Tetrafluoromethane, also known as carbon tetrafluoride or R-14, is the simplest perfluorocarbon (C F₄). As its IUPAC name indicates, tetrafluoromethane is the perfluorinated counterpart to the hydrocarbon methane. It can also be classified as a haloalkane or halomethane. Tetrafluoromethane is a useful refrigerant but also a potent greenhouse gas.^[3] It has a very high bond strength due to the nature of the carbon–fluorine bond.

Bonding

Because of the multiple carbon–fluorine bonds, and the high electronegativity of fluorine, the carbon in tetrafluoromethane has a significant positive partial charge which strengthens and shortens the four carbon–fluorine bonds by providing additional ionic character. Carbon–fluorine bonds are the strongest single bonds in organic chemistry.^[4] Additionally, they strengthen as more carbon–fluorine bonds are added to the same carbon. In the one carbon organofluorine compounds represented by molecules of fluoromethane, difluoromethane, trifluoromethane, and tetrafluoromethane, the carbon–fluorine bonds are strongest in tetrafluoromethane.^[5] This effect is due to the increased coulombic attractions between the fluorine atoms and the carbon because the carbon has a positive partial charge of 0.76.^[5]

Preparation

Tetrafluoromethane is the product when any carbon compound, including carbon itself, is burned in an atmosphere of fluorine. With hydrocarbons, hydrogen fluoride is a coproduct. It was first reported in 1926.^[6] It can also be prepared by the fluorination of carbon dioxide, carbon monoxide or phosgene with sulfur tetrafluoride. Commercially it is manufactured by the reaction of hydrogen fluoride with dichlorodifluoromethane or chlorotrifluoromethane; it is also produced during the electrolysis of metal fluorides MF, MF₂ using a carbon electrode.

Although it can be made from a myriad of precursors and fluorine, elemental fluorine is expensive and difficult to handle. Consequently, CF
₄ is prepared on an industrial scale using hydrogen fluoride:^[3]

CCl₂F₂ + 2 HF → CF₄ + 2 HCl

Laboratory synthesis

Tetrafluoromethane and silicon tetrafluoride can be prepared in the laboratory by the reaction of silicon carbide with fluorine.

SiC + 4 F₂ → CF₄ + SiF₄

Reactions

Tetrafluoromethane, like other fluorocarbons, is very stable due to the strength of its carbon–fluorine bonds. The bonds in tetrafluoromethane have a bonding energy of 515 kJ⋅mol⁻¹. As a result, it is inert to acids and hydroxides. However, it reacts explosively with alkali metals. Thermal decomposition or combustion of CF₄ produces toxic gases (carbonyl fluoride and carbon monoxide) and in the presence of water will also yield hydrogen fluoride.

It is very slightly soluble in water (about 20 mg⋅L⁻¹), but miscible with organic solvents.

Uses

Tetrafluoromethane is sometimes used as a low temperature refrigerant (R-14). It is used in electronics microfabrication alone or in combination with oxygen as a plasma etchant for silicon, silicon dioxide, and silicon nitride.^[7] It also has uses in neutron detectors.^[8]

Environmental effects

Tetrafluoromethane is a potent greenhouse gas that contributes to the greenhouse effect. It is very stable, has an atmospheric lifetime of 50,000 years, and a high greenhouse warming potential 6,500 times that of CO₂.^[9]

Tetrafluoromethane is the most abundant perfluorocarbon in the atmosphere, where it is designated as PFC-14. Its atmospheric concentration is growing.^[10] As of 2019, the man-made gases CFC-11 and CFC-12 continue to contribute a stronger radiative forcing than PFC-14.^[11]

Although structurally similar to chlorofluorocarbons (CFCs), tetrafluoromethane does not deplete the ozone layer^[12] because the carbon–fluorine bond is much stronger than that between carbon and chlorine.^[13]

Main industrial emissions of tetrafluoromethane besides hexafluoroethane are produced during production of aluminium using Hall-Héroult process. CF₄ also is produced as product of the breakdown of more complex compounds such as halocarbons.^[14]

Health risks

Due to its density, tetrafluoromethane can displace air, creating an asphyxiation hazard in inadequately ventilated areas. Otherwise, it is normally harmless due to its stability.

References

External links

Halomethanes

Inorganic compounds of carbon and related ions

Inorganic compounds of carbon and related ions
Compounds	CF CO CO₂ CO₃ CO₄ CO₅ CO₆ COS CS C₂S₂ CS₂ CSe₂ C₃O₂ C₃S₂ SiC
Carbon ions	Carbides [:C≡C:]²⁻, [::C::]⁴⁻, [:C=C=C:]⁴⁻ Cyanide [:C≡N:]⁻ Cyanate [:O−C≡N:]⁻ Thiocyanate [:S−C≡N:]⁻ Fulminate [:C≡N−O:]⁻ Thiofulminate [:C≡N−S:]⁻
Nanostructures	Graphite intercalation compounds Fullerides
Oxides and related	Oxides Nitrides Metal carbonyls Carbonic acid Bicarbonates Carbonates

Fluorine compounds

Authority control databases: National

Halomethanes
Unsubstituted	CH₄
Monosubstituted	CH₃F CH₃Cl CH₃Br CH₃I CH₃At
Disubstituted	CH₂F₂ CH₂ClF CH₂BrF CH₂FI CH₂Cl₂ CH₂BrCl CH₂ClI CH₂Br₂ CH₂BrI CH₂I₂
Trisubstituted	CHF₃ CHClF₂ CHBrF₂ CHF₂I CHCl₂F CHBrClF CHClFI CHBr₂F CHBrFI CHFI₂ CHCl₃ CHBrCl₂ CHCl₂I CHBr₂Cl CHBrClI CHClI₂ CHBr₃ CHBr₂I CHBrI₂ CHI₃
Tetrasubstituted	CF₄ CClF₃ CBrF₃ CF₃I CCl₂F₂ CBrClF₂ CClF₂I CBr₂F₂ CBrF₂I CF₂I₂ CCl₃F CBrCl₂F CCl₂FI CBr₂ClF C*BrClFI CClFI₂ CBr₃F CBr₂FI CBrFI₂ CFI₃ CCl₄ CBrCl₃ CCl₃I CBr₂Cl₂ CBrCl₂I CCl₂I₂ CBr₃Cl CBr₂ClI CBrClI₂ CClI₃ CBr₄ CBr₃I CBr₂I₂ CBrI₃ CI₄
* Chiral compound.