Barium chloride

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 barium salts, it is white, 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
Names
Other names
Barium muriate
Muryate of Barytes[1]
Barium dichloride
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.030.704
EC Number
  • 233-788-1
RTECS number
  • CQ8750000 (anhydrous)
    CQ8751000 (dihydrate)
UNII
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 (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
Std enthalpy of
formation (ΔfH⦵298)
 −858.56 kJ/mol
Hazards
Main hazards Acute Toxic
Safety data sheet See: data page
NIH BaCl
GHS pictograms
GHS Signal word Danger
GHS hazard statements
 H301, H332
P261, P264, P270, P271, P301+310, P304+312, P304+340, P312, P321, P330, P405, P501
NFPA 704 (fire diamond)
Flammability code 0: Will not burn. E.g. waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasReactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
0
3
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
78 mg/kg (rat, oral)
50 mg/kg (guinea pig, oral)[3]
112 mg Ba/kg (rabbit, oral)
59 mg Ba/kg (dog, oral)
46 mg Ba/kg (mouse, oral)[3]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 0.5 mg/m3[4]
REL (Recommended)
 TWA 0.5 mg/m3[4]
IDLH (Immediate danger)
 50 mg/m3[4]
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 constant (εr), etc.
Thermodynamic
data
 Phase behaviour
solid–liquid–gas
UV, IR, NMR, MS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

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.[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+ + 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

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.

Uses

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] Its toxicity limits its applicability.

Safety

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.

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References
  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. "Barium (soluble compounds, as Ba)". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. NIOSH Pocket Guide to Chemical Hazards. "#0045". National Institute for Occupational Safety and Health (NIOSH).
  5. 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.
Salts and covalent derivatives of the chloride ion
HCl He
LiCl BeCl2 BCl3
B2Cl4		 CCl4 NCl3
ClN3		 Cl2O
ClO2
Cl2O7		 ClF
ClF3
ClF5		 Ne
NaCl MgCl2 AlCl
AlCl3		 SiCl4 P2Cl4
PCl3
PCl5		 S2Cl2
SCl2
SCl4		 Cl2 Ar
KCl CaCl
CaCl2		 ScCl3 TiCl2
TiCl3
TiCl4		 VCl2
VCl3
VCl4
VCl5		 CrCl2
CrCl3
CrCl4		 MnCl2 FeCl2
FeCl3		 CoCl2
CoCl3		 NiCl2 CuCl
CuCl2		 ZnCl2 GaCl2
GaCl3		 GeCl2
GeCl4		 AsCl3
AsCl5		 Se2Cl2
SeCl4		 BrCl KrCl
RbCl SrCl2 YCl3 ZrCl3
ZrCl4		 NbCl3
NbCl4
NbCl5		 MoCl2
MoCl3
MoCl4
MoCl5
MoCl6		 TcCl4 RuCl3 RhCl3 PdCl2 AgCl CdCl2 InCl
InCl2
InCl3		 SnCl2
SnCl4		 SbCl3
SbCl5		 Te3Cl2
TeCl4		 ICl
ICl3		 XeCl
XeCl2
XeCl4
CsCl BaCl2   HfCl4 TaCl5 WCl2
WCl3
WCl4
WCl5
WCl6		 Re3Cl9
ReCl4
ReCl5
ReCl6		 OsCl4 IrCl2
IrCl3
IrCl4		 PtCl2
PtCl4		 AuCl
AuCl3		 Hg2Cl2,
HgCl2		 TlCl PbCl2,
PbCl4		 BiCl3 PoCl2,
PoCl4		 AtCl RnCl2
FrCl RaCl2   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
LaCl3 CeCl3 PrCl3 NdCl2,
NdCl3		 PmCl3 SmCl2,
SmCl3		 EuCl2,
EuCl3		 GdCl3 TbCl3 DyCl2,
DyCl3		 HoCl3 ErCl3 TmCl2
TmCl3		 YbCl2
YbCl3		 LuCl3
AcCl3 ThCl4 PaCl5 UCl3
UCl4
UCl5
UCl6		 NpCl4 PuCl3 AmCl2
AmCl3		 CmCl3 BkCl3 CfCl3 EsCl3 Fm Md No LrCl3
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