Dioxygen difluoride

Dioxygen difluoride is a compound of fluorine and oxygen with the molecular formula O₂F₂. It can exist as an orange-colored solid which melts into a red liquid at . It is an extremely strong oxidant and decomposes into oxygen and fluorine even at at a rate of 4% per its lifetime at room temperature is thus extremely short. Dioxygen difluoride reacts vigorously with nearly every chemical it encounters (including ordinary ice) leading to its onomatopoeic nickname "FOOF" (a play on its chemical structure and its explosive tendencies).

Preparation

Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17  mmHg (0.9–2.3 kPa) is optimal) to an electric discharge of 25–30  mA at 2.1–2.4  kV.
A similar method was used for the first synthesis by Otto Ruff in 1933. Another synthesis involves mixing and in a stainless steel vessel cooled to , followed by exposing the elements to bremsstrahlung for several hours. A third method requires heating a mix of fluorine and oxygen to , and then rapidly cooling it using liquid oxygen. All of these methods involve synthesis according to the equation

: + →

It also arises from the thermal decomposition of ozone difluoride:
: → +
or
: 2 → 2 +

Structure and properties

In , oxygen is assigned the unusual oxidation state of +1. In most of its other compounds, oxygen has an oxidation state of −2.

The structure of dioxygen difluoride resembles that of hydrogen peroxide, , in its large dihedral angle, which approaches 90° and C₂ symmetry. This geometry conforms with the predictions of VSEPR theory.

The bonding within dioxygen difluoride has been the subject of considerable speculation, particularly because of the very short O−O distance and the long O−F distances. The O−O bond length is within 2 pm of the 120.7 pm distance for the O=O double bond in the dioxygen molecule, . Several bonding systems have been proposed to explain this, including an O−O ''triple'' bond with O−F single bonds destabilised and lengthened by repulsion between the lone pairs on the fluorine atoms and the Ï€ orbitals of the O−O bond. Repulsion involving the fluorine lone pairs is also responsible for the long and weak covalent bonding in the fluorine molecule.
Computational chemistry indicates that dioxygen difluoride has an exceedingly high barrier to rotation of 81.17 kJ/mol around the O−O bond (in hydrogen peroxide the barrier is 29.45 kJ/mol), this is close to the O−F bond disassociation energy of 81.59 kJ/mol.

The ¹⁹F NMR chemical shift of dioxygen difluoride is 865 ppm, which is by far the highest chemical shift recorded for a fluorine nucleus, thus underlining the extraordinary electronic properties of this compound. Despite its instability, thermochemical data for have been compiled.

Reactivity

The compound readily decomposes into oxygen and fluorine. Even at a temperature of , 4% decomposes each day by this process:

: → +

The other main property of this unstable compound is its oxidizing power, although most experimental reactions have been conducted near . Several experiments with the compound resulted in a series of fires and explosions. Some of the compounds that produced violent reactions with include ethyl alcohol, methane, ammonia, and even water ice.

With and , it gives the corresponding dioxygenyl salts:

:2 + 2 → 2 +

Uses

The compound currently has no practical applications, but has been of theoretical interest. One laboratory used it to synthesize plutonium hexafluoride at unprecedentedly low temperatures, which was significant because previous methods for preparation needed temperatures so high that the plutonium hexafluoride created would decompose rapidly.

See also

* Chlorine trifluoride
* A. G. Streng

References

External links

*

{{fluorides

Fluorides
Nonmetal halides
Oxygen compounds
Oxidizing agents
Chalcohalides
Substances discovered in the 1930s