Lithium peroxide

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Lithium peroxide
Names
Other names
Dilithium peroxide, Lithium (I) peroxide
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.031.585
Properties
Li2O2
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
soluble
Solubility insoluble in alcohol
Structure
hexagonal
Thermochemistry
-13.82 kJ/g
Hazards
not listed
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 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard OX: Oxidizer. E.g., potassium perchlorateNFPA 704 four-colored diamond
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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]

Preparation[edit]

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]

Uses[edit]

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 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]

References[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, FL. (pp: 4-72).
  2. ^ Speight, James G. (2005). Lange's Handbook of Chemistry (16th Edition). (pp: 1.40). McGraw-Hill. Online version available at: http://www.knovel.com/web/portal/browse/display?_EXT_KNOVEL_DISPLAY_bookid=1347&VerticalID=0
  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]