Barium chloride

From Wikipedia, the free encyclopedia

Categories: Chlorides | Metal halides | Barium compounds | Inorganic compounds | Pyrotechnic colorants

Barium chloride

Identifiers
CAS number	[10361-37-2]
Properties
Molecular formula	Ba Cl₂ (anhydrous) BaCl₂·2H₂O
Molar mass	208.2324 g/mol
Appearance	White solid
Density	3.856 g/cm³, solid
Melting point	962 °C
Boiling point	1560 °C
Solubility in water	37.5 g/100 ml (26°C)
Structure
Crystal structure	monoclinic or orthorhombic
Coordination geometry	7-9
Thermochemistry
Std enthalpy of formation Δ_fH^o₂₉₈	−858.56 kJ/mol
Hazards
MSDS	External MSDS
EU classification	Toxic (T)
R-phrases	R20, R25
S-phrases	(S1/2), S45
Flash point	Non-flammable
Related compounds
Other anions	Barium fluoride Barium bromide Barium iodide
Other cations	Calcium chloride Strontium chloride Lead chloride
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Barium chloride is the ionic 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.

1 Structure and properties
2 Preparation
3 Uses
4 Safety
5 References
6 External links

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. Barium Chloride is also used (with Hydrochloric acid) as a test for sulfates. When these two chemicals are mixed with a sulfate salt, a white precipitate forms, which is barium sulfate.

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

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

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