Cl1.5 exist. The aluminium trialkyls and triaryls are reactive, volatile, and colorless liquids or low-melting solids. They catch fire spontaneously in air and react with water, thus necessitating precautions when handling them. They often form dimers, unlike their boron analogues, but this tendency diminishes for branched-chain alkyls (e.g. Pri, Bui, Me3CCH2); for example, triisobutylaluminium exists as an equilibrium mixture of the monomer and dimer. These dimers, such as trimethylaluminium (Al2Me6), usually feature tetrahedral Al centers formed by dimerization with some alkyl group bridging between both aluminium atoms. They are hard acids and react readily with ligands, forming adducts. In industry, they are mostly used in alkene insertion reactions, as discovered by Karl Ziegler, most importantly in "growth reactions" that form long-chain unbranched primary
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
alkenes and alcohols, and in the low-pressure polymerization of ethene and propene. There are also some heterocyclic and cluster organoaluminium compounds involving Al–N bonds. The industrially most important aluminium hydride is lithium aluminium hydride (LiAlH4), which is used as a reducing agent in organic chemistry. It can be produced from lithium hydride and aluminium trichloride. The simplest hydride, aluminium hydride or alane, is not as important. It is a polymer with the formula (AlH3)n, in contrast to the corresponding boron hydride that is a dimer with the formula (BH3)2. == Natural occurrence == === Space === Aluminium's per-particle abundance in the Solar System is 3.15 ppm (parts per million). It is the twelfth most abundant of all elements and third most abundant among the elements that have odd atomic numbers, after hydrogen and nitrogen. The only stable isotope of aluminium, 27Al, is the eighteenth most abundant nucleus in the universe. It is created almost entirely after fusion of carbon in massive stars that will later become Type II supernovas: this fusion creates 26Mg, which upon capturing free protons and neutrons, becomes aluminium. Some smaller quantities of 27Al are created in hydrogen burning shells of evolved stars, where 26Mg can capture free protons. Essentially all aluminium now in existence is 27Al. 26Al was present in the early Solar System with abundance of 0.005% relative to 27Al but its half-life of 728,000 years is too short for any original nuclei to survive; 26Al is therefore extinct. Unlike for 27Al, hydrogen burning is the primary source of 26Al, with the nuclide emerging after a nucleus of 25Mg catches a free proton. However, the trace quantities of 26Al that do exist are the most common gamma ray emitter in the interstellar gas; if the original 26Al were still present, gamma ray maps of the
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
Milky Way would be brighter. === Earth === Overall, the Earth is about 1.59% aluminium by mass (seventh in abundance by mass). Aluminium occurs in greater proportion in the Earth's crust than in the universe at large. This is because aluminium easily forms the oxide and becomes bound into rocks and stays in the Earth's crust, while less reactive metals sink to the core. In the Earth's crust, aluminium is the most abundant metallic element (8.23% by mass) and the third most abundant of all elements (after oxygen and silicon). A large number of silicates in the Earth's crust contain aluminium. In contrast, the Earth's mantle is only 2.38% aluminium by mass. Aluminium also occurs in seawater at a concentration of 0.41 μg/kg. Because of its strong affinity for oxygen, aluminium is almost never found in the elemental state; instead it is found in oxides or silicates. Feldspars, the most common group of minerals in the Earth's crust, are aluminosilicates. Aluminium also occurs in the minerals beryl, cryolite, garnet, spinel, and turquoise. Impurities in Al2O3, such as chromium and iron, yield the gemstones ruby and sapphire, respectively. Native aluminium metal is extremely rare and can only be found as a minor phase in low oxygen fugacity environments, such as the interiors of certain volcanoes. Native aluminium has been reported in cold seeps in the northeastern continental slope of the South China Sea. It is possible that these deposits resulted from bacterial reduction of tetrahydroxoaluminate Al(OH)4−. Although aluminium is a common and widespread element, not all aluminium minerals are economically viable sources of the metal. Almost all metallic aluminium is produced from the ore bauxite (AlOx(OH)3–2x). Bauxite occurs as a weathering product of low iron and silica bedrock in tropical climatic conditions. In 2017, most bauxite was mined in Australia, China,
|
{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
Guinea, and India. == History == The history of aluminium has been shaped by usage of alum. The first written record of alum, made by Greek historian Herodotus, dates back to the 5th century BCE. The ancients are known to have used alum as a dyeing mordant and for city defense. After the Crusades, alum, an indispensable good in the European fabric industry, was a subject of international commerce; it was imported to Europe from the eastern Mediterranean until the mid-15th century. The nature of alum remained unknown. Around 1530, Swiss physician Paracelsus suggested alum was a salt of an earth of alum. In 1595, German doctor and chemist Andreas Libavius experimentally confirmed this. In 1722, German chemist Friedrich Hoffmann announced his belief that the base of alum was a distinct earth. In 1754, German chemist Andreas Sigismund Marggraf synthesized alumina by boiling clay in sulfuric acid and subsequently adding potash. Attempts to produce aluminium date back to 1760. The first successful attempt, however, was completed in 1824 by Danish physicist and chemist Hans Christian Ørsted. He reacted anhydrous aluminium chloride with potassium amalgam, yielding a lump of metal looking similar to tin. He presented his results and demonstrated a sample of the new metal in 1825. In 1827, German chemist Friedrich Wöhler repeated Ørsted's experiments but did not identify any aluminium. (The reason for this inconsistency was only discovered in 1921.) He conducted a similar experiment in the same year by mixing anhydrous aluminium chloride with potassium (the Wöhler process) and produced a powder of aluminium. In 1845, he was able to produce small pieces of the metal and described some physical properties of this metal. For many years thereafter, Wöhler was credited as the discoverer of aluminium. As Wöhler's method could not yield great quantities of aluminium, the
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
|
metal remained rare; its cost exceeded that of gold. The first industrial production of aluminium was established in 1856 by French chemist Henri Etienne Sainte-Claire Deville and companions. Deville had discovered that aluminium trichloride could be reduced by sodium, which was more convenient and less expensive than potassium, which Wöhler had used. Even then, aluminium was still not of great purity and produced aluminium differed in properties by sample. Because of its electricity-conducting capacity, aluminium was used as the cap of the Washington Monument, completed in 1885, the tallest building in the world at the time. The non-corroding metal cap was intended to serve as a lightning rod peak. The first industrial large-scale production method was independently developed in 1886 by French engineer Paul Héroult and American engineer Charles Martin Hall; it is now known as the Hall–Héroult process. The Hall–Héroult process converts alumina into metal. Austrian chemist Carl Joseph Bayer discovered a way of purifying bauxite to yield alumina, now known as the Bayer process, in 1889. Modern production of aluminium is based on the Bayer and Hall–Héroult processes. As large-scale production caused aluminium prices to drop, the metal became widely used in jewelry, eyeglass frames, optical instruments, tableware, and foil, and other everyday items in the 1890s and early 20th century. Aluminium's ability to form hard yet light alloys with other metals provided the metal with many uses at the time. During World War I, major governments demanded large shipments of aluminium for light strong airframes; during World War II, demand by major governments for aviation was even higher. From the early 20th century to 1980, the aluminium industry was characterized by cartelization, as aluminium firms colluded to keep prices high and stable. The first aluminium cartel, the Aluminium Association, was founded in 1901 by the Pittsburgh Reduction
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{
"page_id": 904,
"source": null,
"title": "Aluminium"
}
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