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Carbon tetrafluoride

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Carbon tetrafluoride
Names
IUPAC names
Tetrafluoromethane
Carbon tetrafluoride
Other names
Carbon tetrafluoride, Perfluoromethane, Tetrafluorocarbon, Freon 14, Halon 14, Arcton 0, CFC 14, PFC 14, R 14, UN 1982
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.000.815 Edit this at Wikidata
EC Number
  • 200-896-5
RTECS number
  • FG4920000
UNII
  • InChI=1S/CF4/c2-1(3,4)5 checkY
    Key: TXEYQDLBPFQVAA-UHFFFAOYSA-N checkY
  • InChI=1/CF4/c2-1(3,4)5
  • FC(F)(F)F
Properties
CF4
Molar mass 88.0043 g/mol
Appearance Colorless gas
Odor odorless
Density 3.72 g/L, gas (15 °C)
Melting point −183.6 °C (−298.5 °F; 89.5 K)
Boiling point −127.8 °C (−198.0 °F; 145.3 K)
Critical point (T, P) −45.55 °C (−50.0 °F; 227.6 K), 36.91 standard atmospheres (3,739.9 kPa; 542.4 psi)[1]
0.005%V at 20 °C
0.0038%V at 25 °C
Solubility soluble in benzene, chloroform
Vapor pressure 106.5 kPa at −127 °C
5.15 atm-cu m/mole
1.0004823[2]
Viscosity 17.32 μPa·s[3]
Structure
Tetragonal
Tetrahedral
0 D
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Simple asphyxiant and greenhouse gas
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 1: Exposure would cause irritation but only minor residual injury. E.g. turpentineFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard SA: Simple asphyxiant gas. E.g. nitrogen, helium
1
0
0
SA
Flash point Non-flammable
Safety data sheet (SDS) ICSC 0575
Related compounds
Other anions
 Tetrachloromethane
Tetrabromomethane
Tetraiodomethane
Other cations
 Silicon tetrafluoride
Germanium tetrafluoride
Tin tetrafluoride
Lead tetrafluoride
Related fluoromethanes
 Fluoromethane
Difluoromethane
Fluoroform
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Tetrafluoromethane, also known as carbon tetrafluoride or R-14, is the simplest perfluorocarbon (CF4). 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.[4] It has a very high bond strength due to the nature of the carbon–fluorine bond.

Bonding

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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.[5] 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.[6] 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.[6]

Preparation

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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.[7] 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:[4]

CCl2F2 + 2 HF → CF4 + 2 HCl

Laboratory synthesis

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

SiC + 4 F2 → CF4 + SiF4

Reactions

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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

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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.[8] It also has uses in neutron detectors.[9]

Environmental effects

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Mauna Loa tetrafluoromethane (CF4) timeseries.
PFC-14 measured by the Advanced Global Atmospheric Gases Experiment (AGAGE) in the lower atmosphere (troposphere) at stations around the world. Abundances are given as pollution free monthly mean mole fractions in parts-per-trillion.
Atmospheric concentration of CF4 (PFC-14) vs. similar man-made gases (right graph). Note the log scale.

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.[10]

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

Although structurally similar to chlorofluorocarbons (CFCs), tetrafluoromethane does not deplete the ozone layer[13] because the carbon–fluorine bond is much stronger than that between carbon and chlorine.[14]

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.[15]

Health risks

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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.

See also

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References

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  1. ^ Lide, David R.; Kehiaian, Henry V. (1994). CRC Handbook of Thermophysical and Themochemical Data (PDF). CRC Press. p. 31.
  2. ^ 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. Bibcode:1990NIMPA.292..593A. doi:10.1016/0168-9002(90)90178-9.
  3. ^ Kestin, J.; Ro, S.T.; Wakeham, W.A. (1971). "Reference values of the viscosity of twelve gases at 25°C". Transactions of the Faraday Society. 67: 2308–2313. doi:10.1039/TF9716702308.
  4. ^ a b Siegemund, Günter; Schwertfeger, Werner; Feiring, Andrew; Smart, Bruce; Behr, Fred; Vogel, Herward; McKusick, Blaine (2002). "Fluorine Compounds, Organic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a11_349. ISBN 978-3527306732.
  5. ^ 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.
  6. ^ a b Lemal, D.M. (2004). "Perspective on Fluorocarbon Chemistry". J. Org. Chem. 69 (1): 1–11. doi:10.1021/jo0302556. PMID 14703372.
  7. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  8. ^ K. Williams, K. Gupta, M. Wasilik. Etch Rates for Micromachining Processing – Part II J. Microelectromech. Syst., vol. 12, pp. 761–777, December 2003.
  9. ^ Moon, Myung-Kook; Nam, Uk-Won; Lee, Chang-Hee; Em, V.T.; Choi, Young-Hyun; Cheon, Jong-Kyu; Kong, Kyung-Nam (2005). "Low efficiency 2-dimensional position-sensitive neutron detector for beam profile measurement". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment. 538 (1–3): 592–596. Bibcode:2005NIMPA.538..592M. doi:10.1016/j.nima.2004.09.020.
  10. ^ Artaxo, Paulo; Berntsen, Terje; Betts, Richard; Fahey, David W.; Haywood, James; Lean, Judith; Lowe, David C.; Myhre, Gunnar; Nganga, John; Prinn, Ronald; Raga, Graciela; Schulz, Michael; van Dorland, Robert (February 2018). "Changes in Atmospheric Constituents and in Radiative Forcing" (PDF). Intergovernmental Panel on Climate Change. p. 212. Retrieved 17 March 2021.
  11. ^ "Climate change indicators - Atmospheric concentration of greenhouse gases - Figure 4". United States Environmental Protection Agency. 27 June 2016. Retrieved 2020-09-26.
  12. ^ Butler J. and Montzka S. (2020). "The NOAA Annual Greenhouse Gas Index (AGGI)". NOAA Global Monitoring Laboratory/Earth System Research Laboratories.
  13. ^ Cicerone, Ralph J. (1979-10-05). "Atmospheric Carbon Tetrafluoride: A Nearly Inert Gas" (PDF). Science. 206 (4414): 59–61. Bibcode:1979Sci...206...59C. doi:10.1126/science.206.4414.59. ISSN 0036-8075. PMID 17812452. S2CID 34911990.
  14. ^ "Bond Energies". www2.chemistry.msu.edu. Retrieved 2023-01-15.
  15. ^ Jubb, Aaron M.; McGillen, Max R.; Portmann, Robert W.; Daniel, John S.; Burkholder, James B. (2015). "An atmospheric photochemical source of the persistent greenhouse gas CF4". Geophysical Research Letters. 42 (21): 9505–9511. Bibcode:2015GeoRL..42.9505J. doi:10.1002/2015GL066193. ISSN 0094-8276.
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