Some of the unregulated byproducts may possibly pose greater health risks than the regulated chemicals.[7][clarification needed]
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, instead of simply urinating in the pool.[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]
Another newly-identified byproduct of chloramine is chloronitramide anions, whose toxicity has not yet been determined.[20]
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%.[21]
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.[21] 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.[21] 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.[22]
Synthesis and chemical reactions
Chloramine is a highly unstable compound in concentrated form. Pure chloramine decomposes violently above −40 °C (−40 °F).[23] 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]
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):
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:
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]
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.
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,[25] but only possesses 0.4% of the biocidal effect of HClO.[26]
^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. PMID16586082. S2CID21688495.
^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. PMID17180964. S2CID41960634.
^ abRichardson, 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. PMID17980649.
^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.
