3D model (JSmol)
|Molar mass||53.9962 g/mol|
|Appearance||colorless gas, pale yellow liquid when condensed|
|Melting point||−223.8 °C|
|Boiling point||−144.8 °C|
|Solubility in other solvents||68 mL gaseous OF2 in 1 L (0 °C)|
|Std enthalpy of
|24.5 kJ mol−1|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Preparation[change | change source]
Oxygen difluoride was first prepared in 1929. It was made by the electrolysis of molten potassium fluoride and hydrofluoric acid. The chemicals had a small amount of moisture.  Nowadays, it is prepared by the reaction of fluorine with a dilute aqueous solution of sodium hydroxide. Sodium fluoride is left as a side-product:
- 2 F2 + 2 NaOH → OF2 + 2 NaF + H2O
Reactions[change | change source]
OF2 reacts with many metals. It produces oxides and fluorides. Nonmetals also react with it. Phosphorus reacts with OF2 to form PF5 and POF3. Sulfur produces SO2 and SF4 with it. A noble gas, xenon also reacts with it. It produces XeF4 and xenon oxyfluorides.
Oxygen difluoride reacts very slowly with water to form hydrofluoric acid:
- OF2 (aq) + H2O (aq) → 2 HF (aq) + O2 (g)
Oxygen difluoride oxidizes sulfur dioxide to sulfur trioxide:
- OF2 + SO2 → SO3 + F2
- OF2 + 2 SO2 → S
Safety[change | change source]
OF2 is a dangerous chemical. This is because it is highly oxidising.
References[change | change source]
- Yost, D. M. "Oxygen Fluoride" Inorganic Syntheses, 1939 volume, 1, pages 109-111.
- Paul Lebeau; 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.