Oxygen fluoride

Short description: Any binary compound of oxygen and fluorine

Oxygen difluoride

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula O

2, where n = 1 to 6. Many different oxygen fluorides are known:

* oxygen difluoride (OF

* dioxygen difluoride (O

* trioxygen difluoride or ozone difluoride (O

2)[1][2]
* tetraoxygen difluoride (O

2)[3]
* pentaoxygen difluoride (O

* hexaoxygen difluoride (O

2)[4]
* dioxygen monofluoride or fluoroperoxyl (O
2F)

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]

## Contents

* 1 Synthesis
* 1.1 Oxygen difluoride (OF2)
* 1.2 Dioxygen difluoride (O2F2)
* 1.3 Trioxygen difluoride or ozone difluoride (O3F2)
* 1.4 Fluoroperoxyl
* 2 General preparation of polyoxygen difluorides
* 3 Effects on ozone
* 4 Hypergolic propellant
* 5 See also
* 6 References
* 7 External links

## Synthesis

Here are some synthesis methods and reactions of the three most common oxygen fluorides – oxygen difluoride (OF
2), dioxygen difluoride (O

2) and ozone difluoride (O

2).

### Oxygen difluoride (OF

Oxygen difluoride

A common preparative method involves fluorination of sodium hydroxide:

2 F
2 \+ 2 NaOH → OF
2 \+ 2 NaF + H

2 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:

2 \+ 4 HCl → 2 Cl
2 \+ 2 HF + 2 H

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 (O

Dioxygen difluoride

2 precipitates as a brown solid upon the UV irradiation of a mixture of liquid O
2 and F
2 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 O

2.[10]

2 \+ F
2 → O

2 (electric discharge, 183 °C)

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

2 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 (O

2 is a viscous, blood-red liquid. It remains liquid at 90 K and so can be differentiated from O

2 which has a melting point of about 109 K.[11][3]

Like the other oxygen fluorides, O

2 is endothermic and decomposes at about 115 K with the evolution of heat, which is given by the following reaction:

2 O

2 → O
2 \+ 2 O

2 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 O
2F and is stable only at low temperature. It has been reported to be produced from atomic fluorine and dioxygen.[12]

2 \+ F → O

## General preparation of polyoxygen difluorides

Reaction equation[6] | O
2:F
2 by volume | Current | Temperature of bath (°C)

|
|
|

2 \+ F
2 ⇌ O

2 | 1:1  | 10 – 50 mA  | ~ -196°
3 O
2 \+ 2 F
2 ⇌ 2 O

2 | 3:2  | 25 – 30 mA  | ~ -196°
2 O
2 \+ F
2 ⇌ O

2 | 2:1  | 4 – 5 mA  | ~ -205°

## Effects on ozone

Oxygen- and fluorine-containing radicals like O
2F 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]

3 \+ F → O
2 \+ OF
O + OF → O
2 \+ F

Net reaction:

3 \+ O → 2 O

## Hypergolic propellant

Despite the low solubility of O

2 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 O

2, 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

* Bromine oxide
* Chlorine oxide
* Iodine oxide
* Ozone

## References

1. ↑ Solomon, I. J. (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. doi:10.1021/jp9921194. Bibcode: 1999JPCA..103.7972M.
3. ↑ 3.0 3.1 3.2 3.3 3.4 3.5 3.6 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 (1986). "Fluorine and Oxygen". Fluorine. Berlin, Heidelberg: Springer. pp. 1–161.
6. ↑ 6.0 6.1 Nikitin, Igor Vasil'evich, and V. Ya Rosolovskii (1971). "Oxygen Fluorides and Dioxygenyl Compounds". Russian Chemical Reviews 40 (11): 889–900. doi:10.1070/rc1971v040n11abeh001981. Bibcode: 1971RuCRv..40..889N.
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. ↑ 9.0 9.1 Solomon, Irvine J. Research on Chemistry of O

2 and O

2. No. IITRI-C227-6. IIT RESEARCH INST CHICAGO IL, 1964.
10. ↑ Goetschel, Charles T. (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. doi:10.1063/1.464199. Bibcode: 1993JChPh..98.2198F.

## External links

* National Pollutant Inventory - Fluoride and compounds fact sheet
* Oxygen Fluoride PubChem
* Center for Disease Control and Prevention - Health Hazards of Oxygen Difluoride

* v
* t
* e

Oxygen compounds

* AgO
* Al2O3
* AmO2
* Am2O3
* As2O3
* As2O5
* Au2O3
* B2O3
* BaO
* BeO
* Bi2O3
* BiO2
* Bi2O5
* BrO2
* Br2O3
* Br2O5
* CO
* CO2
* C2O3
* CaO
* CaO2
* CdO
* CeO2
* Ce3O4
* Ce2O3
* ClO2
* Cl2O
* Cl2O3
* Cl2O4
* Cl2O6
* Cl2O7
* CoO
* Co2O3
* Co3O4
* CrO3
* Cr2O3
* Cr2O5
* Cr5O12
* CsO2
* Cs2O3
* CuO
* D2O
* Dy2O3
* Er2O3
* Eu2O3
* FeO
* Fe2O3
* Fe3O4
* Ga2O
* Ga2O3
* GeO
* GeO2
* H2O
* H218O
* H2O2
* HfO2
* HgO
* Hg2O
* Ho2O3
* I2O4
* I2O5
* I2O6
* I4O9
* In2O3
* IrO2
* KO2
* K2O2
* La2O3
* Li2O
* Li2O2
* Lu2O3
* MgO
* Mg2O3
* MnO
* MnO2
* Mn2O3
* Mn2O7
* MoO2
* MoO3
* Mo2O3
* NO
* NO2
* N2O
* N2O3
* N2O4
* N2O5
* NaO2
* Na2O
* Na2O2
* NbO
* NbO2
* Nd2O3
* OF2
* O2F2
* O3F2
* O4F2
* O2PtF6

Chemical formulas

* v
* t
* e

Salts and covalent derivatives of the fluoride ion

HF |  |  |  | He

|
|
|
|

LiF | BeF2 | BF
BF3
B2F4 | CF4
other compounds | NF3
FN3
N2F2
N2F4
NF5§  | F2O
F2O2
other compounds | F2 | Ne
NaF | MgF2 | AlF
AlF3 | SiF4 | P2F4
PF3
PF5 | S2F2
SF2
SF4
SF6 | ClF
ClF3
ClF5 | Ar
KF | CaF
CaF2 | ScF3 | TiF2
TiF3
TiF4 | VF2
VF3
VF4
VF5 | CrF2
CrF3
CrF4
CrF5
CrF6§  | MnF2 | FeF2
FeF3 | CoF2
CoF3 | NiF2 | CuF
CuF2 | ZnF2 | GaF2
GaF3 | GeF2
GeF4 | AsF3
AsF5 | Se2F2
SeF4
SeF6 | BrF
BrF3
BrF5 | KrF2
RbF | SrF
SrF2 | YF3 | ZrF3
ZrF4 | NbF4
NbF5 | MoF2
MoF3
MoF4
MoF5
MoF6 | TcF4
TcF6 | RuF3
RuF5
RuF6 | RhF3
RhF5
RhF6 | PdF2 | AgF | CdF2 | InF
InF2
InF3 | SnF2
SnF4 | SbF3
SbF5 | Te3F2
TeF4
TeF6 | IF
IF3</br>IF5</br>IF7 | XeF2
XeF4
XeF6
CsF | BaF2 |  | HfF4 | TaF5 | WF2
WF3
WF4
WF5
WF6 | Re3F9
ReF4
ReF5
ReF6
ReF7 | OsF4
OsF5
OsF6 | IrF2
IrF3
IrF4
IrF5
IrF6 | PtF2
PtF4
PtF5
PtF6 | AuF
AuF3
Au2F10
AuF5•F2 | Hg2F2
HgF2
HgF4 | TlF | PbF2
PbF4 | BiF3
BiF5 | PoF2
PoF4
PoF6 | AtF  | RnF2
FrF | RaF2 |  | Rf  | Db  | Sg  | Bh  | Hs  | Mt  | Ds  | Rg  | Cn  | Nh  | Fl  | Mc  | Lv  | Ts  | Og
| ↓  |  |
LaF3 | CeF3
CeF4 | PrF3 | NdF2,
NdF3 | PmF3 | SmF2,
SmF3 | EuF2,
EuF3 | GdF3 | TbF3 | DyF2,
DyF3 | HoF3 | ErF3 | TmF2
TmF3 | YbF2
YbF3 | LuF3
AcF3 | ThF4 | PaF5 | UF3
UF4
UF5
UF6 | NpF4
NpF6 | PuF3
PuF6 | AmF2
AmF3 | CmF3 | BkF3 | CfF3 | EsF3 | Fm  | Md  | No  | LrF3

§ means the substance has not been made.

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