Lithium oxide
| |
| Names | |
|---|---|
| IUPAC name
Lithium oxide
| |
| Other names
Lithium monoxide
Lithia | |
| Identifiers | |
| ECHA InfoCard | 100.031.823 |
PubChem CID
|
|
| RTECS number |
|
CompTox Dashboard (EPA)
|
|
| Properties | |
| Li 2O | |
| Molar mass | 29.88 g/mol |
| Appearance | white solid |
| Density | 2.013 g/cm3 |
| Melting point | 1,570 °C (2,860 °F; 1,840 K) |
| decomposes 6.67 g/100 mL (0 °C) 10.02 g/100 mL (100 °C) | |
| log P | 9.23 |
Refractive index (nD)
|
1.644 [1] |
| Structure | |
| Antifluorite (cubic), cF12 | |
| Fm3m, No. 225 | |
| Tetrahedral (Li+); cubic (O2–) | |
| Thermochemistry | |
Heat capacity (C)
|
18.105 J/g K |
Std enthalpy of
formation (ΔfH⦵298) |
-20.01 kJ/g |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards
|
Corrosive, reacts violently with water |
| Flash point | Non-flammable |
| Related compounds | |
Other anions
|
Lithium sulfide |
Other cations
|
Sodium oxide Potassium oxide Rubidium oxide Caesium oxide |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
| |
Lithium oxide (Li
2O) or lithia is an inorganic chemical compound. Lithium oxide is formed along with small amounts of lithium peroxide when lithium metal is burned in the air and combines with oxygen[2]:
- 4Li+O
2 → 2Li
2O.
Pure Li
2O can be produced by the thermal decomposition of lithium peroxide, Li
2O
2 at 450°C[2]
- 2Li
2O
2 → 2Li
2O + O
2
Structure
In the solid state lithium oxide adopts an antifluorite structure which is related to the CaF
2, fluorite structure with Li cations substituted for fluoride anions and oxide anions substituted for calcium cations.[3]
The ground state gas phase Li
2O molecule is linear with a bond length consistent with strong ionic bonding.[4][5] VSEPR theory would predict a bent shape similar to H
2O.
Uses
Lithium oxide is used as a flux in ceramic glazes; and creates blues with copper and pinks with cobalt. Lithium oxide reacts with water and steam, forming lithium hydroxide and should be isolated from them.
Its usage is also being investigated for non-destructive emission spectroscopy evaluation and degradation monitoring within thermal barrier coating systems. It can be added as a co-dopant with yttria in the zirconia ceramic top coat, without a large decrease in expected service life of the coating. At high heat, lithium oxide emits a very detectable spectral pattern, which increases in intensity along with degradation of the coating. Implementation would allow in situ monitoring of such systems, enabling an efficient means to predict lifetime until failure or necessary maintenance.
A possible new use is as a replacement for lithium cobalt oxide as the cathode in the lithium ion batteries used to power electronic devices from mobile phones to laptop computers to battery-powered cars.[6]
See also
References
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0070494398
- ^ a b Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 97–99. ISBN 978-0-08-022057-4.
- ^ Zintl, E.; Harder, A.; Dauth B. (1934). "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums". Z. Elektrochem. Angew. Phys. Chem. 40: 588–93.
{{cite journal}}: line feed character in|journal=at position 68 (help)CS1 maint: multiple names: authors list (link) - ^ Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford Science Publications ISBN 0-19-855370-6
- ^ A spectroscopic determination of the bond length of the LiOLi molecule: Strong ionic bonding, D. Bellert, W. H. Breckenridge, J. Chem. Phys. 114, 2871 (2001); doi:10.1063/1.1349424
- ^ "Air power". The Economist, Technology Quarterly. September 3, 2009. Retrieved September 9, 2009.
External links
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