Chemistry:Oxygen difluoride

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Oxygen difluoride
Structure and dimensions of the oxygen difluoride molecule
Space-filling model of the oxygen difluoride molecule
Names
Other names
  • Oxygen fluoride
  • Hypofluorous anhydride
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
EC Number
  • 231-996-7
RTECS number
  • RS2100000
UNII
Properties
OF
2
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 48.9 atm (at −58.0 °C or −72.4 °F or 215.2 K[lower-alpha 1])
Thermochemistry
43.3 J/mol K
246.98 J/mol K
24.5 kJ mol−1
42.5 kJ/mol
Hazards
GHS pictograms GHS03: OxidizingGHS04: Compressed GasGHS05: CorrosiveGHS06: Toxic
GHS Signal word danger
Template:PPhrases
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
  • 2.6 ppm (rat, 1 hour)
  • 1.5 ppm (mouse, 1 hour)
  • 26 ppm (dog, 1 hour)
  • 16 ppm (monkey, 1 hour)
[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
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Oxygen difluoride is a chemical compound with the formula OF
2
. As predicted by VSEPR theory, the molecule adopts a "bent" molecular geometry similar to that of water. However, 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.[5][6] The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product:

2 F
2
+ 2 NaOH → OF
2
+ 2 NaF + H
2
O

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, OF
2
decomposes to oxygen and fluorine via a radical mechanism.

OF
2
reacts with many metals to yield oxides and fluorides. Nonmetals also react: phosphorus reacts with OF
2
to form PF
5
and POF
3
; sulfur gives SO
2
and SF
4
; and unusually for a noble gas, xenon reacts (at elevated temperatures) yielding XeF
4
and xenon oxyfluorides.

Oxygen difluoride reacts very slowly with water to form hydrofluoric acid:

OF
2

(aq)
+ H
2
O
(l)
→ 2 HF
(aq)
+ O
2

(g)

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

OF
2
+ SO
2
→ SO
3
+ F
2

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

OF
2
+ 2 SO
2
→ S
2
O
5
F
2

Safety

Oxygen difluoride is considered an unsafe gas due to its oxidizing properties. Hydrofluoric acid produced by the hydrolysis of OF
2
with water is highly corrosive and toxic, capable of causing necrosis, leaching calcium from the bones and causing 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 OF
2
would be a solid at 30 K, the fictional alien lifeforms were described as endothermic, maintaining elevated body temperatures and liquid OF
2
blood by radiothermal heating.

Notes

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

References

External links