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Lithium chloride

Lithium chloride
CAS number [7447-41-8] (anhydrous)
[85144-11-2] (hydrate)
RTECS number OJ5950000 (anhydrous)
Molecular formula LiCl
Molar mass 42.39 g/mol
Appearance White crystalline solid
Density 2.07 g/cm³, solid
Melting point

605 °C (878 K)

Boiling point

>1300 °C (>1570 K)

Solubility in water 67.2 g/100 ml (0 °C)
Dipole moment 7.13 D (gas)
MSDS External MSDS
R-phrases 22-36/37/38
S-phrases 26-36/37/39
Related Compounds
Other anions lithium fluoride; lithium bromide; lithium iodide
Other cations sodium chloride; magnesium chloride
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Lithium chloride is a chemical compound with the formula LiCl. It behaves as a fairly typical ionic compound, although the Li+ ion is very small. The salt is hygroscopic and highly soluble in water, and is highly polar. It is more soluble in polar organic solvents such as methanol and acetone than is sodium chloride or potassium chloride.


Chemical properties

Lithium chloride can react as a source of chloride ion. As with any other soluble ionic chloride, it will precipitate insoluble chlorides when added to a solution of an appropriate metal salt such as lead(II) nitrate:

2 LiCl(aq) + Pb(NO3)2(aq) → PbCl2(s) + 2 LiNO3(aq)

The Li+ ion acts as a weak Lewis acid under certain circumstances; for example one mole of lithium chloride is capable of absorbing up to four moles of ammonia.


Lithium chloride may be prepared most simply by reaction of lithium hydroxide or lithium carbonate with hydrochloric acid. It may also be prepared by the highly exothermic reaction of lithium metal with either chlorine or anhydrous hydrogen chloride gas. Anhydrous LiCl is prepared from the hydrate by gently heating under an atmosphere of hydrogen chloride, used to prevent hydrolysis.


Lithium chloride is used for the production of lithium metal, by electrolysis of a LiCl/KCl melt at 450 °C. LiCl is also used as a brazing flux for aluminium in automobile parts. It can be used to improve the efficiency of the Stille reaction. Its desiccant properties can be used to generate potable water by absorbing moisture from the air, which is then released by heating the salt. For a short time in the 1940s lithium chloride was manufactured as a substitute for salt, but this was prohibited after the toxic effects of the compound were recognised.[1] [2][3]


Irritant. Avoid swallowing. Ingestion of this compound can result in poisoning or effects on the central nervous system due to its lithium content; see lithium pharmacology for more details.


  • Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  • N. N. Greenwood, A. Earnshaw, Chemistry of the Elements, 2nd ed., Butterworth-Heinemann, Oxford, UK, 1997.
  • R. Vatassery, titration analysis of LiCl, sat'd in Ethanol by AgNO3 to precipitate AgCl(s). EP of this titration gives%Cl by mass.
  • H. Nechamkin, The Chemistry of the Elements, McGraw-Hill, New York, 1968.
  1. ^ Talbott J. H. (1950). "Use of lithium salts as a substitute for sodium chloride.". Arch Med Interna. 85 (1): 1-10. PMID 15398859.
  2. ^ L. W. Hanlon, M. Romaine, F. J. Gilroy. (1949). "Lithium Chloride as a Substitute for Sodium Chloride in the Diet". Journal of the American Medical Association 139 (11): 688-692.
  3. ^ Case of trie Substitute Salt. TIME (28 Feb 1949).
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Lithium_chloride". A list of authors is available in Wikipedia.
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