Yttrium is a chemical element with symbol Y and atomic number 39. It is a silvery-metallic transition metal chemically similar to the lanthanides and has been classified as a "rare-earth element". Yttrium is always found in combination with lanthanide elements in rare-earth minerals, is never found in nature as a free element. 89Y is the only stable isotope, the only isotope found in the Earth's crust. In 1787, Carl Axel Arrhenius found a new mineral near Ytterby in Sweden and named it ytterbite, after the village. Johan Gadolin discovered yttrium's oxide in Arrhenius' sample in 1789, Anders Gustaf Ekeberg named the new oxide yttria. Elemental yttrium was first isolated in 1828 by Friedrich Wöhler; the most important uses of yttrium are LEDs and phosphors the red phosphors in television set cathode ray tube displays. Yttrium is used in the production of electrodes, electronic filters, superconductors, various medical applications, tracing various materials to enhance their properties. Yttrium has no known biological role.
Exposure to yttrium compounds can cause lung disease in humans. Yttrium is a soft, silver-metallic and crystalline transition metal in group 3; as expected by periodic trends, it is less electronegative than its predecessor in the group and less electronegative than the next member of period 5, zirconium. Yttrium is the first d-block element in the fifth period; the pure element is stable in air in bulk form, due to passivation of a protective oxide film that forms on the surface. This film can reach a thickness of 10 µm; when finely divided, yttrium is unstable in air. Yttrium nitride is formed; the similarities of yttrium to the lanthanides are so strong that the element has been grouped with them as a rare-earth element, is always found in nature together with them in rare-earth minerals. Chemically, yttrium resembles those elements more than its neighbor in the periodic table, if physical properties were plotted against atomic number, it would have an apparent number of 64.5 to 67.5, placing it between the lanthanides gadolinium and erbium.
It also falls in the same range for reaction order, resembling terbium and dysprosium in its chemical reactivity. Yttrium is so close in size to the so-called'yttrium group' of heavy lanthanide ions that in solution, it behaves as if it were one of them. Though the lanthanides are one row farther down the periodic table than yttrium, the similarity in atomic radius may be attributed to the lanthanide contraction. One of the few notable differences between the chemistry of yttrium and that of the lanthanides is that yttrium is exclusively trivalent, whereas about half the lanthanides can have valences other than three; as a trivalent transition metal, yttrium forms various inorganic compounds in the oxidation state of +3, by giving up all three of its valence electrons. A good example is yttrium oxide known as yttria, a six-coordinate white solid. Yttrium forms a water-insoluble fluoride and oxalate, but its bromide, iodide and sulfate are all soluble in water; the Y3+ ion is colorless in solution because of the absence of electrons in the d and f electron shells.
Water reacts with yttrium and its compounds to form Y2O3. Concentrated nitric and hydrofluoric acids do not attack yttrium, but other strong acids do. With halogens, yttrium forms trihalides such as yttrium fluoride, yttrium chloride, yttrium bromide at temperatures above 200 °C. Carbon, selenium and sulfur all form binary compounds with yttrium at elevated temperatures. Organoyttrium chemistry is the study of compounds containing carbon–yttrium bonds. A few of these are known to have yttrium in the oxidation state 0; some trimerization reactions were generated with organoyttrium compounds as catalysts. These syntheses use YCl3 as a starting material, obtained from Y2O3 and concentrated hydrochloric acid and ammonium chloride. Hapticity is a term to describe the coordination of a group of contiguous atoms of a ligand bound to the central atom. Yttrium complexes were the first examples of complexes where carboranyl ligands were bound to a d0-metal center through a η7-hapticity. Vaporization of the graphite intercalation compounds graphite–Y or graphite–Y2O3 leads to the formation of endohedral fullerenes such as Y@C82.
Electron spin resonance studies indicated the formation of 3 − ion pairs. The carbides Y3C, Y2C, YC2 can be hydrolyzed to form hydrocarbons. Yttrium in the Solar System was created through stellar nucleosynthesis by the s-process, but by the r-process; the r-process consists of rapid neutron capture of lighter elements during supernova explosions. The s-process is a slow neutron capture of lighter elements inside pulsating red giant stars. Yttrium isotopes are among the most common products of the nuclear fission of uranium in nuclear explosions and nuclear reactors. In the context of nuclear waste management, the most important isotopes of yttrium
Isotopes of magnesium
Magnesium occurs in three stable isotopes, 24Mg, 25Mg, 26Mg. There are 18 radioisotopes; the longest-lived radioisotope is 28Mg with a half-life of 20.915 hours. The lighter isotopes decay to isotopes of sodium while the heavier isotopes decay to isotopes of aluminium; the shortest-lived is 19Mg with a half-life of 5 picoseconds. Values marked # are not purely derived from experimental data, but at least from systematic trends. Spins with weak assignment arguments are enclosed in parentheses. Uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, which use expanded uncertainties. Isotope masses from: Audi, Georges. P.. "Atomic weights of the elements. Review 2000". Pure and Applied Chemistry. 75: 683–800. Doi:10.1351/pac200375060683. Wieser, Michael E.. "Atomic weights of the elements 2005". Pure and Applied Chemistry. 78: 2051–2066. Doi:10.1351/pac200678112051.
Lay summary. Half-life and isomer data selected from the following sources. See editing notes on this article's talk page. Audi, Georges. "NuDat 2.x database". Brookhaven National Laboratory. Holden, Norman E.. "11. Table of the Isotopes". In Lide, David R. CRC Handbook of Physics. Boca Raton, Florida: CRC Press. ISBN 978-0-8493-0485-9. Magnesium isotopes data from The Berkeley Laboratory Isotopes Project's
Barium is a chemical element with symbol Ba and atomic number 56. It is a soft, silvery alkaline earth metal; because of its high chemical reactivity, barium is never found in nature as a free element. Its hydroxide, known in pre-modern times as baryta, does not occur as a mineral, but can be prepared by heating barium carbonate; the most common occurring minerals of barium are barite and witherite, both insoluble in water. The name barium originates from the alchemical derivative "baryta", from Greek βαρύς, meaning "heavy." Baric is the adjectival form of barium. Barium was identified as a new element in 1774, but not reduced to a metal until 1808 with the advent of electrolysis. Barium has few industrial applications, it was used as a getter for vacuum tubes and in oxide form as the emissive coating on indirectly heated cathodes. It is a component of YBCO and electroceramics, is added to steel and cast iron to reduce the size of carbon grains within the microstructure. Barium compounds are added to fireworks to impart a green color.
Barium sulfate is used as an insoluble additive to oil well drilling fluid, as well as in a purer form, as X-ray radiocontrast agents for imaging the human gastrointestinal tract. The soluble barium ion and soluble compounds are poisonous, have been used as rodenticides. Barium is a silvery-white metal, with a slight golden shade when ultrapure; the silvery-white color of barium metal vanishes upon oxidation in air yielding a dark gray oxide layer. Barium has good electrical conductivity. Ultrapure barium is difficult to prepare, therefore many properties of barium have not been measured yet. At room temperature and pressure, barium has a body-centered cubic structure, with a barium–barium distance of 503 picometers, expanding with heating at a rate of 1.8×10−5/°C. It is a soft metal with a Mohs hardness of 1.25. Its melting temperature of 1,000 K is intermediate between those of the lighter strontium and heavier radium; the density is again intermediate between those of radium. Barium is chemically similar to magnesium and strontium, but more reactive.
It always exhibits the oxidation state of +2, except in a few rare and unstable molecular species that are only characterised in the gas phase such as BaF. Reactions with chalcogens are exothermic. Reactions with other nonmetals, such as carbon, phosphorus and hydrogen, are exothermic and proceed upon heating. Reactions with water and alcohols are exothermic and release hydrogen gas: Ba + 2 ROH → Ba2 + H2↑ Barium reacts with ammonia to form complexes such as Ba6; the metal is attacked by most acids. Sulfuric acid is a notable exception because passivation stops the reaction by forming the insoluble barium sulfate on the surface. Barium combines with several metals, including aluminium, zinc and tin, forming intermetallic phases and alloys. Barium salts are white when solid and colorless when dissolved, barium ions provide no specific coloring, they are denser than the calcium analogs, except for the halides. Barium hydroxide was known to alchemists. Unlike calcium hydroxide, it absorbs little CO2 in aqueous solutions and is therefore insensitive to atmospheric fluctuations.
This property is used in calibrating pH equipment. Volatile barium compounds burn with a green to pale green flame, an efficient test to detect a barium compound; the color results from spectral lines at 455.4, 493.4, 553.6, 611.1 nm. Organobarium compounds are a growing field of knowledge: discovered are dialkylbariums and alkylhalobariums. Barium found in the Earth's crust is a mixture of seven primordial nuclides, barium-130, 132, 134 through 138. Barium-130 undergoes slow radioactive decay to xenon-130 by double beta plus decay, barium-132 theoretically decays to xenon-132, with half-lives a thousand times greater than the age of the Universe; the abundance is ≈ 0.1 %. The radioactivity of these isotopes is so weak. Of the stable isotopes, barium-138 composes 71.7% of all barium. In total, barium has about 40 known isotopes, ranging in mass between 114 and 153; the most stable artificial radioisotope is barium-133 with a half-life of 10.51 years. Five other isotopes have half-lives longer than a day.
Barium has 10 meta states, of which barium-133m1 is the most stable with a half-life of about 39 hours. Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral baryte were found in volcanic rock near Bologna, so were called "Bologna stones." Alchemists were attracted to them. The phosphorescent properties of baryte heated with organics were described by V. Casciorolus in 1602. Carl Scheele determined that baryte contained a new element in 1774, but could not isolate barium, only barium oxide. Johan Gottlieb Gahn isolated barium oxide two year
Lead is a chemical element with symbol Pb and atomic number 82. It is a heavy metal, denser than most common materials. Lead is soft and malleable, has a low melting point; when freshly cut, lead is silvery with a hint of blue. Lead has the highest atomic number of any stable element and three of its isotopes each include a major decay chain of heavier elements. Lead is a unreactive post-transition metal, its weak metallic character is illustrated by its amphoteric nature. Compounds of lead are found in the +2 oxidation state rather than the +4 state common with lighter members of the carbon group. Exceptions are limited to organolead compounds. Like the lighter members of the group, lead tends to bond with itself. Lead is extracted from its ores. Galena, a principal ore of lead bears silver, interest in which helped initiate widespread extraction and use of lead in ancient Rome. Lead production declined after the fall of Rome and did not reach comparable levels until the Industrial Revolution. In 2014, the annual global production of lead was about ten million tonnes, over half of, from recycling.
Lead's high density, low melting point and relative inertness to oxidation make it useful. These properties, combined with its relative abundance and low cost, resulted in its extensive use in construction, batteries and shot, solders, fusible alloys, white paints, leaded gasoline, radiation shielding. In the late 19th century, lead's toxicity was recognized, its use has since been phased out of many applications. However, many countries still allow the sale of products that expose humans to lead, including some types of paints and bullets. Lead is a toxin that accumulates in soft tissues and bones, it acts as a neurotoxin damaging the nervous system and interfering with the function of biological enzymes, causing neurological disorders, such as brain damage and behavioral problems. A lead atom has 82 electrons, arranged in an electron configuration of 4f145d106s26p2; the sum of lead's first and second ionization energies—the total energy required to remove the two 6p electrons—is close to that of tin, lead's upper neighbor in the carbon group.
This is unusual. The similarity of ionization energies is caused by the lanthanide contraction—the decrease in element radii from lanthanum to lutetium, the small radii of the elements from hafnium onwards; this is due to poor shielding of the nucleus by the lanthanide 4f electrons. The sum of the first four ionization energies of lead exceeds that of tin, contrary to what periodic trends would predict. Relativistic effects, which become significant in heavier atoms, contribute to this behavior. One such effect is the inert pair effect: the 6s electrons of lead become reluctant to participate in bonding, making the distance between nearest atoms in crystalline lead unusually long. Lead's lighter carbon group congeners form stable or metastable allotropes with the tetrahedrally coordinated and covalently bonded diamond cubic structure; the energy levels of their outer s- and p-orbitals are close enough to allow mixing into four hybrid sp3 orbitals. In lead, the inert pair effect increases the separation between its s- and p-orbitals, the gap cannot be overcome by the energy that would be released by extra bonds following hybridization.
Rather than having a diamond cubic structure, lead forms metallic bonds in which only the p-electrons are delocalized and shared between the Pb2+ ions. Lead has a face-centered cubic structure like the sized divalent metals calcium and strontium. Pure lead has a silvery appearance with a hint of blue, it tarnishes on contact with moist air and takes on a dull appearance, the hue of which depends on the prevailing conditions. Characteristic properties of lead include high density, malleability and high resistance to corrosion due to passivation. Lead's close-packed face-centered cubic structure and high atomic weight result in a density of 11.34 g/cm3, greater than that of common metals such as iron and zinc. This density is the origin of the idiom to go over like a lead balloon; some rarer metals are denser: tungsten and gold are both at 19.3 g/cm3, osmium—the densest metal known—has a density of 22.59 g/cm3 twice that of lead. Lead is a soft metal with a Mohs hardness of 1.5. It is somewhat ductile.
The bulk modulus of lead—a measure of its ease of compressibility—is 45.8 GPa. In comparison, that of aluminium is 75.2 GPa. Lead's tensile strength, at 12–17 MPa, is low; the melting point of lead—at 327.5 °C —is low compared to most metals. Its boiling point of 1749 °C is the lowest among the carbon group elements; the electrical resistivity of lead at 20 °C is 192 nanoohm-meters an order of magnitude higher than those of other industrial metals. Lead is a superconductor at temperatures lower than 7.19 K.
Table of nuclides
A table of nuclides or chart of nuclides is a two-dimensional graph in which one axis represents the number of neutrons and the other represents the number of protons in an atomic nucleus. Each point plotted on the graph thus represents the nuclide of a real or hypothetical chemical element; this system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements instead of each of their isotopes. A chart or table of nuclides is a simple map to the nuclear, or radioactive, behaviour of nuclides, as it distinguishes the isotopes of an element, it contrasts with a periodic table, which only maps their chemical behavior, since isotopes do not differ chemically to any significant degree, with the exception of hydrogen. Nuclide charts organize nuclides along the X axis by their numbers of neutrons and along the Y axis by their numbers of protons, out to the limits of the neutron and proton drip lines; this representation was first published by Kurt Guggenheimer in 1934 and expanded by Giorgio Fea in 1935, Emilio Segrè in 1945 or G. Seaborg.
In 1958, Walter Seelmann-Eggebert and Gerda Pfennig published the first edition of the Karlsruhe Nuclide Chart. Its 7th edition was made available in 2006. Today, there are several nuclide charts, four of which have a wide distribution: the Karlsruhe Nuclide Chart, the Strasbourg Universal Nuclide Chart, the Chart of the Nuclides from the JAEA and the Nuclide Chart from Knolls Atomic Power Laboratory, it has become a basic tool of the nuclear community. The nuclide table below shows nuclides, including all with half-life of at least one day, they are arranged with increasing atomic numbers from left to right and increasing neutron numbers from top to bottom. Cell color denotes the half-life of each nuclide. In graphical browsers, each nuclide has a tool tip indicating its half-life; each color represents a certain range of length of half-life, the color of the border indicates the half-life of its nuclear isomer state. Some nuclides have multiple nuclear isomers, this table notes the longest one.
Dotted borders mean that a nuclide has a nuclear isomer, their color is represented the same way as for their normal counterparts. Isotopes are nuclides with the same number of protons but differing numbers of neutrons. Isotopes neighbor each other vertically, e.g. carbon-12, carbon-13, carbon-14 or oxygen-15, oxygen-16, oxygen-17. Isotones are nuclides with the same number of neutrons but differing number of protons. Isotones neighbor each other horizontally. Example: carbon-14, nitrogen-15, oxygen-16 in the sample table above. Isobars are nuclides with the same number of nucleons, i.e. mass number, but different numbers of protons and different number of neutrons. Isobars neighbor each other diagonally from lower-left to upper-right. Example: carbon-14, nitrogen-14, oxygen-14 in the sample table above. Isodiaphers are nuclides with the same difference between protons. Like isobars, they at right angles to the isobar lines. Examples: boron-10, carbon-12, nitrogen-14 where N−Z=0. Beyond the neutron drip line along the lower left, nuclides decay by neutron emission.
Beyond the proton drip line along the upper right, nuclides decay by proton emission. Drip lines have only been established for some elements; the island of stability is a hypothetical region of the table of nuclides that contains isotopes far more stable than other transuranic elements. There are no stable nuclides having an equal number of protons and neutrons in their nuclei with atomic number greater than 20 as can be "read" from the chart. Nuclei of greater atomic number require an excess of neutrons for stability; the only stable nuclides having an odd number of protons and an odd number of neutrons are hydrogen-2, lithium-6, boron-10, nitrogen-14 and tantalum-180m. This is because the mass-energy of such atoms is higher than that of their neighbors on the same isobaric chain, so most of them are unstable to beta decay. There are no stable nuclides with mass numbers 5 or 8. There are stable nuclides with all other mass numbers up to 208 with the exceptions of 147 and 151. With the possible exception of the pair tellurium-123 and antimony-123, odd mass numbers are never represented by more than one stable nuclide.
This is because the mass-energy is a convex function of atomic number, so all nuclides on an odd isobaric chain except one have a lower-energy neighbor to which they can decay by beta decay. There are no stable nuclides having atomic number greater than Z=82, although bismuth is stable for all practical human purposes. Elements with atomic numbers from 1 to 82 all have stable isotopes, with the exceptions of technetium and promethium. Interactive Chart of Nuclides app for mobiles: Android or Apple - for PC use The Live Chart of Nuclides - IAEA Another example of a Chart of Nuclides from Korea Data up to Jan 1999 only
In nuclear physics, beta decay is a type of radioactive decay in which a beta ray is emitted from an atomic nucleus. For example, beta decay of a neutron transforms it into a proton by the emission of an electron accompanied by an antineutrino, or conversely a proton is converted into a neutron by the emission of a positron with a neutrino, thus changing the nuclide type. Neither the beta particle nor its associated neutrino exist within the nucleus prior to beta decay, but are created in the decay process. By this process, unstable atoms obtain a more stable ratio of protons to neutrons; the probability of a nuclide decaying due to beta and other forms of decay is determined by its nuclear binding energy. The binding energies of all existing nuclides form what is called the nuclear band or valley of stability. For either electron or positron emission to be energetically possible, the energy release or Q value must be positive. Beta decay is a consequence of the weak force, characterized by lengthy decay times.
Nucleons are composed of up quarks and down quarks, the weak force allows a quark to change type by the exchange of a W boson and the creation of an electron/antineutrino or positron/neutrino pair. For example, a neutron, composed of two down quarks and an up quark, decays to a proton composed of a down quark and two up quarks. Decay times for many nuclides that are subject to beta decay can be thousands of years. Electron capture is sometimes included as a type of beta decay, because the basic nuclear process, mediated by the weak force, is the same. In electron capture, an inner atomic electron is captured by a proton in the nucleus, transforming it into a neutron, an electron neutrino is released; the two types of beta decay are known as beta beta plus. In beta minus decay, a neutron is converted to a proton, the process creates an electron and an electron antineutrino. Β+ decay is known as positron emission. Beta decay conserves a quantum number known as the lepton number, or the number of electrons and their associated neutrinos.
These particles have lepton number +1, while their antiparticles have lepton number −1. Since a proton or neutron has lepton number zero, β+ decay must be accompanied with an electron neutrino, while β− decay must be accompanied by an electron antineutrino. An example of electron emission is the decay of carbon-14 into nitrogen-14 with a half-life of about 5,730 years: 146C → 147N + e− + νeIn this form of decay, the original element becomes a new chemical element in a process known as nuclear transmutation; this new element has an unchanged mass number A, but an atomic number Z, increased by one. As in all nuclear decays, the decaying element is known as the parent nuclide while the resulting element is known as the daughter nuclide. Another example is the decay of hydrogen-3 into helium-3 with a half-life of about 12.3 years: 31H → 32He + e− + νeAn example of positron emission is the decay of magnesium-23 into sodium-23 with a half-life of about 11.3 s: 2312Mg → 2311Na + e+ + νeβ+ decay results in nuclear transmutation, with the resulting element having an atomic number, decreased by one.
The beta spectrum, or distribution of energy values for the beta particles, is continuous. The total energy of the decay process is divided between the electron, the antineutrino, the recoiling nuclide. In the figure to the right, an example of an electron with 0.40 MeV energy from the beta decay of 210Bi is shown. In this example, the total decay energy is 1.16 MeV, so the antineutrino has the remaining energy: 1.16-0.40=0.76 MeV. An electron at the far right of the curve would have the maximum possible kinetic energy, leaving the energy of the neutrino to be only its small rest mass. Radioactivity was discovered in 1896 by Henri Becquerel in uranium, subsequently observed by Marie and Pierre Curie in thorium and in the new elements polonium and radium. In 1899, Ernest Rutherford separated radioactive emissions into two types: alpha and beta, based on penetration of objects and ability to cause ionization. Alpha rays could be stopped by thin sheets of paper or aluminium, whereas beta rays could penetrate several millimetres of aluminium.
In 1900, Paul Villard identified a still more penetrating type of radiation, which Rutherford identified as a fundamentally new type in 1903 and termed gamma rays. Alpha and gamma are the first three letters of the Greek alphabet. In 1900, Becquerel measured the mass-to-charge ratio for beta particles by the method of J. J. Thomson used to identify the electron, he found that m/e for a beta particle is the same as for Thomson's electron, therefore suggested that the beta particle is in fact an electron. In 1901, Rutherford and Frederick Soddy showed that alpha and beta radioactivity involves the transmutation of atoms into atoms of other chemical elements. In 1913, after the products of more radioactive decays were known and Kazimierz Fajans independently proposed their radioactive displacement law, which states that beta emission from one element produces another element one place to the right in the periodic table, while alpha emission produces an element two places to the left; the study of beta decay provided the first physical evidence for the existence of the neutrino.
In both alpha and gamma decay, the resulting particle has a narrow energy distribution, since the particle carries the energy from the diffe
Polonium is a chemical element with symbol Po and atomic number 84. A rare and radioactive metal with no stable isotopes, polonium is chemically similar to selenium and tellurium, though its metallic character resembles that of its horizontal neighbors in the periodic table: thallium and bismuth. Due to the short half-life of all its isotopes, its natural occurrence is limited to tiny traces of the fleeting polonium-210 in uranium ores, as it is the penultimate daughter of natural uranium-238. Though longer-lived isotopes exist, they are much more difficult to produce. Today, polonium is produced in milligram quantities by the neutron irradiation of bismuth. Due to its intense radioactivity, which results in the radiolysis of chemical bonds and radioactive self-heating, its chemistry has been investigated on the trace scale only. Polonium was discovered in 1898 by Marie and Pierre Curie, when it was extracted from the uranium ore pitchblende and identified by its strong radioactivity: it was the first element to be so discovered.
Polonium was named after Marie Curie's homeland of Poland. Polonium has few applications, those are related to its radioactivity: heaters in space probes, antistatic devices, sources of neutrons and alpha particles. Besides being radioactive, polonium is toxic. 210Po is an alpha emitter. A milligram of 210Po emits about as many alpha particles per second as 5 grams of 226Ra. A few curies of 210Po emit a blue glow, caused by ionisation of the surrounding air. About one in 100,000 alpha emissions causes an excitation in the nucleus which results in the emission of a gamma ray with a maximum energy of 803 keV. Polonium is a radioactive element; the alpha form is the only known example of a simple cubic crystal structure in a single atom basis at STP, with an edge length of 335.2 picometers. The structure of polonium has been characterized by X-ray diffraction and electron diffraction.210Po has the ability to become airborne with ease: if a sample is heated in air to 55 °C, 50% of it is vaporized in 45 hours to form diatomic Po2 molecules though the melting point of polonium is 254 °C and its boiling point is 962 °C.
More than one hypothesis exists for. The chemistry of polonium is similar to that of tellurium, although it shows some similarities to its neighbor bismuth due to its metallic character. Polonium dissolves in dilute acids but is only soluble in alkalis. Polonium solutions are first colored in pink by the Po2+ ions, but rapidly become yellow because alpha radiation from polonium ionizes the solvent and converts Po2+ into Po4+; this process is accompanied by bubbling and emission of heat and light by glassware due to the absorbed alpha particles. At pH about 1, polonium ions are hydrolyzed and complexed by acids such as oxalic acid, citric acid, tartaric acid. Polonium has no common compounds, all of its compounds are synthetically created; the most stable class of polonium compounds are polonides, which are prepared by direct reaction of two elements. Na2Po has the antifluorite structure, the polonides of Ca, Ba, Hg, Pb and lanthanides form a NaCl lattice, BePo and CdPo have the wurtzite and MgPo the nickel arsenide structure.
Most polonides decompose upon heating to about 600 °C, except for HgPo that decomposes at ~300 °C and the lanthanide polonides, which do not decompose but melt at temperatures above 1000 °C. For example, PrPo melts at 1250 °C and TmPo at 2200 °C. PbPo is one of the few occurring polonium compounds, as polonium alpha decays to form lead. Polonium hydride is a volatile liquid at room temperature prone to dissociation. Water is the only other known hydrogen chalcogenide, a liquid at room temperature; the two oxides PoO2 and PoO3 are the products of oxidation of polonium. Halides of the structure PoX2, PoX4 and PoF6 are known, they are soluble in the corresponding hydrogen halides, i.e. PoClX in HCl, PoBrX in HBr and PoI4 in HI. Polonium dihalides are formed by direct reaction of the elements or by reduction of PoCl4 with SO2 and with PoBr4 with H2S at room temperature. Tetrahalides can be obtained by reacting polonium dioxide with HCl, HBr or HI. Other polonium compounds include potassium polonite as a polonite, acetate, carbonate, chromate, cyanide and hydroxides, selenate, monosulfide, sulfate and sulfite.
Polonium has 33 known isotopes. They have atomic masses that range from 188 to 220 u. 210Po is the most available and is made via neutron capture by natural bismuth. The longer-lived 209Po and 208Po can be made through the alpha, proton, or deuteron bombardment of lead or bismuth in a cyclotron. Tentatively called "radium F", polonium was discovered by Marie and Pierre Curie in 1898, was named after Marie Curie's native land of Poland. Poland at the time was under Russian and Austro-Hungarian partition, did not exist as an independent country, it was Curie's hope that naming the element after her native land woul