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Barium chloride

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Barium chloride
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Other names
Barium muriate
Muryate of Barytes[1]
Barium dichloride
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10361-37-2 7pxY
10326-27-9 (dihydrate) 7pxN
ChemSpider 23540 7pxY
EC number 233-788-1
Jmol-3D images Image
PubChem Template:Chembox PubChem/format
RTECS number CQ8750000 (anhydrous)
CQ8751000 (dihydrate)
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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)
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Boiling point Script error: No such module "convert".
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]
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Crystal structure orthogonal (anhydrous)
monoclinic (dihydrate)
7-9
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−858.56 kJ/mol
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EU Index 056-004-00-8
EU classification Toxic (T)
Harmful (Xn)
R-phrases R20, R25
S-phrases (S1/2), S45
NFPA 704

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0
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Flash point Non-flammable
US health exposure limits (NIOSH):

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Other anions
Barium fluoride
Barium bromide
Barium iodide
Other cations
Beryllium chloride
Magnesium chloride
Calcium chloride
Strontium chloride
Radium chloride
Lead chloride
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Thermodynamic
data

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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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

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.[4] The coordination of Ba2+ is 8 in the fluorite structure[5] and 9 in the cotunnite structure.[6] 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.[7]

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.

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):[8]

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

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.[9] 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.

References

  1. ^ https://play.google.com/books/reader?printsec=frontcover&output=reader&id=nKQ-AAAAYAAJ&pg=GBS.PA64
  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. ^ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  5. ^ Haase, A.; Brauer, G. (1978). "Hydratstufen und Kristallstrukturen von Bariumchlorid". Z. anorg. allg. Chem. 441: 181–195. doi:10.1002/zaac.19784410120.  edit
  6. ^ 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.  edit
  7. ^ 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.  edit
  8. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 
  9. ^ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.

External links

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