Lithium peroxide

From Wikipedia, the free encyclopedia
Jump to navigation Jump to search
Lithium peroxide
Lithium peroxide.svg
Other names
Dilithium peroxide, Lithium (I) peroxide
3D model (JSmol)
ECHA InfoCard 100.031.585
Molar mass 45.881 g/mol
Appearance fine, white powder
Odor odorless
Density 2.31 g/cm3[1][2]
Melting point 195 °C (383 °F; 468 K)
Boiling point Decomposes to Li2O
Solubility insoluble in alcohol
-13.82 kJ/g
not listed
NFPA 704
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is ☑Y☒N ?)
Infobox references

Lithium peroxide is the inorganic compound with the formula Li2O2. It is a white, nonhygroscopic solid. Because of its high oxygen:mass and oxygen:volume ratios, the solid has been used to remove CO2 from the atmosphere in spacecraft.[3]


It is prepared by the reaction of hydrogen peroxide and lithium hydroxide. This reaction initially produces lithium hydroperoxide:[3][4]

LiOH + H2O2 → LiOOH + 2 H2O

This lithium hydroperoxide has also been described as lithium peroxide monoperoxohydrate trihydrate (Li2O2·H2O2·3H2O). Dehydration of this material gives the anhydrous peroxide salt:

2 LiOOH → Li2O2 + H2O2 + 2 H2O

Li2O2 decomposes at about 450 °C to give lithium oxide:

2 Li2O2 → 2 Li2O + O2

The structure of solid Li2O2 has been determined by X-ray crystallography and density functional theory. The solid features an eclipsed "ethane-like" Li6O2 subunits with an O-O distance of around 1.5 Å.[5]


It is used in air purifiers where weight is important, e.g., spacecraft to absorb carbon dioxide and release oxygen in the reaction:[3]

2 Li2O2 + 2 CO2 → 2 Li2CO3 + O2

It absorbs more CO2 than does the same weight of lithium hydroxide and offers the bonus of releasing oxygen.[6] Furthermore, unlike most other alkali metal peroxides, it is not hygroscopic.

The reversible lithium peroxide reaction is the basis for a prototype lithium–air battery. Using oxygen from the atmosphere allows the battery to eliminate storage of oxygen for its reaction, saving battery weight and size.[7]

The successful combination of a lithium-air battery overlain with an air-permeable mesh solar cell was announced by The Ohio State University in 2014.[8] The combination of two functions in one device (a "solar battery") is expected to reduce costs significantly compared to separate devices and controllers as are currently employed.

See also[edit]


  1. ^ "Physical Constants of Inorganic Compounds," in CRC Handbook of Chemistry and Physics, 91st Edition (Internet Version 2011), W. M. Haynes, ed., CRC Press/Taylor and Francis, Boca Raton, Florida. (pp: 4-72).
  2. ^ Speight, James G. (2005). Lange's Handbook of Chemistry (16th Edition). (pp: 1.40). McGraw-Hill. Online version available at:
  3. ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. p. 98. ISBN 0-08-022057-6.
  4. ^ E. Dönges "Lithium and Sodium Peroxides" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 979.
  5. ^ L. G. Cota and P. de la Mora "On the structure of lithium peroxide, Li2O2" Acta Crystallogr. 2005, vol. B61, pages 133-136. doi:10.1107/S0108768105003629
  6. ^ 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.pub2
  7. ^ Girishkumar, G.; B. McCloskey; AC Luntz; S. Swanson; W. Wilcke (July 2, 2010). "Lithium- air battery: Promise and challenges". The Journal of Physical Chemistry Letters. 1 (14): 2193–2203. doi:10.1021/jz1005384.
  8. ^ [1] Patent-pending device invented at The Ohio State University: the world’s first solar battery.

External links[edit]