Oxygen fluorides arecompounds ofelementsoxygen andfluorine with the general formulaOnF2, wheren = 1 to 6. Many different oxygen fluorides are known:
Oxygen fluorides are strongoxidizing agents with high energy and can release their energy either instantaneously or at a controlled rate. Thus, these compounds attracted much attention as potential oxidizers injet propulsion systems.[5]
A common preparative method involves fluorination ofsodium hydroxide:
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 freehalogens:
It can also displace halogens from their salts.[3] It is both an effectivefluorinating agent and a strongoxidizing agent. When reacted with unsaturatednitrogen fluorides with electrical discharge, it results in the formation ofnitrogen trifluoride, oxide fluorides and other oxides.[6][7]
O2F2 precipitates as a brown solid upon theUV irradiation of a mixture of liquidO2 andF2 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 agas-phase electric discharge in cold containers includingO2F2.[10]
It is typically an orange-yellow solid which rapidly decomposes toO2 andF2 close to its normal boiling point of about 216 K.[3]
O2F2 reacts violently withred phosphorus, even at −196 °C. Explosions can also occur ifFreon-13 is used to moderate the reaction.[9]
O3F2 is a viscous, blood-red liquid. It remains liquid at 90 K and so can be differentiated fromO2F2 which has a melting point of about 109 K.[11][3]
Like the other oxygen fluorides,O3F2 isendothermic and decomposes at about 115 K with the evolution of heat, which is given by the following reaction:
O3F2 is safer to work with thanozone, 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 anhydrousammonia will result in a mild explosion, when they are both at 90 K each.[3]
Fluoroperoxyl is a molecule such as O–O–F, whosechemical formula isO2F and is stable only at low temperature. It has been reported to be produced from atomic fluorine and dioxygen.[12]
| 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° |
Oxygen- and fluorine-containing radicals likeO2F and OF occur in the atmosphere. These along with other halogen radicals have been implicated in thedestruction of ozone in the atmosphere. However, theoxygen monofluorideradicals are assumed to not play as big a role in the ozone depletion because free fluorine atoms in the atmosphere are believed to react withmethane to producehydrofluoric acid which precipitates in rain. This decreases the availability of free fluorine atoms for oxygen atoms to react with and destroy ozone molecules.[13]
Net reaction:
Despite the low solubility ofO3F2 in liquid oxygen, it has been shown to behypergolic with most rocket propellant fuels. The mechanism involves the boiling off oxygen from the solution containingO3F2, 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]