Monochloramine, often called simply chloramine, is the chemical compound with the fomula NH2Cl. Together with dichloramine (NHCl2) and nitrogen trichloride (NCl3), it is one of the three chloramines of ammonia.[1] It is an 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]

Other names
  • Monochloramine
  • Chloramide
  • Chloroazane
3D model (JSmol)
ECHA InfoCard 100.031.095
EC Number
  • 234-217-9
MeSH chloramine
UN number 3093
Molar mass 51.476 g mol−1
Appearance Colorless gas
Melting point −66 °C (−87 °F; 207 K)
Acidity (pKa) 14
Basicity (pKb) 15
Related compounds
Related amines
GHS pictograms
GHS Signal word Danger
H290, H314, H315, H319, H335, H372, H412
P234, P260, P261, P264, P270, P271, P273, P280, P301+330+331, P302+352, P303+361+353, P304+340, P305+351+338, P310, P312, P314, P321, P332+313, P337+313, P362, P363, P390, P403+233, P404, P405
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

The wholesale cost in the developing world is about US$13.80 to US$18.41 per 500 grams.[3]

Water treatment

Chloramine is used as a disinfectant for water. It is less aggressive than chlorine and more stable against light than hypochlorites.[4]

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, but still active, 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.[5]

Some of the unregulated byproducts may possibly pose greater health risks than the regulated chemicals.[6]

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.[7]

Swimming pool disinfection

In swimming pools, chloramines are formed by the reaction of free chlorine with amine groups present in organic substances, mainly bio-derived. 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 responsible for the distinctive "chlorine" smell of swimming pools.[8][9] Some pool test kits designed for use by homeowners do not distinguish free chlorine and chloramines, which can be misleading and lead to non-optimal levels of chloramines in the pool water.[10] There is also evidence that exposure to chloramine can contribute to respiratory problems, including asthma, among swimmers.[11] Respiratory problems related to chloramine exposure are common and prevalent among competitive swimmers.[12]


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.[13][14] Both iodinated disinfection by-products and N-nitrosodimethylamine have been shown to be genotoxic.[14]

Synthesis and chemical reactions

NH2Cl is a highly unstable compound in concentrated form. Pure NH2Cl decomposes violently above −40 °C (−40 °F).[15] Gaseous chloramine at low pressures and low concentrations of chloramine in aqueous solution are thermally slightly more stable. Chloramine is readily soluble in water and ether, but less soluble in chloroform and carbon tetrachloride.[4]


In dilute aqueous solution, chloramine is prepared by the reaction of ammonia with sodium hypochlorite:[4]

NH3 + NaOCl → NH2Cl + NaOH

This reaction is also the first step of the Olin Raschig process for 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 chloramine solution can be concentrated by vacuum distillation and by passing the vapor through potassium carbonate which absorbs the water. Chloramine can be extracted with ether.

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

2 NH3 + Cl2 NH2Cl + NH4Cl

Pure chloramine 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:[16]


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, chloramine is pH-neutral. It is an oxidizing agent (acidic solution: E° = −1.48 V, in basic solution E° = −0.81 V):[4]

NH2Cl + 2 H+ + 2 eNH+
+ Cl

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

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

Nu + NH3Cl+ → NuCl + NH3

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

Chloramine may also aminate nucleophiles (electrophilic amination):

Nu + NH2Cl → NuNH2 + Cl

The amination of ammonia with chloramine to form hydrazine is an example of this mechanism seen in the Olin 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,[17] but only possesses 0.4% of the biocidal effect of HClO.[18]

See also


  1. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  2. Lawrence, Stephen A. (2004). Amines: Synthesis, Properties and Applications. Cambridge University Press. p. 172. ISBN 9780521782845.
  3. "Chloramine". International Drug Price Indicator Guide. Archived from the original on 22 January 2018. Retrieved 8 December 2016.
  4. Hammerl, Anton; Klapötke, Thomas M. (2005), "Nitrogen: Inorganic Chemistry", Encyclopedia of Inorganic Chemistry (2nd ed.), Wiley, pp. 55–58
  6. Stuart W. Krasner (2009-10-13). "The formation and control of emerging disinfection by-products of health concern". Philosophical Transactions of the Royal Society A: Mathematical, Physical and Engineering Sciences. Philosophical Transactions of the Royal Society. 367 (1904): 4077–95. Bibcode:2009RSPTA.367.4077K. doi:10.1098/rsta.2009.0108. PMID 19736234.
  7. 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 1817676. PMID 17384768.
  8. Donegan, Fran J.; David Short (2011). Pools and Spas. Upper Saddle River, New Jersey: Creative Homeowner. ISBN 978-1-58011-533-9.
  9. "Controlling Chloramines in Indoor Swimming Pools". NSW Government. Archived from the original on 2011-04-03. Retrieved 2013-02-15.
  10. Hale, Chris (20 April 2016). "Pool Service Information". Into The Blue Pools. Retrieved 22 April 2016.
  11. Bougault, Valérie; et al. (2009). "The Respiratory Health of Swimmers". Sports Medicine. 39 (4): 295–312. doi:10.2165/00007256-200939040-00003. PMID 19317518.
  12. Lévesque, Benoit; Duchesne, Jean-François; Gingras, Suzanne; Lavoie, Robert; Prud'Homme, Denis; Bernard, Emmanuelle; Boulet, Louis-Philippe; Ernst, Pierre (2006-10-01). "The determinants of prevalence of health complaints among young competitive swimmers". International Archives of Occupational and Environmental Health. 80 (1): 32–39. doi:10.1007/s00420-006-0100-0. PMID 16586082.
  13. 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. PMID 17180964.
  14. 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.
  15. Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  16. Ura, Yasukazu; Sakata, Gozyo (2007). "Chloroamines". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. p. 5.
  17. Jacangelo, J. G.; Olivieri, V. P.; Kawata, K. (1987). "Oxidation of sulfhydryl groups by monochloramine". Water Res. 21 (11): 1339–1344. doi:10.1016/0043-1354(87)90007-8.
  18. Morris, J. C. (1966). "Future of chlorination". J. Am. Water Works Assoc. 58 (11): 1475–1482. doi:10.1002/j.1551-8833.1966.tb01719.x.
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