| Other names
|3D model (Jmol)||Interactive image|
|Molar mass||53.9962 g/mol|
|Appearance||colorless gas, pale yellow liquid when condensed|
|Density|| 1.90 g/cm3 (-224° C, liquid),|
1.719 g/cm3 (-183° C, liquid), 1.521 g/cm3 (liquid at −145 °C), 1.88 g/l (gas at room temperature)
|Melting point||−223.8 °C (−370.8 °F; 49.3 K)|
|Boiling point||−144.75 °C (−228.55 °F; 128.40 K)|
|Vapor pressure||>1 atm (20°C)|
|43.3 J/mol K|
|246.98 J/mol K|
Std enthalpy of
|24.5 kJ mol−1|
Gibbs free energy (ΔfG˚)
|Lethal dose or concentration (LD, LC):|
LC50 (median concentration)
| 2.6 ppm (rat, 1 hr)|
1.5 ppm (mouse, 1 hr)
26 ppm (dog, 1 hr)
16 ppm (monkey, 1 hr)
|US health exposure limits (NIOSH):|
|TWA 0.05 ppm (0.1 mg/m3)|
|C 0.05 ppm (0.1 mg/m3)|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|(what is ?)|
Oxygen difluoride is the chemical compound with the formula OF2. As predicted by VSEPR theory, the molecule adopts a "bent" molecular geometry similar to that of water, but it has very different properties, being a strong oxidizer.
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. The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product:
- 2 F2 + 2 NaOH → OF2 + 2 NaF + H2O
Its powerful oxidizing properties are suggested by the oxidation number of +2 for the oxygen atom instead of its normal -2. 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:
- OF2 (aq) + H2O (l) → 2 HF (aq) + O2 (g)
Oxygen difluoride oxidizes sulfur dioxide to sulfur trioxide:
- OF2 + SO2 → SO3 + F2
- OF2 + 2 SO2 → S
OF2 is a dangerous chemical, as is the case for any strongly oxidizing gas.
- "NIOSH Pocket Guide to Chemical Hazards #0475". National Institute for Occupational Safety and Health (NIOSH).
- "Oxygen difluoride". Immediately Dangerous to Life and Health. National Institute for Occupational Safety and Health (NIOSH).
- Lebeau, P.; Damiens, A. (1929). "Sur un nouveau mode de préparation du fluorure d'oxygène" [A new method of preparation of oxygen fluoride]. Comptes rendus hebdomadaires des séances de l’Académie des sciences (in French). 188: 1253–1255. Retrieved February 21, 2013.
- Lebeau, P.; Damiens, A. (1927). "Sur l'existence d'un composé oxygéné du fluor" [The existence of an oxygen compound of fluorine]. Comptes rendus hebdomadaires des séances de l’Académie des sciences (in French). 185: 652–654. Retrieved February 21, 2013.
- National Pollutant Inventory - Fluoride and compounds fact sheet
- WebBook page for OF2
- CDC - NIOSH Pocket Guide to Chemical Hazards