Lithium oxide

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Lithium oxide
Lithium-oxide-unit-cell-3D-balls-B.png
Lithium-oxide-unit-cell-3D-ionic.png
CaF2 polyhedra.png
__ Li+     __ O2−
Li2O.jpg
Names
IUPAC name
Lithium oxide
Other names
Lithia, Kickerite
Identifiers
CAS Number
  • 12057-24-8 checkY
3D model (JSmol)
ChemSpider
ECHA InfoCard 100.031.823 Edit this at Wikidata
RTECS number
  • OJ6360000
UNII
CompTox Dashboard (EPA)
InChI
  • InChI=1S/2Li.O/q2*+1;-2 checkY
    Key: FUJCRWPEOMXPAD-UHFFFAOYSA-N checkY
  • InChI=1S/2Li.O/q2*+1;-2
    Key: FUJCRWPEOMXPAD-UHFFFAOYAW
  • Key: FUJCRWPEOMXPAD-UHFFFAOYSA-N
  • [Li+].[Li+].[O-2]
Properties
Chemical formula
 Li
2O
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
Refractive index (nD)
 1.644 [1]
Structure
Crystal structure
 Antifluorite (cubic), cF12
Space group
 Fm3m, No. 225
Tetrahedral (Li+); cubic (O2−)
Thermochemistry
Heat capacity (C)
 1.8105 J/g K or 54.1 J/mol K
37.89 J/mol K
Std enthalpy of
formation (ΔfH⦵298)
 -20.01 kJ/g or -595.8 kJ/mol
Gibbs free energy (ΔfG˚)
 -562.1 kJ/mol
Hazards
Main hazards Corrosive, reacts violently with water
NFPA 704 (fire diamond)
3
0
1
W
Flash point Non-flammable
Related compounds
Other anions
 Lithium sulfide
Lithium selenide
Lithium telluride
Lithium polonide
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).
checkY  (what is checkY☒N ?)
Infobox references

Lithium oxide (Li
2O) 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[]

Burning lithium metal produces lithium oxide.

Lithium oxide is produced by thermal decomposition of lithium peroxide at 300-400 °C.[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
2O.

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

2Li
2O
2 → 2Li
2O + O
2

Structure[]

Solid lithium oxide adopts an antifluorite structure with four-coordinated Li+ centers and eight-coordinated oxides.[4]

The ground state gas phase Li
2O molecule is linear with a bond length consistent with strong ionic bonding.[5][6] 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.

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

See also[]

References[]

  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 978-0-08-022057-4.
  4. ^ E. Zintl; A. Harder; B. Dauth (1934). "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums". . 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[]

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