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Oxygen difluoride
CAS number 7783-41-7
PubChem 24547
RTECS number RS2100000
Molecular formula OF2
Molar mass 53.9962 g/mol
Density 1.9 g/cm3
Melting point

−223.8 °C

Boiling point

−144.8 °C

Solubility in other solvents 68 mL gaseous OF2 in 1 L (0 °C)[1]
Std enthalpy of
24.5 kJ mol−1
Related compounds
Related compounds O2F2
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
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.



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:

2 F2 + 2 NaOH → OF2 + 2 NaF + H2O


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, yielding 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

However, in the presence of UV radiation the products are sulfuryl fluoride, SO2F2, and pyrosulfuryl fluoride, S2O5F2:

OF2 + 2SO2 + UV radiation → S2O5F2

Popular culture

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.


  1. ^ Yost, D. M. "Oxygen Fluoride" Inorganic Syntheses, 1939 volume, 1, pages 109-111.
  2. ^ Paul Lebeau; Damiens, A. "A New Method for the Preparation of the Fluorine Oxide”Compt. rend. 1929, volume 188, 1253-5.
  3. ^ Lebeau, P.; Damiens, A. "The Existence of an Oxygen Compound of Fluorine"Compt. rend. 1927, volume 185, pages 652-4.

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