Tetrafluoromethane

Tetrafluoromethane, also known as carbon tetrafluoride, is the simplest fluorocarbon (C F₄). It has a very high bond strength due to the nature of the carbon–fluorine bond. It can also be classified as a haloalkane or halomethane. 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. Tetrafluoromethane is a potent greenhouse gas.

Bonding

Carbon–fluorine bonds are the strongest in organic chemistry.^[2] 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.^[3] 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.^[3]

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.^[4] 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 fluorine 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 myriad precursors and fluorine, elemental fluorine is expensive and difficult to handle. Consequently, CF
4 is prepared on an industrial scale using hydrogen fluoride:^[5]

CCl₂F₂ + 2 HF → CF₄ + 2 HCl

Laboratory synthesis

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

SiC + 2 F₂ → CF₄ + Si

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. It is used in electronics microfabrication alone or in combination with oxygen as a plasma etchant for silicon, silicon dioxide, and silicon nitride.^[6] It also has uses in neutron detectors ^[7]

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 of 6500 (which is given for the first 100 years thereof, CO₂ has a factor of 1); however, the low amount in the atmosphere restricts the overall radiative forcing effect.

Although structurally similar to chlorofluorocarbons (CFCs), tetrafluoromethane does not deplete the ozone layer. This is because the depletion is caused by the chlorine atoms in CFCs, which dissociate when struck by UV radiation. Carbon–fluorine bonds are stronger and less likely to dissociate. According to Guinness World Records Tetrafluoromethane is the most persistent greenhouse gas.

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.^[8]

Health risks

Due to its density, tetrafluoromethane can displace air, creating an asphyxiation hazard in inadequately ventilated areas.

References

↑ Abjean, R.; A. Bideau-Mehu; Y. Guern (15 July 1990). "Refractive index of carbon tetrafluoride (CF4) in the 300-140 nm wavelength range". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment. 292 (3): 593–594. doi:10.1016/0168-9002(90)90178-9.
↑ O'Hagan D (February 2008). "Understanding organofluorine chemistry and in cations. An introduction to the C–F bond". Chemical Society Reviews. 37 (2): 308–19. doi:10.1039/b711844a. PMID 18197347.
1 2 Lemal, D.M. (2004). "Perspective on Fluorocarbon Chemistry". J. Org. Chem. 69 (1): 1–11. doi:10.1021/jo0302556. PMID 14703372.
↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
↑ G. Siegemund, W. Schwertfeger, A. Feiring, B. Smart, F. Behr, H. Vogel, B. McKusick "Fluorine Compounds, Organic" in "Ullmann’s Encyclopedia of Industrial Chemistry" 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a11_349
↑ K. Williams, K. Gupta, M. Wasilik. Etch Rates for Micromachining Processing – Part II J. Microelectromech. Syst., vol. 12, pp. 761–777, December 2003.
↑ "Low efficiency 2-dimensional position-sensitive neutron detector for beam profile measurement". Retrieved 19 July 2012.
↑ Jubb, Aaron M.; M. R. McGillen, R. W. Portmann, J. S. Daniel, and J. B. Burkholder (2015). "An atmospheric photochemical source of the persistent greenhouse gas CF4". Geophys. Res. Lett. (42). doi:10.1002/2015GL066193. Cite uses deprecated parameter |coauthors= (help)

External links

Halomethanes

Monosubstituted	CH₃F CH₃Cl CH₃Br CH₃I

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.

Inorganic compounds of carbon and related ions

Compounds	CBr₄ CCl₄ CF CF₄ CI₄ CO CO₂ CO₃ COS CS CS₂ CSe₂ C₃O₂ C₃S₂ SiC

Carbon ions	Carbides [:C≡C:]^2–, [::C::]^4–, [:C=C=C:]^4– Cyanides [:C≡N:]^– Cyanates [:O-C≡N:]^– Thiocyanates [:S-C≡N:]^– Fulminates [:C≡N-O:]^– Thiofulminates [:C≡N-S:]^–

Oxides and related	Oxides Metal carbonyls Carbonic acid Bicarbonates Carbonates

Fluorine compounds

AgF AgF₂ Ag₂F AlF₃ AmF₃ AmF₄ AsF₃ AsF₅ AuF₃ BF₃ B₂F₄ BaF₂ BeF₂ BiF₃ BiF₅ BrF₃ BrF₅ CF₄ C₂F₆ CaF₂ CdF₂ CeF₃ ClF₃ ClF₅ CoF₂ CoF₃ CrF₂ CrF₃ CrF₄ CrF₅ CrF₆ CsF CuF DyF₃ ErF₃ EuF₂ EuF₃ DF HF FI KF LiF NaF RbF TlF FeF₂ HgF₂ Hg₂F₂ KrF₂ MgF₂ MnF₂ NiF₂ F₂O F₂O₄ PbF₂ PdF₂ SnF₂ SrF₂ XeF₂ ZnF₂ ZrF₂ FeF₃ GaF₃ GdF₃ HoF₃ InF₃ LaF₃ LuF₃ MnF₃ F₃N NdF₃ F₃P PrF₃ PuF₃ SbF₃ ScF₃ SmF₃ TbF₃ TiF₃ TlF₃ TmF₃ UF₃ VF₃ YF₃ YbF₃ GeF₄ HfF₄ MoF₄ F₄N₂ PbF₄ PuF₄ SF₄ SeF₄ SiF₄ SnF₄ TeF₄ ThF₄ UF₄ VF₄ WF₄ XeF₄ ZrF₄ IF₅ NbF₅ PF₅ SbF₅ TaF₅ UF₅ VF₅ WF₅ MoF₆ OsF₆ PtF₆ ReF₆ SF₆ SeF₆ TeF₆ UF₆ WF₆ XeF₆ IF₇ ReF₇ IrF₃ IrF₆ OsF₄ OsF₅ PmF₃ PrF₄ PuF₆ ReF₄ ReF₅ RhF₆ RuF₃ RuF₄ RuF₅ RuF₆ TcF₅ TcF₆ VF₂ YbF₂ AgBF₄ AgPF₆ Cs₂AlF₅ K₃AlF₆ Na₃AlF₆ KAsF₆ LiAsF₆ NaAsF₆ HBF₄ KBF₄ LiBF₄ NaBF₄ RbBF₄ Ba(BF₄)₂ Ni(BF₄)₂ Pb(BF₄)₂ Sn(BF₄)₂ BaClF BaSiF₆ BaGeF₆ BrOF₃ BrO₂F CBrF₃ CBr₂F₂ CBr₃F CClF₃ CCl₂F₂ CCl₃F CFN CF₂O CF₃I CHF₃ CH₂F₂ CH₃F C₂Cl₃F₃ C₂H₃F C₆H₅F C₇H₅F₃ C₁₅F₃₃N ClFO₂ CrFO₄ CrF₂O₂ CsBF₄ NH₄F FNO FNO₂ FNO₃ KHF₂ NaHF₂ ThOF₂ NH₅F₂ F₂OS F₃OP F₃PS HPF₆ HSbF₆ KPF₆ KSbF₆ LiPF₆ NaPF₆ NaSbF₆ Na₂SiF₆ Na₂TiF₆ Na₂ZrF₆ TlPF₆ IOF₃ K₂NbF₇ K₂TaF₇ IO₃F

Chemical formulas

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