Barium chloride

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Barium chloride
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UN number 1564
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Molar mass 208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)
Appearance White powder, or colourless or white crystals (anhydrous)
Colourless rhomboidal crystals (dihydrate)[1][2]
Odor Odourless
Density 3.856 g/cm3 (anhydrous)
3.0979 g/cm3 (dihydrate)
Melting point Template:Chembox CalcTemperatures
Boiling point Template:Chembox CalcTemperatures
Solubility Soluble in methanol, insoluble ethyl acetate, slightly soluble in hydrochloric acid and nitric acid, soluble in ethanol.[3][2] The dihydrate of barium chloride is soluble in methanol, almost insoluble in ethanol, acetone and ethyl acetate.[2]
Template:Longitem −72.6·10−6 cm3/mol
Template:Longitem PbCl2-type orthorhombic (anhydrous)
monoclinic (dihydrate)
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Template:Longitem 123.9 J/(mol·K)
Template:Longitem −858.56 kJ/mol
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Barium chloride is an inorganic compound with the formula Template:Chem2. It is one of the most common water-soluble salts of barium. Like most other water-soluble barium salts, it is a white powder, highly toxic, and imparts a yellow-green coloration to a flame. It is also hygroscopic, converting to the dihydrate Template:Chem2, which are colourless crystals with a bitter salty taste. It has limited use in the laboratory and industry.[4][2]

Preparation

On an industrial scale, barium chloride is prepared via a two step process from barite (barium sulfate).[5] The first step requires high temperatures.

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The second step requires reaction between barium sulfide and hydrogen chloride:

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or between barium sulfide and calcium chloride:

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In place of HCl, chlorine can be used.[4] Barium chloride is extracted out from the mixture with water. From water solutions of barium chloride, its dihydrate (Template:Chem2) can be crystallized as colorless crystals.[1]

Barium chloride can in principle be prepared by the reaction between barium hydroxide or barium carbonate with hydrogen chloride. These basic salts react with hydrochloric acid to give hydrated barium chloride.

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Structure and properties

Template:Chem2 crystallizes in two forms (polymorphs). At room temperature, the compound is stable in the orthorhombic cotunnite ([[Lead(II) chloride|Template:Chem2]]) structure, whereas the cubic fluorite structure ([[calcium fluoride|Template:Chem2]]) is stable between 925 and 963 °C.[6] Both polymorphs accommodate the preference of the large Template:Chem2 ion for coordination numbers greater than six.[7] The coordination of Template:Chem2 is 8 in the fluorite structure[8] and 9 in the cotunnite structure.[9] When cotunnite-structure Template:Chem2 is subjected to pressures of 7–10 GPa, it transforms to a third structure, a monoclinic post-cotunnite phase. The coordination number of Template:Chem2 increases from 9 to 10.[10]

In aqueous solution Template:Chem2 behaves as a simple salt; in water it is a 1:2 electrolyteTemplate:Cln and the solution exhibits a neutral pH. Its solutions react with sulfate ion to produce a thick white solid precipitate of barium sulfate.

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This precipitation reaction is used in chlor-alkali plants to control the sulfate concentration in the feed brine for electrolysis.

Oxalate effects a similar reaction:

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

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Template:Chem2 is stable in the air at room temperature, but loses one water of crystallization above Template:Cvt, becoming Template:Chem2, and becomes anhydrous above Template:Cvt.[1] Template:Chem2 may be formed by shaking the dihydrate with methanol.[2]

Template:Chem2 readily forms eutectics with alkali metal chlorides.[2]

Uses

Although inexpensive, barium chloride finds limited applications in the laboratory and industry.

Its main laboratory use is as a reagent for the gravimetric determination of sulfates. The sulfate compound being analyzed is dissolved in water and hydrochloric acid is added. When barium chloride solution is added, the sulfate present precipitates as barium sulfate, which is then filtered through ashless filter paper. The paper is burned off in a muffle furnace, the resulting barium sulfate is weighed, and the purity of the sulfate compound is thus calculated.

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.[4] It is also used to make red pigments such as Lithol red and Red Lake C. Its toxicity limits its applicability.Script error: No such module "Unsubst".

Toxicity

Barium chloride, along with other water-soluble barium salts, is highly toxic.[11] It irritates eyes and skin, causing redness and pain. It damages kidneys. Fatal dose of barium chloride for a human has been reported to be about 0.8-0.9 g. Systemic effects of acute barium chloride toxicity include abdominal pain, diarrhea, nausea, vomiting, cardiac arrhythmia, muscular paralysis, and death. The Template:Chem2 ions compete with the Template:Chem2 ions, causing the muscle fibers to be electrically unexcitable, thus causing weakness and paralysis of the body.[2] Sodium sulfate and magnesium sulfate are potential antidotes because they form barium sulfate Template:Chem/link, which is relatively non-toxic because of its insolubility in water.

Barium chloride is not classified as a human carcinogen.[2]

References

  1. ↑ a b c d Script error: No such module "citation/CS1".
  2. ↑ a b c d e f g h Script error: No such module "citation/CS1".
  3. ↑ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  4. ↑ a b c Script error: No such module "citation/CS1".
  5. ↑ Script error: No such module "citation/CS1".Script error: No such module "Unsubst".
  6. ↑ Script error: No such module "Citation/CS1".
  7. ↑ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. Template:ISBN.Script error: No such module "Unsubst".
  8. ↑ Script error: No such module "Citation/CS1".
  9. ↑ Script error: No such module "Citation/CS1".
  10. ↑ Script error: No such module "Citation/CS1".
  11. ↑ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.Script error: No such module "Unsubst".

External links

Template:Barium compounds Template:Chlorides

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