Oxygen difluoride

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Oxygen difluoride
Other names difluorine monoxide
fluorine monoxide
oxygen fluoride
hypofluorous anhydride
CAS number 7783-41-7
PubChem 24547
RTECS number RS2100000
Molecular formula OF2
Molar mass 53.9962 g/mol
Appearance colorless gas, pale yellow liquid when condensed
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 HFO
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

Oxygen difluoride is the chemical compound with the formula OF2. It has a V-shaped molecular structure like H2O, but has different properties. It is a strong oxidizer.

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. [2][3] 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]

It is a very powerful oxidizing agent. This is proved by the oxidation number of +2 for the oxygen atom. This oxidising number is quite unusual. Above 200 °C, OF2 decomposes to oxygen and fluorine.

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

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

OF2 + 2 SO2S2O5F2

Safety[change | change source]

OF2 is a dangerous chemical. This is because it is highly oxidising.

References[change | change source]

  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.

Other websites[change | change source]