caesium ← barium → lanthanum Sr ↑ Ba ↓ Ra 56Ba Periodic table Appearance silvery gray General properties Name, symbol, number barium, Ba, 56 Pronunciation /ˈbɛəriəm/ BAIR-ee-əm Element category alkaline earth metals Group, period, block 2, 6, s Standard atomic weight 137.33g·mol−1 Electron configuration Xe 6s2 Electrons per shell 2, 8, 18, 18, 8, 2 (Image) Physical properties Phase solid Density (near r.t.) 3.51 g·cm−3 Liquid density at m.p. 3.338 g·cm−3 Melting point 1000 K, 727 °C, 1341 °F Boiling point 2170 K, 1897 °C, 3447 °F Heat of fusion 7.12 kJ·mol−1 Heat of vaporization 140.3 kJ·mol−1 Specific heat capacity (25 °C) 28.07 J·mol−1·K−1 Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 911 1038 1185 1388 1686 2170 Atomic properties Oxidation states 2 (strongly basic oxide) Electronegativity 0.89 (Pauling scale) Ionization energies 1st: 502.9 kJ·mol−1 2nd: 965.2 kJ·mol−1 3rd: 3600 kJ·mol−1 Atomic radius 222 pm Covalent radius 215±11 pm Van der Waals radius 268 pm Miscellanea Crystal structure body-centered cubic Magnetic ordering paramagnetic Electrical resistivity (20 °C) 332 nΩ·m Thermal conductivity (300 K) 18.4 W·m−1·K−1 Thermal expansion (25 °C) 20.6 µm·m−1·K−1 Speed of sound (thin rod) (20 °C) 1620 m/s Young's modulus 13 GPa Shear modulus 4.9 GPa Bulk modulus 9.6 GPa Mohs hardness 1.25 CAS registry number 7440-39-3 Most stable isotopes Main article: Isotopes of barium iso NA half-life DM DE (MeV) DP 130Ba 0.106% 130Ba is stable with 74 neutrons 132Ba 0.101% 132Ba is stable with 76 neutrons 133Ba syn 10.51 y ε 0.517 133Cs 134Ba 2.417% 134Ba is stable with 78 neutrons 135Ba 6.592% 135Ba is stable with 79 neutrons 136Ba 7.854% 136Ba is stable with 80 neutrons 137Ba 11.23% 137Ba is stable with 81 neutrons 138Ba 71.7% 138Ba is stable with 82 neutrons v · d · e

Barium ( /ˈbɛəriəm/ BAIR-ee-əm) is a chemical element with the symbol Ba and atomic number 56. It is the fifth element in Group 2, a soft silvery metallic alkaline earth metal. Barium is never found in nature in its pure form due to its reactivity with air. Its oxide is historically known as baryta but it reacts with water and carbon dioxide and is not found as a mineral. The most common naturally occurring minerals are the very insoluble barium sulfate, BaSO4 (barite), and barium carbonate, BaCO3 (witherite). Barium's name originates from Greek barys (βαρύς), meaning "heavy", describing the high density of some common barium-containing ores. Barium has few industrial uses, but the metal has been historically used to scavenge air in vacuum tubes. Barium compounds impart a green color to flames and have been used in fireworks. Barium sulfate is used for its density, insolubility, and X-ray opacity. It is used as an insoluble heavy mud-like paste when drilling oil wells, and in purer form, as an X-ray radiocontrast agent for imaging the human gastrointestinal tract. Soluble barium compounds are poisonous due to release of the soluble barium ion, and have been used as rodenticides. New uses for barium continue to be sought. For example, it is a component of some "high temperature" YBCO superconductors. Contents 1 Characteristics 1.1 Physical properties 1.2 Chemical properties 1.3 Isotopes 1.4 Occurrence 2 Production 3 Compounds 4 History 5 Applications 5.1 Applications of barium compounds 5.1.1 Applications of other barium compounds 6 Precautions 6.1 Non-toxicity of barium sulfate 7 See also 8 References 9 External links // Characteristics Barite Physical properties

Barium is a soft and ductile metal. Its simple compounds are notable for their relatively high (for an alkaline earth element) specific gravity. This is true of the most common barium-bearing mineral, its sulfate barite BaSO4, also called 'heavy spar' due to the high density (4.5 g/cm³). Chemical properties Barium is a highly reducing metal. It reacts exothermically with oxygen at room temperature to form barium oxide and peroxide. The reaction is violent if barium is powdered. It also reacts violently with dilute acids, alcohol and water Ba + 2 H2O → Ba(OH)2 + H2↑ Barium combines with several metals, including aluminium, zinc, lead and tin, forming intermetallic compounds and alloys.1 Isotopes Main article: Isotopes of barium Naturally occurring barium is a mix of seven stable isotopes, the most abundant being 138Ba (71.7 %). There are twenty-two isotopes known, but most of these are highly radioactive and have half-lives in the several millisecond to several day range. The only notable exceptions are 133Ba which has a half-life of 10.51 years, and 137mBa (2.55 minutes).2 133Ba is a standard calibrant for gamma-ray detectors in nuclear physics studies. Occurrence The abundance of barium is 0.0425 % in the Earth's crust and 13 µg/L in sea water. It occurs in the minerals barite (as the sulfate) and witherite (as the carbonate).1 Athough witherite deposits were mined from the 17th century till 19693 in northern England, for example in the Settlingstones Mine near Newbrough,4 today nearly all barium is mined as barite. A rare gem containing barium is known, called benitoite. Large deposits of barite are found in China, Germany, India, Morocco, and in the United States.5 Production Trend in world production of barite

Because barium quickly oxidizes in air, it is difficult to obtain the free metal and it is never found free in nature. The metal is primarily found in, and extracted from, barite. Because barite is so insoluble, it cannot be used directly for the preparation of other barium compounds, or barium metal. Instead, the ore is heated with carbon to reduce it to barium sulfide:6 BaSO4 + 2 C → BaS + 2 CO2 The barium sulfide is then hydrolyzed or treated with acids to form other barium compounds, such as the chloride, nitrate, and carbonate. Barium is commercially produced through the electrolysis of molten barium chloride (BaCl2): (cathode) Ba2+ + 2 e− → Ba (anode) 2 Cl– → Cl2 + 2 e− Barium metal is also obtained by the reduction of barium oxide with finely divided aluminium at temperatures between 1100 and 1200 °C: 4 BaO + 2 Al → BaO·Al2O3 + 3 Ba The barium vapor is cooled and condensed to give the solid metal, which can be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags.1 Compounds Ba2+ is the dominant oxidation state throughout the chemistry of barium. Its properties generally resemble those of other alkaline earth ions such as strontium and calcium. All halides, pseudohalides and chalcogenides are known, usually as colourless solids. The sulfate is famously insoluble. BaO forms a peroxide when heated in air. The oxide is basic and reacts with acids to give salts. Barium reduces oxides, chlorides and sulfides of less active metals. For example: Ba + CdO → BaO + Cd Ba + ZnCl2 → BaCl2 + Zn 3 Ba + Al2S3 → 3 BaS + 2 Al At elevated temperatures, barium combines with nitrogen and hydrogen to produce the nitride Ba3N2 and hydride BaH2, respectively. When heated with nitrogen and carbon, it forms the cyanide: Ba + N2 + 2 C → Ba(CN)2 History

Barium's name originates from Greek βαρύς barys, meaning "heavy", describing the density of some common barium-containing ores. Alchemists in the early Middle Ages knew about some barium minerals. Smooth pebble-like stones of mineral barite found in Bologna, Italy were known as "Bologna stones". The fact that after exposed to light, they would glow for years, attracted witches and alchemists to them.7 Carl Scheele identified barite as containing a new element in 1774, but could not isolate barium. Oxidized barium was at first called barote, by Guyton de Morveau, a name which was changed by Antoine Lavoisier to baryta. Barium was first isolated by electrolysis of molten barium salts in 1808, by Sir Humphry Davy in England.8 Davy, by analogy with calcium named "barium" after baryta, with the "-ium" ending signifying a metallic element.7 The production of pure oxygen in the Brin process was a large scale application of barium peroxide before electrolysis and fractionally distill liquefied air became the dominant ways to produce oxygen. In this process the barium oxide reacts at 500–600°C with air to form barium peroxide which decomposes at above 700°C by releasing oxygen.910 2 BaO + O2 ⇌ 2 BaO2 Applications Amoebiasis as seen in radiograph of barium-filled colon Green barium fireworks The dominating application of elemental barium is as a scavenger or “getter” removing the last traces of oxygen and other gases in electronic vacuum tubes such as television cathode ray tubes.1 An alloy of barium with nickel is used in spark plug wire. Applications of barium compounds Barium sulfate (the mineral barite, BaSO4) is important to the petroleum industry, e.g. as drilling mud, a weighting agent in drilling new oil wells.5 It is also a filler in a variety of products such as rubber. Taking advantage of its opacity to X-rays, the sulfate is used as a radiocontrast agent for X-ray imaging of the digestive system ("barium meals" and "barium enemas").5 Lithopone, a pigment that contains barium sulfate and zinc sulfide, is a permanent white that has good covering power, and does not darken when exposed to sulfides.11 Applications of other barium compounds

Aside from the sulfate, other compounds of barium find only niche applications. Applications are limited by the toxicity of Ba2+ ions (Barium carbonate is a rat poison), which is not a problem for the insoluble BaSO4. Barium oxide is used in a coating for the electrodes of fluorescent lamps, which facilitates the release of electrons. Barium carbonate is also used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass.5 Barium, commonly as barium nitrate, is used to give green colors in fireworks.12 The species responsible for the brilliant green is barium monochloride; in the absence of a source of chlorine a yellow or "apple" green is produced instead. Barium peroxide can be used as a catalyst to start an aluminothermic reaction when welding rail tracks together. It can also be used in green tracer ammunition and as a bleaching agent.13 Barium titanate is a promising electroceramic. Barium fluoride is used for optics in infrared applications, since it is transparent from about 0.15 to 12 micrometres.14 The ratio of Barium:Aluminium (Biogenic Barium) within marine cores is used as a proxy for surface ocean export production in the past Precautions Soluble barium compounds are poisonous. At low doses, barium acts as a muscle stimulant, whereas higher doses affect the nervous system, causing cardiac irregularities, tremors, weakness, anxiety, dyspnea and paralysis. This may be due to its ability to block potassium ion channels which are critical to the proper function of the nervous system.15 However, individual responses to barium salts vary widely, with some being able to handle barium nitrate casually without problems, and others becoming ill from working with it in small quantities. Barium acetate was used by Marie Robards to poison her father in Texas in 1993. She was tried and convicted in 1996.16 Non-toxicity of barium sulfate

Because it is highly insoluble in water as well as stomach acids, barium sulfate can be taken orally. It is eliminated completely from the digestive tract. Unlike other heavy metals, barium does not bioaccumulate.1718 However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.19 See also Book: Barium Wikipedia Books are collections of articles that can be downloaded or ordered in print. References ^ a b c d Robert Kresse, Ulrich Baudis, Paul Jäger, H. Hermann Riechers, Heinz Wagner, Jochen Winkler, Hans Uwe Wolf, "Barium and Barium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2007 Wiley-VCH, Weinheim. doi:10.1002/14356007.a03_325.pub2 ^ David R. Lide, Norman E. Holden (2005). "Section 11, Table of the Isotopes". CRC Handbook of Chemistry and Physics, 85th Edition. Boca Raton, Florida: CRC Press.  ^ Industrial minerals. 1969. p. 28. http://books.google.com/?id=JjEmAQAAIAAJ&dq=Settlingstones+Witherite+Mine&q=Settlingstones+#search_anchor.  ^ "Alston Moor Cumbria, UK". Steetley Minerals. http://www.rock-site.co.uk/EZ/rs/rs/page151.php.  ^ a b c d C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN 0849304814.  ^ "Toxicological Profile for Barium and Barium Compounds. Agency for Toxic Substances and Disease Registry". CDC. 2007.. http://www.atsdr.cdc.gov/toxprofiles/tp24.pdf.  ^ a b Robert E. Krebs (2006). The history and use of our earth's chemical elements: a reference guide. Greenwood Publishing Group. p. 80. ISBN 0313334382. http://books.google.com/?id=yb9xTj72vNAC.  ^ Davy, H. (1808) "Electro-chemical researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia," Philosophical Transactions of the Royal Society of London, vol. 98, pages 333-370. ^ Jensen, William B. (2009). "The Origin of the Brin Process for the Manufacture of Oxygen". Journal of Chemical Education 86: 1266. doi:10.1021/ed086p1266.  ^ Ihde, Aaron John (1984-04-01). The development of modern chemistry. p. 681. ISBN 9780486642352. http://books.google.de/books?id=34KwmkU4LG0C&pg=PA681.  ^ Chris J. Jones, John Thornback (2007). Medicinal applications of coordination chemistry. Royal Society of Chemistry. p. 102. ISBN 0854045961. http://books.google.com/?id=uEJHsZWyO-EC.  ^ Michael S. Russell, Kurt Svrcula (2008). Chemistry of Fireworks. Royal Society of Chemistry. p. 110. ISBN 0854041273. http://books.google.com/?id=yxRyOf8jFeQC.  ^ Brent, G. F.; Harding, M. D. (1995). "Surfactant coatings for the stabilization of barium peroxide and lead dioxide in pyrotechnic compositions". Propellants Explosives Pyrotechnics 20: 300. doi:10.1002/prep.19950200604.  ^ "Crystran Ltd. Optical Component Materials". http://www.crystran.co.uk/barium-fluoride-baf2.htm. Retrieved 2010-12-29.  ^ Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds. McGraw-Hill. pp. 77–78. ISBN 0070494398. http://books.google.com/?id=Xqj-TTzkvTEC&pg=PA243. Retrieved 2009-06-06. ^ "Boyfriend fight preceded Roanoke mom's slaying". http://www.buffalo.edu/news/pdf/October08/DallanMorningNewsEwingSlaying.pdf. Retrieved 2009-06-06.  ^ "Toxicity Profiles, Ecological Risk Assessment". US EPA. http://www.epa.gov/region5/superfund/ecology/html/toxprofiles.htm#ba. Retrieved 2009-06-06.  ^ Moore, J. W. (1991). Inorganic Contaminants of Surface Waters, Research and Monitoring Priorities. New York: Springer-Verlag.  ^ Doig AT (February 1976). "Baritosis: a benign pneumoconiosis". Thorax 31 (1): 30–9. doi:10.1136/thx.31.1.30. PMID 1257935.  External links Wikimedia Commons has media related to: Barium Look up barium in Wiktionary, the free dictionary. WebElements.com – Barium Elementymology & Elements Multidict 3-D Holographic Display Using Strontium Barium Niobate  This article incorporates text from a publication now in the public domain: Chisholm, Hugh, ed (1911). Encyclopædia Britannica (Eleventh ed.). Cambridge University Press.  v · d · e  Barium compounds

BaBr2 | BaCO3 | BaC2O4 | Ba(ClO3)2 | BaCl2 | BaCrO4 | BaF2 | BaFeO4 | BaFe2O4 | BaI2 | Ba(NO3)2 | BaO | BaO2 | Ba(OH)2 | BaS | BaSO4 | BaTiO3 v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties Large version 

BaBr2 | BaCO3 | BaC2O4 | Ba(ClO3)2 | BaCl2 | BaCrO4 | BaF2 | BaFeO4 | BaFe2O4 | BaI2 | Ba(NO3)2 | BaO | BaO2 | Ba(OH)2 | BaS | BaSO4 | BaTiO3 v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties Large version 

BaBr2 | BaCO3 | BaC2O4 | Ba(ClO3)2 | BaCl2 | BaCrO4 | BaF2 | BaFeO4 | BaFe2O4 | BaI2 | Ba(NO3)2 | BaO | BaO2 | Ba(OH)2 | BaS | BaSO4 | BaTiO3 v · d · e Periodic table H   He Li Be   B C N O F Ne Na Mg   Al Si P S Cl Ar K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases Unknown chem. properties Large version