In chemistry, trihalomethanes (THMs) are chemical compounds in which three of the four hydrogen atoms of methane (CH4) are replaced by halogen atoms. Many trihalomethanes find uses in industry as solvents or refrigerants. THMs are also environmental pollutants, and many are considered carcinogenic. Trihalomethanes with all the same halogen atoms are called haloforms.

Table of common trihalomethanes

Common trihalomethanes (ordered by molecular weight)
Molecular

formula

 IUPAC name CAS registry number Common name Other names Molecule
CHF3 trifluoromethane 75-46-7 fluoroform Freon 23, R-23, HFC-23 Fluoroform
CHClF2 chlorodifluoromethane 75-45-6 chlorodifluoromethane R-22, HCFC-22 Chlorodifluoromethane
CHCl3 trichloromethane 67-66-3 chloroform R-20, methyl trichloride Chloroform
CHBrCl2 bromodichloromethane 75-27-4 bromodichloromethane dichlorobromomethane, BDCM Bromodichloromethane
CHBr2Cl dibromochloromethane 124-48-1 dibromochloromethane chlorodibromomethane, CDBM Dibromochloromethane
CHBr3 tribromomethane 75-25-2 bromoform methyl tribromide Bromoform
CHI3 triiodomethane 75-47-8 iodoform methyl triiodide Iodoform

Industrial uses

Only chloroform has significant applications of the haloforms. In the predominant application, chloroform is required for the production of tetrafluoroethylene (TFE), precursor to teflon.[1] Chloroform is fluorinated by reaction with hydrogen fluoride to produce chlorodifluoromethane (R-22). Pyrolysis of chlorodifluoromethane (at 550-750 °C) yields TFE, with difluorocarbene as an intermediate.

Refrigerants and solvents

Trihalomethanes released to the environment break down faster than chlorofluorocarbons (CFCs), thereby doing much less damage to the ozone layer. Trifluoromethane and chlorodifluoromethane are both used as refrigerants. Chlorodifluoromethane is a refrigerant HCFC, or hydrochlorofluorocarbon, while fluoroform is an HFC, or hydrofluorocarbon. Fluoroform is not ozone depleting.

Chloroform is a common solvent in organic chemistry.

Occurrence and production

The total global flux of chloroform through the environment is approximately 660000 tonnes per year,[2] and about 90% of emissions are natural in origin. Many kinds of seaweed produce chloroform, and fungi are believed to produce chloroform in soil.[3]

Most of the haloforms — specifically, chloroform (CHCl3), bromoform (CHBr3), and iodoform (CHI3) — are easy to prepare through the haloform reaction, although this method does not lend itself to bulk syntheses. (Fluoroform (CHF3) cannot be prepared in this manner.)

Chloroform is produced by heating mixtures of methane or methyl chloride with chlorine. Dichloromethane is a coproduct.[4]

Bromochlorofluoromethane is one of the simplest possible stable chiral compounds, and is used for studies.

Regulation

Trihalomethanes were the subject of the first drinking water regulations issued after passage of the U.S. Safe Drinking Water Act in 1974.[5]

The EPA limits the total concentration of the four chief constituents (chloroform, bromoform, bromodichloromethane, and dibromochloromethane), referred to as total trihalomethanes (TTHM), to 80 parts per billion in treated water.[6]

Traces of chloroform are produced in swimming pools.[7][8][9][10]

References

  1. ↑ Dae Jin Sung; Dong Ju Moon; Yong Jun Lee; Suk-In Hong (2004). "Catalytic Pyrolysis of Difluorochloromethane to Produce Tetrafluoroethylene". International Journal of Chemical Reactor Engineering. 2: A6. doi:10.2202/1542-6580.1065. S2CID 97895482.
  2. ↑ Gribble, Gordon W. (2004). "Natural Organohalogens: A New Frontier for Medicinal Agents?". Journal of Chemical Education. 81 (10): 1441. Bibcode:2004JChEd..81.1441G. doi:10.1021/ed081p1441.
  3. ↑ Cappelletti, M. (2012). "Microbial degradation of chloroform". Applied Microbiology and Biotechnology. 96 (6): 1395–409. doi:10.1007/s00253-012-4494-1. PMID 23093177. S2CID 12429523.
  4. ↑ Rossberg, Manfred; Lendle, Wilhelm; Pfleiderer, Gerhard; Tögel, Adolf; Dreher, Eberhard-Ludwig; Langer, Ernst; Rassaerts, Heinz; Kleinschmidt, Peter; Strack, Heinz; Cook, Richard; Beck, Uwe; Lipper, Karl-August; Torkelson, Theodore R.; Löser, Eckhard; Beutel, Klaus K.; Mann, Trevor (2006). "Chlorinated Hydrocarbons". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a06_233.pub2. ISBN 3527306730.
  5. ↑ EPA Alumni Association: Senior EPA officials discuss early implementation of the Safe Drinking Water Act of 1974, Video, Transcript (see pages 12-13).
  6. ↑ "EPA | Envirofacts | ICR | Regulations". archive.epa.gov. Retrieved 2021-10-11.
  7. ↑ Lindstrom, A B; Pleil, J.D.; Berkoff, D.C. (1997). "Alveolar breath sampling and analysis to assess trihalomethane exposures during competitive swimming training". Environmental Health Perspectives. 105 (6): 636–642. doi:10.1289/ehp.97105636. ISSN 0091-6765. PMC 1470079. PMID 9288498.
  8. ↑ Drobnic, Franchek; Freixa, Assumpci??; Casan, Pere; Sanchis, Joaqu??N; Guardino, Xavier (1996). "Assessment of chlorine exposure in swimmers during training". Medicine & Science in Sports & Exercise. 28 (2): 271–274. doi:10.1097/00005768-199602000-00018. ISSN 0195-9131. PMID 8775165.
  9. ↑ Aiking, Harry; van Ackert, Manila B.; Schölten, Rob J.P.M.; Feenstra, Jan F.; Valkenburg, Hans A. (1994). "Swimming pool chlorination: a health hazard?". Toxicology Letters. 72 (1–3): 375–380. doi:10.1016/0378-4274(94)90051-5. ISSN 0378-4274. PMID 7911264.
  10. ↑ Nickmilder, M.; Bernard, A. (2011). "Associations between testicular hormones at adolescence and attendance at chlorinated swimming pools during childhood". International Journal of Andrology. 34 (5pt2): e446–e458. doi:10.1111/j.1365-2605.2011.01174.x. ISSN 0105-6263. PMC 3229674. PMID 21631527.
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