Dioxygen difluoride

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Dioxygen difluoride
Preferred IUPAC name

Dioxygen difluoride

Systematic name

Fluorooxy hypofluorite

Other names

Monofluorooxygenyl hypofluorite
Difluorine dioxide
Fluorine dioxide
Perfluoroperoxide
Fluorine peroxide
Difluorine peroxide
FOOF
Fluoroperoxyl fluoride

Identifiers
Abbreviations FOOF
CAS number 7783-44-0 YesY
PubChem 123257 YesY
ChemSpider 109870 YesY
ChEBI CHEBI:47866 YesY
Gmelin Reference 1570
Jmol-3D images Image 1
Properties
Molecular formula O
2F
2
Molar mass 69.996 g·mol−1
Melting point −154 °C; −245 °F; 119 K
Boiling point −57 °C; −71 °F; 216 K (extrapolated)
Solubility in other solvents decomp.
Related compounds
Related compounds O
3F
2
H
2O
2
OF
2
FClO
2
Cl
2O
2
S
2Cl
2
S
2F
2
 YesY (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references

Dioxygen difluoride is a compound of fluorine and oxygen with the molecular formula O
2F
2. It exists as an orange solid that melts into a red liquid at −163 °C. It is an extremely strong oxidant and decomposes into OF
2 and oxygen even at −160 °C (4% per day); its lifetime at room temperature is thus extremely short.[1] Dioxygen difluoride reacts with nearly every chemical it encounters – even ordinary ice – leading to its onomatopoeic nickname "FOOF" (a play on its chemical formula).

The material has no practical applications but has been of theoretical interest. One laboratory use of it was the synthesis of plutonium hexafluoride at unprecedentedly low temperatures, which was significant because previous methods for its preparation needed temperatures so high that the plutonium hexafluoride created would rapidly decompose.[2]

Preparation

Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17 mmHg is optimal) to an electric discharge of 25–30 mA at 2.1–2.4 kV.[3] A similar method was used for the first synthesis by Otto Ruff in 1933.[4] Another synthesis involves mixing O2 and F2 in a stainless steel vessel cooled to −196 °C, followed by exposing the elements to 3 MeV bremsstrahlung for several hours. A third method requires heating a mix of fluorine and oxygen to 700 °C (1,292 °F), and then rapidly cooling it using liquid oxygen.[5] All of these methods involve synthesis according to the equation:

O
2 + F
2 → O
2F
2

It also arises from the thermal decomposition of ozone difluoride:[6]

O3F2 → O2F2 + 1/2 O2

Structure and properties

In O
2F
2, oxygen is assigned the unusual oxidation state of +1. In most of its other compounds, oxygen has an oxidation state of −2.

The structure of dioxygen difluoride resembles that of hydrogen peroxide, H
2O
2, in its large dihedral angle, which approaches 90°. This geometry conforms with the predictions of VSEPR theory. The O−O bond length is within 2 pm of the 120.7 pm distance for the O=O double bond in dioxygen, O2.

The bonding within dioxygen difluoride has been the subject of considerable speculation over the years, particularly because of the very short O–O distance and the long O–F distances. Bridgeman and Rothery have proposed a scheme which essentially has an O–O triple bond and an O–F single bond that is destabilised and lengthened by repulsion between the lone pairs on the fluorine atoms and the π-orbitals of the O–O bond.[7] Repulsion involving the fluorine lone pairs is also responsible for the long and weak covalent bonding in the fluorine molecule.

The 19F NMR chemical shift of dioxygen difluoride is 865 ppm, which is by far the highest chemical shift recorded for a fluorine nucleus, thus underlining the extraordinary electronic properties of this compound. Despite its instability, thermochemical data for O
2F
2 have been compiled.[8]

Reactivity

The compound readily decomposes into oxygen difluoride and oxygen. Even at a temperature of −160 °C, 4% decomposes each day[1] by this process:

2 O
2F
2 → 2 OF
2 + O
2

The other main property of this unstable compound is its oxidizing power, despite the fact that all reactions are conducted near −100 °C.[9] Several series of experiments with the compound resulted in a series of fires and explosions. Some of the compounds that produced violent reactions with O2F2 include ethyl alcohol, methane, ammonia, and even with water ice.[9]

With BF3 and PF5, it gives the corresponding dioxygenyl salts:[1][10]

2 O
2F
2 + 2 PF
5 → 2 [O
2]+
[PF
6]−
 + F
2

See also

References

  1. ↑ 1.0 1.1 1.2 Holleman, A. F.; Wiberg, E. (2001). Inorganic Chemistry. Academic Press. ISBN 0-12-352651-5. 
  2. ↑ Malm, J. G., P. G. Eller, and L. B. Asprey. "Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride." Journal of the American Chemical Society 106, no. 9 (1984): 2726-2727.
  3. ↑ W. Kwasnik "Dioxygen Difluoride" in Handbook of Preparative Inorganic Chemistry, 2nd Ed. Edited by G. Brauer, Academic Press, 1963, NY. Vol. 1. p. 162.
  4. ↑ Ruff, O.; Mensel, W. (1933). "Neue Sauerstofffluoride: O
    2    F
    2     und OF". Zeitschrift für anorganische und allgemeine Chemie 211 (1–2): 204–208. doi:10.1002/zaac.19332110122.     
  5. ↑ Thomas Mills (1991). "Direct synthesis of liquid-phase dioxygen difluoride". Journal of Fluorine Chemistry 52 (3): 267–276. doi:10.1016/S0022-1139(00)80341-3. 
  6. ↑ Kirshenbaum, A. D.; Grosse, A. V. (1959). "Ozone Fluoride or Trioxygen Difluoride, O
    3    F
    2    ". Journal of the American Chemical Society 81 (6): 1277. doi:10.1021/ja01515a003.     
  7. ↑ Bridgeman, A. J.; Rothery, J. (1999). "Bonding in mixed halogen and hydrogen peroxides". Journal of the Chemical Society, Dalton Transactions 1999 (22): 4077–4082. doi:10.1039/a904968a. 
  8. ↑ Lyman, John L. (1989). Thermodynamic Properties of Dioxygen Difluoride (O2F2) and Dioxygen Fluoride (O2F). American Chemical Society and the American Institute of Physics for the National Institute of Standards and Technology. Retrieved 5 August 2013. 
  9. ↑ 9.0 9.1 Streng, A. G. (1963). "The Chemical Properties of Dioxygen Difluoride". Journal of the American Chemical Society 85 (10): 1380–1385. doi:10.1021/ja00893a004. 
  10. ↑ Solomon, I. J.; et al. (1964). "New Dioxygenyl Compounds". Inorganic Chemistry 3 (3): 457. doi:10.1021/ic50013a036. 

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