Oxygen difluoride

Oxygen difluoride
Names
Other names
oxygen fluoride
hypofluorous anhydride
Identifiers
7783-41-7 YesY
ChEBI CHEBI:30494 YesY
ChemSpider 22953 YesY
EC Number 231-996-7
Jmol interactive 3D Image
PubChem 24547
RTECS number RS2100000
Properties
OF2
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)
hydrolyzes[1]
Vapor pressure >1 atm (20°C)[2]
Thermochemistry
43.3 J/mol K
246.98 J/mol K
24.5 kJ mol−1
42.5 kJ/mol
Hazards
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]
US health exposure limits (NIOSH):
TWA 0.05 ppm (0.1 mg/m3)[2]
C 0.05 ppm (0.1 mg/m3)[2]
0.5 ppm[2]
Related compounds
Related compounds
 HFO
O2F2
NHF2
NF3
SCl2
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.

Preparation

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

Reactions

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

However, in the presence of UV radiation the products are sulfuryl fluoride, SO
2F
2, and pyrosulfuryl fluoride, S
2O
5F
2:

OF2 + 2 SO2 → S
2O
5F
2

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.

Safety

OF2 is a dangerous chemical, as is the case for any strongly oxidizing gas.

References

  1. ↑ http://www.chemyq.com/En/xz/xz1/2818mqnrv.htm
  2. 1 2 3 4 "NIOSH Pocket Guide to Chemical Hazards #0475". National Institute for Occupational Safety and Health (NIOSH).
  3. ↑ "Oxygen difluoride". Immediately Dangerous to Life and Health. 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 February 21, 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 February 21, 2013.

External links

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