Barium chloride
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.030.704 |
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CompTox Dashboard (EPA)
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| Properties | |
| BaCl2 | |
| 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 960 °C (dihydrate) |
| Boiling point | 1560 °C |
| 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 [1] |
| Structure | |
| orthogonal (anhydrous) monoclinic (dihydrate) | |
| 7-9 | |
| Thermochemistry | |
Std enthalpy of
formation (ΔfH⦵298) |
−858.56 kJ/mol |
| Hazards | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Related compounds | |
Other anions
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Barium fluoride Barium bromide Barium iodide |
Other cations
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Beryllium chloride Magnesium chloride Calcium chloride Strontium chloride Radium chloride Lead chloride |
| Supplementary data page | |
| Barium chloride (data page) | |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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
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.[2] The coordination of Ba2+ is 8 in the fluorite structure[3] and 9 in the cotunnite structure.[4] 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.[5]
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+ + SO42- → BaSO4
Oxalate effects a similar reaction:
- Ba2+ + C2O42- → BaC2O4
When it is mixed with sodium hydroxide, it gives the dihydroxide, which is moderately soluble in water.
Preparation
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):[6]
This first step requires high temperatures.
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
Uses
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.
Safety
Barium chloride, along with other water-soluble barium salts, is highly toxic.[7] Sodium sulfate and magnesium sulfate are potential antidotes because they form the insoluble solid barium sulfate BaSO4, which is nontoxic.
References
- ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
- ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi: 10.1002/zaac.19784410120 , please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
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|doi= 10.1021/j100804a038instead. - ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi: 10.1107/S0021889895001580 , please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with
|doi= 10.1107/S0021889895001580instead. - ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.