Barium chloride

Summary

Barium chloride is an inorganic compound with the formula BaCl2. It is one of the most common water-soluble salts of barium. Like most other water-soluble barium salts, it is white, highly toxic, and imparts a yellow-green coloration to a flame. It is also hygroscopic, converting first to the dihydrate BaCl2(H2O)2. It has limited use in the laboratory and industry.[5]

Barium chloride
Cotunnite structure.png
Barium chloride.jpg
Names
Other names
Barium muriate
Muryate of Barytes[1]
Barium dichloride crystals
Neutral barium chloride
Identifiers
  • 10361-37-2 checkY
  • 10326-27-9 (dihydrate) checkY
3D model (JSmol)
  • Interactive image
ChemSpider
  • 23540 checkY
ECHA InfoCard 100.030.704 Edit this at Wikidata
EC Number
  • 233-788-1
  • 25204
RTECS number
  • CQ8750000 (anhydrous)
    CQ8751000 (dihydrate)
UNII
  • 0VK51DA1T2 checkY
  • EL5GJ3U77E (dihydrate) checkY
UN number 1564
  • DTXSID7044508 Edit this at Wikidata
  • InChI=1S/Ba.2ClH/h;2*1H/q+2;;/p-2 checkY
    Key: WDIHJSXYQDMJHN-UHFFFAOYSA-L checkY
  • InChI=1/Ba.2ClH/h;2*1H/q+2;;/p-2
    Key: WDIHJSXYQDMJHN-NUQVWONBAL
  • [Ba+2].[Cl-].[Cl-]
Properties
BaCl2
Molar mass 208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)
Appearance White solid
Odor Odourless
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
Structure
orthogonal (anhydrous)
monoclinic (dihydrate)
7–9
Thermochemistry
123.9 J/(k mol)
−858.56 kJ/mol
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
Highly toxic, corrosive
GHS labelling:
GHS06: Toxic
Danger
H301, H302, H332
P261, P264, P270, P271, P301+P310, P304+P312, P304+P340, P312, P321, P330, P405, P501
NFPA 704 (fire diamond)
3
0
0
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]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 0.5 mg/m3[3]
REL (Recommended)
TWA 0.5 mg/m3[3]
IDLH (Immediate danger)
50 mg/m3[3]
Safety data sheet (SDS) NIH BaCl
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
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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Infobox references

Structure and propertiesEdit

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

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+ + SO2−
4
→ BaSO4

Oxalate effects a similar reaction:

Ba2+ + C
2
O2−
4
BaC2O4

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

PreparationEdit

On an industrial scale, it is prepared via a two step process from barite (barium sulfate):[10]

BaSO4 + 4 C → BaS + 4 CO

This first step requires high temperatures.

BaS + 2 HCl → BaCl2 + H2S

In place of HCl, chlorine can be used.[5]

Barium chloride can in principle be prepared from barium hydroxide or barium carbonate. These basic salts react with hydrochloric acid to give hydrated barium chloride.

UsesEdit

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

[5] It is also used to make red pigments such as Lithol red and Red Lake C. Its toxicity limits its applicability.

SafetyEdit

Barium chloride, along with other water-soluble barium salts, is highly toxic.[11] Sodium sulfate and magnesium sulfate are potential antidotes because they form barium sulfate BaSO4, which is relatively non-toxic because of its insolubility.

ReferencesEdit

  1. ^ Chemical Recreations: A Series of Amusing and Instructive Experiments, which May be Performed with Ease, Safety, Success, and Economy ; to which is Added, the Romance of Chemistry : An Inquiry into the Fallacies of the Prevailing Theory of Chemistry : With a New Theory and a New Nomenclature. R. Griffin & Company. 1834.
  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 or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  5. ^ a b c Kresse, Robert; Baudis, Ulrich; Jäger, Paul; Riechers, H. Hermann; Wagner, Heinz; Winkler, Jocher; Wolf, Hans Uwe (2007). "Barium and Barium Compounds". In Ullman, Franz (ed.). Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a03_325.pub2. ISBN 978-3527306732.
  6. ^ Wells, A. F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  7. ^ Haase, A.; Brauer, G. (1978). "Hydratstufen und Kristallstrukturen von Bariumchlorid". Z. anorg. allg. Chem. 441: 181–195. doi:10.1002/zaac.19784410120.
  8. ^ 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.
  9. ^ 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.
  10. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  11. ^ The Merck Index, 7th edition, Merck & Co., Rahway, New Jersey, 1960.

External linksEdit

  • International Chemical Safety Card 0614. (anhydrous)
  • International Chemical Safety Card 0615. (dihydrate)
  • Barium chloride's use in industry.
  • ChemSub Online: Barium chloride.