Dioxygen difluoride

Dioxygen difluoride
Names
Preferred IUPAC name
Dioxygen difluoride
Systematic IUPAC name
Fluorooxy hypofluorite
Other names
  • Monofluorooxygenyl hypofluorite
  • Difluorine dioxide
  • Fluorine dioxide
  • Perfluoroperoxide
  • Fluorine peroxide
  • Difluorine peroxide
  • FOOF
  • Fluoroperoxyl fluoride
Identifiers
Abbreviations FOOF
7783-44-0 Yes
ChEBI CHEBI:47866 Yes
ChemSpider 109870 Yes
1570
Jmol-3D images Image
PubChem 123257
Properties
O
2F
2
Molar mass 69.996 g·mol−1
Appearance orange as a solid
Density 1.45 g/cm3 (at b.p.)
Melting point −154 °C (−245 °F; 119 K)
Boiling point −57 °C (−71 °F; 216 K) extrapolated
Solubility in other solvents decomposes
Thermochemistry
Specific
heat capacity (C)
 62.1 J/mol K
277.2 J/mol K
Std enthalpy of
formation (ΔfHo298)
 19.2 kJ/mol
58.2 kJ/mol
Related compounds
Related compounds
Except where noted otherwise, data is given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
 Yes verify (what is: Yes/?)
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 (110 K). It is an extremely strong oxidant and decomposes into oxygen and fluorine even at −160 °C (113 K) at a rate of 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's 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 O
2 and F
2 in a stainless steel vessel cooled to −196 °C (77.1 K), 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]

O
3F
2 → O
2F
2 + ½ O
2

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° and C2 symmetry. This geometry conforms with the predictions of VSEPR theory.

The bonding within dioxygen difluoride has been the subject of considerable speculation, particularly because of the very short O–O distance and the long O–F distances. The O−O bond length is within 2 pm of the 120.7 pm distance for the O=O double bond in dioxygen, O
2. Several bonding systems have been proposed to explain this, including a O–O triple bond with O–F single bonds which are 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. Computational chemistry indicates that dioxygen difluoride has an exceedingly high barrier to rotation of 81.17 kJ/mol around the O-O bond (in hydrogen peroxide its 29.45 kJ/mol), this is close to the O-F bond disassociation energy of 81.59 kJ/mol.[8]

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.[9]

Reactivity

The compound readily decomposes into oxygen and fluorine. Even at a temperature of −160 °C (113 K), 4% decomposes each day[1] by this process:

O
2F
2 → O
2 + F
2

The other main property of this unstable compound is its oxidizing power, despite the fact that all reactions are conducted near −100 °C (173 K).[10] Several series of experiments with the compound resulted in a series of fires and explosions. Some of the compounds that produced violent reactions with O
2F
2 include ethyl alcohol, methane, ammonia, and even with water ice.[10]

With BF
3 and PF
5, it gives the corresponding dioxygenyl salts:[1][11]

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.; Eller, P. G.; Asprey, L. B. (1984). "Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride". Journal of the American Chemical Society 106 (9): 2726–2727. doi:10.1021/ja00321a056.
  3. ↑ Kwasnik, W. (1963). "Dioxygen Difluoride". In Brauer, G. Handbook of Preparative Inorganic Chemistry 1 (2nd ed.). NY: Academic Press. 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. ↑ Mills, Thomas (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. ↑ Kraka, Elfi; He, Yuan; Cremer, Dieter. "Quantum Chemical Descriptions of FOOF: The Unsolved Problem of Predicting Its Equilibrium Geometry". The Journal of Physical Chemistry A 105 (13): 3269–3276. doi:10.1021/jp002852r.
  9. ↑ Lyman, John L. (1989). "Thermodynamic Properties of Dioxygen Difluoride (O2F2) and Dioxygen Fluoride (O2F)" (PDF). American Chemical Society and the American Institute of Physics for the National Institute of Standards and Technology. Retrieved 5 August 2013.
  10. ↑ 10.0 10.1 Streng, A. G. (1963). "The Chemical Properties of Dioxygen Difluoride". Journal of the American Chemical Society 85 (10): 1380–1385. doi:10.1021/ja00893a004.
  11. ↑ Solomon, Irvine J.; Brabets, Robert I.; Uenishi, Roy K.; Keith, James N.; McDonough, John M. (1964). "New Dioxygenyl Compounds". Inorganic Chemistry 3 (3): 457. doi:10.1021/ic50013a036.

External links