Encyclopedia of Chemistry

Barium chloride is the chemical compound with the formula BaCl₂. It is one of the most important water-soluble salts of barium. Like other barium salts, it is toxic and imparts a yellow-green coloration to a flame. It is also hygroscopic.

Structure and properties

BaCl₂ crystallizes in both the fluorite and lead chloride motifs, both of which accommodate the preference of the large Ba²⁺ ion for coordination numbers greater than six.^[1] In aqueous solution BaCl₂ behaves as a simple salt; In water it is a 1:2 electrolyte and the solution exhibits a neutral pH.

Barium chloride reacts with sulfate ion to produce a thick white precipitate of barium sulfate.

BaCl₂(aq) + SO₄^2- → BaSO₄(s) + 2 Cl^-(aq)

Oxalate effects a similar reaction:

BaCl₂(aq) + Na₂C₂O₄(aq) → BaC₂O₄ (s) + 2 NaCl(aq)

Preparation

Although inexpensively available, barium chloride can be prepared from barium hydroxide or barium carbonate, the latter being found naturally as the mineral witherite. These basic salts react with hydrochloric acid to give hydrated barium chloride. On an industrial scale, it is prepared via a two step process from barite (barium sulfate)^[4]:

BaSO₄ + 4 C → BaS + 4 CO

This first step requires high temperatures.

BaS + CaCl₂ → BaCl₂ + CaS

The second step reqiures fusion of the reactants. The BaCl₂ can then be leached out from the mixture with water.

Uses

As a cheap, soluble salt of barium, barium chloride finds wide application in the laboratory. It is commonly used as a test for sulfate ion (see chemical properties above). In industry, barium chloride is mainly used in the purification of brine solution in caustic chlorine plants and also in the manufacture of heat treatment salts, case hardening of steel, in the manufacture of pigments, and in the manufacture of other barium salts. BaCl₂ is also used in fireworks to give a bright green color. However, its toxicity limits its applicability.

Safety

Barium chloride, along with other water-soluble barium salts, is toxic. Sodium sulfate is a potential antidote because it forms the insoluble solid BaSO₄.

References

1. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.

1. Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford: Butterworth-Heinemann. ISBN 0-7506-3365-4.
2. Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
3. The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.
4. H. Nechamkin, The Chemistry of the Element, McGraw-Hill, New York, 1968.