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Astatine

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 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. 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 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 $300 pm$, 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 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 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 AtO2- has since been determined to be AtO(OH)2-, a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH). The well characterized AtO3- 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 AtO3-, such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion AtO4-; 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 AtBr2- and AtI2- ions, or in a chloride solution, they may produce species like AtCl2- 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, AtOCl2- or AtCl2- may be produced. The polyhalides PdAtI2, CsAtI2, TlAtI2, and PbAtI are known or presumed to have been precipitated. In a plasma ion source mass spectrometer, the ions [AtI]+, [AtBr]+, and [AtCl]+ have been formed by introducing lighter halogen vapors into a 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 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 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 αστατος (astatos) meaning unstable, because of its propensity for radioactive decay, with the ending "-ine", found in the names of the four previously discovered halogens. The name was also chosen to continue the tradition of the four stable halogens, where the name referred to a property of the element.

Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry. Subsequent investigators reported 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 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).

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.

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 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 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. In order to eliminate undesired nuclides, the maximum energy of the particle accelerator is set to a value (optimally 29.17 MeV) above that for the reaction producing astatine-211 (to produce the desired isotope) and below the one producing astatine-210 (to avoid producing other astatine isotopes).

Separation methods
Since astatine is the main product of the synthesis, after its formation it must only be separated from the target and any significant contaminants. Several methods are available, "but they generally follow one of two 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 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 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 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 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.

title: Astatine


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Atom

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 99.94% 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 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.

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". 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 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 (Fe2O2 and Fe2O3).

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 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 are not a form of light but made of negatively-charged particles because they can be deflected by electric and magnetic fields. He measured these particles to be at least a thousand 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 this positive charge was everywhere in the atom, the atom being in the shape of a sphere. Following from this, he imagined 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 along with most of the atom's mass is concentrated in a tiny nucleus at the center of the atom. Only such an intense concentration of positive charge, anchored by its high mass and separated from the negative charge, 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 down into the 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 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. 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 1917 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 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. An element's atomic number, which is defined as the element's position on the periodic table, is also the number of protons it has in its nucleus. 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 Schroedinger used this idea to develop the Schroedinger 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 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.

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 $9.11 kg$, 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 $1.673 kg$. The number of protons in an atom is 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 $1.675 kg$. 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 $2.5 m$—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 equal to $$1.07 \sqrt[3]{A}$$ femtometres, where $$A$$ 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 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 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 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.

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 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.

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 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 $1.66 kg$. 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 $207.977 Da$.

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 $6.022$). This number was chosen so that if an element has an atomic mass of 1 u, a 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 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 ($2$) atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of $2 kg$ 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 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 is measured in units of the reduced Planck constant (ħ), with electrons, protons and neutrons all having spin $6$ 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, 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 $\ell$ 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, 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 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. 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 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 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 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 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 $1/2$ 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 with molecular matter.

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.

license: Creative Commons - Attribution Share-Alike - https://creativecommons.org/licenses/by-sa/4.0/

title: Atom


text:
Arable land

Arable land (from the arabilis, "able to be ploughed") is any land capable of being ploughed and used to grow crops. Alternatively, for the purposes of agricultural statistics, the term often has a more precise definition: "Arable land is the land under temporary agricultural crops (multiple-cropped areas are counted only once), temporary meadows for mowing or pasture, land under market and kitchen gardens and land temporarily fallow (less than five years). The abandoned land resulting from shifting cultivation is not included in this category. Data for 'Arable land' are not meant to indicate the amount of land that is potentially cultivable." A more concise definition appearing in the Eurostat glossary similarly refers to actual rather than potential uses: "land worked (ploughed or tilled) regularly, generally under a system of crop rotation". In Britain, arable land has traditionally been contrasted with pasturable land such as heaths, which could be used for sheep-rearing but not as farmland.

Arable land is vulnerable to land degradation and some types of un-arable land can be enriched to create useful land. Climate change and biodiversity loss, are driving pressure on arable land.

By country

According to the Food and Agriculture Organization of the United Nations, in 2013, the world's arable land amounted to 1.407 billion hectares, out of a total of 4.924 billion hectares of land used for agriculture.

Non-arable land
Agricultural land that is not arable according to the FAO definition above includes:
 * Meadows and pastures – land used as pasture and grazed range, and those natural grasslands and sedge meadows that are used for hay production in some regions.
 * Permanent crop – land that produces crops from woody vegetation, e.g. orchard land, vineyards, coffee plantations, rubber plantations, and land producing nut trees;

Other non-arable land includes land that is not suitable for any agricultural use. Land that is not arable, in the sense of lacking capability or suitability for cultivation for crop production, has one or more limitations – a lack of sufficient freshwater for irrigation, stoniness, steepness, adverse climate, excessive wetness with the impracticality of drainage, excessive salts, or a combination of these, among others. Although such limitations may preclude cultivation, and some will in some cases preclude any agricultural use, large areas unsuitable for cultivation may still be agriculturally productive. For example, United States NRCS statistics indicate that about 59 percent of US non-federal pasture and unforested rangeland is unsuitable for cultivation, yet such land has value for grazing of livestock. In British Columbia, Canada, 41 percent of the provincial Agricultural Land Reserve area is unsuitable for the production of cultivated crops, but is suitable for uncultivated production of forage usable by grazing livestock. Similar examples can be found in many rangeland areas elsewhere.

Land conversion
Land incapable of being cultivated for the production of crops can sometimes be converted to arable land. New arable land makes more food and can reduce starvation. This outcome also makes a country more self-sufficient and politically independent, because food importation is reduced. Making non-arable land arable often involves digging new irrigation canals and new wells, aqueducts, desalination plants, planting trees for shade in the desert, hydroponics, fertilizer, nitrogen fertilizer, pesticides, reverse osmosis water processors, PET film insulation or other insulation against heat and cold, digging ditches and hills for protection against the wind, and installing greenhouses with internal light and heat for protection against the cold outside and to provide light in cloudy areas. Such modifications are often prohibitively expensive. An alternative is the seawater greenhouse, which desalinates water through evaporation and condensation using solar energy as the only energy input. This technology is optimized to grow crops on desert land close to the sea.

The use of artifices does not make the land arable. Rock still remains rock, and shallow – less than 6 feet – turnable soil is still not considered toilable. The use of artifice is an open-air none recycled water hydroponics relationship. The below described circumstances are not in perspective, have limited duration, and have a tendency to accumulate trace materials in soil that either there or elsewhere cause deoxygenation. The use of vast amounts of fertilizer may have unintended consequences for the environment by devastating rivers, waterways, and river endings through the accumulation of non-degradable toxins and nitrogen-bearing molecules that remove oxygen and cause non-aerobic processes to form.

Examples of infertile non-arable land being turned into fertile arable land include:
 * Aran Islands: These islands off the west coast of Ireland (not to be confused with the Isle of Arran in Scotland's Firth of Clyde) were unsuitable for arable farming because they were too rocky. The people covered the islands with a shallow layer of seaweed and sand from the ocean. Today, crops are grown there, even though the islands are still considered non-arable.
 * Israel: The construction of desalination plants along Israel's coast allowed agriculture in some areas that were formerly desert. The desalination plants, which remove the salt from ocean water, have produced a new source of water for farming, drinking, and washing.
 * Slash and burn agriculture uses nutrients in wood ash, but these expire within a few years.
 * Terra preta, fertile tropical soils produced by adding charcoal.

Examples
Examples of fertile arable land being turned into infertile land include:
 * Droughts such as the "Dust Bowl" of the Great Depression in the US turned farmland into desert.
 * Each year, arable land is lost due to desertification and human-induced erosion. Improper irrigation of farmland can wick the sodium, calcium, and magnesium from the soil and water to the surface. This process steadily concentrates salt in the root zone, decreasing productivity for crops that are not salt-tolerant.
 * Rainforest deforestation: The fertile tropical forests are converted into infertile desert land. For example, Madagascar's central highland plateau has become virtually totally barren (about ten percent of the country) as a result of slash-and-burn deforestation, an element of shifting cultivation practiced by many natives.

title: Arable land