| Oxygen difluoride | |
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Other names
difluorine monoxide |
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| Identifiers | |
| CAS number | 7783-41-7 |
| PubChem | 24547 |
| ChemSpider | 22953 |
| ChEBI | CHEBI:30494 |
| RTECS number | RS2100000 |
| Jmol-3D images | Image 1 |
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| Properties | |
| Molecular formula | OF2 |
| Molar mass | 53.9962 g/mol |
| Appearance | colorless gas, pale yellow liquid when condensed |
| Density | 1.9 g/cm3 as liquid at -145°C |
| Melting point |
−223.8 °C |
| Boiling point |
−144.8 °C |
| Solubility in other solvents | 68 mL gaseous OF2 in 1 L (0 °C)[1] |
| Thermochemistry | |
| Std enthalpy of formation ΔfH |
24.5 kJ mol−1 |
| Related compounds | |
| Related compounds | HFO O2F2 NHF2 NF3 SCl2 |
| (verify) (what is: /?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
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| Infobox references | |
Oxygen difluoride is the chemical compound with the formula OF2. As predicted by VSEPR theory, the molecule adopts a "V" shaped structure like H2O, but it has very different properties, being a strong oxidizer.
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Oxygen difluoride was first reported in 1929; it was obtained by the electrolysis of molten potassium fluoride and hydrofluoric acid containing small quantities of water.[2][3] The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product:
Its powerful oxidizing properties are suggested by the oxidation number of +2 for the oxygen atom, which is unusual. Above 200 °C, OF2 decomposes to oxygen and fluorine via a radical mechanism.
OF2 reacts with many metals to yield oxides and fluorides. Nonmetals also react: phosphorus reacts with OF2 to form PF5 and POF3; sulfur gives SO2 and SF4; and unusually for a noble gas, xenon reacts, at elevated temperatures, yielding XeF4 and xenon oxyfluorides.
Oxygen difluoride reacts very slowly with water to form hydrofluoric acid:
Oxygen difluoride oxidizes sulfur dioxide to sulfur trioxide:
However, in the presence of UV radiation the products are sulfuryl fluoride, SO2F2, and pyrosulfuryl fluoride, S2O5F2:
In Robert L. Forward's science fiction novel Camelot 30K, oxygen difluoride was used as a biochemical solvent by fictional life forms living in the solar system's Kuiper belt.
OF2 is a dangerous chemical, as is the case for any strongly oxidizing gas.