Chloramine

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Chloramine
Chloramine
Chloramine
Other names Monochloramine
Chloramide
Identifiers
CAS number 10599-90-3
Properties
Molecular formula NH2Cl
Molar mass 51.48 g/mol
Appearance colorless
Melting point

−66 °C
Boiling point

°C
Solubility in other solvents Soluble
Related compounds
Related compounds Dichloramine
Nitrogen trichloride
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)
Infobox disclaimer and references

Chloramine (monochloramine) is a chemical compound with the formula NH2Cl. It is usually used as a dilute solution where it is used as a disinfectant. The term chloramine also refers to a family of organic compounds with the formulas R2NCl and RNCl2 (R is an organic group). Dichloramine, NHCl2, and nitrogen trichloride, NCl3, are also well known.

Contents

[edit] Synthesis and chemical reactions

NH2Cl is a highly unstable compound in concentrated form, much less as a pure liquid. Pure NH2Cl decomposes violently above −40 °C.[1] NH2Cl is, however, quite stable in dilute solution, and this considerable stability is the basis of its applications.

NH2Cl is prepared by the chemical reaction between ammonia and hypochlorous acid[2] under mildly alkaline conditions:

NH3 + HOCl → NH2Cl + H2O

The synthesis is conducted in dilute solution. In this reaction HOCl undergoes attack by the nucleophile NH3. At lower pH's, further chlorination occurs.

The above syntheses are useful but do not deliver NH2Cl in pure form. The pure compound can be prepared by contacting fluoroamine with calcium chloride:

NH2F + CaCl2 → NH2Cl + CaClF

NH2Cl is a key intermediate in the traditional synthesis of hydrazine.

Monochloramine oxidizes sulfhydrals and disulfides in the same manner as HClO,[3] but only possesses 0.4% of the biocidal effect of HClO.[4]

[edit] Uses in water treatment

NH2Cl is commonly used in low concentrations as a disinfectant in municipal water systems as an alternative to chlorination. This application is increasing. Chlorine (sometimes referred to as Free Chlorine) is being displaced by chloramine, which is much more stable and does not dissipate from the water before it reaches consumers. NH2Cl also exhibits less tendency to convert organic materials into chlorocarbons such as chloroform and carbon tetrachloride. Such compounds have been identified as carcinogens and in 1979 the U.S. EPA began regulating their levels in U.S. drinking water. Furthermore, water treated with chloramine lacks the distinct chlorine odour of the gaseous treatment and so has improved taste.

Chloramine in tap water gives a greenish cast to the water in bulk, versus the normally bluish cast to pure water or water containing only free chlorine disinfectant. This greenish color may be observed by filling a white polyethylene bucket with chloraminated tap water and comparing it to chloramine-free water such as distilled water or a sample from a swimming pool.

New swimming pool initially filled with chloramine-treated tap water, showing greenish color of chloramine in bulk water. The color is less apparent in smaller volumes, but is noticeable in a white 5-gallon bucket, or even faintly detectable in a glass tumbler on careful inspection.
New swimming pool initially filled with chloramine-treated tap water, showing greenish color of chloramine in bulk water. The color is less apparent in smaller volumes, but is noticeable in a white 5-gallon bucket, or even faintly detectable in a glass tumbler on careful inspection.

Chloramine can be removed from tap water by treatment with superchlorination (10 ppm or more of free chlorine, such as from a dose of sodium hypochlorite bleach or pool sanitizer) while maintaining a pH of about 7 (such as from a dose of hydrochloric acid). Hypochlorous acid from the free chlorine strips the ammonia from the chloramine, and the ammonia outgasses from the surface of the bulk water. This process takes about 24 hours for normal tap water concentrations of a few ppm of chloramine. Residual free chlorine can then be removed by exposure to bright sunlight for about 4 hours.

[edit] Situations where NH2Cl should be removed

Aquarium owners must remove the chloramine from their tap water because it is toxic to fish. Aging the water for a few days removes chlorine but not the more stable chloramine, which can be neutralised using products available at pet stores.

Many animals are sensitive to chloramine and it must be removed from water given to many animals in zoos.

Chloramine must also be removed from the water prior to use in kidney dialysis machines, as it would come in contact with the bloodstream across a permeable membrane. However, since chloramine is neutralized by the digestive process, kidney dialysis patients can still safely drink chloramine-treated water.

Home brewers use reducing agents such as sodium metabisulfite or potassium metabisulfite to remove chloramine from brewing liquor as it, unlike chlorine, cannot be removed by boiling (A.J. DeLange). Residual sodium can cause off flavors in beer (See Brewing, Michael Lewis) so potassium metabisulfite is preferred.

In swimming pools, chloramines are formed by the reaction of free chlorine with organic substances. Chloramines, compared to free chlorine, are both less effective as a sanitizer and more irritating to the eyes of swimmers. When swimmers complain of eye irritation from "too much chlorine" in a pool, the problem is typically a high level of chloramines, caused by too little chlorine in relation to the amount of organic matter. Pool test kits designed for use by homeowners are sensitive to both free chlorine and chloramines, which can be misleading.

[edit] Organic chloramines

A variety of organic chloramines are known and proven useful in organic synthesis. One example is N-chloromorpholine ClN(CH2CH2)2O, N-chloropiperidine, and N-chloroquinuclidinium chloride.[5]

[edit] Safety

NH2Cl is toxic in large quantities. US EPA regulations limit chloramine concentration to 4 parts per million (ppm). A typical target level in US public water supplies is 3 ppm.

[edit] References

  1. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  2. ^ Fair, G. M., J. C. Morris, S. L. Chang, I. Weil, and R. P. Burden. 1948. The behavior of chlorine as a water disinfectant. J. Am. Water Works Assoc. 40:1051-1061.
  3. ^ Jacangelo, J. G., V. P. Olivieri, and K. Kawata. 1987. Oxidation of sulfhydryl groups by monochloramine. Water Res. 21:1339-1344.
  4. ^ Morris, J. C. 1966. Future of chlorination. J. Am. Water Works Assoc. 58:1475-1482.
  5. ^ Lindsay Smith, J. R.; McKeer, L. C.; Taylor, J. M. "4-Chlorination of Electron-Rich Benzenoid Compounds: 2,4-Dichloromethoxybenzene" Organic Syntheses, CollectedVolume 8, p.167 (1993)..http://www.orgsyn.org/orgsyn/pdfs/CV8P0167.pdf describes several N-chloramines

[edit] External links