Oxygen Difluoride
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Oxygen Difluoride
Oxygen difluoride
Structure and dimensions of the oxygen difluoride molecule
Space-filling model of the oxygen difluoride molecule
Other names
oxygen fluoride
hypofluorous anhydride
3D model (JSmol)
ECHA InfoCard 100.029.087 Edit this at Wikidata
EC Number
  • 231-996-7
RTECS number
  • RS2100000
Molar mass 53.9962 g/mol
Appearance colorless gas, pale yellow liquid when condensed
Odor peculiar, foul
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 48.9 atm (at -58.0 °C or -72.4 °F or 215.2 K[a])
43.3 J/mol K
246.98 J/mol K
24.5 kJ mol-1
42.5 kJ/mol
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
2.6 ppm (rat, 1 hr)
1.5 ppm (mouse, 1 hr)
26 ppm (dog, 1 hr)
16 ppm (monkey, 1 hr)[3]
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 0.05 ppm (0.1 mg/m3)[2]
REL (Recommended)
C 0.05 ppm (0.1 mg/m3)[2]
IDLH (Immediate danger)
0.5 ppm[2]
Related compounds
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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.[4][5] 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)

It can oxidize sulphur dioxide to sulfur trioxide and elemental fluorine:

OF2 + SO2 -> SO3 + F2

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

OF2 + 2 SO2 ->


Oxygen difluoride is considered an unsafe gas due to its oxidizing properties. Hydrofluoric acid produced by the hydrolysis of OF2 with water is highly corrosive and toxic, capable of causing necrosis, leaching calcium from the bones and cardiovascular damage, among a host of other insidious effects.

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. While would be a solid at 30 K, the fictional alien lifeforms were described as endothermic, maintaining elevated body temperatures and liquid blood by radiothermal heating.


  1. ^ This is its critical temperature, which is below ordinary room temperature


  1. ^ "difluorine monoxide;oxygen difluoride,physical properties,suppliers,CAS,MSDS,structure,Molecular Formula, Molecular Weight ,Solubility,boiling point, melting point". www.chemyq.com.
  2. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0475". National Institute for Occupational Safety and Health (NIOSH).
  3. ^ "Oxygen difluoride". Immediately Dangerous to Life and Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  4. ^ 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 2013.
  5. ^ 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 2013.

External links

  This article uses material from the Wikipedia page available here. It is released under the Creative Commons Attribution-Share-Alike License 3.0.



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