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56 caesiumbariumlanthanum


Name, Symbol, Number barium, Ba, 56
Chemical series alkaline earth metals
Group, Period, Block 2, 6, s
Appearance silvery white
Atomic mass 137.327(7) g/mol
Electron configuration [Xe] 6s2
Electrons per shell 2, 8, 18, 18, 8, 2
Physical properties
Phase solid
Density (near r.t.) 3.51 g/cm³
Liquid density at m.p. 3.338 g/cm³
Melting point 1000 K
(727 °C, 1341 °F)
Boiling point 2170 K
(1897 °C, 3447 °F)
Heat of fusion 7.12 kJ/mol
Heat of vaporization 140.3 kJ/mol
Heat capacity (25 °C) 28.07 J/(mol·K)
Vapor pressure
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 911 1038 1185 1388 1686 2170
Atomic properties
Crystal structure cubic body centered
Oxidation states 2
(strongly basic oxide)
Electronegativity 0.89 (Pauling scale)
Ionization energies 1st: 502.9 kJ/mol
2nd: 965.2 kJ/mol
3rd: 3600 kJ/mol
Atomic radius 215 pm
Atomic radius (calc.) 253 pm
Covalent radius 198 pm
Magnetic ordering paramagnetic
Electrical resistivity (20 °C) 332 nΩ·m
Thermal conductivity (300 K) 18.4 W/(m·K)
Thermal expansion (25 °C) 20.6 µm/(m·K)
Speed of sound (thin rod) (20 °C) 1620 m/s
Speed of sound (thin rod) (r.t.) 13 m/s
Shear modulus 4.9 GPa
Bulk modulus 9.6 GPa
Mohs hardness 1.25
CAS registry number 7440-39-3
Notable isotopes
Main article: Isotopes of barium
iso NA half-life DM DE (MeV) DP
130Ba 0.106% Ba is stable with 74 neutrons
132Ba 0.101% Ba is stable with 76 neutrons
133Ba syn 10.51 y ε 0.517 133Cs
134Ba 2.417% Ba is stable with 78 neutrons
135Ba 6.592% Ba is stable with 79 neutrons
136Ba 7.854% Ba is stable with 80 neutrons
137Ba 11.23% Ba is stable with 81 neutrons
138Ba 71.7% Ba is stable with 82 neutrons

Barium (chemical symbol Ba, atomic number 56) is a soft, silvery chemical element classified as an alkaline earth metal. Given its reactivity with air, it is never found in its pure form in nature. In addition, its oxide reacts with water and carbon dioxide and is not found as a mineral. The most common naturally occurring minerals of barium are barite (barium sulfate, BaSO4) and witherite (barium carbonate, BaCO3).

Barium and its compounds have a variety of applications. For instance, metallic barium is used to remove traces of oxygen from vacuum tubes. Barium sulfate is useful for X-ray diagnostics of the digestive system and as a weighting agent in drilling oil wells. Barium carbonate is used in rat poisons and in the manufacture of glass, porcelain, bricks, and cement. Barium oxide is used for coating cathodes in fluorescent lamps, and the hydroxide, a chemical base, is used to clean up acid spills. The salts of barium (particularly its nitrate, chloride, and chlorate) may be used in fireworks to produce green colors. In industry, barium chloride is used mainly to purify brine solutions in chlorine plants and to manufacture heat-treatment salts, pigments, and other barium salts. It should be noted, however, that barium and its water-soluble compounds are toxic.


Occurrence and production

It is difficult to find barium in its pure, metallic form in nature, as it rapidly becomes oxidized in air. It is primarily found in and extracted from the mineral barite, a crystalline form of barium sulfate (BaSO4).

Barium is commercially produced through the electrolysis of molten barium chloride (BaCl2). Barium ions (Ba2+) migrate to the cathode, where they gain electrons (e) and are converted to metallic barium. At the same time, chloride ions (Cl) migrate to the anode, where they lose electrons and are converted to chlorine gas. The reactions at the electrodes can be written as follows:

At the cathode: Ba2+ + 2e → Ba
At the anode: 2Cl → Cl2 (gas) + 2e

Discovery and etymology

Barium (from the Greek word barys, meaning "heavy") was first identified in 1774 by Carl Scheele and extracted in 1808 by Sir Humphry Davy in England. The oxide was initially called barote, by Guyton de Morveau. Antoine Lavoisier changed the name to baryta, from which "barium" was derived to describe the metal.

Notable characteristics

As a member of the family of alkaline earth metals, barium lies in group two (former group 2A) of the periodic table, between strontium and radium. In addition, it is placed in period six, between cesium and lanthanum.

Barium is chemically similar to calcium but is more reactive. This metal readily oxidizes when exposed to air and is highly reactive with water or alcohol, producing hydrogen gas. Upon burning in air or oxygen, it produces not just barium oxide (BaO) but also barium peroxide.

To store barium in its pure form, protecting it from oxidation by the air, it should be kept under a petroleum-based fluid (such as kerosene) or other suitable oxygen-free liquid that excludes air.

The compounds of barium, particularly the water-soluble ones, are toxic. In addition, as barium is a heavy element, its compounds are notable for their high specific gravity (density). For example, barite, the most common barium-bearing mineral, is also called "heavy spar," based on its high density (4.5 g/cm3).


Naturally occurring barium is a mix of seven stable isotopes. There are 22 known isotopes, but most of them are highly radioactive, with half-lives in the range of several milliseconds to several minutes. The only notable exceptions are 133Ba, with a half-life of 10.51 years, and 137mBa (2.6 minutes).


Some of the important compounds of barium are noted below. Their uses are mentioned under Applications.

Barium carbonate

As noted above, barium carbonate (BaCO3) can be found in nature in the form of the mineral witherite. It is the chief source of barium salts.

Barium carbonate reacts with various acids to form soluble barium salts. For example, it reacts with hydrochloric acid (HCl) to produce barium chloride:

BaCO3(s) + 2 HCl(aq) → BaCl2(aq) + CO2(g) + H2O(l)

The reaction with sulfuric acid, however, is poor, because barium sulfate is highly insoluble.

Barium chloride

Barium chloride (BaCl2) is a highly toxic, ionic, water-soluble salt of barium. It imparts a yellow-green coloration to a flame.

In water solution, it can react with sulfate ions (SO42-) to produce a thick white precipitate of barium sulfate. The reaction is as follows.

BaCl2(aq) + SO42-(aq) → BaSO4(s) + 2Cl-(aq)

Barium hydroxide

Barium hydroxide, also known as baryta, is a strong, corrosive chemical base. It has a white granular or powdery appearance. It may be prepared by dissolving barium oxide (BaO) in water. It crystallizes as the octahydrate, which may be converted to the monohydrate by heating in air or completely dehydrated at 100 °C in a vacuum. The monohydrate is the usual commercial form.

Barium sulfate

Barium sulfate (or barium sulphate) is a white crystalline solid with the formula BaSO4. It is very insoluble in water and other potential solvents. As mentioned above, the mineral barite is composed largely of barium sulfate.


Barium and its compounds have a number of applications:

  • Metallic barium is a "getter" in vacuum tubes, to remove the last traces of oxygen.
  • An alloy of barium with nickel is used in sparkplug wire.
  • Barium compounds, and especially barite (barium sulfate), are extremely important to the petroleum industry. Barite is used as a weighting agent in drilling new oil wells. A weighting agent is a material that adds body to petroleum.
  • Barium sulfate is also a good absorber of X rays and is used in medical diagnostic work to obtain X-ray images of the digestive system (through "barium meals" and "barium enemas"). Although the water-soluble compounds of barium are often highly toxic, the extremely low solubility of barium sulfate protects the patient from absorbing harmful amounts of barium.
  • 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.
  • Barium sulfate is used as a filler in plastics.
  • Barium carbonate is used for the preparation of rat poison and in the manufacture of glass, porcelain, bricks, and cement.
  • Barium oxide is used in coating the cathodes of fluorescent lamps and cathode ray tubes, as it facilitates the release of electrons.
  • As a chemical base, barium hydroxide is used clean up acid spills and make them significantly less harmful. It is also used in the laboratory for the titration of weak acids, and for reactions such as the hydrolysis of esters and nitriles.
  • Under the name baryta, barium hydroxide is used in homeopathic remedies.
  • In the laboratory, a barium chloride solution (in water) is used to test for the presence of sulfate ions. If sulfate is present, a white precipitate of barium sulfate is produced. By a similar reaction, barium chloride can be used to prepare other insoluble salts, such as barium oxalate (BaC2O4).
  • In industry, barium chloride is mainly used in the purification of brine solution in caustic chlorine plants and in the manufacture of heat-treatment salts, pigments, and other barium salts.
  • The salts of barium (particularly its nitrate, chloride, and chlorate) may be used in fireworks to produce green colors.
  • Impure barium sulfide phosphoresces after exposure to light.
  • 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 tracers for bullets.


Barium dust, if inhaled, can accumulate in the lungs, leading to a condition called baritosis. In addition, barium compounds that are soluble in water or acid are extremely 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 the result of its ability to block potassium ion channels, which are critical to the proper functioning of the nervous system.

Barium sulfate can be used in medicine only because it does not dissolve and is eliminated completely from the digestive tract. Unlike other heavy metals, barium does not bioaccumulate (accumulate in the bodies of living systems).[1]


  1. Barium Info Retrieved December 4, 2007.


  • Greenwood, N.N. and A. Earnshaw. Chemistry of the Elements. Oxford, UK: Pergamon Press, 1984.
  • Handbook of Chemistry and Physics, 71st edition. Ann Arbor, Michigan: CRC Press, 1990.
  • Nechamkin, H. The Chemistry of the Element. New York: McGraw-Hill, 1968.

External links

All links retrieved December 18, 2012.


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