Dioxygen difluoride
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| Dioxygen difluoride | |
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| Other names | dioxygen difluoride fluorine dioxide difluorine dioxide perfluoroperoxide |
| Identifiers | |
| CAS number | [7783-44-0] |
| Properties | |
| Molecular formula | F2O2 |
| Molar mass | 69.996 g mol−1 |
| Melting point |
−154 °C |
| Boiling point |
−57 °C (extrapolated) |
| Solubility in other solvents | decomp. |
| Related compounds | |
| Related compounds | OF2, FClO2 S2Cl2 |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Dioxygen difluoride is a compound with the formula O2F2. This yellow compound is a strong oxidant and decomposes into OF2 and oxygen even at -160 °C (4% per day).[1]
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[edit] Preparation
Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7-17 mmgHg is optimal)[citation needed] to an electric discharge of 25-30 mA at 2.1-2.4 kV. This is basically the reaction used for the first synthesis by Otto Ruff in 1933.[2]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.[citation needed]
[edit] Structure and electronic description
In O2F2 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, H2O2, 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 has 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.[3] Repulsion involving the fluorine lone pairs is also responsible for the long and weak covalent bonding in the fluorine molecule.
[edit] Reactivity
The overarching property of this unstable compound is its oxidizing power, despite the fact that all reactions must be conducted near −100 °C.[4] With BF3 and PF5, it gives the corresponding dioxygenyl salts:[1]
- 2O2F2 + 2PF5 → 2[O2+]PF6− + F2
It converts uranium and plutonium oxides into the corresponding hexafluorides.[5]
[edit] References
- ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Otto Ruff, W. Menzel (1933)). "Neue Sauerstofffluoride: O2F2 und OF". Zeitschrift für anorganische und allgemeine Chemie 211 (1-2): 204–208. doi:.
- ^ Bridgeman, A. J., Rothery, J. "Bonding in mixed halogen and hydrogen peroxides" Journal of the Chemical Society, Dalton Transactions, 1999, 4077-4082. DOI:10.1039/a904968a
- ^ Streng, A. G. "The Chemical Properties of Dioxygen Difluoride" Journal of the American Chemical Society 85 (10), 1380 - 1385. DOI:10.1021/ja00893a004
- ^ Atwood, D. A. "Fluorine: Inorganic Chemistry" in Encyclopedia of Inorganic Chemistry 2006 John Wiley & Sons, New York. DOI:10.1002/0470862106.ia076
