Lithium carbonate

From Wikipedia, the free encyclopedia
Jump to: navigation, search
Lithium carbonate
Lithium carbonate
Lithium-carbonate-xtal-1979-Mercury-3D-sf.png
Identifiers
CAS number 554-13-2 YesY
PubChem 11125
ChemSpider 10654 YesY
UNII 2BMD2GNA4V YesY
KEGG D00801 YesY
ChEBI CHEBI:6504 YesY
ChEMBL CHEMBL1200826 N
RTECS number OJ5800000
Jmol-3D images Image 1
Properties
Molecular formula CLi2O3
Molar mass 73.89 g mol−1
Appearance Odorless white powder
Density 2.11 g/cm3
Melting point 723 °C (1,333 °F; 996 K)
Boiling point 1,310 °C (2,390 °F; 1,580 K)
decomposes from ~1300 °C
Solubility in water 1.54 g/100 mL (0 °C)
1.43 g/100 mL (10 °C)
1.29 g/100 mL (25 °C)
1.08 g/100 mL (40 °C)
0.69 g/100 mL (100 °C)[1]
Solubility Insoluble in acetone, ammonia, alcohol[2]
Refractive index (nD) 1.428[3]
Viscosity 4.64 cP (777 °C)
3.36 cP (817 °C)[2]
Thermochemistry
Specific
heat capacity
C
97.4 J/mol·K[2]
Std molar
entropy
So298
90.37 J/mol·K[2]
Std enthalpy of
formation
ΔfHo298
-1215.6 kJ/mol[2]
Gibbs free energy ΔG -1132.4 kJ/mol[2]
Hazards
MSDS ICSC 1109
GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[4]
GHS signal word Warning
GHS hazard statements H302, H319[4]
GHS precautionary statements P305+351+338[4]
EU Index Not listed
EU classification Harmful Xn Irritant Xi
R-phrases R22, R36
S-phrases S26, S36/37
Main hazards Irritant
Flash point Non-flammable
LD50 525 mg/kg (oral, rat)[5]
Related compounds
Other cations Sodium carbonate
Potassium carbonate
Rubidium carbonate
Caesium carbonate
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 N (verify) (what is: YesY/N?)
Infobox references

Lithium carbonate is an inorganic compound, the lithium salt of carbonate with the formula Li2CO3. This white salt is widely used in the processing of metal oxides. For the treatment of bipolar disorder, it is on the World Health Organization's List of Essential Medicines, a list of the most important medication needed in a basic health system.[6]

Uses[edit]

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.[7] It is also used in the manufacture of most lithium-ion battery cathodes, which are made of lithium cobalt oxide.

Medical uses[edit]

Main article: Lithium pharmacology

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 1949, John Cade discovered the anti-manic 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.[8] 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.[9] Despite these findings, a great deal remains unknown regarding lithium's mood controlling properties.[citation needed]

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.[10]

Properties and reactions[edit]

Unlike sodium carbonate, which forms at least three hydrates, lithium carbonate exists only in the anhydrous form.[11] 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 tenfold under a mild pressure of carbon dioxide; this effect is due to the formation of the metastable bicarbonate, which is more soluble:[7]

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.[12]

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

Production[edit]

Lithium is extracted from primarily two sources: pegmatite crystals and lithium salt from brine pools. Approximately 30,000 tons were produced in 1989. It also exists as the rare mineral zabuyelite.[13]

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:[11]

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2

References[edit]

  1. ^ Seidell, Atherton; Linke, William F. (1952). Solubilities of Inorganic and Organic Compounds. Van Nostrand. 
  2. ^ a b c d e f http://chemister.ru/Database/properties-en.php?dbid=1&id=608
  3. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  4. ^ a b c Sigma-Aldrich Co., Lithium carbonate. Retrieved on 2014-06-03.
  5. ^ http://chem.sis.nlm.nih.gov/chemidplus/rn/554-13-2
  6. ^ "WHO Model List of Essential Medicines" (PDF). World Health Organization. October 2013. Retrieved 22 April 2014. 
  7. ^ 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
  8. ^ Medical use
  9. ^ Yildiz, A; Guleryuz, S; Ankerst, DP; Ongür, D; Renshaw, PF (2008). "Protein kinase C inhibition in the treatment of mania: a double-blind, placebo-controlled trial of tamoxifen". Archives of General Psychiatry 65 (3): 255–63. doi:10.1001/archgenpsychiatry.2007.43. PMID 18316672. 
  10. ^ Simard, M; Gumbiner, B; Lee, A; Lewis, H; Norman, D (1989). "Lithium carbonate intoxication. A case report and review of the literature". Archives of Internal Medicine 149 (1): 36–46. doi:10.1001/archinte.149.1.36. PMID 2492186. 
  11. ^ 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.
  12. ^ E. R. Caley, P. J. Elving "Purification of Lithium Carbonate" Inorganic Syntheses, 1939, vol. 1, p. 1. doi:10.1002/9780470132326.ch1
  13. ^ David Barthelmy. "Zabuyelite Mineral Data". Mineralogy Database. Retrieved 2010-02-07. 

External links[edit]