Lithium superoxide

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.

Lithium superoxide

Identifiers
CAS Number	12136-56-0 Y
3D model (JSmol)	Interactive image
InChI InChI=1S/Li.O2/c;1-2/q+1;-1 Key: GMZGUKXTXROVJB-UHFFFAOYSA-N
SMILES [Li+].O=[O-]
Properties
Chemical formula	LiO2
Molar mass	38.94 g·mol−1
Density	g/cm3, solid[clarification needed]
Melting point	<25 °C (decomposes)
Related compounds
Other cations	Sodium superoxide Potassium superoxide Rubidium superoxide Caesium superoxide
Related Lithium oxides	Lithium oxide Lithium peroxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Y verify (what is YN ?) Infobox references

Structure

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 (O−2) 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⁺][O−2](am)

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

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 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^[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 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

• Lithium oxide
• Lithium peroxide

References

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. 50 (10). AIP Publishing: 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. 115 (47). American Chemical Society (ACS): 23625–23633. doi:10.1021/jp206796h. ISSN 1932-7447.
3. 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. 114 (31). American Chemical Society (ACS): 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. 91 (24). American Chemical Society (ACS): 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. 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. 5 (5). American Chemical Society (ACS): 813–819. doi:10.1021/jz500084e. ISSN 1948-7185. PMID 26274072.
9. 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. 116 (26). American Chemical Society (ACS): 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:

    • 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. doi:10.1021/j100345a052. 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. doi:10.1021/j100324a041. ISSN 0022-3654.

    For an argument that the different photoionization rate of lithium should produce a dissimilar equilibrium, see:

    • Swider, William (1987). "Chemistry of mesospheric potassium and its different seasonal behavior as compared to sodium". Journal of Geophysical Research. 92 (D5): 5621. Bibcode:1987JGR....92.5621S. doi:10.1029/jd092id05p05621. ISSN 0148-0227.