Tetrafluoromethane

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Tetrafluoromethane
Carbon-tetrafluoride-2D-dimensions.png Carbon-tetrafluoride-3D-balls-B.png
Identifiers
CAS number 75-73-0 YesY
PubChem 6393
ChemSpider 6153 YesY
UNII 94WG9QG0JN YesY
EC number 200-896-5
ChEBI CHEBI:38825 YesY
RTECS number FG4920000
Jmol-3D images Image 1
Properties
Molecular formula 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)
Solubility in water 0.005%V at 20 °C
0.0038%V at 25 °C
Solubility soluble in benzene, chloroform
Vapor pressure 3.65 MPa at 15 °C
106.5 kPa at −127 °C
kH 5.15 atm-cu m/mole
Refractive index (nD) 1.0004823

[1]

Hazards
MSDS ICSC 0575
EU Index Not listed
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
Flash point Non-flammable
Autoignition temperature 1100 °C
Related compounds
Other cations Silicon tetrafluoride
Germanium tetrafluoride
Tin tetrafluoride
Lead tetrafluoride
Related fluoromethanes Fluoromethane
Difluoromethane
Fluoroform
Related compounds Tetrachloromethane
Tetrabromomethane
Tetraiodomethane
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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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.

Bonding[edit]

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[edit]

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[edit]

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

SiC + 2 F2 → CF4 + Si

Reactions[edit]

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[edit]

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[edit]

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.

Health risks[edit]

Depending on the concentration, inhalation of tetrafluoromethane can cause headaches, 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.

See also[edit]

References[edit]

  1. ^ 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. 
  2. ^ 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. 
  3. ^ a b Lemal, D.M. (2004). "Perspective on Fluorocarbon Chemistry". J. Org. Chem. 69 (1): 1–11. doi:10.1021/jo0302556. PMID 14703372. 
  4. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 
  5. ^ 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
  6. ^ K. Williams, K. Gupta, M. Wasilik. Etch Rates for Micromachining Processing – Part II J. Microelectromech. Syst., vol. 12, pp. 761–777, December 2003.
  7. ^ "Low efficiency 2-dimensional position-sensitive neutron detector for beam profile measurement". Retrieved 19 July 2012. 

External links[edit]