 | |||
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| Names | |||
|---|---|---|---|
| Preferred IUPAC name Dioxygen difluoride | |||
Systematic IUPAC name
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Other names
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| Identifiers | |||
3D model (JSmol) | |||
| Abbreviations | FOOF | ||
| ChEBI | |||
| ChemSpider |
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| 1570 | |||
| UNII | |||
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| 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) | |||
Std molar entropy(S⦵298) | 277.2 J/(mol·K) | ||
Std enthalpy of formation(ΔfH⦵298) | 19.2 kJ/mol | ||
Gibbs free energy(ΔfG⦵) | 58.2 kJ/mol | ||
| Related compounds | |||
Related compounds | |||
| Hazards | |||
| GHS labelling: | |||
    | |||
| Danger | |||
| NFPA 704 (fire diamond) | |||
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |||
Dioxygen difluoride is acompound offluorine andoxygen with themolecular formula O2F2. It can exist as an orange-red colored solid which melts into a red liquid at −163 °C (110 K). It is an extremely strongoxidant anddecomposes into oxygen and fluorine even at −160 °C (113 K) at a rate of 4% perday — its lifetime at room temperature is thus extremely short.[1] Dioxygen difluoride reacts vigorously with nearly every chemical it encounters (including ordinaryice) leading to itsonomatopoeic nicknameFOOF (a play on its chemical structure and its explosive tendencies).[2]
Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17 mmHg (0.9–2.3 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 byOtto Ruff in 1933.[4] Another synthesis involves mixingO
2 andF
2 in astainless steel vessel cooled to −196 °C (77.1 K), followed by exposing the elements to3 MeVbremsstrahlung 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 usingliquid oxygen.[5] All of these methods involve synthesis according to the equation
It also arises from thethermal decomposition ofozone difluoride:[6]
InO
2F
2, oxygen is assigned the unusualoxidation state of +1. In most of its other compounds, oxygen has an oxidation state of −2.
The structure of dioxygen difluoride resembles that ofhydrogen peroxide,H
2O
2, in its largedihedral angle, which approaches 90° and C2symmetry. This geometry conforms with the predictions ofVSEPR 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 120.7pm distance for the O=O double bond in thedioxygen molecule,O
2. Several bonding systems have been proposed to explain this, including an O−Otriple bond with O−F single bonds destabilised and lengthened by repulsion between thelone pairs on the fluorine atoms and theπ orbitals of the O−O bond.[7] Repulsion involving the fluorine lone pairs is also responsible for the long and weakcovalent 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]
The19F NMRchemical 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 forO
2F
2 have been compiled.[9]
The compound readilydecomposes into oxygen and fluorine. Even at a temperature of −160 °C (113 K), 4% decomposes each day[1] by this process:
The other main property of this unstable compound is itsoxidizing power, although most experimental reactions have been conducted near −100 °C (173 K).[10] Several experiments with the compound resulted in a series of fires and explosions. Some of the compounds that produced violent reactions withO
2F
2 includeethyl alcohol,methane,ammonia, and evenwater ice.[10]
WithBF
3 andPF
5, it gives the correspondingdioxygenyl salts:[1][11]
The compound currently has no practical applications, but has been of theoretical interest.Los Alamos National Laboratory used it to synthesizeplutonium 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]
