Lithium hydroxide

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Lithium hydroxide
Lithium hydroxide
IUPAC name
Lithium hydroxide
Other names
1310-65-2 YesY
1310-66-3 (monohydrate) N
ChEBI CHEBI:33979 YesY
ChemSpider 3802 YesY
Jmol interactive 3D Image
PubChem 3939
RTECS number OJ6307070
UN number 2680
Molar mass 23.95 g/mol (anhydrous)
41.96 g/mol (monohydrate)
Appearance hygroscopic white solid
Density 1.46 g/cm3 (anhydrous)
1.51 g/cm3 (monohydrate)
Melting point 462 °C (864 °F; 735 K)
Boiling point 924 °C (1,695 °F; 1,197 K) decomposes
12.7 g/100 mL (0 °C)
12.8 g/100 mL (20 °C)
17.5 g/100 mL (100 °C)
22.3 g/100 mL (10 °C)
26.8 g/100 mL (80 °C)[1]
Solubility in methanol anhydrous:
9.76 g/100 g (20 °C, 48 hours mixing)
13.69 g/100 g (20 °C, 48 hours mixing)[2]
Solubility in ethanol anhydrous:
2.36 g/100 g (20 °C, 48 hours mixing)
2.18 g/100 g (20 °C, 48 hours mixing)[2]
Solubility in isopropanol anhydrous:
0 g/100 g (20 °C, 48 hours mixing)
0.11 g/100 g (20 °C, 48 hours mixing)[2]
Basicity (pKb) -0.63[3]
1.464 (anhydrous)
1.460 (monohydrate)
2.071 J/g K
-20.36 kJ/g
Main hazards Corrosive
Safety data sheet ICSC 0913
ICSC 0914 (monohydrate)
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
210 mg/kg (oral, rat)[4]
Related compounds
Other anions
Lithium amide
Other cations
Sodium hydroxide
Potassium hydroxide
Rubidium hydroxide
Caesium hydroxide
Related compounds
Lithium oxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Lithium hydroxide is an inorganic compound with the formula LiOH. It is a white hygroscopic crystalline material. It is soluble in water and slightly soluble in ethanol, and is the weakest base among the alkali metal hydroxides. It is available commercially in anhydrous form and as the monohydrate (LiOH.H2O), both of which are strong bases.

Production and reactions[edit]

Lithium hydroxide is produced in a metathesis reaction between lithium carbonate and calcium hydroxide:[5]

Li2CO3 + Ca(OH)2 → 2 LiOH + CaCO3

The initially produced hydrate is dehydrated by heating under vacuum up to 180 °C.

In the laboratory, lithium hydroxide arises by the action of water on lithium or lithium oxide. The equations for these processes follow:

2 Li + 2 H2O → 2 LiOH + H2
Li2O + H2O → 2 LiOH

Typically, these reactions are avoided.

Although lithium carbonate is more widely used, the hydroxide is an effective precursor to lithium salts, e.g.

LiOH + HF → LiF + H2O.


Lithium hydroxide is mainly consumed for the production of lithium greases. A popular lithium grease is lithium stearate, which is a general-purpose lubricating grease due to its high resistance to water and usefulness at both high and low temperatures.

Carbon dioxide scrubbing[edit]

Further information: carbon dioxide scrubber

Lithium hydroxide is used in breathing gas purification systems for spacecraft, submarines, and rebreathers to remove carbon dioxide from exhaled gas by producing lithium carbonate and water:[6]

2 LiOH·H2O + CO2 → Li2CO3 + 3 H2O


2LiOH + CO2 → Li2CO3 + H2O

The latter, anhydrous hydroxide, is preferred for its lower mass and lesser water production for respirator systems in spacecraft. One gram of anhydrous lithium hydroxide can remove 450 cm3 of carbon dioxide gas. The monohydrate loses its water at 100–110 °C.

Other uses[edit]

It is used as a heat transfer medium and as a storage-battery electrolyte. It is also used in ceramics and some Portland cement formulations. Lithium hydroxide (isotopically enriched in lithium-7) is used to alkalize the reactor coolant in pressurized water reactors for corrosion control.

See also[edit]


  1. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3. 
  2. ^ a b c Khosravi, Javad (2007). "9: Results". PRODUCTION OF LITHIUM PEROXIDE AND LITHIUM OXIDE IN AN ALCOHOL MEDIUM. ISBN 978-0-494-38597-5. 
  3. ^ Lew. Kristi., Acids and Bases (Essential Chemistry). Infobase Publishing (2009). p43.
  4. ^
  5. ^ Wietelmann, U; Bauer, RJ (2000). "Lithium and Lithium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a15_393. ISBN 3-527-30673-0. 
  6. ^ Jaunsen, JR (1989). "The Behavior and Capabilities of Lithium Hydroxide Carbon Dioxide Scrubbers in a Deep Sea Environment". US Naval Academy Technical Report. USNA-TSPR-157. Retrieved 2008-06-17. 

External links[edit]