Barium chloride

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Barium chloride
Cotunnite structure.png
Barium chloride.jpg
Other names
Barium muriate
Muryate of Barytes[1]
Barium dichloride
3D model (Jmol)
ECHA InfoCard 100.030.704
EC Number 233-788-1
RTECS number CQ8750000 (anhydrous)
CQ8751000 (dihydrate)
Molar mass 208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)
Appearance White solid
Density 3.856 g/cm3 (anhydrous)
3.0979 g/cm3 (dihydrate)
Melting point 962 °C (1,764 °F; 1,235 K) (960 °C, dihydrate)
Boiling point 1,560 °C (2,840 °F; 1,830 K)
31.2 g/100 mL (0 °C)
35.8 g/100 mL (20 °C)
59.4 g/100 mL (100 °C)
Solubility soluble in methanol, insoluble in ethanol, ethyl acetate[2]
-72.6·10−6 cm3/mol
orthogonal (anhydrous)
monoclinic (dihydrate)
−858.56 kJ/mol
Safety data sheet See: data page
Toxic (T)
Harmful (Xn)
R-phrases R20, R25
S-phrases (S1/2), S45
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
78 mg/kg (rat, oral)
50 mg/kg (guinea pig, oral)[4]
112 mg Ba/kg (rabbit, oral)
59 mg Ba/kg (dog, oral)
46 mg Ba/kg (mouse, oral)[4]
US health exposure limits (NIOSH):
PEL (Permissible)
TWA 0.5 mg/m3[3]
REL (Recommended)
TWA 0.5 mg/m3[3]
IDLH (Immediate danger)
50 mg/m3[3]
Related compounds
Other anions
Barium fluoride
Barium bromide
Barium iodide
Other cations
Beryllium chloride
Magnesium chloride
Calcium chloride
Strontium chloride
Radium chloride
Lead chloride
Supplementary data page
Refractive index (n),
Dielectric constantr), etc.
Phase behaviour
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Barium chloride is the inorganic compound with the formula BaCl2. It is one of the most common 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[edit]

BaCl2 crystallizes in two forms (polymorphs). One form has the cubic fluorite (CaF2) structure and the other the orthorhombic cotunnite (PbCl2) structure. Both polymorphs accommodate the preference of the large Ba2+ ion for coordination numbers greater than six.[5] The coordination of Ba2+ is 8 in the fluorite structure[6] and 9 in the cotunnite structure.[7] When cotunnite-structure BaCl2 is subjected to pressures of 7–10 GPa, it transforms to a third structure, a monoclinic post-cotunnite phase. The coordination number of Ba2+ increases from 9 to 10.[8]

In aqueous solution BaCl2 behaves as a simple salt; in water it is a 1:2 electrolyte and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white precipitate of barium sulfate.

Ba2+(aq) + SO42−(aq) → BaSO4(s)

Oxalate effects a similar reaction:

Ba2+(aq) + C2O42−(aq)BaC2O4(s)

When it is mixed with sodium hydroxide, it gives the dihydroxide, which is moderately soluble in water.


Barium chloride can be prepared from barium hydroxide or barium carbonate, with barium carbonate 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):[9]

BaSO4(s) + 4 C(s) → BaS(s) + 4 CO(g)

This first step requires high temperatures.

BaS + CaCl2 → BaCl2 + CaS

The second step requires fusion of the reactants. The BaCl2 can then be leached out from the mixture with water. From water solutions of barium chloride, the dihydrate can be crystallized as white crystals: BaCl2·2H2O


As an inexpensive, 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, along with other water-soluble barium salts, is highly toxic.[10] Sodium sulfate and magnesium sulfate are potential antidotes because they form the insoluble solid barium sulfate BaSO4, which is relatively non-toxic because of its insolubility.


  1. ^
  2. ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  3. ^ a b c "NIOSH Pocket Guide to Chemical Hazards #0045". National Institute for Occupational Safety and Health (NIOSH). 
  4. ^ a b "Barium (soluble compounds, as Ba)". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH). 
  5. ^ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  6. ^ Haase, A.; Brauer, G. (1978). "Hydratstufen und Kristallstrukturen von Bariumchlorid". Z. anorg. allg. Chem. 441: 181–195. doi:10.1002/zaac.19784410120. 
  7. ^ Brackett, E. B.; Brackett, T. E.; Sass, R. L. (1963). "The Crystal Structures of Barium Chloride, Barium Bromide, and Barium Iodide". J. Phys. Chem. 67 (10): 2132. doi:10.1021/j100804a038. 
  8. ^ Léger, J. M.; Haines, J.; Atouf, A. (1995). "The Post-Cotunnite Phase in BaCl2, BaBr2 and BaI2 under High Pressure". J. Appl. Cryst. 28 (4): 416. doi:10.1107/S0021889895001580. 
  9. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9. 
  10. ^ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.

External links[edit]