Oxygen fluoride

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula O_nF₂, where n = 1 to 6. Many different oxygen fluorides are known:

oxygen difluoride (OF₂)
dioxygen difluoride (O₂F₂)
trioxygen difluoride or ozone difluoride (O₃F₂)[1][2]
tetraoxygen difluoride (O₄F₂)[3]
pentaoxygen difluoride (O₅F₂)
hexaoxygen difluoride (O₆F₂)[4]
dioxygen monofluoride of fluoroperoxyl (O₂F)

tetraoxygen difluoride

Oxygen difluoride

Oxygen fluorides are strong oxidizing agents with high energy and can release their energy either instantaneously or at a controlled rate. Thus, these compounds attracted much attention as potential fuels in jet propulsion systems.[5]

Synthesis

Here are some synthesis methods and reactions of the three most common oxygen fluorides – oxygen difluoride (OF₂), dioxygen difluoride (O₂F₂) and ozone difluoride (O₃F₂).

Oxygen difluoride (OF₂)

Oxygen difluoride

A common preparative method involves fluorination of sodium hydroxide:

{\ce {2F2 + 2NaOH -> OF2 + 2NaF + H2O}}

OF₂ is a colorless gas at room temperature and a yellow liquid below 128 K. Oxygen difluoride has an irritating odor and is poisonous.[3] It reacts quantitatively with aqueous haloacids to give free halogens:

{\ce {OF2 + 4HCl -> 2Cl2 + 2HF + 2H2O}}

It can also displace halogens from their salts.[3] It is both an effective fluorinating agent and a strong oxidizing agent. When reacted with unsaturated nitrogen fluorides with electrical discharge, it results in the formation of nitrogen trifluoride, oxide fluorides and other oxides.[6][7]

Dioxygen difluoride (O₂F₂)

Dioxygen difluoride

O₂F₂ precipitates as a brown solid upon the UV irradiation of a mixture of liquid O₂ and F₂ at -196 °C.[8] It also only appears to be stable below -160 °C.[9] The general method of preparation of many oxygen fluorides is a gas-phase electric discharge in cold containers including O₂F₂.[10]

{\ce {{O2}+F2->[{\text{Electric Discharge}}][{\text{183 °C}}]O2F2}}

It is typically an orange-yellow solid which rapidly decomposes to O₂ and F₂ close to its normal boiling point of about 216 K.[3]

O₂F₂ reacts violently with red phosphorus, even at -196 °C. Explosions can also occur if Freon-13 is used to moderate the reaction.[9]

Ozone difluoride (O₃F₂)

O₃F₂ is a viscous, blood-red liquid. It remains liquid at 90 K and so can be differentiated from O₂F₂ which has a melting point of about 109 K.[11][3]

Like the other oxygen fluorides, O₃F₂ is endothermic and decomposes at about 115 K with the evolution of heat, which is given by the following reaction:

{\ce {2 O3F2 -> O2 + 2O2F2}}

O₃F₂ is safer to work with than ozone, and can be evaporated, or thermally decomposed, or exposed to electric sparks, without any explosions. But on contact with organic matter or oxidizable compounds, it can detonate or explode. Thus, the addition of even one drop of ozone difluoride to solid anhydrous ammonia will result in a mild explosion, when they are both at 90 K each.[3]

Fluoroperoxyl

Fluoroperoxyl is a molecule such as O-O-F, whose chemical formula is O₂F and is stable only at low temperature. It has been reported to be produced from atomic fluorine and dioxygen.[12]

{\rm {O_{2}+F\longrightarrow O_{2}F}}

General preparation of polyoxygen difluorides

Reaction equation[6]	F₂:O₂ by volume	Current	Temperature of bath (°C)
${\ce {O2 + F2 <=> O2F2}}$	1:1	10–50 mA	~ -196°
${\ce {3O2 + 2F2 <=> 2O3F2}}$	2:3	25–30 mA	~ -196°
${\ce {2O2 + F2 <=> O4F2}}$	1:2	4 – 5 mA	~ -205°

Effects on ozone

Oxygen- and fluorine-containing radicals like FO₂ and FO occur in the atmosphere. These along with other halogen radicals have been implicated in the destruction of ozone in the atmosphere. However, the fluorine oxide radicals are assumed to not play as big a role in the ozone depletion because free fluorine atoms in the atmosphere are believed to react with methane to produce hydrofluoric acid which precipitates in rain. This decreases the availability of free fluorine atoms for oxygen atoms to react with and destroy ozone molecules.[13]

{\begin{array}{lr}&{\ce {F + O3 -> FO + O2}}\\&{\ce {FO + O -> F + O2}}\\\hline {\text{Net Reaction:}}&{\ce {O3 + O -> 2 O2}}\end{array}}

Hypergolic propellant

Despite the low solubility of O₃F₂ in liquid oxygen, it has been shown to be hypergolic with most rocket propellant fuels. The mechanism involves the boiling off oxygen from the solution containing O₃F₂, making it more reactive to have a spontaneous reaction with the rocket fuel. The degree of reactivity is also dependent on the type of fuel used.[3]

References

Solomon, I. J. et al. (1968). "Additional Studies Concerning the Existence of O₃F₂". Journal of the American Chemical Society. 90 (20): 5408–5411. doi:10.1021/ja01022a014.CS1 maint: uses authors parameter (link)
Misochko, Eugenii Ya, Alexander V. Akimov, Charles A. Wight (1999). "Infrared spectroscopic observation of the stabilized Intermediate complex FO₃ formed by reaction of mobile Fluorine atoms with ozone molecules Trapped in an Argon Matrix". The Journal of Physical Chemistry A. 103 (40): 7972–7977. Bibcode:1999JPCA..103.7972M. doi:10.1021/jp9921194.CS1 maint: uses authors parameter (link)
Streng, A. G. (1963). "The Oxygen Fluorides". Chemical Reviews. 63 (6): 607–624. doi:10.1021/cr60226a003.
Streng, A. G., A. V. Grosse (1966). "Two New Fluorides of Oxygen, O₅F₂ and O₆F₂". Journal of the American Chemical Society. 88: 169–170. doi:10.1021/ja00953a035.CS1 maint: uses authors parameter (link)
Jäger, Susanne et al. (1986). "Fluorine and Oxygen". Fluorine. Berlin, Heidelberg: Springer. pp. 1–161.CS1 maint: uses authors parameter (link)
Nikitin, Igor Vasil'evich, and V. Ya Rosolovskii (1971). "Oxygen Fluorides and Dioxygenyl Compounds". Russian Chemical Reviews. 40 (11): 889–900. Bibcode:1971RuCRv..40..889N. doi:10.1070/rc1971v040n11abeh001981.CS1 maint: uses authors parameter (link)
Lawless, Edward W., Ivan C. Smith (1968). Inorganic high-energy oxidizers: synthesis, structure, and properties. M. Dekker.CS1 maint: uses authors parameter (link)
Marx, Rupert, Konrad Seppelt (2015). "Structure investigations on oxygen fluorides". Dalton Transactions. 44 (45): 19659–19662. doi:10.1039/c5dt02247a. PMID 26351980.CS1 maint: uses authors parameter (link)
Solomon, Irvine J. Research on Chemistry of ${\ce {O3F2}}$ and ${\ce {O2F2}}$ . No. IITRI-C227-6. IIT RESEARCH INST CHICAGO IL, 1964.
Goetschel, Charles T. et al. (1969). "Low-Temperature Radiation Chemistry. I. Preparation of Oxygen Fluorides and Dioxygenyl Tetrafluoroborate". Journal of the American Chemical Society. 91 (17): 4702–4707. doi:10.1021/ja01045a020.CS1 maint: uses authors parameter (link)
De Marco, Ronald A., and Jean'ne M. Shreeve . "Fluorinated Peroxides." Advances in Inorganic Chemistry and Radiochemistry. Vol. 16. Academic Press, 1974. 109-176.
J.L.Lyman and R. Holland, J. Phys. Chem.,1988,92, 7232.
Francisco J. S. (1993). "An ab initio investigation of the significance of the HOOF intermediate in coupling reactions involving FOO x and HO x species". The Journal of Chemical Physics. 98 (3): 2198–2207. Bibcode:1993JChPh..98.2198F. doi:10.1063/1.464199.

External links

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[1] Solomon, I. J. et al. (1968). "Additional Studies Concerning the Existence of O₃F₂". Journal of the American Chemical Society. 90 (20): 5408–5411. doi:10.1021/ja01022a014.CS1 maint: uses authors parameter (link)

[2] Misochko, Eugenii Ya, Alexander V. Akimov, Charles A. Wight (1999). "Infrared spectroscopic observation of the stabilized Intermediate complex FO₃ formed by reaction of mobile Fluorine atoms with ozone molecules Trapped in an Argon Matrix". The Journal of Physical Chemistry A. 103 (40): 7972–7977. Bibcode:1999JPCA..103.7972M. doi:10.1021/jp9921194.CS1 maint: uses authors parameter (link)

[:0-3] Streng, A. G. (1963). "The Oxygen Fluorides". Chemical Reviews. 63 (6): 607–624. doi:10.1021/cr60226a003.

[4] Streng, A. G., A. V. Grosse (1966). "Two New Fluorides of Oxygen, O₅F₂ and O₆F₂". Journal of the American Chemical Society. 88: 169–170. doi:10.1021/ja00953a035.CS1 maint: uses authors parameter (link)

[5] Jäger, Susanne et al. (1986). "Fluorine and Oxygen". Fluorine. Berlin, Heidelberg: Springer. pp. 1–161.CS1 maint: uses authors parameter (link)

[:1-6] Nikitin, Igor Vasil'evich, and V. Ya Rosolovskii (1971). "Oxygen Fluorides and Dioxygenyl Compounds". Russian Chemical Reviews. 40 (11): 889–900. Bibcode:1971RuCRv..40..889N. doi:10.1070/rc1971v040n11abeh001981.CS1 maint: uses authors parameter (link)

[7] Lawless, Edward W., Ivan C. Smith (1968). Inorganic high-energy oxidizers: synthesis, structure, and properties. M. Dekker.CS1 maint: uses authors parameter (link)

[8] Marx, Rupert, Konrad Seppelt (2015). "Structure investigations on oxygen fluorides". Dalton Transactions. 44 (45): 19659–19662. doi:10.1039/c5dt02247a. PMID 26351980.CS1 maint: uses authors parameter (link)

[:2-9] Solomon, Irvine J. Research on Chemistry of ${\ce {O3F2}}$ and ${\ce {O2F2}}$ . No. IITRI-C227-6. IIT RESEARCH INST CHICAGO IL, 1964.

[10] Goetschel, Charles T. et al. (1969). "Low-Temperature Radiation Chemistry. I. Preparation of Oxygen Fluorides and Dioxygenyl Tetrafluoroborate". Journal of the American Chemical Society. 91 (17): 4702–4707. doi:10.1021/ja01045a020.CS1 maint: uses authors parameter (link)

[11] De Marco, Ronald A., and Jean'ne M. Shreeve . "Fluorinated Peroxides." Advances in Inorganic Chemistry and Radiochemistry. Vol. 16. Academic Press, 1974. 109-176.

[12] J.L.Lyman and R. Holland, J. Phys. Chem.,1988,92, 7232.

[13] Francisco J. S. (1993). "An ab initio investigation of the significance of the HOOF intermediate in coupling reactions involving FOO x and HO x species". The Journal of Chemical Physics. 98 (3): 2198–2207. Bibcode:1993JChPh..98.2198F. doi:10.1063/1.464199.