Oxygen difluoride

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. The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product: :

Structure and bonding

It is a covalently bonded molecule with a bent molecular geometry and a F-O-F bond angle of 103 degrees. Its powerful oxidizing properties are suggested by the oxidation number of +2 for the oxygen atom instead of its normal −2.

Reactions

Above 200 °C, decomposes to oxygen and fluorine by a radical mechanism. :

reacts with many metals to yield oxides and fluorides. Nonmetals also react: phosphorus reacts with to form and ; sulfur gives and ; and unusually for a noble gas, xenon reacts (at elevated temperatures) yielding and xenon oxyfluorides.

Reactions of oxygen difluoride and hydrogen halides or halide salts produce the free halogen. For example:Oxygen difluoride reacts with water to form hydrofluoric acid: :

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

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

Safety

Oxygen difluoride is considered an unsafe gas due to its oxidizing properties. It reacts explosively with water, hydrogen sulfide, diborane, and nitrogen oxides. Hydrofluoric acid produced by the hydrolysis of 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 highly toxic effects. Other acute poisoning effects include: pulmonary edema, bleeding lungs, headaches, etc. Chronic exposure to oxygen difluoride, like that of other chemicals that release fluoride ions, can lead to fluorosis and other symptoms of chronic fluoride poisoning. Oxygen difluoride may be associated with kidney damage. The maximum workplace exposure limit is 0.05 ppm.

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 30K, the fictional alien lifeforms were described as endothermic, maintaining elevated body temperatures and liquid blood by radiothermal heating.

Notes

References

References

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