Dioxygen difluoride

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Dioxygen difluoride
Stick model of dioxygen difluoride
Stick model of dioxygen difluoride
Spacefill model of dioxygen difluoride
Spacefill model of dioxygen difluoride
Names
Preferred IUPAC name
Dioxygen difluoride
Systematic IUPAC name
Fluorooxygen hypofluorite
Other names
  • Monofluorooxygenyl hypofluorite
  • Peroxydifluoride
  • Oxygen perfluoride
  • Oxyfluoryl hypofluorite
  • Fluorine peroxide
  • FOOF
Identifiers
3D model (
JSmol
)
Abbreviations FOOF
ChEBI
ChemSpider
1570
UNII
  • InChI=1S/F2O2/c1-3-4-2 checkY
    Key: REAOZOPEJGPVCB-UHFFFAOYSA-N checkY
  • InChI=1/F2O2/c1-3-4-2
    Key: REAOZOPEJGPVCB-UHFFFAOYAK
  • FOOF
Properties
O
2F
2
Molar mass 69.996 g·mol−1
Appearance orange as a solid
red as a liquid
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
62.1 J/(mol·K)
277.2 J/(mol·K)
Std enthalpy of
formation
(ΔfH⦵298)
 19.2 kJ/mol
58.2 kJ/mol
Related compounds
Related compounds
Hazards
GHS labelling:
GHS03: OxidizingGHS05: CorrosiveGHS06: ToxicGHS09: Environmental hazard
Danger
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 4: Readily capable of detonation or explosive decomposition at normal temperatures and pressures. E.g. nitroglycerinSpecial hazard OX: Oxidizer. E.g. potassium perchlorate
4
0
4
OX
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Dioxygen difluoride is a

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 vigorously with nearly every chemical it encounters (including ordinary ice) leading to its onomatopoeic nickname FOOF (a play on its chemical structure and its explosive tendencies).[2]

Preparation

Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17 

kPa) 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

ozone difluoride:[6]

2 O
3F
2 → 2 O
2F
2 + 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.

Dioxygen difluoride's structure

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

dioxygen molecule, O
2. Several bonding systems have been proposed to explain this, including an O−O triple bond with O−F single bonds 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 the barrier is 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

ethyl alcohol, methane, ammonia, and even 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

Uses

The compound currently has no practical applications, but has been of theoretical interest. One laboratory used it to synthesize plutonium hexafluoride at unprecedentedly low temperatures, which was significant because previous methods for preparation needed temperatures so high that the plutonium hexafluoride created would decompose rapidly.[12]

See also

References

  1. ^
    ISBN 978-0-12-352651-9
    .
  2. ^ Lowe, Derek (2010-02-23). "Things I Won't Work With: Dioxygen Difluoride". www.science.org. Retrieved 2022-05-26.
  3. ^ Kwasnik, W. (1963). "Dioxygen Difluoride". In Brauer, G. (ed.). Handbook of Preparative Inorganic Chemistry. Vol. 1 (2nd ed.). NY: Academic Press. p. 162.
  4. doi:10.1002/zaac.19332110122
    .
  5. doi:10.1016/S0022-1139(00)80341-3
    .
  6. doi:10.1021/ja01515a003
    .
  7. doi:10.1039/a904968a
    .
  8. doi:10.1021/jp002852r
    .
  9. doi:10.1063/1.555830. Archived from the original
    (PDF) on 4 March 2016. Retrieved 5 August 2013.
  10. ^
    doi:10.1021/ja00893a004
    .
  11. doi:10.1021/ic50013a036
    .
  12. doi:10.1021/ja00321a056
    .

External links

  • Perfluoroperoxide in Linstrom, Peter J.; Mallard, William G. (eds.); NIST Chemistry WebBook, NIST Standard Reference Database Number 69, National Institute of Standards and Technology, Gaithersburg (MD)
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