Chloramines are derivatives of ammonia by substitution of one, two or three hydrogen atoms with chlorine atoms: monochloramine (chloroamine, NH2Cl), dichloramine (NHCl2), and nitrogen trichloride (NCl3).[1][full citation needed] The term chloramine also refers to a family of organic compounds with the formulas R2NCl and RNCl2 (where R is an organic group).

Monochloramine (chloramine) is an inorganic compound with the formula NH2Cl. It is an unstable colorless liquid at its melting point of −66 °C (−87 °F), but it is usually handled as a dilute aqueous solution, in which form it is sometimes used as a disinfectant. Chloramine is too unstable to have its boiling point measured.[2] It is listed as a tumorigen and mutagen.[3][better source needed]

The wholesale cost in the developing world is about 13.80 to 18.41 U.S. dollars per 500 grams.[4]

Water treatment

Chloramine is used as a disinfectant for water because it is less aggressive than chlorine and more stable against light than hypochlorites.[5]

Drinking water disinfection

NH2Cl is commonly used in low concentrations as a secondary disinfectant in municipal water distribution systems as an alternative to chlorination. This application is increasing. Chlorine (referred to in water treatment as free chlorine) is being displaced by chloramine—to be specific monochloramine—which is much more stable and does not dissipate as rapidly as free chlorine. NH2Cl also has a much lower, however still present, tendency than free chlorine to convert organic materials into chlorocarbons such as chloroform and carbon tetrachloride. Such compounds have been identified as carcinogens and in 1979 the United States Environmental Protection Agency began regulating their levels in U.S. drinking water.[6]

Some of the unregulated byproducts may possibly pose greater health risks than the regulated chemicals.[7]

Adding chloramine to the water supply may increase exposure to lead in drinking water, especially in areas with older housing; this exposure can result in increased lead levels in the bloodstream, which may pose a significant health risk.[8]

Swimming pool disinfection

In swimming pools, chloramines are formed by the reaction of free chlorine with amine groups present in organic substances, such as urine, sweat and shed skin cells. Chloramines, compared to free chlorine, are both less effective as a sanitizer and, if not managed correctly, more irritating to the eyes of swimmers. Chloramines are also responsible for the distinctive "chlorine" smell of swimming pools.[9][10] Some pool test kits designed for use by homeowners are not able to distinguish free chlorine and chloramines, which can be misleading and lead to non-optimal levels of chloramines in the pool water.[11] There is also evidence that exposure to chloramine can contribute to respiratory problems, including asthma, among swimmers.[12] Respiratory problems related to chloramine exposure are common and prevalent among competitive swimmers.[13]


US EPA drinking water quality standards limit chloramine concentration for public water systems to 4 parts per million (ppm) based on a running annual average of all samples in the distribution system. In order to meet EPA-regulated limits on halogenated disinfection by-products, many utilities are switching from chlorination to chloramination. While chloramination produces fewer regulated total halogenated disinfection by-products, it can produce greater concentrations of unregulated iodinated disinfection byproducts and N-nitrosodimethylamine.[14][15] Both iodinated disinfection by-products and N-nitrosodimethylamine have been shown to be genotoxic.[15]

Synthesis and chemical reactions

NH2Cl is a highly unstable compound in concentrated form. Pure NH2Cl decomposes violently above −40 °C (−40 °F).[16] Gaseous chloroamine at low pressures and low concentrations of chloroamine in aqueous solution are thermally slightly more stable. Chloroamine is readily soluble in water and ether, but less soluble in chloroform and carbon tetrachloride.[5]


In dilute aqueous solution, chloroamine is prepared by the reaction of ammonia with sodium hypochlorite:[5]

NH3 + ClO → NH2Cl + OH

This is also the first step of the Raschig hydrazine synthesis. The reaction has to be carried out in a slightly alkaline medium (pH 8.5–11). The acting chlorinating agent in this reaction is hypochlorous acid (HOCl), which has to be generated by protonation of hypochlorite, and then reacts in a nucleophilic substitution of the hydroxyl against the amino group. The reaction occurs quickest at around pH 8. At higher pH values the concentration of hypochlorous acid is lower, at lower pH values ammonia is protonated to form ammonium ions NH+
, which do not react further.

The chloroamine solution can be concentrated by vacuum distillation and by passing the vapor through potassium carbonate which absorbs the water. Chloroamine can be extracted with ether.

Gaseous chloroamine can be obtained from the reaction of gaseous ammonia with chlorine gas (diluted with nitrogen gas):

2 NH3(g) + Cl2(g) ⇌ NH2Cl(g) + NH4Cl(s)

Pure chloroamine can be prepared by passing fluoroamine through calcium chloride:

2 NH2F + CaCl2 → 2 NH2Cl + CaF2


The covalent N−Cl bonds of chloramines are readily hydrolyzed with release of hypochlorous acid:[17]

RR′NCl + H2O ⇌ RR′NH + HOCl

The quantitative hydrolysis constant (K value) is used to express the bactericidal power of chloramines, which depends on their generating hypochlorous acid in water. It is expressed by the equation below, and is generally in the range 10−4 to 10−10 (2.8×10−10 for monochloramine):

In aqueous solution, chloramine slowly decomposes to dinitrogen and ammonium chloride in a neutral or mildly alkaline (pH ≤ 11) medium:

3 NH2Cl → N2 + NH4Cl + 2 HCl

However, only a few percent of a 0.1 M chloramine solution in water decomposes according to the formula in several weeks. At pH values above 11, the following reaction with hydroxide ions slowly occurs:

3 NH2Cl + 3 OH → NH3 + N2 + 3 Cl + 3 H2O

In an acidic medium at pH values of around 4, chloramine disproportionates to form dichloramine, which in turn disproportionates again at pH values below 3 to form nitrogen trichloride:

2 NH2Cl + H+ ⇌ NHCl2 + NH+
3 NHCl2 + H+ ⇌ 2 NCl3 + NH+

At low pH values, nitrogen trichloride dominates and at pH 3–5 dichloramine dominates. These equilibria are disturbed by the irreversible decomposition of both compounds:

NHCl2 + NCl3 + 2 H2O → N2 + 3 HCl + 2 HOCl


In water, chloroamine is pH-neutral. It is an oxidizing agent (acidic solution: E° = −1.48 V, in basic solution E° = −0.81 V):[5]

NH2Cl + 2 H+ + 2 eNH+
+ Cl

Reactions of chloroamine include radical, nucleophilic, and electrophilic substitution of chlorine, electrophilic substitution of hydrogen, and oxidative additions.

Chloroamine can, like hypochlorous acid, donate positively charged chlorine in reactions with nucleophiles (Nu):

Nu + NH3Cl+ → NuCl + NH3

Examples of chlorination reactions include transformations to dichloroamine and nitrogen trichloride in acidic medium, as described in the decomposition section.

Chloroamine may also aminate nucleophiles (electrophilic amination):

Nu + NH2Cl → NuNH2 + Cl

The amination of ammonia with chloroamine to form hydrazine is an example of this mechanism (the Raschig process):

NH2Cl + NH3 + NaOH → N2H4 + NaCl + H2O

Chloramine electrophilically aminates itself in neutral and alkaline media to start its decomposition:

2 NH2Cl → N2H3Cl + HCl

The chlorohydrazine (N2H3Cl) formed during self-decomposition is unstable and decomposes itself, which leads to the net decomposition reaction:

3 NH2Cl → N2 + NH4Cl + 2 HCl

Monochloramine oxidizes sulfhydryls and disulfides in the same manner as hypochlorous acid,[18] but only possesses 0.4% of the biocidal effect of HClO.[19]

Removing from water

Chloramines should be removed from water for dialysis, aquariums, hydroponic applications, and homebrewing beer.[citation needed]. Chloramine must be removed from water prior to use in kidney dialysis machines because it can cause hemolytic anemia if it enters the blood stream.[20] In hydroponic applications, chloramine stunts the growth of plants.[21]

When a chemical or biological process that changes the chemistry of chloramines is used, it falls under reductive dechlorination. Other techniques use physical—not chemical—methods for removing chloramines.[citation needed]

Ultraviolet light

The use of ultraviolet light for chlorine or chloramine removal is an established technology that has been widely accepted in pharmaceutical, beverage, and dialysis applications.[22] UV is also used for disinfection at aquatic facilities [23].

Ascorbic acid and sodium ascorbate

Ascorbic acid (vitamin C) and sodium ascorbate completely neutralize both chlorine and chloramine, but degrade in a day or two, which makes them usable only for short-term applications. SFPUC determined that 1000 mg of vitaminC  tablets, crushed and mixed in with bath water, completely remove chloramine in a medium-size bathtub without significantly depressing pH.[24]

Activated carbon

Activated carbon has been used for chloramine removal long before catalytic carbon, a form of activated carbon, became available[citation needed]; standard activated carbon requires a very long contact time, which means a large volume of carbon is needed. For thorough removal, up to four times the contact time of catalytic carbon may be required.[citation needed]

Most dialysis units now depend on granular activated carbon (GAC) filters, two of which should be placed in series so that chloramine breakthrough can be detected after the first one, before the second one fails.[25] Additionally, sodium metabisulfite injection may be used in certain circumstances.[26][full citation needed]

Campden tablets

Home brewers use reducing agents such as sodium metabisulfite or potassium metabisulfite (both proprietary sold as Campden tablets) to remove chloramine from brewing fermented beverages. However, residual sodium can cause off flavors in beer[27][full citation needed] so potassium metabisulfite is preferred.

Sodium thiosulfate

Sodium thiosulfate is used to dechlorinate tapwater for aquariums or treat effluent from wastewater treatments prior to release into rivers[citation needed]. The reduction reaction is analogous to the iodine reduction reaction. Treatment of tapwater requires between 0.1 and 0.3 grams of pentahydrated (crystalline) sodium thiosulfate per 10 L of water[citation needed]. Many animals are sensitive to chloramine, and it must be removed from water given to many animals in zoos[citation needed].

Other methods

Chloramine, like chlorine, can be removed by boiling and aging. However, time required to remove chloramine is much longer than that of chlorine. The time required to remove half of the chloramine (half-life) from 10 US gallons (38 l; 8.3 imp gal) of water by boiling is 26.6 hours, whereas the half-life of free chlorine in boiling 10 gallons of water is only 1.8 hours.[28][better source needed]

Organic chloramines

A variety of organic chloramines are known and proven useful in organic synthesis. Examples include N-chloromorpholine (ClN(CH2CH2)2O), N-chloropiperidine, and N-chloroquinuclidinium chloride.[29]

Reduction of organic chloramines

Chloramines are often an unwanted side-product of oxidation reactions of organic compounds (with amino groups) with bleach. The reduction of chloramines back into amines can be carried out through a mild hydride donor. Sodium borohydride will reduce chloramines, but this reaction is greatly accelerated with acid catalysis.[citation needed]

See also


  1. ^ Clause 2.4 Chloramines ISO 7393-2
  2. ^ Lawrence, Stephen A. (2004). Amines: Synthesis, Properties and Applications. Cambridge University Press. p. 172. ISBN 9780521782845. 
  3. ^ https://www.wolframalpha.com/input/?i=chloramine
  4. ^ "Chloramine". International Drug Price Indicator Guide. Archived from the original on 22 January 2018. Retrieved 8 December 2016. 
  5. ^ a b c d Hammerl, Anton; Klapötke, Thomas M. (2005), "Nitrogen: Inorganic Chemistry", Encyclopedia of Inorganic Chemistry (2nd ed.), Wiley, pp. 55–58 
  6. ^ http://www.epa.gov/fedrgstr/EPA-WATER/2006/January/Day-04/w03.pdf
  7. ^ Stuart W. Krasner (2009-10-13). "The formation and control of emerging disinfection by-products of health concern". 367 (1904). Philosophical Transactions of the Royal Society: 4077–95. doi:10.1098/rsta.2009.0108. 
  8. ^ Marie Lynn Miranda; et al. (February 2007). "Changes in Blood Lead Levels Associated with Use of Chloramines in Water Treatment Systems". Environmental Health Perspectives. 115 (2): 221–5. doi:10.1289/ehp.9432. PMC 1817676Freely accessible. PMID 17384768. 
  9. ^ Donegan, Fran J.; David Short (2011). Pools and Spas. Upper Saddle River, New Jersey: Creative Homeowner. ISBN 978-1-58011-533-9. 
  10. ^ "Controlling Chloramines in Indoor Swimming Pools". NSW Government. Archived from the original on 2011-04-03. Retrieved 2013-02-15. 
  11. ^ Hale, Chris (20 April 2016). "Pool Service Information". Into The Blue Pools. Retrieved 22 April 2016. 
  12. ^ Bougault, Valérie; et al. (2009). "The Respiratory Health of Swimmers". Sports Medicine. 39 (4): 295–312. doi:10.2165/00007256-200939040-00003. 
  13. ^ "The determinants of prevalence of health complaints among young competitive swimmers". International Archives of Occupational and Environmental Health. 80 (1): 32–39. 2006-10-01. doi:10.1007/s00420-006-0100-0. 
  14. ^ Krasner, Stuart W.; Weinberg, Howard S.; Richardson, Susan D.; Pastor, Salvador J.; Chinn, Russell; Sclimenti, Michael J.; Onstad, Gretchen D.; Thruston, Alfred D. (2006). "Occurrence of a New Generation of Disinfection Byproducts". Environmental Science & Technology. 40 (23): 7175–7185. doi:10.1021/es060353j. 
  15. ^ a b Richardson, Susan D.; Plewa, Michael J.; Wagner, Elizabeth D.; Schoeny, Rita; DeMarini, David M. (2007). "Occurrence, genotoxicity, and carcinogenicity of regulated and emerging disinfection by-products in drinking water: A review and roadmap for research". Mutation Research/Reviews in Mutation Research. 636 (1–3): 178–242. doi:10.1016/j.mrrev.2007.09.001. PMID 17980649. 
  16. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  17. ^ Ura, Yasukazu; Sakata, Gozyo (2007). "Chloroamines". Ullmann's Encyclopedia of Industrial Chemistry (7th ed.). Wiley. p. 5. 
  18. ^ Jacangelo, J. G.; Olivieri, V. P.; Kawata, K. (1987). "Oxidation of sulfhydryl groups by monochloramine". Water Res. 21: 1339–1344. 
  19. ^ Morris, J. C. (1966). "Future of chlorination". J. Am. Water Works Assoc. 58: 1475–1482. 
  20. ^ Hakim, Nadey (2009). Artificial Organs. London: Springer-Verlag. p. 51. ISBN 9781848822818. Retrieved 2014-06-14. Water that contains chloramine is safe for people to drink, bathe, and cook in because the digestive process neutralizes it. Chloramine can, however, easily harm patients if it enters the blood stream during the dialysis process causing hemolytic anemia. 
  21. ^ Date, S.; Terabayashi, S.; Kobayashi, Y.; Fujime, Y. (2005), "Effects of chloramines concentration in nutrient solution and exposure time on plant growth in hydroponically cultured lettuce", Scientia Horticulturae, 103 (3): 257–265, doi:10.1016/j.scienta.2004.06.019 
  22. ^ Adelstein, Ben (2010-10-13). "Considering UV technology in water bottling". Watertechonline.com. Archived from the original on 2013-02-09. Retrieved 2013-11-23. 
  23. ^ "dechloraminator". UVgermi.com. 2017-08-30. 
  24. ^ "Questions Regarding Chlorine and Chloramine Removal From Water (Updated June 2013)". San Francisco Public Utilities Commission. Retrieved 2013-11-23. 
  25. ^ Ward, D. M. (Oct 1996). "Chloramine removal from water used in hemodialysis". Adv. Ren. Replace Ther. 3 (4): 337–347. PMID 8914698. 
  26. ^ Handbook of Dialysis, page 81
  27. ^ Brewing, Michael Lewis
  28. ^ "Experiments in Removing Chlorine and Chloramine From Brewing Water" (PDF). 1998-11-03. Archived from the original (PDF) on 2013-11-10. Retrieved 2013-11-23. 
  29. ^ Lindsay Smith, J. R.; McKeer, L. C.; Taylor, J. M. "4-Chlorination of Electron-Rich Benzenoid Compounds: 2,4-Dichloromethoxybenzene". Organic Syntheses. 67: 222. doi:10.15227/orgsyn.067.0222. ; Collective Volume, 8, p. 167 

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