Lithium carbonate

Lithium carbonate


Lithium carbonate is an health system.[1]

Contents

  • Uses 1
    • Medical uses 1.1
  • Properties and reactions 2
    • Production 2.1
  • References 3
  • External links 4

Uses

Lithium carbonate is an important industrial chemical. It forms low-melting fluxes with silica and other materials. Glasses derived from lithium carbonate are useful in ovenware. Lithium carbonate is a common ingredient in both low-fire and high-fire ceramic glaze. Its alkaline properties are conducive to changing the state of metal oxide colorants in glaze particularly red iron oxide (Fe2O3). Cement sets more rapidly when prepared with lithium carbonate, and is useful for tile adhesives. When added to aluminium trifluoride, it forms LiF which gives a superior electrolyte for the processing of aluminium.[2] It is also used in the manufacture of most lithium-ion battery cathodes, which are made of lithium cobalt oxide.

Medical uses

In 1843, lithium carbonate was used as a new solvent for stones in the bladder. In 1859, some doctors recommended a therapy with lithium salts for a number of ailments, including gout, urinary calculi, rheumatism, mania, depression, and headache. In 1948, John Cade discovered the antimanic effects of lithium ions. This finding led lithium, specifically lithium carbonate, to be used to treat mania associated with bipolar disorder.

Lithium carbonate is used to treat mania, the elevated phase of bipolar disorder. Lithium ions interfere with ion transport processes (see “sodium pump”) that relay and amplify messages carried to the cells of the brain.[3] Mania is associated with irregular increases in protein kinase C (PKC) activity within the brain. Lithium carbonate and sodium valproate, another drug traditionally used to treat the disorder, act in the brain by inhibiting PKC’s activity and help to produce other compounds that also inhibit the PKC.[4] Despite these findings, a great deal remains unknown regarding lithium's mood-controlling properties.

Use of lithium salts exhibit a number of risks and side effects, especially at higher doses. Lithium intoxication affects the central nervous and renal systems and is potentially lethal.[5]

Properties and reactions

Unlike sodium carbonate, which forms at least three hydrates, lithium carbonate exists only in the anhydrous form.[6] Its solubility in water is low relative to other lithium salts. The isolation of lithium from aqueous extracts of lithium ores capitalizes on this poor solubility. Its apparent solubility increases 10-fold under a mild pressure of carbon dioxide; this effect is due to the formation of the metastable bicarbonate, which is more soluble:[2]

Li2CO3 + CO2 + H2O \overrightarrow{\leftarrow} 2 LiHCO3

The extraction of lithium carbonate at high pressures of CO2 and its precipitation upon depressuring is the basis of the Quebec process.

Lithium carbonate can also be purified by exploiting its diminished solubility in hot water. Thus, heating a saturated aqueous solution causes crystallization of Li2CO3.[7]

Lithium carbonate, and other carbonates of group 1, do not decarboxylate readily. Li2CO3 decomposes at temperatures around 1300°C.

Production

Lithium is extracted from primarily two sources: pegmatite crystals and lithium salt from brine pools. About 30,000 tons were produced in 1989. It also exists as the rare mineral zabuyelite.[8]

Lithium carbonate is generated by combining lithium peroxide with carbon dioxide. This reaction is the basis of certain air purifiers, e.g., in spacecraft, used to absorb carbon dioxide:[6]

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2

References

  1. ^
  2. ^ a b Ulrich Wietelmann, Richard J. Bauer "Lithium and Lithium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH: Weinheim. doi:10.1002/14356007.a15 393
  3. ^ Medical use
  4. ^
  5. ^
  6. ^ a b Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. Pages=84-85 ISBN 0-7506-3365-4.
  7. ^ E. R. Caley, P. J. Elving "Purification of Lithium Carbonate" Inorganic Syntheses, 1939, vol. 1, p. 1. doi:10.1002/9780470132326.ch1
  8. ^

External links

  • Official FDA information published by Drugs.com