Astatine is a radioactive chemical element with 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. A sample of the pure element has never been assembled, because any macroscopic specimen would be immediately vaporized by the heat of its own radioactivity.

The bulk properties of astatine are not known with any certainty. Many of them have been estimated based on the element's position on the periodic table as a heavier analog of iodine, and a member of the halogens (the group of elements including fluorine, chlorine, bromine, and iodine). Astatine is likely to have a dark or lustrous appearance and may be a semiconductor or possibly a metal; it probably has a higher melting point than that of iodine. Chemically, several anionic species of astatine are known and most of its compounds resemble those of iodine. It also shows some metallic behavior, including being able to form a stable monatomic cation in aqueous solution (unlike the lighter halogens).

The first synthesis of the element was in 1940 by Dale R. Corson, Kenneth Ross MacKenzie, and Emilio G. Segrè at the University of California, Berkeley, who named it from the Greek astatos (ἄστατος), meaning "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 can only be produced synthetically, usually by bombarding bismuth-209 with alpha particles.

Astatine,  85At
Pronunciation/ˈæstətiːn, -tɪn/ (AS-tə-teen, -⁠tin)
Appearanceunknown, probably metallic
Mass number210 (most stable isotope)
Astatine in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson


Atomic number (Z)85
Groupgroup 17 (halogens)
Periodperiod 6
Element category  metalloid, sometimes classified as a nonmetal, or a metal[1][2]
Electron configuration[Xe] 4f14 5d10 6s2 6p5
Electrons per shell
2, 8, 18, 32, 18, 7
Physical properties
Phase at STPsolid
Melting point575 K ​(302 °C, ​576 °F)
Boiling point610 K ​(337 °C, ​639 °F)
Density (near r.t.)(At2) 6.35±0.15 g/cm3 (predicted)[3]
Molar volume(At2) 32.94 cm3/mol (predicted)[3]
Heat of vaporization(At2) 54.39 kJ/mol[4]
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 361 392 429 475 531 607
Atomic properties
Oxidation states−1, +1, +3, +5, +7[5]
ElectronegativityPauling scale: 2.2
Ionization energies
  • 1st: 899.003 kJ/mol[6]
Covalent radius150 pm
Van der Waals radius202 pm
Other properties
Natural occurrencefrom decay
Crystal structureface-centered cubic (fcc)
Face-centered cubic crystal structure for astatine

Thermal conductivity1.7 W/(m·K)
CAS Number7440-68-8
Namingafter Greek astatos (αστατος), meaning "unstable"
DiscoveryDale R. Corson, Kenneth Ross MacKenzie, Emilio Segrè (1940)
Main isotopes of astatine
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct
209At syn 5.41 h β+ 209Po
α 205Bi
210At syn 8.1 h β+ 210Po
α 206Bi
211At syn 7.21 h ε 211Po
α 207Bi


Astatine is an extremely radioactive element; all its isotopes have short 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 one second or less. Of the first 101 elements in the periodic table, only francium is less stable, and all the astatine isotopes more stable than francium are in any case synthetic and do not occur in nature.[7]

The bulk properties of astatine are not known with any certainty.[8] Research is limited by its short half-life, which prevents the creation of weighable quantities.[9] A visible piece of astatine would immediately vaporize itself because of the heat generated by its intense radioactivity.[10] It remains to be seen if, with sufficient cooling, a macroscopic quantity of astatine could be deposited as a thin film.[2] Astatine is usually classified as either a nonmetal or a metalloid;[11][12] metal formation has also been predicted.[2][13]


Most of the physical properties of astatine have been estimated (by interpolation or extrapolation), using theoretically or empirically derived methods.[14] 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).[15][16][17] The melting and boiling points of astatine are also expected to follow the trend seen in the halogen series, increasing with atomic number. On this basis they are estimated to be 575 and 610 K (302 and 337 °C; 575 and 638 °F), respectively.[18] Some experimental evidence suggests astatine may have lower melting and boiling points than those implied by the halogen trend.[19] Astatine sublimes less readily than does iodine, having a lower vapor pressure.[9] 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.[a] 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.[21][22]

The structure of solid astatine is unknown.[23] As an analogue of iodine it may have an orthorhombic crystalline structure composed of diatomic astatine molecules, and be a semiconductor (with a band gap of 0.7 eV).[24] 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.[2] Evidence for (or against) the existence of diatomic astatine (At2) is sparse and inconclusive.[25][26][27][28][29] Some sources state that it does not exist, or at least has never been observed,[30][31] while other sources assert or imply its existence.[19][32][33] Despite this controversy, many properties of diatomic astatine have been predicted;[34] for example, its bond length would be 300±10 pm, dissociation energy 83.7±12.5 kJ/mol,[35] and heat of vaporization (∆Hvap) 54.39 kJ/mol.[4] The latter figure means that astatine may (at least) be metallic in the liquid state on the basis that elements with a heat of vaporization greater than ~42 kJ/mol are metallic when liquid;[36] diatomic iodine, with a value of 41.71 kJ/mol,[37] falls just short of the threshold figure.[b]


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."[24] Many of its apparent chemical properties have been observed using tracer studies on extremely dilute astatine solutions,[33][40] typically less than 10−10 mol·L−1.[41] Some properties – such as anion formation – align with other halogens.[9] Astatine has some metallic characteristics as well, such as plating onto a cathode,[c] coprecipitating with metal sulfides in hydrochloric acid,[43] and forming a stable monatomic cation in aqueous solution.[43][44] It forms complexes with EDTA, a metal chelating agent,[45] 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.[46]

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,[47][48][49][50] That would be consistent with the electronegativity of astatine on the Allred–Rochow scale (1.9) being less than that of hydrogen (2.2).[51][d] 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 is predicted to be reduced by one-third because of spin–orbit interactions.[41]


Less reactive than iodine, astatine is the least reactive of the halogens,[53] although its compounds have been synthesized in microscopic 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.[54][55][e] 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,[10] palladium, silver, thallium, and lead.[58] 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.[59]

The formation of an astatine compound with hydrogen – usually referred to as hydrogen astatide – was noted by the pioneers of astatine chemistry.[60] 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.[9][61]

Astatine is known to bind to boron,[62] carbon, and nitrogen.[63] Various boron cage compounds have been prepared with At–B bonds, these being more stable than At–C bonds.[64] 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.[65] The dipyridine-astatine(I) cation, [At(C5H5N)2]+, forms ionic compounds with perchlorate[63] (a non-coordinating anion[66]) and with nitrate, [At(C5H5N)2]NO3.[63] 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.[63]

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.[9][67] The species previously thought to be AtO
has since been determined to be AtO(OH)
, a hydrolysis product of AtO+ (another such hydrolysis product being AtOOH).[68] The well characterized AtO
anion can be obtained by, for example, the oxidation of astatine with potassium hypochlorite in a solution of potassium hydroxide.[65][69] Preparation of lanthanum triastatate La(AtO3)3, following the oxidation of astatine by a hot Na2S2O8 solution, has been reported.[70] Further oxidation of AtO
, such as by xenon difluoride (in a hot alkaline solution) or periodate (in a neutral or alkaline solution), yields the perastatate ion AtO
; this is only stable in neutral or alkaline solutions.[71] 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.[65][72]

Astatine may form bonds to the other chalcogens; these include S7At+ and At(CSN)
with sulfur, a coordination selenourea compound with selenium, and an astatine–tellurium colloid with tellurium.[73]

Structure of astatine monoiodide, one of the astatine interhalogens and the heaviest known diatomic interhalogen.

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.[56] 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
and AtI
ions,[56] or in a chloride solution, they may produce species like AtCl
or AtBrCl
via equilibrium reactions with the chlorides.[57] 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
or AtCl
may be produced.[56] The polyhalides PdAtI2, CsAtI2, TlAtI2,[74][75][76] and PbAtI[77] 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.[56] 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.[f] 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.[56][71]


Periodic table by Mendeleev (1971), with astatine missing below chlorine, bromine and iodine ("J")
Dmitri Mendeleev's table of 1871, with an empty space at the eka-iodine position

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).[81] Scientists tried to find it in nature; given its extreme rarity, these attempts resulted in several false discoveries.[82]

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.[83][84][85] In 1934, H. G. MacPherson of University of California, Berkeley disproved Allison's method and the validity of his discovery.[86] 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; moreover, astatine is not found in the thorium series, and the true identity of dakin is not known.[87]

In 1936, a team of Romanian physicist Horia Hulubei and French physicist Yvette Cauchois claimed to have discovered element 85 via X-ray analysis. In 1939, they published another paper which supported and extended previous data. In 1944, Hulubei published a summary of data he had obtained up to that time, claiming it was supported by the work of other researchers. He chose the name "dor", presumably from the Romanian for "longing" [for peace], as World War II had started five years earlier. As Hulubei was writing in French, a language which does not accommodate the "ine" suffix, dor would likely have been rendered in English as "dorine", had it been adopted. In 1947, Hulubei's claim was effectively rejected by the Austrian chemist Friedrich Paneth, who would later chair the IUPAC committee responsible for recognition of new elements. Even though Hulubei's samples did contain astatine, his means to detect it were too weak, by current standards, to enable correct identification.[88] 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.[89]

Emilio Segrè, one of the discoverers of the main-group element astatine

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). Karlik and 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).[90][g] 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",[91] but Karlik and Bernert were again unable to reproduce these results.[54]

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.[1] 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 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.[92] 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.[93][94] (Since then, astatine has been determined in a third decay chain, the neptunium series.[95]) 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"[92] coming from the Greek astatos (αστατος) 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.[96]

Corson and his colleagues classified astatine as a metal on the basis of its analytical chemistry.[97] Subsequent investigators reported iodine-like,[98][99] cationic,[100][101] or amphoteric behavior.[102][103] 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."[96]


Alpha decay characteristics for sample astatine isotopes[h]
Half-life[7] Probability
of alpha
207 −13.243 MeV 1.80 h 8.6% 20.9 h
208 −12.491 MeV 1.63 h 0.55% 12.3 d
209 −12.880 MeV 5.41 h 4.1% 5.5 d
210 −11.972 MeV 8.1 h 0.175% 193 d
211 −11.647 MeV 7.21 h 41.8% 17.2 h
212 −8.621 MeV 0.31 s ≈100% 0.31 s
213 −6.579 MeV 125 ns 100% 125 ns
214 −3.380 MeV 558 ns 100% 558 ns
219 10.397 MeV 56 s 97% 58 s
220 14.350 MeV 3.71 min 8% 46.4 min
221[i] 16.810 MeV 2.3 min experimentally
alpha stable

There are 39 known isotopes of astatine, with atomic masses (mass numbers) of 191–229. Theoretical modeling suggests that 37 more isotopes could exist.[104] No stable or long-lived astatine isotope has been observed, nor is one expected to exist.[105]

Astatine's alpha decay energies follow the same trend as for other heavy elements.[105] 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%.[7][j] 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.[105] Even though heavier astatine isotopes release less energy, no long-lived astatine isotope exists, because of the increasing role of beta decay (electron emission).[105] This decay mode is especially important for astatine; as early as 1950 it was postulated that all isotopes of the element undergo beta decay,[106] though nuclear mass measurements reveal that 215At is in fact beta-stable, as it has the lowest mass of all isobars with A = 215.[7] A beta decay mode has been found for all other astatine isotopes except for astatine-213, astatine-214, and astatine-216m.[7] Astatine-210 and lighter isotopes exhibit beta plus decay (positron emission), astatine-216 and heavier isotopes exhibit beta (minus) decay, and astatine-212 decays via both modes, while astatine-211 undergoes electron capture.[7]

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.[7]

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,[k] 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-214m1; its half-life of 265 nanoseconds is shorter than those of all ground states except that of astatine-213.[7][104]

Natural occurrence

Decay Chain(4n+1, Neptunium Series)
Neptunium series, showing the decay products, including astatine-217, formed from neptunium-237

Astatine is the rarest naturally occurring element.[l] The total amount of astatine in the Earth's crust (quoted mass 2.36 × 1025 grams)[107] is estimated to be less than one gram at any given time.[9]

Any astatine present at the formation of the Earth has long since disappeared; the four naturally occurring isotopes (astatine-215, -217, -218 and -219)[108] 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).[109] 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.[110] Astatine-218 was the first astatine isotope discovered in nature.[111] Astatine-219, with a half-life of 56 seconds, is the longest lived of the naturally occurring isotopes.[7]

Isotopes of astatine are sometimes not listed as naturally occurring because of misconceptions[102] that there are no such isotopes,[112] or discrepancies in the literature. Astatine-216 has been counted as a naturally occurring isotope but reports of its observation[113] (which were described as doubtful) have not been confirmed.[114]



Possible reactions after bombarding bismuth-209 with alpha particles
Reaction[m] Energy of alpha particle
+ 4
+ 2 1
26 MeV[54]
+ 4
+ 3 1
40 MeV[54]
+ 4
+ 4 1
60 MeV[115]

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 giga becquerels[116] (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.[116][117][n] Solvent radiolysis due to the cumulative effect of astatine decay[119] 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.[120]

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.[121] The target is kept under a chemically neutral nitrogen atmosphere,[122] and is cooled with water to prevent premature astatine vaporization.[121] In a particle accelerator, such as a cyclotron,[123] 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)[124] 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).[121]

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.[125][o] 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.[116]


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[127] (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.[117][128]


The bismuth (or sometimes bismuth trioxide) target is dissolved in, for example, concentrated nitric or perchloric acid. Astatine is extracted using an organic solvent such as butyl or isopropyl ether, or thiosemicarbazide. 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).[129] Wet methods involve "multiple radioactivity handling steps" and are not well suited for isolating larger quantities of astatine. They can enable the production of astatine in a specific oxidation state and may have greater applicability in experimental radiochemistry.[116]

Uses and precautions

Several 211At-containing molecules and their experimental uses[130]
Agent Applications
[211At]astatine-tellurium colloids Compartmental tumors
6-[211At]astato-2-methyl-1,4-naphtaquinol diphosphate Adenocarcinomas
211At-labeled methylene blue Melanomas
Meta-[211At]astatobenzyl guanidine Neuroendocrine tumors
5-[211At]astato-2'-deoxyuridine Various
211At-labeled biotin conjugates Various pretargeting
211At-labeled octreotide Somatostatin receptor
211At-labeled monoclonal antibodies and fragments Various
211At-labeled bisphosphonates Bone metastases

Newly formed astatine-211 is the subject of ongoing research in nuclear medicine.[130] 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 radiotherapy, since it decays either via emission of an alpha particle (to bismuth-207),[131] 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.[130] 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.[132]

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.[121] 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."[116] Significant morbidity in cell culture models of human cancers has been achieved with from one to ten astatine-211 atoms bound per cell.[133]

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[p] (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.[120][135][q]

Animal studies show that astatine, similarly to iodine, although to a lesser extent, is preferentially 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+.[46] 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.[136] Early research suggested that injection of astatine into female rodents caused morphological changes in breast tissue;[137] 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.[134] Trace amounts of astatine can be handled safely in fume hoods if they are well-aerated; biological uptake of the element must be avoided.[138]

See also


  1. ^ This half-vaporization period grows to 16 hours if it is instead put on a gold or a platinum surface; this may be caused by poorly understood interactions between astatine and these noble metals.[20]
  2. ^ The extrapolated molar refractivity of diatomic astatine is 41.4 cm3, using the method given by Johnson[38] (simple plot of the values for F, Cl, Br and I vs the cube of their covalent radii). This indicates astatine may be a metal in its condensed state, based on the Goldhammer-Herzfeld criterion, which predicts metallic behavior if the ratio of molar refractivity to molar volume is ≥1.[39]
  3. ^ It is also possible that this is sorption on a cathode.[42]
  4. ^ The algorithm used to generate the Allred-Rochow scale fails in the case of hydrogen, providing a value that is close to that of oxygen (3.5). Hydrogen is instead assigned a value of 2.2. Despite this shortcoming, the Allred-Rochow scale has achieved a relatively high degree of acceptance.[52]
  5. ^ Iodine can act as a carrier despite it reacting with astatine in water because these reactions require iodide (I), not (only) I2.[56][57]
  6. ^ An initial attempt to fluoridate astatine using chlorine trifluoride resulted in formation of a product which became stuck to the glass. Chlorine monofluoride, chlorine, and tetrafluorosilane were formed. The authors called the effect "puzzling", admitting they had expected formation of a volatile fluoride.[78] Ten years later, the compound was predicted to be non-volatile, out of line with the other halogens but similar to radon fluoride;[79] by this time, the latter had been shown to be ionic.[80]
  7. ^ In other words, some other substance was undergoing beta decay (to a different end element), not polonium-218.
  8. ^ In the table, under the words "mass excess", the energy equivalents are given rather than the real mass excesses; "mass excess daughter" stands for the energy equivalent of the mass excess sum of the daughter of the isotope and the alpha particle; "alpha decay half-life" refers to the half-life if decay modes other than alpha are omitted.
  9. ^ The value for mass excess of astatine-221 is calculated rather than measured.
  10. ^ This means that, if decay modes other than alpha are omitted, then astatine-210 has an alpha decay half-life of 4,628.6 hours (128.9 days) and astatine-211 has one of only 17.2 hours (0.7 days). Therefore, astatine-211 is very much less stable toward alpha decay than the previous isotope.
  11. ^ "m1" means that this state of the isotope is the next possible one above – with an energy greater than – the ground state. "m2" and similar designations refer to further higher energy states. The number may be dropped if there is only one well-established meta state, such as astatine-216m. Other designation techniques are sometimes used.
  12. ^ Emsley[10] states that this title has been lost to berkelium, "a few atoms of which can be produced in very-highly concentrated uranium-bearing deposits"; however, his assertion is not corroborated by any primary source.
  13. ^ A nuclide is commonly denoted by a symbol of the chemical element this nuclide belongs to, preceded by a non-spaced superscript mass number and a subscript atomic number of the nuclide located directly under the mass number. (Neutrons may be considered as nuclei with the atomic mass of 1 and the atomic charge of 0, with the symbol being n.) With the atomic number omitted, it is also sometimes used as a designation of an isotope of an element in isotope-related chemistry.
  14. ^ See however Nagatsu et al.[118] who encapsulate the bismuth target in a thin aluminium foil and place it in a niobium holder capable of holding molten bismuth.
  15. ^ See also Lavrukhina and Pozdnyakov.[126]
  16. ^ In other words, where carbon's one s atomic orbital and three p orbitals hybridize to give four new orbitals shaped as intermediates between the original s and p orbitals.
  17. ^ "Unfortunately, the conundrum confronting the … field is that commercial supply of 211At awaits the demonstration of clinical efficacy; however, the demonstration of clinical efficacy requires a reliable supply of 211At."[116]


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External links


-ine is a suffix used in chemistry to denote two kinds of substance. The first is a chemically basic and alkaloidal substance. It was proposed by Joseph Louis Gay-Lussac in an editorial accompanying a paper by Friedrich Sertürner describing the isolation of the alkaloid "morphium", which was subsequently renamed to "morphine". Examples include quinine, morphine and guanidine. The second usage is to denote a hydrocarbon of the second degree of unsaturation. Examples include hexine and heptine. With simple hydrocarbons, this usage is identical to the IUPAC suffix -yne.

The suffix is usually pronounced either or depending on the word it appears in and the accent of the speaker. In a few words (for example, quinine and strychnine), the sound is normal in some accents. Gasoline ends with ; glycerine more often with than with .

It is noteworthy also that some elements of the periodic table (namely the halogens, in the Group 17) have this suffix: fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At), ending which was continued in the artificially created tennessine (Ts).

The suffix -in () is etymologically related and overlaps in usage with -ine. Many proteins and lipids have names ending with -in: for example, the enzymes pepsin and trypsin, the hormones insulin and gastrin, and the lipids stearin (stearine) and olein.

Astatine monobromide

Astatine monobromide is an interhalogen compound with the chemical formula AtBr.

Astatine monoiodide

Astatine monoiodide is an interhalogen compound with the chemical formula AtI.

Berta Karlik

Berta Karlik was an Austrian physicist. She worked for the University of Vienna, eventually becoming the first female professor at the institution. While working with Ernst Foyn she published a paper on the radioactivity of seawater.

She discovered that the element 85 astatine is a product of the natural decay processes. The element was first synthesized in 1940 by Dale R. Corson, K. R. MacKenzie, and Emilio Segrè, after several scientists in vain searched for it in radioactive minerals.

Emilio Segrè

Emilio Gino Segrè (1 February 1905 – 22 April 1989) was an Italian-American physicist and Nobel laureate, who discovered the elements technetium and astatine, and the antiproton, a subatomic antiparticle, for which he was awarded the Nobel Prize in Physics in 1959.

From 1943 to 1946 he worked at the Los Alamos National Laboratory as a group leader for the Manhattan Project. He found in April 1944 that Thin Man, the proposed plutonium gun-type nuclear weapon, would not work because of the presence of plutonium-240 impurities.

Born in Tivoli, near Rome, Segrè studied engineering at the University of Rome La Sapienza before taking up physics in 1927. Segrè was appointed assistant professor of physics at the University of Rome in 1932 and worked there until 1936, becoming one of the Via Panisperna boys. From 1936 to 1938 he was director of the Physics Laboratory at the University of Palermo. After a visit to Ernest O. Lawrence's Berkeley Radiation Laboratory, he was sent a molybdenum strip from the laboratory's cyclotron deflector in 1937, which was emitting anomalous forms of radioactivity. After careful chemical and theoretical analysis, Segrè was able to prove that some of the radiation was being produced by a previously unknown element, named technetium, which was the first artificially synthesized chemical element that does not occur in nature.

In 1938, Benito Mussolini's fascist government passed anti-Semitic laws barring Jews from university positions. As a Jew, Segrè was now rendered an indefinite émigré. At the Berkeley Radiation Lab, Lawrence offered him a job as a research assistant. While at Berkeley, Segrè helped discover the element astatine and the isotope plutonium-239, which was later used to make the Fat Man nuclear bomb dropped on Nagasaki.

In 1944, he became a naturalized citizen of the United States. On his return to Berkeley in 1946, he became a professor of physics and of history of science, serving until 1972. Segrè and Owen Chamberlain were co-heads of a research group at the Lawrence Radiation Laboratory that discovered the antiproton, for which the two shared the 1959 Nobel Prize in Physics.

Segrè was also active as a photographer and took many photos documenting events and people in the history of modern science, which were donated to the American Institute of Physics after his death. The American Institute of Physics named its photographic archive of physics history in his honor.


Francium is a chemical element with symbol Fr and atomic number 87. It used to be known as eka-caesium. It is extremely radioactive; its most stable isotope, francium-223 (originally called actinium K after the natural decay chain it appears in), has a half-life of only 22 minutes. It is the second-most electropositive element, behind only caesium, and is the second rarest naturally occurring element (after astatine). The isotopes of francium decay quickly into astatine, radium, and radon. The electronic structure of a francium atom is [Rn] 7s1, and so the element is classed as an alkali metal.

Bulk francium has never been viewed. Because of the general appearance of the other elements in its periodic table column, it is assumed that francium would appear as a highly reactive metal, if enough could be collected together to be viewed as a bulk solid or liquid. Obtaining such a sample is highly improbable, since the extreme heat of decay caused by its short half-life would immediately vaporize any viewable quantity of the element.

Francium was discovered by Marguerite Perey in France (from which the element takes its name) in 1939. It was the last element first discovered in nature, rather than by synthesis. Outside the laboratory, francium is extremely rare, with trace amounts found in uranium and thorium ores, where the isotope francium-223 continually forms and decays. As little as 20–30 g (one ounce) exists at any given time throughout the Earth's crust; the other isotopes (except for francium-221) are entirely synthetic. The largest amount produced in the laboratory was a cluster of more than 300,000 atoms.


The halogens () are a group in the periodic table consisting of five chemically related elements: Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). The artificially created element 117 (Tennessine, Ts) may also be a halogen. In the modern IUPAC nomenclature, this group is known as group 17. The symbol X is often used generically to refer to any halogen.

The name "halogen" means "salt-producing". When halogens react with metals they produce a wide range of salts, including calcium fluoride, sodium chloride (common table salt), silver bromide and potassium iodide.

The group of halogens is the only periodic table group that contains elements in three of the main states of matter at standard temperature and pressure. All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from minerals or salts. The middle halogens, that is chlorine, bromine and iodine, are often used as disinfectants. Organobromides are the most important class of flame retardants. Elemental halogens are dangerous and can be lethally toxic.


Helvetium was the suggested name of chemical element number 85, now known as astatine, given to it by the Swiss chemist Walter Minder. Walter Minder announced the discovery in 1940. He chose the name based on "Helvetia", the Latin name for Switzerland, to honor his country of birth.

In the year 1942 he together with Alice Leigh-Smith announced a second time the discovery of element number 85. This time he proposed the name anglohelvetium to honor also England, the home of Alice Leigh-Smith.

Later it was proven that in fact he had not discovered element 85.

Hydrogen astatide

Hydrogen astatide, also known as astatine hydride, astatane, astidohydrogen or hydroastatic acid, is a chemical compound with the chemical formula HAt, consisting of an astatine atom covalently bonded to a hydrogen atom. It thus is a hydrogen halide.

This chemical compound can dissolve in water to form hydroastatic acid, which exhibits properties very similar to the other four binary acids, and is in fact the strongest among them. However, it is limited in use due to its ready decomposition into elemental hydrogen and astatine, as well as the short half-life of the various isotopes of astatine. Because the atoms have a nearly equal electronegativity, and as the At+ ion has been observed, dissociation could easily result in the hydrogen carrying the negative charge. Thus, a hydrogen astatide sample can undergo the following reaction:

2 HAt → H+ + At− + H− + At+ → H2 + At2This results in elemental hydrogen gas and astatine precipitate. Further, a trend for hydrogen halides, or HX, is that enthalpy of formation becomes less negative, i.e., decreases in magnitude but increases in absolute terms, as the halide becomes larger. Whereas hydroiodic acid solutions are stable, the hydronium-astatide solution is clearly less stable than the water-hydrogen-astatine system. Finally, radiolysis from astatine nuclei could sever the H-At bonds.

Additionally, astatine has no stable isotopes. The most stable is astatine-210, which has a half-life of approximately 8.1 hours, making its chemical compounds especially difficult to work with, as the astatine will quickly decay into other elements.

Hydrogen halide

Hydrogen halides are diatomic inorganic compounds with the formula HX where X is one of the halogens: fluorine, chlorine, bromine, iodine, or astatine. Hydrogen halides are gases that dissolve in water to give acids which are commonly known as hydrohalic acids.


An interhalogen compound is a molecule which contains two or more different halogen atoms (fluorine, chlorine, bromine, iodine, or astatine) and no atoms of elements from any other group.

Most interhalogen compounds known are binary (composed of only two distinct elements). Their formulae are generally XYn, where n = 1, 3, 5 or 7, and X is the less electronegative of the two halogens. They are all prone to hydrolysis, and ionize to give rise to polyhalogen ions. Those formed with astatine have a very short half-life due to astatine being intensely radioactive.

No interhalogen compounds containing three or more different halogens are definitely known, although a few books claim that IFCl2 and IF2Cl have been obtained, and theoretical studies seem to indicate that some compounds in the series BrClFn are barely stable.

Isotopes of astatine

Astatine (85At) has 39 known isotopes, all of which are radioactive; the range of their mass numbers is from 191 to 229. There also exist 23 metastable excited states. The longest-lived isotope is 210At, which has a half-life of 8.1 hours; the longest-lived isotope existing in naturally occurring decay chains is 219At with a half-life of 56 seconds.

Major actinide

Major actinides is a term used in the nuclear power industry that refers to the plutonium and uranium present in used nuclear fuel, as opposed to the minor actinides neptunium, americium, curium, berkelium, and californium.


A metalloid is a type of chemical element which has properties in between, or that are a mixture of, those of metals and nonmetals. There is neither a standard definition of a metalloid nor complete agreement on the elements appropriately classified as such. Despite the lack of specificity, the term remains in use in the literature of chemistry.

The six commonly recognised metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Five elements are less frequently so classified: carbon, aluminium, selenium, polonium, and astatine. On a standard periodic table, all eleven elements are located in a diagonal region of the p-block extending from boron at the upper left to astatine at lower right. Some periodic tables include a dividing line between metals and nonmetals and the metalloids may be found close to this line.

Typical metalloids have a metallic appearance, but they are brittle and only fair conductors of electricity. Chemically, they behave mostly as nonmetals. They can form alloys with metals. Most of their other physical properties and chemical properties are intermediate in nature. Metalloids are usually too brittle to have any structural uses. They and their compounds are used in alloys, biological agents, catalysts, flame retardants, glasses, optical storage and optoelectronics, pyrotechnics, semiconductors, and electronics.

The electrical properties of silicon and germanium enabled the establishment of the semiconductor industry in the 1950s and the development of solid-state electronics from the early 1960s.The term metalloid originally referred to nonmetals. Its more recent meaning, as a category of elements with intermediate or hybrid properties, became widespread in 1940–1960. Metalloids are sometimes called semimetals, a practice that has been discouraged, as the term semimetal has a different meaning in physics than in chemistry. In physics, it specifically refers to the electronic band structure of a substance.


In chemistry, a nonmetal (or non-metal) is a chemical element that mostly lacks the characteristics of a metal. Physically, a nonmetal tends to have a relatively low melting point, boiling point, and density. A nonmetal is typically brittle when solid and usually has poor thermal conductivity and electrical conductivity. Chemically, nonmetals tend to have relatively high ionization energy, electron affinity, and electronegativity. They gain or share electrons when they react with other elements and chemical compounds. Seventeen elements are generally classified as nonmetals: most are gases (hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon and radon); one is a liquid (bromine); and a few are solids (carbon, phosphorus, sulfur, selenium, and iodine). Metalloids such as boron, silicon, and germanium are sometimes counted as nonmetals.

The nonmetals are divided into two categories reflecting their relative propensity to form chemical compounds: reactive nonmetals and noble gases. The reactive nonmetals vary in their nonmetallic character. The less electronegative of them, such as carbon and sulfur, mostly have weak to moderately strong nonmetallic properties and tend to form covalent compounds with metals. The more electronegative of the reactive nonmetals, such as oxygen and fluorine, are characterised by stronger nonmetallic properties and a tendency to form predominantly ionic compounds with metals. The noble gases are distinguished by their great reluctance to form compounds with other elements.

The distinction between categories is not absolute. Boundary overlaps, including with the metalloids, occur as outlying elements in each category show or begin to show less-distinct, hybrid-like, or atypical properties.

Although five times more elements are metals than nonmetals, two of the nonmetals—hydrogen and helium—make up over 99 percent of the observable universe. Another nonmetal, oxygen, makes up almost half of the Earth's crust, oceans, and atmosphere. Living organisms are composed almost entirely of nonmetals: hydrogen, oxygen, carbon, and nitrogen. Nonmetals form many more compounds than metals.

Period 6 element

A period 6 element is one of the chemical elements in the sixth row (or period) of the periodic table of the elements, including the lanthanides. The periodic table is laid out in rows to illustrate recurring (periodic) trends in the chemical behaviour of the elements as their atomic number increases: a new row is begun when chemical behaviour begins to repeat, meaning that elements with similar behaviour fall into the same vertical columns. The sixth period contains 32 elements, tied for the most with period 7, beginning with caesium and ending with radon. Lead is currently the last stable element; all subsequent elements are radioactive. For bismuth, however, its only primordial isotope, 209Bi, has a half-life of more than 1019 years, over a billion times longer than the current age of the universe. As a rule, period 6 elements fill their 6s shells first, then their 4f, 5d, and 6p shells, in that order; however, there are exceptions, such as gold.

Periodic table (crystal structure)

For elements that are solid at standard temperature and pressure the table gives the crystalline structure of the most thermodynamically stable form(s) in those conditions. In all other cases the structure given is for the element at its melting point. Data is presented only for the first 114 elements as well as the 118th (hydrogen through flerovium and oganesson), and predictions are given for elements that have never been produced in bulk (astatine, francium, and elements 100–114 and 118).

Post-transition metal

Post-transition metals are a set of metallic elements in the periodic table located between the transition metals to their left, and the metalloids to their right. Depending on where these adjacent groups are judged to begin and end, there are at least five competing proposals for which elements to include: the three most common contain six, ten and thirteen elements, respectively (see image). All proposals include gallium, indium, tin, thallium, lead, and bismuth.

Physically, post-transition metals are soft (or brittle), have poor mechanical strength, and have melting points lower than those of the transition metals. Being close to the metal-nonmetal border, their crystalline structures tend to show covalent or directional bonding effects, having generally greater complexity or fewer nearest neighbours than other metallic elements.

Chemically, they are characterised—to varying degrees—by covalent bonding tendencies, acid-base amphoterism and the formation of anionic species such as aluminates, stannates, and bismuthates (in the case of aluminium, tin, and bismuth, respectively). They can also form Zintl phases (half-metallic compounds formed between highly electropositive metals and moderately electronegative metals or metalloids).

The name is universally used, but not officially sanctioned by any organization such as the IUPAC. The origin of the term is unclear: one early use was in 1940 in a chemistry text. Alternate names for this group are B-subgroup metals, other metals, and p-block metals; and at least eleven other labels.


Tennessine is a synthetic chemical element with symbol Ts and atomic number 117. It is the second-heaviest known element and the penultimate element of the 7th period of the periodic table.

The discovery of tennessine was officially announced in Dubna, Russia, by a Russian–American collaboration in April 2010, which makes it the most recently discovered element as of 2019. One of its daughter isotopes was created directly in 2011, partially confirming the results of the experiment. The experiment itself was repeated successfully by the same collaboration in 2012 and by a joint German–American team in May 2014. In December 2015, the Joint Working Party of the International Union of Pure and Applied Chemistry (IUPAC) and the International Union of Pure and Applied Physics, which evaluates claims of discovery of new elements, recognized the element and assigned the priority to the Russian–American team. In June 2016, the IUPAC published a declaration stating that the discoverers had suggested the name tennessine after Tennessee, United States. In November 2016, they officially adopted the name "tennessine".

Tennessine may be located in the "island of stability", a concept that explains why some superheavy elements are more stable compared to an overall trend of decreasing stability for elements beyond bismuth on the periodic table. The synthesized tennessine atoms have lasted tens and hundreds of milliseconds. In the periodic table, tennessine is expected to be a member of group 17, all other members of which are halogens. Some of its properties may significantly differ from those of the halogens due to relativistic effects. As a result, tennessine is expected to be a volatile metal that neither forms anions nor achieves high oxidation states. A few key properties, such as its melting and boiling points and its first ionization energy, are nevertheless expected to follow the periodic trends of the halogens.

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