3D model (Jmol)
|Molar mass||86.845(3) g/mol|
|Melting point||552 °C (1,026 °F; 825 K)|
|Boiling point||1,265 °C (2,309 °F; 1,538 K)|
|143 g/100 mL (0 °C)
166.7 g/100 mL (20 °C)
266 g/100 mL (100 °C)
|Solubility||soluble in methanol, ethanol, ether, acetone
slightly soluble in pyridine
Refractive index (nD)
|51.88 J/mol K|
|66.9 J/mol K|
Std enthalpy of
Gibbs free energy (ΔfG˚)
Std enthalpy of
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|1800 mg/kg (oral, rat)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Production and properties
LiBr is prepared by treatment of lithium carbonate with hydrobromic acid. The salt forms several crystalline hydrates, unlike the other alkali metal bromides. The anhydrous salt forms cubic crystals similar to common salt (sodium chloride).
Lithium hydroxide and hydrobromic acid (aqueous solution of hydrogen bromide) will precipitate lithium bromide in the presence of water.
LiOH + HBr → LiBr + H2O
Lithium bromide is used in air-conditioning systems as desiccant.
Lithium bromide is used as a salt in absorption chilling along with water (see absorption refrigerator). Otherwise the salt is useful as a reagent in organic synthesis. For example, it reversibly forms adducts with some pharmaceuticals.
Lithium bromide was used as a sedative, beginning in the early 1900s, but it fell into disfavor in the 1940s when some heart patients died after using it as a salt substitute.
Like lithium carbonate and lithium chloride, it was used as treatment for bipolar disorder.
Doses as low as 225 mg/day of LiBr can lead to bromism.
- Ulrich Wietelmann, Richard J. Bauer "Lithium and Lithium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH: Weinheim.
- Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5
- Bipolar disorder
- "A PDF file from GFS Chemicals, a supplier of lithium bromide" (PDF). Retrieved 2005-09-15.