Oxygen fluoride

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Oxygen difluoride Oxygen difluoride.svg
Oxygen difluoride

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:

Contents

Tetraoxygen difluoride Tetraoxygen difluoride.svg
Tetraoxygen 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, properties and reactions

Oxygen difluoride (OF2)

Oxygen difluoride Oxygen-difluoride-3D-vdW.png
Oxygen difluoride

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)

Dioxygen difluoride Dioxygen-difluoride-3D-vdW.png
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]

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)

Trioxygen difluoride structural formula V.1.svg

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 volumeCurrentTemperature of bath (°C)
O2 + F2 ⇌ O2F21:110 – 50 mA~ -196°
3 O2 + 2 F2 ⇌ 2 O3F23:225 – 30 mA~ -196°
2 O2 + F2 ⇌ O4F22:14 – 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]

See also

Related Research Articles

<span class="mw-page-title-main">Oxidizing agent</span> Chemical compound used to oxidize another substance in a chemical reaction

An oxidizing agent is a substance in a redox chemical reaction that gains or "accepts"/"receives" an electron from a reducing agent. In other words, an oxidizer is any substance that oxidizes another substance. The oxidation state, which describes the degree of loss of electrons, of the oxidizer decreases while that of the reductant increases; this is expressed by saying that oxidizers "undergo reduction" and "are reduced" while reducers "undergo oxidation" and "are oxidized". Common oxidizing agents are oxygen, hydrogen peroxide, and the halogens.

In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

Chlorine trifluoride is an interhalogen compound with the formula ClF3. This colorless, poisonous, corrosive, and extremely reactive gas condenses to a pale-greenish yellow liquid, the form in which it is most often sold. Despite being famous for its extreme oxidation properties and igniting many things, chlorine trifluoride is not combustible itself. The compound is primarily of interest in plasmaless cleaning and etching operations in the semiconductor industry, in nuclear reactor fuel processing, historically as a component in rocket fuels, and various other industrial operations owing to its corrosive nature.

<span class="mw-page-title-main">Oxygen difluoride</span> Chemical compound

Oxygen difluoride is a chemical compound with the formula OF2. As predicted by VSEPR theory, the molecule adopts a bent molecular geometry. It is a strong oxidizer and has attracted attention in rocketry for this reason. With a boiling point of −144.75 °C, OF2 is the most volatile (isolable) triatomic compound. The compound is one of many known oxygen fluorides.

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 −163 °C (110 K). It is an extremely strong oxidant and decomposes into oxygen and fluorine even at −160 °C (113 K) at a rate of 4% per day — 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).

<span class="mw-page-title-main">Silver(II) fluoride</span> Chemical compound

Silver(II) fluoride is a chemical compound with the formula AgF2. It is a rare example of a silver(II) compound - silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.

<span class="mw-page-title-main">Xenon difluoride</span> Chemical compound

Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF
2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.

<span class="mw-page-title-main">Dioxygenyl</span> Chemical compound

The dioxygenyl(or dioxyl) ion, O+
2, is a rarely-encountered oxycation in which both oxygen atoms have a formal oxidation state of +1/2. It is formally derived from oxygen by the removal of an electron:

Dioxygenyl hexafluoroplatinate is a compound with formula O2PtF6. It is a hexafluoroplatinate of the unusual dioxygenyl cation, O2+, and is the first known compound containing this cation. It can be produced by the reaction of dioxygen with platinum hexafluoride. The fact that PtF
6 is strong enough to oxidise O
2, whose first ionization potential is 12.2 eV, led Neil Bartlett to correctly surmise that it might be able to oxidise xenon (first ionization potential 12.13 eV). This led to the discovery of xenon hexafluoroplatinate, which proved that the noble gases, previously thought to be inert, are able to form chemical compounds.

<span class="mw-page-title-main">Chromyl fluoride</span> Chemical compound

Chromyl fluoride is an inorganic compound with the formula CrO2F2. It is a violet-red colored crystalline solid that melts to an orange-red liquid.

In chemistry, molecular oxohalides (oxyhalides) are a group of chemical compounds in which both oxygen and halogen atoms are attached to another chemical element A in a single molecule. They have the general formula AOmXn, where X is a halogen. Known oxohalides have fluorine (F), chlorine (Cl), bromine (Br), and/or iodine (I) in their molecules. The element A may be a main group element, a transition element, a rare earth element or an actinide. The term oxohalide, or oxyhalide, may also refer to minerals and other crystalline substances with the same overall chemical formula, but having an ionic structure.

<span class="mw-page-title-main">Neptunium(VI) fluoride</span> Chemical compound

Neptunium(VI) fluoride (NpF6) is the highest fluoride of neptunium, it is also one of seventeen known binary hexafluorides. It is an orange volatile crystalline solid. It is relatively hard to handle, being very corrosive, volatile and radioactive. Neptunium hexafluoride is stable in dry air but reacts vigorously with water.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

<span class="mw-page-title-main">Chlorine trifluoride oxide</span> Chemical compound

Chlorine oxide trifluoride or chlorine trifluoride oxide is a corrosive liquid molecular compound with formula ClOF3. It was developed secretly as a rocket fuel oxidiser.

<span class="mw-page-title-main">Xenon oxydifluoride</span> Chemical compound

Xenon oxydifluoride is an inorganic compound with the molecular formula XeOF2. The first definitive isolation of the compound was published on 3 March 2007, producing it by the previously-examined route of partial hydrolysis of xenon tetrafluoride.

Tetraoxygen difluoride is an inorganic chemical compound of oxygen, belonging to the family of oxygen fluorides. It consists of two O2F units bound together with a weak O-O bond, and is the dimer of the O2F radical.

<span class="mw-page-title-main">Seleninyl fluoride</span> Chemical compound

Seleninyl fluoride is an oxyfluoride of selenium with the chemical formula SeOF2.

Oxygen monofluoride is an unstable binary inorganic compound radical of fluorine and oxygen with the chemical formula OF. This is the simplest of many oxygen fluorides.

Dioxygen monofluoride is a binary inorganic compound radical of fluorine and oxygen with the chemical formula O2F. The compound is stable only at low temperature. This is one of many known oxygen fluorides.

Hexaoxygen difluoride is a binary inorganic compound of fluorine and oxygen with the chemical formula O6F2. The compound is one of many known oxygen fluorides.

References

  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.
  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.
  3. 1 2 3 4 5 6 7 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.
  5. ↑ Jäger, Susanne; et al. (1986). "Fluorine and Oxygen". Fluorine. Berlin, Heidelberg: Springer. pp. 1–161.
  6. 1 2 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.
  7. ↑ Lawless, Edward W.; Ivan C. Smith (1968). Inorganic high-energy oxidizers: synthesis, structure, and properties. M. Dekker.
  8. ↑ Marx, Rupert; Konrad Seppelt (2015). "Structure investigations on oxygen fluorides". Dalton Transactions. 44 (45): 19659–19662. doi:10.1039/c5dt02247a. PMID   26351980.
  9. 1 2 Solomon, Irvine J. Research on Chemistry of O3F2 and 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.
  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.
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