Lithium superoxide
3D model (JSmol)
  • InChI=1S/Li.O2/c;1-2/q+1;-1
  • [Li+].O=[O-]
Molar mass 38.94 g·mol−1
Density g/cm3, solid[clarification needed]
Melting point <25 °C (decomposes)
Related compounds
Other cations
Related Lithium oxides
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 inorganic salt with formula LiO2. A radical compound, it can be produced at low temperature in matrix isolation experiments, or in certain nonpolar, non-protic solvents. Lithium superoxide is also a transient species during the reduction of oxygen in a lithium–air galvanic cell, and serves as a main constraint on possible solvents for such a battery. For this reason, it has been investigated thoroughly using a variety of methods, both theoretical and spectroscopic.


The LiO2 molecule is a misnomer: the bonds between lithium and oxygen are highly ionic, with almost complete electron-transfer.[1] The force constant between the two oxygen atoms matches the constants measured for the superoxide anion (O2) in other contexts. The bond length for the O-O bond was determined to be 1.34 Å. Using a simple crystal structure optimization, the Li-O bond was calculated to be approximately 2.10 Å.[2]

There have been quite a few studies regarding the clusters formed by LiO2 molecules. The most common dimer has been found to be the cage isomer. Second to it is the singlet bypyramidal structure. Studies have also been done on the chair complex and the planar ring, but these two are less favorable, though not necessarily impossible.[3]

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 ozonating lithium peroxide (Li2O2) in freon 12:

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+][O2](am)

Lithium superoxide is, however, only metastable in ammonia, gradually oxidizing the solvent to water and nitrogen gas:

2 O2 + 2 NH3 → N2 + 2 H2O + 2 OH

Unlike other known decompositions of LiO2, this reaction bypasses lithium peroxide.[6]


Like other superoxides, lithium superoxide is the product of a one-electron reduction of an oxygen molecule. It thus appears whenever oxygen is mixed with single-electron redox catalysts, such as p-benzoquinone.[7]

In batteries

Lithium superoxide also appears at the cathode of a lithium-air galvanic cell during discharge, as in the following reaction:[8]

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.[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 mesosphere contains a persistent layer of alkali metal cations ablated from meteors. For sodium and potassium, many of the ions bond to form particles of the corresponding superoxide. It is currently unclear whether lithium should react analogously.[11]

See also


  1. ^ Andrews, Lester (1969-05-15). "Infrared Spectrum, Structure, Vibrational Potential Function, and Bonding in the Lithium Superoxide Molecule LiO2". The Journal of Chemical Physics. AIP Publishing. 50 (10): 4288–4299. Bibcode:1969JChPh..50.4288A. doi:10.1063/1.1670893. ISSN 0021-9606.
  2. ^ Lau, Kah Chun; Curtiss, Larry A.; Greeley, Jeffrey (2011-11-09). "Density Functional Investigation of the Thermodynamic Stability of Lithium Oxide Bulk Crystalline Structures as a Function of Oxygen Pressure". The Journal of Physical Chemistry C. American Chemical Society (ACS). 115 (47): 23625–23633. doi:10.1021/jp206796h. ISSN 1932-7447.
  3. ^ a b Bryantsev, Vyacheslav S.; Blanco, Mario; Faglioni, Francesco (2010-07-16). "Stability of Lithium Superoxide LiO2 in the Gas Phase: Computational Study of Dimerization and Disproportionation Reactions". The Journal of Physical Chemistry A. American Chemical Society (ACS). 114 (31): 8165–8169. Bibcode:2010JPCA..114.8165B. doi:10.1021/jp1047584. ISSN 1089-5639. PMID 20684589.
  4. ^ Lindsay, D. M.; Garland, D. A. (1987). "ESR spectra of matrix-isolated lithium superoxide". The Journal of Physical Chemistry. American Chemical Society (ACS). 91 (24): 6158–6161. doi:10.1021/j100308a020. ISSN 0022-3654.
  5. ^ Vol'nov, I. I.; Tokareva, S. A.; Belevskii, V. N.; Klimanov, V. I. (1967-07-01). "Investigation of the nature of the interaction of lithium peroxide with ozone". Bulletin of the Academy of Sciences of the USSR, Division of Chemical Science. 16 (7): 1369–1371. doi:10.1007/BF00905329. ISSN 1573-9171.
  6. ^ Zhang, Xinmin; Guo, Limin; Gan, Linfeng; Zhang, Yantao; Wang, Jin; Johnson, Lee R.; Bruce, Peter G.; Peng, Zhangquan (2017-05-18). "LiO 2 : Cryosynthesis and Chemical/Electrochemical Reactivities". The Journal of Physical Chemistry Letters. 8 (10): 2334–2338. doi:10.1021/acs.jpclett.7b00680. ISSN 1948-7185. PMID 28481552. S2CID 46818521.
  7. ^ Nava, Matthew; Zhang, Shiyu; Pastore, Katharine S.; Feng, Xiaowen; Lancaster, Kyle M.; Nocera, Daniel G.; Cummins, Christopher C. (2021-12-13). "Lithium superoxide encapsulated in a benzoquinone anion matrix". Proceedings of the National Academy of Sciences. 118 (51). Bibcode:2021PNAS..11819392N. doi:10.1073/pnas.2019392118. ISSN 0027-8424. PMC 8713792. PMID 34903644.
  8. ^ a b Das, Ujjal; Lau, Kah Chun; Redfern, Paul C.; Curtiss, Larry A. (2014-02-13). "Structure and Stability of Lithium Superoxide Clusters and Relevance to Li–O2 Batteries". The Journal of Physical Chemistry Letters. American Chemical Society (ACS). 5 (5): 813–819. doi:10.1021/jz500084e. ISSN 1948-7185. PMID 26274072.
  9. ^ a b Bryantsev, Vyacheslav S.; Faglioni, Francesco (2012-06-21). "Predicting Autoxidation Stability of Ether- and Amide-Based Electrolyte Solvents for Li–Air Batteries". The Journal of Physical Chemistry A. American Chemical Society (ACS). 116 (26): 7128–7138. Bibcode:2012JPCA..116.7128B. doi:10.1021/jp301537w. ISSN 1089-5639. PMID 22681046.
  10. ^ Lu, Jun (2016). "A lithium - oxygen battery based on lithium superoxide". Nature. 529 (7586): 377–381. Bibcode:2016Natur.529..377L. doi:10.1038/nature16484. PMID 26751057. S2CID 4452883.
  11. ^ For arguments claiming (or assuming) similarity, see: For an argument that the different photoionization rate of lithium should produce a dissimilar equilibrium, see: