Lithium oxide

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Lithium oxide
Lithium-oxide-unit-cell-3D-balls-B.png
Names
IUPAC name
Lithium oxide
Other names
Kickerite
Identifiers
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.031.823
RTECS number OJ6360000
Properties
Li
2
O
Molar mass 29.88 g/mol
Appearance white solid
Density 2.013 g/cm3
Melting point 1,438 °C (2,620 °F; 1,711 K)
Boiling point 2,600 °C (4,710 °F; 2,870 K)
reacts violently to form LiOH
log P 9.23
1.644 [1]
Structure
Antifluorite (cubic), cF12
Fm3m, No. 225
Tetrahedral (Li+); cubic (O2−)
Thermochemistry
1.8105 J/g K or 54.1 J/mol K
37.89 J/mol K
-20.01 kJ/g or -595.8 kJ/mol
-562.1 kJ/mol
Hazards
Main hazards Corrosive, reacts violently with water
NFPA 704
Flammability code 0: Will not burn. E.g., waterHealth code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gasReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calciumSpecial hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
0
3
1
Flash point Non-flammable
Related compounds
Other anions
Lithium sulfide
Other cations
Sodium oxide
Potassium oxide
Rubidium oxide
Caesium oxide
Related lithium oxides
Lithium peroxide
Lithium superoxide
Related compounds
Lithium hydroxide
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 oxide (Li
2
O) or lithia is an inorganic chemical compound. It is a white solid. Although not specifically important, many materials are assessed on the basis of their Li2O content. For example, the Li2O content of the principal lithium mineral spodumene (LiAlSi2O6) is 8.03%.[2]

Production[edit]

Lithium oxide is produced by thermal dehydration of lithium hydroxide.[2]

Lithium oxide forms along with small amounts of lithium peroxide when lithium metal is burned in the air and combines with oxygen:[3]

4Li + O
2
→ 2Li
2
O
.

Pure Li
2
O
can be produced by the thermal decomposition of lithium peroxide, Li
2
O
2
, at 450 °C[3]

2Li
2
O
2
→ 2Li
2
O
+ O
2

Structure[edit]

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.[4]

The ground state gas phase Li
2
O
molecule is linear with a bond length consistent with strong ionic bonding.[5][6] VSEPR theory would predict a bent shape similar to H
2
O
.

Uses[edit]

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.

Lithium metal might be obtained from lithium oxide by electrolysis, releasing oxygen as by-product.

See also[edit]

References[edit]

  1. ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
  2. ^ a b Wietelmann, Ulrich and Bauer, Richard J. (2005) "Lithium and Lithium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH: Weinheim. doi:10.1002/14356007.a15_393.
  3. ^ a b Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 97–99. ISBN 0-08-022057-6.
  4. ^ E. Zintl; A. Harder; B. Dauth (1934). "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums". Zeitschrift für Elektrochemie und Angewandte Physikalische Chemie. 40: 588–93.
  5. ^ Wells A.F. (1984) Structural Inorganic Chemistry 5th edition Oxford Science Publications ISBN 0-19-855370-6
  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

External links[edit]