Lithium peroxide
| Names | |
|---|---|
| Other names
Dilithium peroxide, Lithium (I) peroxide
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| Identifiers | |
12031-80-0  |
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| ChemSpider | 23787 |
| Jmol 3D model | Interactive image |
| PubChem | 25489 |
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| 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 | |
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Std enthalpy of
formation (ΔfH |
-13.82 kJ/g |
| Vapor pressure | {{{value}}} |
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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 low density, the solid has been used to remove CO2 from the atmosphere in spacecraft.[3]
Preparation
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 funcitonal theory. The solid features an eclipsed "ethane-like" Li6O2 subunits with an O-O distance of around 1.5 Å.[5]
Uses
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-permiable 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
References
- ↑ "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).
- ↑ 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
- ↑ 3.0 3.1 3.2 Lua error in package.lua at line 80: module 'strict' not found.
- ↑ 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.
- ↑ 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
- ↑ 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
- ↑ Lua error in package.lua at line 80: module 'strict' not found.
- ↑ [1] Patent-pending device invented at The Ohio State University: the world’s first solar battery.
