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Tetrafluoromethane: Wikis


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IUPAC name
Other names Carbon tetrafluoride, Perfluoromethane, Tetrafluorocarbon, Freon 14, Halon 14, Arcton 0, CFC 14, PFC 14, R 14, UN 1982
CAS number 75-73-0 Yes check.svgY
PubChem 6393
EC number 200-896-5
RTECS number FG4920000
Molecular formula CF4
Molar mass 88.0043 g/mol
Appearance Colorless odorless gas
Density 3.72 g/l, gas (15 °C)
Melting point

-183.6 °C (89.6 K)

Boiling point

-127.8 °C (145.4 K)

Solubility in water 0.005 %V at 20 °C
0.0038 %V at 25 °C
Vapor pressure 3.65 MPa at 15 °C
106.5 kPa at -127 °C
Refractive index (nD) 1.113
EU Index Not listed
NFPA 704
NFPA 704.svg
Flash point Non-flammable
1100 °C
Related compounds
Related fluoromethanes Fluoromethane
Related compounds Tetrachloromethane
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Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Tetrafluoromethane, also known as carbon tetrafluoride, is the simplest fluorocarbon (CF4). 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 highest 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.



Carbon–fluorine bonds are the strongest in organic chemistry.[1] 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.[2] 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.[2]


Pure tetrafluoromethane was first synthesised in 1926.[3]

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

SiC + 2 F2 → CF4 + Si

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.

Tetrafluoromethane, like other fluorocarbons, is very stable due to the strength of its carbon-fluorine bonds. The bonds in tetrafluoromethane are with 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 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 ethanol, ether, benzene.


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.[4]

Environmental effects

Tetrafluoromethane is a potent greenhouse gas that contributes to the greenhouse effect. It is very stable, has an atmospheric lifespan of 50,000 years, and a high greenhouse warming potential of 6500 (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.

Health risks

Inhalation of tetrafluoromethane can cause, depending on concentration, headache, nausea, dizziness and damage to the cardiovascular system (mainly the heart). Long-term exposure can cause severe heart damage.

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


  1. ^ O'Hagan D (February 2008). "Understanding organofluorine chemistry. An introduction to the C–F bond". Chem Soc Rev 37 (2): 308–19. doi:10.1039/b711844a. PMID 18197347.  
  2. ^ a b Lemal, D.M. (2004), "Perspective on Fluorocarbon Chemistry", J. Org. Chem. 69: 1–11, doi:10.1021/jo0302556, PMID 14703372  
  3. ^ Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0-7506-3365-4  
  4. ^ K. Williams, K. Gupta, M. Wasilik. Etch Rates for Micromachining Processing - Part II J. Microelectromech. Syst., vol. 12, pp. 761-777, Dec. 2003.

See also

External links



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