Lithium chloride
| Lithium chloride | |
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Lithium chloride |
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Lithium(1+) chloride |
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| Identifiers | |
| CAS number | 7447-41-8  |
| PubChem | 4933294 |
| ChemSpider | 22449 = ? |
| UNII | G4962QA067 |
| UN number | 2056 |
| MeSH | Lithium+chloride |
| ChEBI | CHEBI:48607 |
| ChEMBL | CHEMBL69710  |
| RTECS number | OJ5950000 |
| Jmol-3D images | Image 1 |
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| Properties | |
| Molecular formula | LiCl |
| Molar mass | 42.394(4) g/mol |
| Appearance | white solid hygroscopic [1] |
| Density | 2.068 g/cm3 (anhydrous) |
| Melting point |
605 °C |
| Boiling point |
1382 °C (decomp) |
| Solubility in water | 83.2 g/100 mL (20 °C) (anhydrous) |
| Solubility | highly soluble in alcohol, pyridine 4.1 g/100 mL (acetone) |
| Refractive index (nD) | 1.662 |
| Structure | |
| Coordination geometry |
Octahedral |
| Molecular shape | Linear (gas) |
| Dipole moment | 7.13 D (gas) |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
-9.638 kJ/g |
| Specific heat capacity, C | 1.132 J/(g K) |
| Hazards | |
| MSDS | External MSDS |
| EU Index | Not listed |
| NFPA 704 |
 0
1
0
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| Flash point | Non-flammable |
| Related compounds | |
| Other anions | Lithium fluoride Lithium bromide Lithium iodide |
| Other cations | Sodium chloride Potassium chloride Rubidium chloride Caesium chloride |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data |
Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
 (verify) (what is:  / ?)Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
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| Infobox references |
Lithium chloride is a chemical compound with the formula LiCl. The salt is a typical ionic compound, although the small size of the Li+ ion gives rise to properties not seen for other alkali metal chlorides, such as extraordinary solubility in polar solvents (83g/100 mL of water at 20 °C) and its hygroscopic properties.[2]
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[edit] Chemical properties
The salt forms crystalline hydrates, unlike the other alkali metal chlorides.[3] Mono-, tri-, and pentahydrates are known.[4] It also absorbs up to four equivalents of ammonia. As with any other ionic chlorides, solutions of lithium chloride can serve as a source of chloride ion, e.g. forming a precipitate upon treatment with silver nitrate:
- LiCl + AgNO3 → AgCl + LiNO3
[edit] Preparation
Lithium chloride is produced by treatment of lithium carbonate with hydrochloric acid. It can in principle also be generated by the highly exothermic reaction of lithium metal with either chlorine or anhydrous hydrogen chloride gas. Anhydrous LiCl is prepared from the hydrate by heating with a stream of hydrogen chloride.
[edit] Uses
Lithium chloride is mainly used for the production of lithium metal by electrolysis of a LiCl/KCl melt at 600 °C. LiCl is also used as a brazing flux for aluminium in automobile parts. It is used as a desiccant for drying air streams.[2] In more specialized applications, lithium chloride finds some use in organic synthesis, e.g. as an additive in the Stille reaction. Also, in biochemical applications, it can be used to precipitate RNA from cellular extracts.[5]
Lithium chloride is also used as a flame colorant to produce dark red flames.
Lithium chloride is used as a relative humidity standard in the calibration of hygrometers. At 25C a saturated solution (45.81%) of the salt will yield an equilibrium relative humidity of 11.30%. Additionally, lithium chloride can itself be used as a hygrometer. This deliquescent salt forms a self solution when exposed to air. The equilibrium LiCl concentration in the resulting solution is directly related to the relative humidity of the air. The relative humidity at 25C, with minimal error in the range 10C to 30C, in percent, can be estimated from the following first order equation: RH=107.93-2.11C, where C is solution LiCl concentration, percent by mass.
[edit] Precautions
Lithium salts affect the central nervous system; see lithium pharmacology for more details. For a short time in the 1940s lithium chloride was manufactured as a salt substitute, but this was prohibited after the toxic effects of the compound were recognized.[6][7][8]
[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.
- 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.
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398
- ^ a b 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. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Andreas Hönnerscheid, Jürgen Nuss, Claus Mühle, Martin Jansen "Die Kristallstrukturen der Monohydrate von Lithiumchlorid und Lithiumbromid" Zeitschrift für anorganische und allgemeine Chemie, 2003, volume 629, p. 312-316.doi: 10.1002/zaac.200390049
- ^ Cathala, G., Savouret, J., Mendez, B., West, B.L., Karin, M., Martial, J.A., and Baxter, J.D. (1983). "A Method for Isolation of Intact, Translationally Active Ribonucleic Acid". DNA 2 (4): 329–335. doi:10.1089/dna.1983.2.329. PMID 6198133.
- ^ Talbott J. H. (1950). "Use of lithium salts as a substitute for sodium chloride". Arch Med Interna. 85 (1): 1–10. PMID 15398859.
- ^ L. J. Stone, M. luton, lu3. J. Gilroy. (1949). "Lithium Chloride as a Substitute for Sodium Chloride in the Diet". Journal of the American Medical Association 139 (11): 688–692. PMID 18128981.
- ^ "Case of trie Substitute Salt". TIME. 28 February 1949. http://www.time.com/time/magazine/article/0,9171,799873,00.html.
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