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Oxygen difluoride
From Wikipedia, the free encyclopedia
| Oxygen difluoride | |
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| Systematic name | oxygen(II) fluoride |
| Other names | difluorine monoxide fluorine monoxide oxygen difluoride oxygen fluoride |
| Molecular formula | OF2 |
| Molar mass | 53.9962 g mol−1 |
| CAS number | [7783-41-7] |
| Density | 2.26 kg dm−3 |
| Solubility (in water) | 68 mL gaseous OF2 in 1 L (0 °C)[1] |
| Melting point | −224 °C |
| Boiling point | −145 °C |
| Color | red-brown (gas) pale yellow (liquid) |
| Heat of formation, ΔHf° | 24.5 kJ mol−1 |
| Average O−F bond energy | 187 kJ mol−1 |
| Related compounds | O2F2 NHF2 NF3 SCl2 |
| Disclaimer and references | |
Oxygen difluoride is the chemical compound with the formula OF2. It is a powerful oxidant and fluorinating agent. The oxygen atom has an oxidation number of +2, unlike almost all other oxygen compounds, where it is −2. The only exceptions are peroxides, with oxygen in the −1 oxidation state (e.g. hydrogen peroxide, H2O2), superoxides, with oxygen in the −½ oxidation state (e.g. sodium superoxide, NaO2) and dioxygen difluoride, O2F2, in which oxygen adopts an oxidation state of +1.
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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:
- 2F2 + 2NaOH → OF2 + 2NaF + H2O
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, yielding XeF4 and xenon oxyfluorides.
Oxygen difluoride reacts very slowly with water to form hydrofluoric acid:
- OF2(aq) + H2O(aq) → 2HF(aq) + O2(g)
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.
- ^ Yost, D. M. "Oxygen Fluoride" Inorganic Syntheses, 1939 volume, 1, pages 109-111.
- ^ Lebeau, P.; Damiens, A. "A New Method for the Preparation of the Fluorine Oxide” Compt. rend. 1929, volume 188, 1253-5.
- ^ Lebeau, P.; Damiens, A. "The Existence of an Oxygen Compound of Fluorine" Compt. rend. 1927, volume 185, pages 652-4.
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