Oxygen fluoride

Any binary compound of oxygen and fluorine

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 monofluoride (OF)
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 or fluoroperoxyl (O₂F)

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, properties and reactions

Oxygen difluoride (OF₂)

A common preparative method involves fluorination of sodium hydroxide:

2 F₂ + 2 NaOH → OF₂ + 2 NaF + 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:

OF₂ + 4 HCl → 2 Cl₂ + 2 HF + 2 H₂O

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)

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]

O₂ + F₂ → O₂F₂ (electric discharge, 183 °C)

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]

Trioxygen difluoride or 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:

2 O₃F₂ → O₂ + 2 O₂F₂

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]

O₂ + F → O₂F

General preparation of polyoxygen difluorides

Reaction equation^[6]	O₂:F₂ by volume	Current	Temperature of bath (°C)
O₂ + F₂ ⇌ O₂F₂	1:1	10 – 50 mA	~ -196°
3 O₂ + 2 F₂ ⇌ 2 O₃F₂	3:2	25 – 30 mA	~ -196°
2 O₂ + F₂ ⇌ O₄F₂	2:1	4 – 5 mA	~ -205°

Effects on ozone

Oxygen- and fluorine-containing radicals like O₂F and OF occur in the atmosphere. These along with other halogen radicals have been implicated in the destruction of ozone in the atmosphere. However, the oxygen monofluoride 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]

O₃ + F → O₂ + OF

O + OF → O₂ + F

Net reaction:

O₃ + O → 2 O₂

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.
^ 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.
^ ^a ^b ^c ^d ^e ^f ^g 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.
^ Jäger, Susanne; et al. (1986). "Fluorine and Oxygen". Fluorine. Berlin, Heidelberg: Springer. pp. 1–161.
^ ^a ^b Nikitin, Igor Vasil'evich; V. Ya Rosolovskii (1971). "Oxygen Fluorides and Dioxygenyl Compounds". Russian Chemical Reviews. 40 (11): 889–900. Bibcode:1971RuCRv..40..889N. doi:10.1070/rc1971v040n11abeh001981. S2CID 250903149.
^ Lawless, Edward W.; Ivan C. Smith (1968). Inorganic high-energy oxidizers: synthesis, structure, and properties. M. Dekker.
^ Marx, Rupert; Konrad Seppelt (2015). "Structure investigations on oxygen fluorides". Dalton Transactions. 44 (45): 19659–19662. doi:10.1039/c5dt02247a. PMID 26351980.
^ ^a ^b Solomon, Irvine J. Research on Chemistry of O₃F₂ and 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.
^ 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.