Dioxygen difluoride
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{{Chembox new
| Name = Dioxygen difluoride
| ImageFile = Dioxygen-difluoride-3D-balls.png
<!-- | ImageSize = 120px -->
| ImageName = Dioxygen difluoride
| ImageFile1 = Dioxygen-difluoride-3D-vdW.png
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| ImageName1 = Dioxygen difluoride
| OtherNames = dioxygen difluoride<br /> fluorine dioxide<br /> difluorine dioxide<br /> perfluoroperoxide
| Section1 = {{Chembox Identifiers
| CASOther = [7783-44-0]
}}
| Section2 = {{Chembox Properties
| Formula = F<sub>2</sub>O<sub>2</sub>
| MolarMass = 69.996 g mol<sup>−1</sup>
| Solvent = other solvents
| SolubleOther = decomp.
| MeltingPt = −154 °C
| BoilingPt = −57 °C (extrapolated)
}}
| Section8 = {{Chembox Related
| OtherCpds = [[Oxygen difluoride|OF<sub>2</sub>]],<br /> [[Chloryl fluoride|FClO<sub>2</sub>]]<br /> [[Disulfur dichloride|S<sub>2</sub>Cl<sub>2</sub>]]}}
}}
'''Dioxygen difluoride''' is a [[Chemical compound|compound]] with the [[chemical formula|formula]] O<sub>2</sub>F<sub>2</sub>. This yellow compound is a strong oxidant and decomposes into OF<sub>2</sub> and oxygen even at -160 °C (4% per day).<ref name=Holl>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.</ref>
==Preparation==
Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7-17 [[Torr|mmgHg]] is optimal){{Fact|date=May 2007}} to an electric discharge of 25-30 [[Ampere|mA]] at 2.1-2.4 [[Volt|kV]]. This is basically the reaction used for the first synthesis by [[Otto Ruff]] in 1933.<ref>{{cite journal
| title = Neue Sauerstofffluoride: O<sub>2</sub>F<sub>2</sub> und OF
| author = [[Otto Ruff]], W. Menzel
| volume = 211
| issue = 1-2
| pages = 204–208
| year = 1933)
| journal = Zeitschrift für anorganische und allgemeine Chemie
| doi = 10.1002/zaac.19332110122}}</ref>Another synthesis involves mixing O<sub>2</sub> and F<sub>2</sub> in a [[stainless steel]] vessel cooled to −196 °C, followed by exposing the elements to 3 [[Electronvolt|MeV]] [[bremsstrahlung]] for several hours.{{Fact|date=May 2007}}
==Structure and electronic description==
In O<sub>2</sub>F<sub>2</sub> 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<sub>2</sub>O<sub>2</sub>, 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 [[1 E-12 m|pm]] distance for the O=O double bond in [[Oxygen|dioxygen]], O<sub>2</sub>.
<center>[[Image:Dioxygen-difluoride-2D.png|300px|Dioxygen difluoride's structure]]</center>
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 pair]]s on the fluorine atoms and the [[pi bond|π-orbital]]s of the O-O bond.<ref>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</ref> Repulsion involving the fluorine lone pairs is also responsible for the long and weak [[covalent radius of fluorine|covalent bonding in the fluorine molecule]].
==Reactivity==
The overarching property of this unstable compound is its oxidizing power, despite the fact that all reactions must be conducted near −100 °C.<ref>Streng, A. G. "The Chemical Properties of Dioxygen Difluoride" ''Journal of the American Chemical Society'' '''85''' (10), 1380 - 1385. DOI:10.1021/ja00893a004</ref> With [[boron trifluoride|BF<sub>3</sub>]] and [[phosphorus pentafluoride|PF<sub>5</sub>]], it gives the corresponding [[dioxygenyl]] salts:<ref name=Holl/>
:2O<sub>2</sub>F<sub>2</sub> + 2PF<sub>5</sub> → 2[O<sub>2</sub><sup>+</sup>][[Hexafluorophosphate|PF<sub>6</sub><sup>−</sup>]] + F<sub>2</sub>
It converts uranium and plutonium oxides into the corresponding hexafluorides.<ref>Atwood, D. A. "Fluorine: Inorganic Chemistry" in Encyclopedia of Inorganic Chemistry 2006 John Wiley & Sons, New York. DOI:10.1002/0470862106.ia076</ref>
==References==
<references/>
==External links==
* [http://www.npi.gov.au/database/substance-info/profiles/44.html National Pollutant Inventory - Flouride and compounds fact sheet]
* [http://webbook.nist.gov/cgi/cbook.cgi?ID=C7783440 WebBook page for O<sub>2</sub>F<sub>2</sub>]
[[Category:Fluorides]]
[[Category:Oxygen compounds]]
[[ru:Диоксидифторид]]
[[zh:二氟化二氧]]