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Barium is a chemical element with the symbol Ba and atomic number 56. It is the fifth element in Group 2. Barium is a soft silvery metallic alkaline earth metal. It 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.

Metallic barium has few industrial uses, but 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 heaviness, 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 found: it is an essential ingredient in "high temperature" YBCO superconductors.



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).


Barium 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 (g)

At elevated temperatures, barium combines with chlorine, nitrogen and hydrogen to produce BaCl2, Ba3N2 and BaH2, respectively. Barium reduces oxides, chlorides and sulfides of less reactive metals. For example:

Ba + CdO → BaO + Cd
Ba + ZnCl2 → BaCl2 + Zn
3 Ba + Al2S3 → 3 BaS + 2 Al

When heated with nitrogen and carbon, it forms the cyanide:

Ba + N2 + 2 C → Ba(CN)2

Barium combines with several metals, including aluminium, zinc, lead and tin, forming intermetallic compounds and alloys.[1]

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]


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.[3]

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. Davy, by analogy with calcium named "barium" after baryta, with the "-ium" ending signifying a metallic element.[3]

Occurrence and production
Trend in world production of barite

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] A rare gem containing barium is known, called benitoite. Large deposits of barite are found in China, Germany, India, Morocco, and in the US.[4]

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:[5]

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 (g) + 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 (g)

The barium vapor is cooled by means of a water jacket and condensed into the solid metal. The solid block may be cast into rods or extruded into wires. Being a flammable solid, it is packaged under argon in steel containers or plastic bags.[1]


The most important use of elemental barium is as a scavenger removing last traces of oxygen and other gases in television and other electronic tubes. Besides, an isotope of barium, 133Ba, is routinely used as a standard source in the calibration of gamma-ray detectors in nuclear physics studies.[1]

Barium is an important component of YBCO superconductors. An alloy of barium with nickel is used in spark plug wire. Barium oxide is used in a coating for the electrodes of fluorescent lamps, which facilitates the release of electrons.

Barium compounds, and especially barite (BaSO4), are extremely important to the petroleum industry.

* Barite is used in rubber production and in drilling mud, a weighting agent in drilling new oil wells.[4]
* Barium sulfate is used as a radiocontrast agent for X-ray imaging of the digestive system ("barium meals" and "barium enemas").[4] 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.[6]
* Barium carbonate is a rat poison and can also be used in making bricks. Unlike the sulfate, the carbonate dissolves in stomach acid, allowing it to be poisonous. Barium carbonate is also used in glassmaking. Being a heavy element, barium increases the refractive index and luster of the glass.[4]
* Barium, commonly as barium nitrate, is used to give green colors in fireworks.[7] 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.[8]
* Barium titanate was proposed in 2007 to be used in next generation battery technology for electric cars.[9]
* Barium fluoride is used for optics in infrared applications, since it is transparent from about 0.15 to 12 microns.[10]


Barium powder is pyrophoric: it can explode in contact with air or oxidizing gases. It is likely to explode when combined with halogenated hydrocarbon solvents. It reacts violently with water. Oxidation occurs very easily and metallic barium should be kept under a petroleum-based fluid (such as kerosene) or other suitable oxygen-free liquids that exclude air.

All water or acid soluble barium compounds are poisonous. At low doses, barium acts as a muscle stimulant, while 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.[1] 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.[11]

Barium sulfate can be taken orally because it is highly insoluble in water, and is eliminated completely from the digestive tract.[1] Unlike other heavy metals, barium does not bioaccumulate.[12][13] However, inhaled dust containing barium compounds can accumulate in the lungs, causing a benign condition called baritosis.[14]

See also


1. ^ a b c d e f Patnaik, Pradyot (2003). Handbook of Inorganic Chemical Compounds. McGraw-Hill. pp. 77–78. ISBN 0070494398. http://books.google.com/books?id=Xqj-TTzkvTEC&pg=PA243. Retrieved 2009-06-06.
2. ^ 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.
3. ^ 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/books?id=yb9xTj72vNAC&.
4. ^ a b c d C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN 0849304814.
5. ^ "Toxicological Profile for Barium and Barium Compounds. Agency for Toxic Substances and Disease Registry". CDC. 2007.. http://www.atsdr.cdc.gov/toxprofiles/tp24.pdf.
6. ^ Chris J. Jones, John Thornback (2007). Medicinal applications of coordination chemistry. Royal Society of Chemistry. p. 102. ISBN 0854045961. http://books.google.co.jp/books?id=uEJHsZWyO-EC&.
7. ^ Michael S. Russell, Kurt Svrcula (2008). Chemistry of Fireworks. Royal Society of Chemistry. p. 110. ISBN 0854041273. http://books.google.co.jp/books?id=yxRyOf8jFeQC&.
8. ^ Brent, G. F. (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.
9. ^ "Battery Breakthrough?". http://www.technologyreview.com/Biztech/18086/. Retrieved 2009-06-06.
10. ^ "Crystran Ltd. Optical Component Materials". http://www.crystran.co.uk/barium-fluoride-baf2.htm. Retrieved 2010-12-29.
11. ^ "Boyfriend fight preceded Roanoke mom's slaying". http://www.buffalo.edu/news/pdf/October08/DallanMorningNewsEwingSlaying.pdf. Retrieved 2009-06-06.
12. ^ "Toxicity Profiles, Ecological Risk Assessment". http://www.epa.gov/region5/superfund/ecology/html/toxprofiles.htm#ba. Retrieved 2009-06-06.
13. ^ Moore, J. W. (1991). Inorganic Contaminants of Surface Waters, Research and Monitoring Priorities. New York: Springer-Verlag.
14. ^ Doig AT (February 1976). "Baritosis: a benign pneumoconiosis". Thorax 31 (1): 30–9. doi:10.1136/thx.31.1.30. PMID 1257935.

External links

* WebElements.com – Barium
* Elementymology & Elements Multidict
* 3-D Holographic Display Using Strontium Barium Niobate