Muryate of Barytes
|Jmol 3D model||Interactive image|
|RTECS number||CQ8750000 (anhydrous)
|Molar mass||208.23 g/mol (anhydrous)
244.26 g/mol (dihydrate)
|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|
Std enthalpy of
|Safety data sheet||See: data page|
EU classification (DSD)
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|78 mg/kg (rat, oral)
50 mg/kg (guinea pig, oral)
LDLo (lowest published)
|112 mg Ba/kg (rabbit, oral)
59 mg Ba/kg (dog, oral)
46 mg Ba/kg (mouse, oral)
|US health exposure limits (NIOSH):|
|TWA 0.5 mg/m3|
|TWA 0.5 mg/m3|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
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. The coordination of Ba2+ is 8 in the fluorite structure and 9 in the cotunnite structure. 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.
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.
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):
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
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.
Barium chloride, along with other water-soluble barium salts, is highly toxic. 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.
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