(Redirected from NH2Cl)

Monochloramine, often called chloramine, is the chemical compound with the formula NH2Cl. Together with dichloramine (NHCl2) and nitrogen trichloride (NCl3), it is one of the three chloramines of ammonia.[3] It is a 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.[4]

Stereo, skeletal formula of chloramine with all explicit hydrogens added
Spacefill model of chloramine
Other names
  • Chloramine
  • Chloramide[1]
3D model (JSmol)
ECHA InfoCard 100.031.095 Edit this at Wikidata
EC Number
  • 234-217-9
MeSH chloramine
UN number 3093
  • InChI=1S/ClH2N/c1-2/h2H2 ☒N
  • NCl
Molar mass 51.476 g mol−1
Appearance Colorless gas
Melting point −66 °C (−87 °F; 207 K)
Acidity (pKa) 14
Basicity (pKb) 15
Occupational safety and health (OHS/OSH):
Main hazards
Corrosive acid
Ingestion hazards
Corrosive; nausea and vomiting
Inhalation hazards
Eye hazards
Skin hazards
GHS labelling:
GHS07: Exclamation mark GHS08: Health hazard GHS05: Corrosive
H290, H314, H315, H319, H335, H372, H412
P234, P260, P261, P264, P270, P271, P273, P280, P301+P330+P331, P302+P352, P303+P361+P353, P304+P340, P305+P351+P338, P310, P312, P314, P321, P332+P313, P337+P313, P362, P363, P390, P403+P233, P404, P405, P501
NFPA 704 (fire diamond)
Lethal dose or concentration (LD, LC):
935 mg/kg (rat, oral)[2]
Related compounds
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Water treatment


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

Drinking water disinfection


Chloramine 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 less reactive and does not dissipate as rapidly as free chlorine. Chloramine 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 (EPA) began regulating their levels in US drinking water.[6]

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

Due to its acidic nature, 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. Fortunately, water treatment plants can add caustic chemicals at the plant which have the dual purpose of reducing the corrosivity of the water, and stabilizing the disinfectant.[8]

Swimming pool disinfection


In swimming pools, chloramines are formed by the reaction of free chlorine with amine groups present in organic substances, mainly those biological in origin (e.g., urea in sweat and urine). 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, which is often misattributed to elemental chlorine by the public.[9][10] 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.[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]

Though chloramine's distinctive smell has been described by some as pleasant and even nostalgic,[14] its formation in pool water as a result of bodily fluids being exposed to chlorine can be minimised by encouraging showering and other hygiene methods prior to entering the pool,[15] as well as refraining from swimming while suffering from digestive illnesses and taking breaks to use the bathroom.[16][17]



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.[18][19] Both iodinated disinfection by-products and N-nitrosodimethylamine have been shown to be genotoxic, causing damage to the genetic information within a cell resulting in mutations which may lead to cancer.[19]

Lead poisoning incidents


In the year 2000, Washington, DC, switched from chlorine to monochloramine, causing lead to leach from unreplaced pipes. The number of babies with elevated blood lead levels rose about tenfold, and by one estimate fetal deaths rose between 32% and 63%.[20]

Trenton, Missouri made the same switch, causing about one quarter of tested households to exceed EPA drinking water lead limits in the period from 2017 to 2019. 20 children tested positive for lead poisoning in 2016 alone.[20] In 2023, Virginia Tech Professor Marc Edwards said lead spikes occur in several water utility system switchovers per year, due to lack of sufficient training and lack of removal of lead pipes.[20] Lack of utility awareness that lead pipes are still in use is also part of the problem; the EPA has required all water utilities in the United States to prepare a complete lead pipe inventory by October 16, 2024.[21]

Synthesis and chemical reactions


Chloramine is a highly unstable compound in concentrated form. Pure chloramine decomposes violently above −40 °C (−40 °F).[22] 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.[5]



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

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:[23]

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, chloramine 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 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,[24] but only possesses 0.4% of the biocidal effect of HClO.[25]

See also



  1. ^ "CHLORAMINE". CAMEO Chemicals. NOAA.
  2. ^ a b "Chloramine T Trihydrate SDS". Fisher.[permanent dead link]
  3. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  4. ^ Lawrence, Stephen A. (2004). Amines: Synthesis, Properties and Applications. Cambridge University Press. p. 172. ISBN 9780521782845.
  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. ^ "Govinfo" (PDF).
  7. ^ 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. 367 (1904). Philosophical Transactions of the Royal Society: 4077–95. Bibcode:2009RSPTA.367.4077K. doi:10.1098/rsta.2009.0108. PMID 19736234.
  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 1817676. 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. PMID 19317518. S2CID 26017985.
  13. ^ 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. S2CID 21688495.
  14. ^ "The smell of chlorine: nostalgic or noxious?". Rheem Thermal Swimming Pool Heating. 2016-08-22. Retrieved 2020-11-22.
  15. ^ "Chloramines: Understanding "Pool Smell"". Retrieved 2020-11-22.
  16. ^ "The Chlorine Smell From Pools May Actually Indicate Bodily Fluids Mixed In The Water, According To The CDC". Bustle. Retrieved 2020-11-22.
  17. ^ "Chemical Irritation of the Eyes and Lungs | Healthy Swimming | Healthy Water | CDC". 2019-05-15. Retrieved 2020-11-22.
  18. ^ 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. S2CID 41960634.
  19. ^ 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.
  20. ^ a b c Allison Kite (July 20, 2022). "'Time bomb' lead pipes will be removed. But first water utilities have to find them". NPR. Midwest Newsroom.
  21. ^ "Lead and Copper Rule Improvements". 4 May 2022.
  22. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  23. ^ Ura, Yasukazu; Sakata, Gozyo (2007). "Chloroamines". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. p. 5. ISBN 978-3527306732.
  24. ^ 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.
  25. ^ 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.