Lithium chloride
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| Lithium chloride | |
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| General | |
| Molecular formula | LiCl |
| Molar mass | 42.39 g/mol |
| Appearance | White crystalline solid |
| CAS number | [7447-41-8] (anhydrous) [85144-11-2] (hydrate) |
| Hydrates | monohydrate, trihydrate |
| Properties | |
| Density and phase | 2.07 g/cm³, solid |
| Solubility in water | 63.7 g/100 ml (0 °C) |
| In ethanol In acetone |
42.4 g/100 ml (25 °C) 4.11 g/100 ml (25 °C) |
| Melting point | 605 °C (878 K) |
| Boiling point | >1300 °C (>1570 K) |
| Structure | |
| Coordination geometry | Octahedral |
| Crystal structure | |
| Dipole moment | 7.13 D (gas) |
| Hazards | |
| MSDS | External MSDS |
| Main hazards | Irritant |
| R/S statement | R: 22-36/37/38 S: 26-36/37/39 |
| RTECS number | OJ5950000 (anhydrous) |
| Supplementary data page | |
| Structure & properties | n, εr, etc. |
| Thermodynamic data | Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Related compounds | |
| Other anions | lithium fluoride |
| Other cations | sodium chloride |
| Except where noted otherwise, data are given for materials in their standard state (at 25°C, 100 kPa) Infobox disclaimer and references |
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Lithium chloride, LiCl, 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.
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[edit] 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.
| Solubility of LiCl in various solvents (g LiCl / 100g of solvent at 25° C) |
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| H2O | 55 |
| Liquid ammonia | 3.02 |
| Liquid sulfur dioxide | 0.012 |
| Methanol | 21 - 41 |
| Formic acid | 27.5 |
| Sulfolane | 1.5 |
| Acetonitrile | 0.14 |
| Acetone | 0.83 |
| Formamide | 28.2 |
| Dimethylformamide | 11 - 28 |
| Reference: Burgess, J. Metal Ions in Solution (Ellis Horwood, New York, 1978) ISBN 0-85312-027-7 |
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[edit] Preparation
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.
[edit] Uses
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] [1]
[edit] Precautions
Irritant. Avoid swallowing.
[edit] References
- 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.
- H. Nechamkin, The Chemistry of the Elements, McGraw-Hill, New York, 1968.
- ^ Talbott J. H. (1950). "Use of lithium salts as a substitute for sodium chloride.". Arch Med Interna. 85 (1): 1-10. PMID 15398859.
- ^ 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.
