Dioxygen difluoride

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.{{cite journal| last1=Ruff | first1=O.| last2=Mensel | first2=W.| year = 1933

: + →

It also arises from the thermal decomposition of ozone difluoride: : 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 π orbitals of the O−O bond.{{cite journal| last1=Bridgeman |first1=A. J.

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 −160 °C, 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 −100 °C. 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:{{cite journal| last1 = Solomon | first1 = Irvine J.| first2 = Robert I. | last2 = Brabets| first3 = Roy K. | last3 = Uenishi| first4 = James N. | last4 = Keith| first5 = John M. | last5 = McDonough| year=1964| title=New Dioxygenyl Compounds| journal=Inorganic Chemistry

:2 + 2 → 2 +

Uses

The compound currently has no practical applications, but has been of theoretical interest. Los Alamos National 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.

References

References

1. (2001). "Inorganic Chemistry". [[Academic Press]].
2. Lowe, Derek. (2010-02-23). "Things I Won't Work With: Dioxygen Difluoride".
3. Kwasnik, W.. (1963). "Handbook of Preparative Inorganic Chemistry". Academic Press.
4. Mills, Thomas. (1991). "Direct synthesis of liquid-phase dioxygen difluoride". Journal of Fluorine Chemistry.
5. (1959). "Ozone Fluoride or Trioxygen Difluoride, {{chem". [[Journal of the American Chemical Society]].
6. (2001). "Quantum Chemical Descriptions of FOOF: The Unsolved Problem of Predicting Its Equilibrium Geometry". The Journal of Physical Chemistry A.
7. Lyman, John L.. (1989). "Thermodynamic Properties of Dioxygen Difluoride (O₂F₂) and Dioxygen Fluoride (O₂F)". American Chemical Society and the American Institute of Physics for the National Institute of Standards and Technology.
8. Streng, A. G.. (1963). "The Chemical Properties of Dioxygen Difluoride". [[Journal of the American Chemical Society]].
9. (1984). "Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride". [[Journal of the American Chemical Society]].