Carbon tetrachloride

Carbon tetrachloride, also known by many other names (see Table) is the chemical compound with the formula CCl₄. It is a reagent in synthetic chemistry and was formerly widely used in fire extinguishers and as a precursor to refrigerants. It is a colorless liquid with a "sweet" smell that can be detected at low levels.

Both carbon tetrachloride and tetrachloromethane are acceptable names under IUPAC nomenclature. Colloquially, it is called "carbon tet".

Additional recommended knowledge

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Production

The production of carbon tetrachloride has steeply declined since the 1980's due to environmental concerns and the decreased demand for CFCs, which were derived from carbon tetrachloride. In 1992, production in the U.S.-Europe-Japan was estimated at 720,000,000 kg.^[1]

Carbon tetrachloride was originally synthesised in 1839 by reaction of chloroform with chlorine, from the french chemist Henri Victor Regnault,^[2] but now it is mainly synthesized from methane:

CH₄ + 4 Cl₂ → CCl₄ + 4HCl

The production often utilizes by-products of other chlorination reactions, such as the syntheses of dichloromethane and chloroform. Higher chlorocarbons are also subjected to "chlorinolysis:"

C₂Cl₆ + Cl₂ → 2 CCl₄

Prior to the 1950's, carbon tetrachloride was manufactured by the chlorination of carbon disulfide at 105 to 130 °C:

CS₂ + 3Cl₂ → CCl₄ + S₂Cl₂^[1]

Properties

In the carbon tetrachloride molecule, four chlorine atoms are positioned symmetrically as corners in a tetrahedral configuration joined to a carbon atom, in the center, by single covalent bonds. Because of this symmetrical geometry, the molecule has no net dipole moment; that is, CCl₄ is non-polar. As a solvent, it is well suited to dissolving other non-polar compounds, fats and oils. It is somewhat volatile, giving off vapors having a smell characteristic of other chlorinated solvents, somewhat similar to the tetrachloroethylene smell reminiscent of dry cleaners' shops.

Solid tetrachloromethane has 2 allotropes: crystaline II below -47.5 °C (225.6 K) and crystaline I above -47.5 °C.^[3]

At -47.3 °C it has monoclinic crystal structure with space group C2/c and lattice constants a = 20.3, b = 11.6, c = 19.9 (.10^-1 nm), β = 111°.^[4]

Uses

In the early 20th century, carbon tetrachloride was widely used as a dry cleaning solvent, as a refrigerant, and in fire extinguishers^[5]. However, once it became apparent that carbon tetrachloride exposure had severe adverse health effects, safer alternatives such as tetrachloroethylene were found for these applications, and its use in these roles declined from about 1940 onward. Carbon tetrachloride persisted as a pesticide to kill insects in stored grain, but in 1970, it was banned in consumer products in the United States.

Prior to the Montreal Protocol, large quantities of carbon tetrachloride were used to produce the freon refrigerants R-11 (trichlorofluoromethane) and R-12 (dichlorodifluoromethane). However, these refrigerants are now believed to play a role in ozone depletion and have been phased out. Carbon tetrachloride is still used to manufacture less destructive refrigerants.

Carbon tetrachloride has also been used in the detection of neutrinos. Carbon tetrachloride is one of the most potent hepatotoxins, and is widly used in scientific research to evaluate hepatoprotective agents 7,8

Reactivity

Carbon tetrachloride has practically no flammability at lower temperatures. Under high temperatures in air, it forms poisonous phosgene.

Because it has no C-H bonds, carbon tetrachloride does not easily undergo free-radical reactions. Hence it is a useful solvent for halogenations either by the elemental halogen, or by a halogenation reagent such as N-bromosuccinimide.

In organic chemistry, carbon tetrachloride serves as a source of chlorine in the Appel reaction.

As a solvent

It is used as a solvent in synthetic chemistry research, but because of its adverse health effects, it is no longer commonly used, and chemists generally try to substitute it with other solvents. It is sometimes useful as a solvent for infrared spectroscopy because there are no significant absorption bands > 1600 cm^-1. Because carbon tetrachloride does not have any hydrogen atoms, it was historically used in proton NMR spectroscopy. However, carbon tetrachloride is toxic, and its dissolving power is low^[6]. Its use has been largely superseded by deuterated solvents, which offer superior solvating properties and allow for deuterium lock by the spectrometer.

Safety

Exposure to high concentrations of carbon tetrachloride (including vapor) can affect the central nervous system, degenerate the liver^[7] and kidneys^[8] and may result (after prolonged exposure) in coma and even death^[9]. Chronic exposure to carbon tetrachloride can cause liver^[10][11] and kidney damage and could result in cancer^[12] More information can be found in Material safety data sheets.

Carbon tetrachloride is also both ozone-depleting^[13] and a greenhouse gas^[14]. However, since 1992^[15] its atmospheric concentrations have been in decline for the reasons described above (see also the atmospheric time-series figure).

See also

• Haloalkane
• Halomethane
• Chloromethane
• Dichloromethane
• Chloroform
• Tetrafluoromethane
• Tetrabromomethane
• Carbon tetraiodide (Tetraiodomethane)

• Dutch standards

References

1. ^ ^{a b} Manfred Rossberg, Wilhelm Lendle, Gerhard Pfleiderer, Adolf Tögel, Eberhard-Ludwig Dreher, Ernst Langer, Heinz Rassaerts, Peter Kleinschmidt, Heinz Strack, Richard Cook, Uwe Beck, Karl-August Lipper, Theodore R. Torkelson, Eckhard Löser, Klaus K. Beutel, “Chlorinated Hydrocarbons” in Ullmann’s Encyclopedia of Chemical Technology, 2007 John Wiley & Sons: New York.
2. ^ V. Regnault (1839). "Ueber die Chlorverbindungen des Kohlenstoffs, C2Cl2 und CCl2". Annalen der Pharmacie 30 (3). doi:10.1002/jlac.18390300310.
3. ^ http://webbook.nist.gov/cgi/cbook.cgi?ID=C56235&Units=SI&Mask=4#Thermo-Phase
4. ^ F. Brezina, J. Mollin, R. Pastorek, Z. Sindelar. Chemicke tabulky anorganickych sloucenin (Chemical tables of inorganic compounds). SNTL, 1986.
5. ^ Doherty R. E. (2000). "A History of the Production and Use of Carbon Tetrachloride, Tetrachloroethylene, Trichloroethylene and 1,1,1-Trichloroethane in the United States: Part 1--Historical Background; Carbon Tetrachloride and Tetrachloroethylene" (1): 69 - 81. doi:10.1006/enfo.2000.0010.
6. ^ Introduction to Nuclear Magnetic Resonance Spectroscopy, Michigan State University
7. ^ WF Seifert, A Bosma, A Brouwer, HF Hendriks, PJ (1994). "Vitamin A deficiency potentiates carbon tetrachloride-induced liver fibrosis in rats". Hepatology 19: 193-201.
8. ^ Liu KX, Kato Y, Yamazaki M, Higuchi O, Nakamura T, Sugiyama Y. (1993). "Decrease in the hepatic clearance of hepatocyte growth factor in carbon tetrachloride-intoxicated rats". Hepatology 17: 651-60.
9. ^ Recknagel R.O., Glende E.A., Dolak J.A., Waller R.L. (1989). "Mechanism of Carbon-tetrachloride Toxicity". Pharmacology Therapeutics (43): 139-154. doi:10.1016/0163-7258(89)90050-8.
10. ^ Recknagel RO (1967). "Carbon tetrachloride Hepatotoxicity". Pharmacological Reviews 19 (2): 145.
11. ^ Masuda Y (2006). "Learning toxicology from carbon tetrachloride-induced hepatotoxicity". Yakugaku Zasshi -Journal of the Pharmaceutical Society of Japan 126 (10): 885-899.
12. ^ Rood AS, McGavran PD, Aavenson JW, et al. (2001). "Stochastic estimates of exposure and cancer risk from carbon tetrachloride released to the air from the Rocky Flats Plant". Risk Analysis 21 (4): 675-695.
13. ^ Fraser P. (1997). "Chemistry of stratospheric ozone and ozone depletion". Australian Meteorological Magazine 46 (3): 185-193.
14. ^ Evans WFJ, Puckrin E (1996). "A measurement of the greenhouse radiation associated with carbon tetrachloride (CCl₄)". Geophysical Research Letters 23 (14): 1769-1772.
15. ^ Walker, S. J., R. F. Weiss & P. K. Salameh (2000). "Reconstructed histories of the annual mean atmospheric mole fractions for the halocarbons CFC-11, CFC-12, CFC-113 and carbon tetrachloride". Journal of Geophysical Research 105: 14285—14296.