Chemistry:Monochloramine

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Monochloramine
Stereo, skeletal formula of chloramine with all explicit hydrogens added
Spacefill model of chloramine
Names
Other names
  • Chloramine
  • Chloramide[1]
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
EC Number
  • 234-217-9
KEGG
MeSH chloramine
UNII
UN number 3093
Properties
NH2Cl
Molar mass 51.476 g mol−1
Appearance Colorless gas
Melting point −66 °C (−87 °F; 207 K)
Acidity (pKa) 14
Basicity (pKb) 15
Hazards
Main hazards Corrosive acid
GHS pictograms GHS07: Harmful GHS08: Health hazard GHS05: Corrosive
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
Ingestion hazard Corrosive; nausea and vomiting
Inhalation hazard Corrosive
Eye hazard Irritation
Skin hazard Irritation
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).
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Infobox references
Tracking categories (test):

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]

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]

Safety

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]

Production

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+4), 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

Decomposition

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):

[math]\displaystyle{ K = \frac{ c_{\text{RR}'\text{NH}} \cdot c_\text{HOCl} }{ c_{\text{RR}'\text{NCl}} } }[/math]

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+4
3 NHCl2 + H+ ⇌ 2 NCl3 + NH+4

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

Reactions

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 e− → NH+4 + 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

References

  1. ↑ "CHLORAMINE". NOAA. https://cameochemicals.noaa.gov/chemical/19968. 
  2. ↑ 2.0 2.1 "Chloramine T Trihydrate SDS". https://www.fishersci.com/shop/msdsproxy?storeId=10652&productName=O1779250. [yes|permanent dead link|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) (in en). Amines: Synthesis, Properties and Applications. Cambridge University Press. p. 172. ISBN 9780521782845. https://books.google.com/books?id=35xwBwjhe2MC&pg=PA172. 
  5. ↑ 5.0 5.1 5.2 5.3 Hammerl, Anton; Klapötke, Thomas M. (2005), "Nitrogen: Inorganic Chemistry", Encyclopedia of Inorganic Chemistry (2nd ed.), Wiley, pp. 55–58 
  6. ↑ "Govinfo". 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". Philosophical Transactions of the Royal Society A: Mathematical, Physical and Engineering Sciences (Philosophical Transactions of the Royal Society) 367 (1904): 4077–95. doi:10.1098/rsta.2009.0108. PMID 19736234. Bibcode: 2009RSPTA.367.4077K. 
  8. ↑ Marie Lynn Miranda (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. 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. https://web.archive.org/web/20110403074224/http://www.health.nsw.gov.au/factsheets/environmental/control_chloramines.html. Retrieved 2013-02-15. 
  11. ↑ Hale, Chris (20 April 2016). "Pool Service Information". http://intothebluesrq.com/newsletter/. 
  12. ↑ Bougault, Valérie (2009). "The Respiratory Health of Swimmers". Sports Medicine 39 (4): 295–312. doi:10.2165/00007256-200939040-00003. PMID 19317518. 
  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. 
  14. ↑ "The smell of chlorine: nostalgic or noxious?" (in en-US). 2016-08-22. https://rheemthermal.com.au/blog/health-and-safety/the-smell-of-chlorine-nostalgic-or-noxious/. 
  15. ↑ "Chloramines: Understanding "Pool Smell"". https://chlorine.americanchemistry.com/Science-Center/Chlorine-Compound-of-the-Month-Library/Chloramines-Understanding-Pool-Smell/. 
  16. ↑ "The Chlorine Smell From Pools May Actually Indicate Bodily Fluids Mixed In The Water, According To The CDC" (in en). https://www.bustle.com/p/the-chlorine-smell-from-pools-may-actually-indicate-bodily-fluids-mixed-in-the-water-according-to-the-cdc-17989837. 
  17. ↑ "Chemical Irritation of the Eyes and Lungs | Healthy Swimming | Healthy Water | CDC" (in en-us). 2019-05-15. https://www.cdc.gov/healthywater/swimming/swimmers/rwi/chemical-irritants.html. 
  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. 
  19. ↑ 19.0 19.1 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. ↑ 20.0 20.1 20.2 Allison Kite (July 20, 2022). "'Time bomb' lead pipes will be removed. But first water utilities have to find them". Midwest Newsroom. NPR. https://www.npr.org/sections/health-shots/2022/07/20/1112049811/lead-pipe-removal. 
  21. ↑ "Lead and Copper Rule Improvements". https://www.epa.gov/ground-water-and-drinking-water/lead-and-copper-rule-improvements. 
  22. ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN:0-12-352651-5.
  23. ↑ Ura, Yasukazu; Sakata, Gozyo (2007). "Ullmann's Encyclopedia of Industrial Chemistry". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. p. 5. 
  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. 

External links



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