Lithium superoxide
| Identifiers | |
|---|---|
3D model (
JSmol ) |
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| Properties | |
| LiO2 | |
| Molar mass | 38.94 g·mol−1 |
| Density | g/cm3, solid[clarification needed] |
| Melting point | <25 °C (decomposes) |
| Related compounds | |
Other cations
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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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Lithium superoxide is an unstable
Structure
The LiO2 molecule is a misnomer: the bonds between lithium and oxygen are highly
There have been quite a few studies regarding the clusters formed by LiO2 molecules. The most common
Production and reactions
Lithium superoxide is extremely reactive because of the odd number of electrons present in the π* molecular orbital of the superoxide anion.[4] Matrix isolation techniques can produce pure samples of the compound, but they are only stable at 15-40 K.[3]
At higher (but still cryogenic) temperatures, lithium superoxide can be produced by
- Li2O2(f12) + 2 O3(g) → 2 LiO2(f12) + 2 O2(g)
The resulting product is only stable up to −35 °C.[5]
Alternatively, lithium electride dissolved in anhydrous ammonia will reduce oxygen gas to yield the same product:
- [Li+][e−](am) + O2(g) → [Li+][O−2](am)
Lithium superoxide is, however, only
- 2 O−2 + 2 NH3 → N2 + 2 H2O + 2 OH−
Unlike other known decompositions of LiO2, this reaction bypasses lithium peroxide.[6]
Occurrence
Like other superoxides, lithium superoxide is the product of a one-electron
In batteries
Lithium superoxide also appears at the
- Li+ + e− + O2 → LiO2
This product typically then reacts and proceed to form lithium peroxide, Li2O2
- 2 LiO2 → Li2O2 + O2
The mechanism for this last reaction has not been confirmed and developing a complete theory of the oxygen reduction process remains a theoretical challenge as of 2022[update].[9] Indeed, recent work suggests that LiO2 can be stabilized via a suitable cathode made of graphene with iridium nanoparticles.[10]
A significant challenge when investigating these batteries is finding an ideal solvent in which to perform these reactions; current candidates are ether- and amide-based, but these compounds readily react with the superoxide and decompose.[9] Nevertheless, lithium-air cells remain the focus of intense research, because of their large energy density—comparable to the internal combustion engine.[8]
In the atmosphere
Lithium superoxide can also form for extended periods of time in low-density, high-energy environments, such as the upper atmosphere. The
See also
References
- ISSN 0021-9606.
- ISSN 1932-7447.
- ^ PMID 20684589.
- ISSN 0022-3654.
- ISSN 1573-9171.
- S2CID 46818521.
- PMID 34903644.
- ^ PMID 26274072.
- ^ PMID 22681046.
- S2CID 4452883.
- ^ For arguments claiming (or assuming) similarity, see:
- Plane, John M. C.; Rajasekhar, B.; Bartolotti, Libero (1989). "Theoretical and experimental determination of the lithium and sodium superoxide bond dissociation energies". The Journal of Physical Chemistry. 93 (8). American Chemical Society (ACS): 3141–3145. ISSN 0022-3654.
- Plane, John M. C.; Rajasekhar, B. (June 1988). "A study of the reaction Li + O2 + M (M = N2, He) over the temperature range 267-1100 K by time-resolved laser-induced fluorescence of Li(22PJ-22S1/2)". The Journal of Physical Chemistry. 92 (13): 3884–3890. ISSN 0022-3654.
- Swider, William (1987). "Chemistry of mesospheric potassium and its different seasonal behavior as compared to sodium". Journal of Geophysical Research. 92 (D5): 5621. ISSN 0148-0227.
- Plane, John M. C.; Rajasekhar, B.; Bartolotti, Libero (1989). "Theoretical and experimental determination of the lithium and sodium superoxide bond dissociation energies". The Journal of Physical Chemistry. 93 (8). American Chemical Society (ACS): 3141–3145.