Precursor Compounds and the Haloform Reaction

Since 1975, many investigators have assumed that the ubiquitous appearance of chloroform (CHCl₃) and other THM's (trihalomethanes, or haloforms) in chlorinated water can be explained by the mechanisms involved in the ''haloform reaction" and that the principal precursors of THM's that are found in natural waters are humic substances. As discussed subsequently in the section pertaining to chlorine chemistry, the haloform reaction will proceed only if specific functional groups are present in the available pool of organic compounds. It is likely that the haloform reaction does occur when natural waters are chlorinated and that humic substances provide the necessary functional groups, but it is not certain that either of these postulates is true. For that reason, both topics—the haloform reaction and humic substances—merit further discussion.

Humic Substances

As was mentioned previously, it is an attractive assumption that naturally occurring humic substances, which are derived from the structural components of living and decaying plants and/or soil dissolution and runoff, provide the most ubiquitous source of haloform precursors in natural water systems. Only limited information is available concerning the structure of these complex natural products, and it is not yet known whether all the major structural features have been identified, if any structural differences exist among the humic substances in waters from different geographic areas, and if these substances are closely or distantly related to soil humic and marine humic materials.

The term "humic acid" is generic and refers to that fraction of soil organic material that is soluble in alkaline solutions but insoluble in acid and ethyl alcohol (Christman and Oglesby, 1971). The fraction that is soluble in acid is commonly labelled "fulvic acid," and that material precipitated by acid but soluble in ethyl alcohol is "hymatomelanic acid." Soils vary widely in their relative compositions of these acids, but aquatic organic material behaves operationally as fulvic acid (Black and Christman, 1963), which typically contains more oxygen and less nitrogen than the humic acid fraction in both soil and aquatic organic matter. Marine organic matter (including sedimentary material) is derived largely from marine organisms and contains more sulfur than its fresh water equivalent (Nissenbaum and Kaplan, 1972; Stuermer and Harvey, 1978). Christman and Oglesby (1971), Steelink (1977), Schnitzer and Kahn (1972), and Dubach et al. (1964) have reported the presence of carboxyl, phenolic and alcoholic hydroxyl, carboxyl, and methoxyl functional groups in humic material. It would appear that the more oxygenated fulvic acid fraction has a greater carboxyl acidity than the humic acid fraction.

Whittaker and Likens (1973) estimated that 90% of the terrestrial biospheric carbon (standing biomass) is tied up in woody tissue. Lignin is a dominant (20%-40%) chemical entity in woody tissue. Because of its refractory nature, it is probably a principal precursor of soil humus, although a myriad of other natural products unquestionably contribute to the complex pool of soil organic matter. Lignin itself is a mixed polymer of guaiacyl (I), syringyl (II), and p-hydroxyphenylpropane (III) aromatic moieties:

No other substitution patterns are known in nature and no other length of alkyl side chain has been found in lignin from any source. Oxidative degradation of lignin produces, therefore, only three aromatic substitution patterns (I, II, and III), although the relative amounts of each vary among the gymnosperms, angiosperms, and the grasses. Intermonomeric linkages in the lignin macromolecule are of both carbon-to-carbon and ether linkage types. The largest single contributor is believed to be the b-4' ether configuration. Side-chain carbon atoms may be in various states of oxygenation or unsaturation, and may contain methyl ketone, allyl, and secondary alcohol configurations.

Significant changes occur in the humification process as reflected by comparative functional group data for lignin and soil humic acid (Table III-1).

TABLE III-1

Comparative Functional Group Analysis of Soil Humic Acid and Spruce Lignin.

This process, which is oxidative in nature, may strongly affect the characteristics of aquatic humic material. Microbial mediation is apparent when there is a marked decrease in methoxyl groups and increases in phenolic hydroxyl and carboxyl acidity.

The contribution of woody tissues to marine humus is not apparent from the results of degradative experiments on marine fulvic acids, which are considered to be autochthonous materials. Degradation of both soil humic acid and aquatic humic material reflects a partial lignitic origin (Table III-2), although a variety of other aromatic patterns (m-dihydroxy) and aliphatic chain lengths (C₂—C₁₇) must result from other natural product sources. The data in Table III-2 indicate key areas of inadequacy in our knowledge of the chemical nature of aquatic humic substances. It is not possible to model natural aquatic humic material with a desirable degree of chemical accuracy, and it certainly is not possible to state that THM's, which appear in chlorinated water containing humic substances, are derived by the classic haloform reaction.

TABLE III-2

Chemical Degradation Products of Humic Substances.

The ultimate concern for public health protection is, of course, the fact that THM's are formed during the chlorination of drinking water sources. Consequently, a discussion of chemical mechanisms may appear to be rather academic. However, a precise understanding of the mechanisms by which the THM's are formed may prove to be truly beneficial by helping water utility personnel avoid the conditions during treatment that promote the appearance of high concentrations of these compounds in finished water. Studies with model compounds under well-defined laboratory conditions have been useful in elucidating these mechanisms and reaction conditions. Examples are given in the section pertaining to chlorine chemistry.

Chlorine

Chlorine has been the principal disinfectant of community water supplies for several decades. Until recently, its use had never been questioned seriously because the health benefits derived from it were so obvious. Although an occasional taste-and-odor problem in finished water was attributable to the reaction of chlorine with some substance in the raw water, the events were usually intermittent, short-lived, and presumably did not affect the public health. However, in 1974, Rook (1974) in the Netherlands and Bellar et al. (1974) in the United States reported that chlorine reacts with organic precursors that are found in many source waters to produce a potential carcinogen, chloroform (CHCl₃).

In December 1974, Congress passed the Safe Drinking Water Act (PL 93-523), and in early 1975, the U.S. Environmental Protection Agency (EPA) began an 80-city water supply survey—the National Organics Reconnaissance Survey (NORS)—to determine the extent of the problem (Symons et al., 1975). As part of NORS, finished waters from five cities (Miami, Florida; Seattle, Washington; Ottumwa, Iowa; Philadelphia, Pennsylvania; and Cincinnati, Ohio), which represented the major types of water sources in the United States, were analyzed as thoroughly as possible for all volatile organic compounds, i.e., those that can be stripped from solution by purging with an inert gas (Coleman et al., 1976). Seventy-two compounds were identified, 53% of them containing one or more halogens. A later study, the EPA National Organic Monitoring Survey (NOMS), included analyses of samples that had been taken from the water supplies of 113 cities (Brass et al., 1977) on four occasions over an 18-month period during 1976 and 1977. The source waters of a few cities were examined, but most of the effort was directed toward an analysis of finished waters for chloroform and 20 other volatile organic compounds. In addition to the 21 compounds that were originally selected, five others appeared frequently and were reported.

Since 1974, there have been numerous other surveys similar to NORS and NOMS, but they have been more restricted in scope. In addition, research activity has been intensified to isolate and identify the precursors, products, and mechanisms that are associated with the presence of potentially toxic organic compounds in both water and wastewater. In December 1976, the EPA published a list of 1,259 compounds that had been identified in a variety of waters (including industrial effluents) in Europe and in the United States (Shackelford and Keith, 1976). The agency is currently compiling a comprehensive register of all data concerning the identification of organic pollutants in water.

Properties of Aqueous Chlorine

Various aspects of chlorine chemistry have been reviewed by Jolley et al. (1978), Miller et al. (1978), Morris (1975, 1978), and Rosenblatt (1975). A synopsis of the basic principles will provide some understanding of the various forms that chlorine can assume in water and the reactions that it can undergo with certain types of compounds.

Reactive Forms of Chlorine in Water

"Aqueous chlorine" is a misleading term because the active form of chlorine that is present in treated water and wastewater is not the gaseous chlorine molecule (Cl₂) but, rather, a hydrolysis product, hypochlorous acid (HOCl), which is formed from the reaction between the chlorine molecule and water:

(1)

Hypochlorous acid, a weak acid, can ionize as follows:

(2)

The degree of ionization depends primarily on the pH and temperature of the water. The concentration of hypochlorous acid and the hypochlorite ion (OCl^-) are approximately equal at pH 7.5 and 25°C.

Another form of chlorine, the hypochloronium acidium ion (H₂OCl⁺), is known to exist (Miller et al., 1978; Rosenblatt, 1975), but its concentration would be extremely low in water at pH's between 5 and 9. Still another form of chlorine, the chloronium (or chlorinium) ion (Cl⁺), has been proposed as an important reactant in aqueous solutions of organic compounds (Carlson and Caple, 1978), although its existence is disputed (Rosenblatt, 1975). Nevertheless, Morris (1978) pointed out that "the reactant behavior of HOCl with organic carbon and amino nitrogen is as an electrophilic agent in which the chlorine atom takes on partially the characteristics of Cl⁺ and combines with an electron pair in the substrate." Finally, Carlson and Caple (1978) mentioned that another form of chlorine, the chlorine radical (Cl·), may react in the light to produce chlorine-substituted organic compounds when the parent chlorine molecule is not lost by any other significant reaction pathway. Rosenblatt (1975), citing others, described this form as "probably the most selective chlorinating species of all."

Free chlorine species (HOCl, OCl^-, Cl₂, H₂OCl⁺, Cl⁺) will oxidize both the bromide ion (Br^-) and iodide ion (I^-) to hypobromous and hypoiodous acids (HOBr and HOI). This reaction, as will be discussed later, is postulated to account for the presence of bromine- and iodine-substituted organic compounds, particularly the mixed-halide haloforms, in waters that had been disinfected by chlorination.

Reactions of Hypochlorous Acid with Organic Compounds

Chlorine reacts in solutions of organic compounds by one or more of three basic mechanisms (Jolley et al., 1978; Miller et al., 1978; Morris, 1975; Morris, 1978), namely, addition, during which chlorine atoms are added to a compound; oxidation; and substitution, during which chlorine atoms are substituted for some other atom that is present in the organic reactant. All three of these reactions involve hypochlorous acid as an electrophile.

Only addition and substitution reactions produce chlorinated organic compounds. Oxidation reactions account for most of the "chlorine demand" of natural waters and waste treatment effluents (Jolley et al., 1978; Morris, 1975), but the end products are not chlorinated organic compounds. That is not to say that those products cannot be harmful. Miller et al. (1978) have mentioned that epoxides can be produced from carbon-chlorinated compounds at pH values that are common in water treatment plants (e.g., pH 9.5-10.5) where softening is practiced. To illustrate, they describe the reaction between ethylene (C₂H₄) and hypochlorous acid, which yields ethylene chlorohydrin (ClCH₂CH₂OH) as an intermediate. This hydrolyzes to form the epoxide, ethylene oxide (C₂H₄O). Carlson and Caple (1978) mentioned one such reaction, in which a mixture of chlorohydrins resulted from the reaction of oleic acid [CH₃(CH₂)₇CH=CH(CH₂)₇COOH] with hypochlorous acid. Presumably, these would be converted to epoxides if the pH were to be increased. Carlson and Caple also showed how a ubiquitous natural compound, a-terpineol [CH₃C₆H₄C(CH₃)₂OH], could form epoxides when reacted with hypochlorous acid. These reactions illustrate how chlorination may result in the development of nonchlorinated products, e.g., the epoxides, which may pose health risks. In instances such as those just discussed, a chlorinated intermediate, which itself should be evaluated toxicologically, is involved.

Chlorine By-Products Found in Drinking Water and Selected Nonpotable Waters

The most frequently mentioned products of aqueous reactions between chlorine and selected types of organic compounds are discussed in this section. Special attention is given to the trihalomethanes (THM's) because of the current interest in them as potentially hazardous byproducts of chlorination in municipal water treatment facilities. The specific reactions by which THM's are produced in chlorinated natural waters are not well understood because the chemical structures of the precursor organic compounds, which are thought to be primarily humic substances, are highly varied and extremely complex. A summary of the relevant facts concerning these ubiquitous, natural organic substances is presented in the section on precursors. The term "haloform reaction" is often mentioned as the mechanism by which THM's are produced when natural waters are chlorinated. This has not been validated definitively in actual water treatment systems. However, the reaction will be discussed in conjunction with THM formation in natural waters because it is one possible mechanism that has been described thoroughly in the literature.

Other products of chlorination that are discussed below include the chlorinated phenols, which have been of concern primarily because they impart offensive tastes to drinking water, and compounds that have been isolated during carefully controlled experiments involving analyses of both chlorinated and unchlorinated samples of freshwater and its analogs. Data derived from in-plant studies are difficult to interpret because raw waters are seldom uniform in composition over time, especially surface waters and raw sewages. Thus, there is always some unavoidable uncertainty as to whether compounds that have been isolated from finished waters were produced by chlorination of precursor organic compounds in the influent or whether they were present before the influent entered the plant. Despite the uncertainty, some published in-plant studies are described in this section. Those related to THM formation are especially noteworthy. Abundant laboratory evidence shows that the high concentrations of THM's that have been observed in many finished drinking waters are produced by reactions between chlorine and precursor organic compounds that are commonly found in natural waters.

Trihalomethanes

The current interest in THM's was prompted by the reports of Rook (1974) and Bellar et al. (1974), which linked the presence of chloroform and other trihalogenated derivatives of methane (CH₄) in finished drinking water to the chlorination of raw waters during treatment. Rook (1974) demonstrated that THM's were produced by the chlorination of aqueous extracts of peat, and he postulated that the precursors were the humic substances. The presence of THM's in laboratory-chlorinated waters that had been taken from a lake in a ''peaty region" helped confirm Rook's hypothesis. Later laboratory studies involving the chlorination of aqueous solutions of humic substances (e.g., Babcock and Singer, 1977; Hoehn et al., 1978; Stevens et al., 1976) further confirmed Rook's hypothesis. Bellar et al. (1974) reported that THM concentrations in Ohio River water increased as it was being treated and chlorinated at several points in the Cincinnati, Ohio, waterworks. Rook (1974) reported similar observations for treated water from the Rhine and Meuse rivers at the Berenplaat treatment plant in the Netherlands.

The prevalence of THM's in U.S. municipal water supplies was demonstrated by the NORS survey in 1975 and the NOMS survey in 1976. According to Symons et al. (1975), the total THM (TTHM) concentrations in the 80-city survey (NORS) were log-normally distributed, while chloroform appeared in the highest concentrations (median, 21 µg/liter), followed by bromodichloromethane (CHCl₂Br) (median, 6 µg/liter) and lesser concentrations of chlorodibromomethane (CHBr₂CI) and tribromomethane (CHBr₃). The distributions are shown in Figure III-1. Brass et al. (1977) reported median concentrations of chloroform and bromodichloromethane of 27 and 10 µg/liter, respectively, in the 113 supplies that were surveyed during the first phase of NOMS. During that phase, sampling and handling procedures were identical to those used during NORS. The THM concentrations in the raw water sources, which included rivers, groundwaters, and lakes, were much lower. Hoehn et al. (1977) and Hoehn and Randall (1977) reported results of an extensive 2-yr THM monitoring program at a water treatment plant and at the distribution system of a northern Virginia water supply that was supplied by a reservoir. Mean finished-water THM concentrations varied seasonally and were from 1 to 2 orders of magnitude greater than those found in the source water. Average chloroform concentrations in finished drinking water were from 5 to 10 times greater than the median reported for NORS and NOMS, but concentrations of bromodichloromethane were comparable.

Figure III-1

Frequency distribution of trihalomethane data (NORS). From Stevens et al., 1978.

As was mentioned earlier, it is difficult to determine whether many organic compounds that are recovered from finished waters during in-plant studies actually were formed during the treatment process, especially when analyses of intermittently collected "grab" samples provide the data base. The uncertainty factor is greater when compounds are seldom recovered from finished water and when they are occasionally detected in untreated water at concentrations that are similar to those found in treated water. In these instances, the origin of a substance that has been isolated from finished water would be determined more accurately by a comparative analysis of samples from both influent and effluent water that has been collected and composited for a period at least equal to the hydraulic detention time within the plant. In the in-plant surveys, the attribution of THM's to chlorination is more certain, although grab samples were used, primarily because THM's have often been found in finished waters in much higher concentrations than in the untreated influent waters.

Mention has already been made of the fact that mixed-halide THM's have been identified on numerous occasions, the most common ones being the bromochloromethanes, CHCl₂Br and CHClBr₂. These, and occasionally iodine-containing, mixed-halide THM's, have been reported but in much lower concentrations than chloroform. The mechanisms that are involved in the formation of these compounds during chlorination are discussed in the bromine-iodine section of this chapter. Fluorinated isomers are not expected to be produced by these mechanisms (Kleopfer, 1976).

The production of chlorinated THM's in concentrations that are typical of those reportedly recovered from finished drinking waters is dependent upon the initial presence of free chlorine. Thus, unnitrified, chlorinated sewage effluents that contain no free chlorine usually contain low (< 10 µg/liter) THM concentrations (Stevens et al., 1978) unless industrial wastes containing THM's are discharged into the sewers. The reason for this phenomenon is that free chlorine species react extremely rapidly with ammonia (NH₃) (Morris, 1967), which is present in sewage in concentrations of parts per million. Under those conditions, the concentration of free chlorine is extremely low (Jolley et al., 1976).

Nonhaloform Products of THM Reactions

While it is not certain that the classic haloform reaction is solely responsible for the appearance of THM's in chlorinated drinking water, it illustrates the fact that organic compounds other than THM's are formed as a natural consequence of reactions involving carbanions that lead to THM formation. Some of these compounds are chlorinated intermediates, while others are terminal products that are formed simultaneously with the THM's. For example, if the classic haloform reaction (involving a ketone) is operative, a carboxylic acid is formed during the final hydrolysis step, which results in the production of a THM:

(3)

Nonhalogenated products, which may contain aromatic or alicyclic structures (e.g., benzene rings or cyclic hydrocarbons, respectively), should not be ignored when assessing potential health risks that are associated with water chlorination.

Another fact to consider is that if the free THM were not produced (e.g., by hydrolysis, in the haloform reaction), some chlorinated intermediate containing a chlorinated methyl group might be present. Suffet et al. (1976) have reported finding 1,1,1-trichloroacetone (Cl₃CCOCH₃) in two different potable water supplies near Philadelphia. One water treatment plant is on the Delaware River; the other is on the Schuylkill River. The 1976 EPA Survey (Shackelford and Keith, 1976) reported no instances of trichloroacetone in rivers, lakes, or groundwaters. Trichloroacetone is readily hydrolyzed at a pH above 5 (Suffet et al., 1978, private communication). Suffet et al. (1976) utilized two different on-line composite isolation methods with dechlorination and adjustment to pH 4 to stabilize the compound. They isolated it with a XAD-2 resin and continuous liquid-liquid extractor. This compound is a likely intermediate in the haloform reaction, especially when acetone (C₃H₆O) is present in the river water. The haloform reaction for acetone is extremely slow (Bell and Lidwell, 1940; Morris and Baum, 1978). Trichloroacetone hydrolysis is much quicker. To be observed, the compound must be extracted immediately or the pH must be lowered to ≤4 to stabilize it.

Morris and Baum (1978) reported that trichloroacetate (CCl₃COO^-) and hexachloroacetone (CCl₃COCCl₃) are possible intermediates of the classic haloform reaction, but they argued that substantial quantities of other chlorinated organic compounds, except chloroform, should not be formed. Obviously, the formation of chlorinated intermediates other than these could occur by mechanisms other than the haloform reaction.

Pfaender et al. (1978) presented corroborating evidence that trichlorinated intermediates do occur in chlorinated waters. They detected more chloroform in water samples that had been analyzed by direct aqueous injections (DAI) of samples into a gas chromatograph than by a conventional technique that involves stripping the volatile THM's from solution with an inert purge gas. They attributed the higher recoveries to thermal decomposition of nonvolatile, chlorinated intermediate compounds that had been introduced into the analytical instrument by DAI but which remained in solution during the analysis involving the purge technique. They did not identify any of the intermediates but did show that chloral (CCl₃CHO), a previously suspected intermediate, was stable under the conditions of the analysis by DAI and could not account for the observed increases in chloroform concentration.

Carbon tetrachloride (CCl₄) is not produced during the haloform reaction, nor should it be produced by any other mechanism involving electrophilic substitution of chlorine on a carbanion. Shackelford and Keith (1976) reported its presence in finished drinking water on numerous occasions when it had not been detected in the untreated water. The most likely source of carbon tetrachloride in these instances is the chlorine itself. Carbon tetrachloride and hexachloroethane (C₂Cl₆) have been identified as contaminants of the chlorine-manufacturing process (Laubusch, 1959). Hexachlorobenzene (C₆Cl₆) has also been identified as a contaminant (Blankenship, 1978).

Reactions Involving Phenolic Substances

The primary concern with phenolic compounds in drinking water supplies has been the offensive tastes that result from their reaction with hypochlorous acid. The presence of the hydroxyl group on the benzene ring activates the ring and permits chlorine to substitute readily. Multiple substitutions can lead to rupture of the parent ring (Burttschell et al., 1959). Morris (1975) has provided a good overview of the available literature of the subject, including that related to the products that have been observed after rupture of the ring. Lee and Morris (1962) and Lee (1967) have described in detail the kinetics of phenol chlorination and have recommended controls for the taste and odor problem that is caused by chlorophenols.

Phenolic substances that are present in natural waters are highly varied. Reactions produced in the laboratory with the simple phenols may not occur when the more complex phenols are present. Therefore, conclusions that are derived from these laboratory studies should be extended with caution to systems that contain the more complex and diverse forms of phenols that exist in nature.

Reaction Products from Chlorinated Surface Waters and Sewage Effluents

Because of the difficulties of interpreting in-plant studies involving analyses of influent and effluent grab samples, most of the investigations discussed below are limited to those involving analyses of pairs of influent water or sewage samples that differed only in that one of each pair was chlorinated in the laboratory. Despite the uncertainties that are associated with data from influent-effluent analyses, a few investigations involving in-plant studies have been reviewed, either because the compounds that were detected in the chlorinated effluents were seldom, if ever, detected in untreated waters or because the study involved a broad survey of many treatment facilities.

Products from Chlorinated Surface Waters

Jolley et al. (1978) chlorinated surface sources that were used to cool water at two Tennessee electric power-generating facilities, one in Kingston (Watts Bar Lake) and the other near Memphis (Mississippi River). Radioisotopic ³⁶Cl in HO³⁶Cl was applied in dosages of 2.1 mg/liter at Kingston and 3.4 mg/liter at Memphis and was allowed to react for 75 and 15 min, respectively. The waters were concentrated by vacuum distillation, and compounds were separated for analysis with a scintillation counter by high-pressure liquid chromatography (HPLC). Chlorination yields (as Cl) of chloroorganic constituents that were separated by HPLC were 0.5% at Kingston and 3.0% at Memphis. A variety of compounds was recovered. The influent water was not analyzed because the presence of the ³⁶Cl isotope in the molecules indicated that the product was derived from the chlorination. There was no attempt to confirm the identity of the reported compounds by gas chromatography/mass spectrometry analysis.

Among the chlorinated compounds that were recovered in concentrations ranging from a few tenths of a part per billion (ppb) to as high as 20 ppb were a nucleoside, three purines, a pyrimidine, seven aromatic acids, and five phenolics. The Mississippi River sample from Memphis contained the highest concentrations (0.7 to 20.0 ppb), because it contained less ammonia—hence more free chlorine to serve as the active oxidizing agent. The analytical methods recovered only nonvolatile, or slightly volatile, compounds and did not permit detection of large polymeric substances such as humic acids or nucleic acids. The pyrimidine, 5-chlorouracil, which was present at concentrations of 0.6 and 7.0 ppb in the two samples, has captured attention because it could be incorporated into cellular genetic material (Gehrs and Southworth, 1978; Jolley et al., 1978). Gehrs and Southworth (1978) commented that one of the chlorophenols, 4-chlororesorcinol [ClC₆H₃-1,3-(OH)₂], was potentially toxic, but they made no mention of the possible impacts of the other chloroderivatives that were recovered.

In a study of chlorinated (2.0 mg/liter, 68-hr contact time) and unchlorinated waters from Lake Zurich, Giger et al. (1976) recovered several a-chloroketones and a variety of THM's that contained chlorine, bromine, and iodine, all of which were absent in the unchlorinated control sample. The ketones included 2,2-dichlorobutanone (CH₃CCl₂COCH₃), 2,2-dichloropentan-3-one (CH₃CCl₂COC₂H₅), 1,1,l-trichloroacetone (also recovered by Suffet et al., 1976), and 3,3-dichlorohexan-4-one (C₂H₅CCl₂COC₂H₅). In addition, they detected minor quantities of chlorinated alkylated benzene compounds. The a-ketones and THM's were identified by mass spectrometry.

Brass et al. (1977) mentioned that two sets of raw and finished waters from an unidentified city were analyzed for volatile organics during NOMS Phase III. They observed that aromatic compounds (unspecified) were found in finished water but not in the raw water samples. Further work was anticipated. A later publication concerning NOMS-III (Munch et al., 1977) lists the relative recoveries of 25 various organic compounds from the raw and finished waters at 11 locations. Compounds other than THM's, which were recovered more often in finished waters than in raw, included benzene (C₆H₆) (4 versus 0), chlorobenzene (C₆H₅Cl) (3 versus 1), ethylbenzene (C₆H₅C₂H₅) (3 versus 0), dichloromethane (CH₂Cl₂) (3 versus 0), toluene (C₆H₅CH₃) (3 versus 1), and a xylene isomer [C₆H₄(CH₃)₂] (4 versus 1). (Compounds where the difference in recoveries was only 1 are not listed.) The ranges of concentrations that were observed when the compounds were quantified included: 0.58-6.10 µg/liter for chlorobenzene, 0.1-1.5 µg/liter for benzene, 0.42-0.57 µg/liter for ethylbenzene, and 0.48-19.00 µg/liter for toluene. The means were 2.7, 0.88, 0.50, and 6.9 µg/liter, respectively.

Products from Chlorinated Sewage Effluents

Jolley (1975) chlorinated primary- and secondary-treated sewage effluents with various forms of ³⁶Cl-tagged chlorine compounds to determine which chlorination products could be found. Chlorine was dispensed as a gas at dosages and contact times that were similar to those expected in actual plant operations. Additional experiments were performed in the laboratory, wherein a ³⁶Cl-tagged hypochlorite solution was used as the chlorinating agent. Of the 44 compounds that were recovered, 17 mentioned previously as having been recovered from cooling waters were identified in the sewage effluents. Less than 1% of the chlorine could be accounted for within organic compounds. Most of it was in the form of chloride ion (Cl^-), indicating that most of the chlorine that is applied to sewage serves as an oxidant. There are many difficulties inherent in the separation and identification of compounds in waters that are heavily laden with organic and inorganic contaminants. It may be that more chloroorganics were present than could be recovered. However, this possibility does not alter the basic conclusion that most of the chlorine is simply reduced to chloride ion during the oxidation of sewage constituents. Nevertheless, the oxidized organic compounds have importance from an environmental perspective whether or not they contain chlorine.

Garrison et al. (1976) identified extractable volatile organics in lime-clarified, tertiary-treated sewage from the Blue Plains pilot plant near Washington, D.C. Chlorination was effected by sodium hypochlorite (NaOCl), which was added to achieve breakpoint and a final, free residual of 3-5 mg/liter. The final pH was between 7.0 and 7.5. Samples were extracted and fractionated by standard analytical procedures, then concentrated and methylated before analysis. The most pronounced effect of the chlorination was evidenced by the appearance of chlorocyclohexane (C₆H₁₁Cl) and tetra-, penta-, and hexachloroethanes (C₂H₂Cl₄; C₂HCl₅; C₂Cl₆) in the effluent. None of the compounds was quantified.

These compounds are unusual constituents of chlorinated waters. The pentachloroethane is not an ordinary commercial material and, therefore, would probably not have an industrial origin. Commenting on his previous work (Garrison et al., 1976), Garrison (1978) provided some insight concerning the appearance of these compounds. Studies at the EPA Laboratory in Athens, Georgia, have shown that cyclohexene (C₆H₁₂) is a common contaminant of methylene chloride (CH₂Cl₂), which is used as the extracting solvent. If the aqueous system being extracted contained residual chlorine, the cyclohexene could be converted to 1,2-dichlorocyclohexane by the addition of chlorine to the double bond. During the original study, the solvent may have contained other alkene impurities that could have reacted with chlorine to form the tetra-, penta-, and hexachloroethanes, but that possibility was not explored. If those impurities had existed, they would have been masked during analysis by the methylene chloride peak on the gas chromatogram, since they are highly volatile substances. No "blanks" (control tests) were analyzed because the possible problem was not recognized. Moreover, both raw and treated sewages were being analyzed and the differences in composition were of prime interest. However, Garrison (1978) emphasized that these chlorinated aliphatic compounds, although unusual, have been identified in studies subsequent to the one reported in the 1976 publication.

Glaze et al. (1976, 1978) and Glaze and Henderson (1975) reported the appearance of 38 compounds that were generated by superchlorination (2,000 mg/liter) of sewage effluent that had received secondary treatment. Most of the chlorinated compounds that were identified were aromatic derivatives. Many of them involved aromatics with no activating substituent groups. Among those that appeared were the chloroderivatives of benzene, toluene, and benzyl alcohol (C₆H₅CH₂OH) and the nonaromatic derivatives such as chlorocyclohexane (reported also by Garrison et al., 1976), a chloroalkyl acetate, and three chlorinated acetone derivatives (tri-, penta-, and hexachloroacetone). The acetone derivatives may be precursors of chloroform (see also Suffet et al., 1976). According to the authors, they may have been produced by an acid-catalyzed reaction between acetone and chlorine.

Glaze and Henderson (1975) and Glaze et al. (1976, 1978) also reported that an increase in the concentration of chlorine-containing compounds, as estimated by the total organic chlorine (TOCl) concentration in nonpurgeable compounds, was observed in the sewage as the chlorine dose was increased from 25 to 2,000 mg/liter. The TOCl concentration increased from 80 to 906 µg/liter, but decreased to 164 µg/liter when the chlorine dose was increased to 3,000 mg/liter, thereby demonstrating the destruction of chloroorganic compounds that were formed at lower chlorine dosages. The chlorine that was bound in organic compounds accounted for less than 0.05% of the applied chlorine. The chlorine dosages far exceeded expected concentrations in routine sewage treatment plant operations. However, there are commercially available treatment units in which sewage and industrial waste sludges can be oxidized by superchlorination within the range of the high dosages used by Glaze et al. Superchlorination would probably be required to reduce the concentrations of chloroorganic compounds that increase when chlorine dosages are increased. Fuchs and Kuhn (1976) demonstrated a marked increase in concentrations of chlorinated organic compounds at a water treatment plant when Rhine River water was chlorinated to breakpoint. Activated carbon treatment reduced these concentrations, the removal ranging from 28% to 69% of the influent TOCl concentrations. Superchlorination was not attempted.

Sievers et al. (1978) found increased levels of aromatic hydrocarbons—such as toluene, xylenes, and styrene (C₆H₅CH = CH₂)—in chlorinated, secondary-treated wastewaters from the Metro-Denver sewage plant in District No. 1. Excess chlorine additions under laboratory conditions resulted in the formation of species of chlorotoluene (ClC₆H₄CH₃) and chloroxylene [ClC₈CH₃(CH₃)₂]. No industrial discharges enter the system. Therefore, the authors believe that the precursors were not of industrial origin. Concentrations of these aromatics ranged from 0.1 to 2.0 ppb.

In the EPA survey (Shackelford and Keith, 1976), toluene was found in finished drinking waters on several occasions, although it has been detected also in river water. Garrison (1978) reported that recent studies by the EPA laboratories in Las Vegas, Nevada, show toluene to be present at concentrations up to 2 µg/liter in pristine mountain streams in the Great Smoky Mountains National Park. The EPA investigators attributed this to automobile exhausts.

The EPA survey document also reported the frequent detection in finished drinking waters of the xylenes (ortho, meta, and para isomers) and, less frequently, styrene. These compounds were present in various raw water sources as well. Sievers et al. (1978) believed the precursors to be nonvolatile but did not postulate what they might be. According to them, toluene, styrene, and xylenes have been identified in 20 U.S. drinking water supplies and in certain Canadian supplies.

Chlorination of Model Compounds

Several recent publications have described the results of laboratory experiments involving chlorination of model compounds in aqueous solutions. The products have been reported, and, in some instances, mechanisms were proposed. A select number of those publications are reported below.

Reactions Resulting in THM Formation

Rook (1976) presented evidence that diketones, which have been described as degradation products of fulvic acids (one component of humic substances), do produce chloroform more rapidly than simple ketones. He chlorinated 1,3-cyclohexanedione [C₆H₈( = O)₂], 5,5-dimethylcyclo-1,3-hexanedione [(CH₃)₂C₆H₆(+O)₂] (dimedone), and 1,3-indandione [C₉H₅(=O)₂] at concentrations (> 200 mg/liter Cl₂) that were high in comparison to dosages used in water treatment plants at pH 7.5 and 11.0 at 10°C. He obtained yields of chloroform ranging from 50% to 100% of the theoretical yields (0.5-1.0 mol of chloroform per mole of precursor) in 4 hr.

In a later publication, Rook (1977) reported that the yields of chloroform from a variety of hydroxylated compounds were quite high. Those producing the highest yields with relative low chlorine dosages were metadihydroxy compounds such as 1,3-benzenediol (resorcinol) [C₆H₄(OH)₂], 1,3-dihydroxynaphthalene [C₁₀H₆(OH)₂], and 3,5-dihydroxybenzoic acid [HOOCC₆H₃(OH)₂]. Resorcinol is known to be a degradation product of fulvic acids (Christman and Ghassemi, 1966). The hydroxyl groups are strongly ''activating." Consequently, the carbon positioned between the two hydroxyl groups is activated from both sides, thereby becoming a strong carbanion, which is readily attacked by chlorine.

Rook (1977) pointed out that chloroform may be just one of many byproducts of reactions between chlorine and these types of compounds. He mentioned that substitution, oxidative ring fissions, or even ring contractions and other fragmentations may occur. Christman et al. (1978a) also chlorinated resorcinol under laboratory conditions and found that the carbon between the hydroxyl substituents was removed from the structure, thereby producing a 1,2-diketocyclopentenedione (3chloro-5,5-dichlorocyclopent-3-ene-1,2-dione). They also identified mono-, di-, and trichlorinated resorcinols that had been produced by simple substitution reactions. Additional work by Christman and coworkers is described below.

Morris and Baum (1978) chlorinated a variety of compounds that could produce reactive carbanions in solutions at pH 7 and 11. Several of the compounds contained the pyrrole ring, which exhibits active carbanion formation and is a component of many natural compounds such as the plant pigments chlorophyll and xanthophyll. The pyrrole ring is also present in tryptophane, proline (both amino acids), and indole (a component of certain protein putrefaction products). They also studied a class of compounds that are described as "acetogenins," which also contain β-ketonic groups in structures that are common to natural pigments.

Other compounds that Morris and Baum investigated as possible haloform precursors were vanillin (4-hydroxy-3-methoxybenzaldehyde and syringaldehyde (4-hydroxy-3,5,dimethoxylbenzaldehyde), both alkaline degradation products of woody materials. Neither compound has the 1,3-dihydroxy configuration that reacts rapidly to produce THM's. In some of the studies, the reactions with chlorine were permitted to proceed in near-neutral solutions for several hours. Then the pH was increased to between 9 and 11. While THM's were produced at the lower pH, greater concentrations were found when the pH was increased. Their major conclusion was that naturally occurring compounds, from which THM's can be produced, are capable of forming chlorinated intermediates as well at neutral pH as at high pH, and that only the final hydrolysis, which results in higher concentrations of THM's, is enhanced by elevated pH. They suggested that the hygienic quality of drinking water might best be monitored by measuring TOCl concentrations.

Christman et al. (1978b) have reacted several phenolic humic model compounds with hypochlorous acid in dilute aqueous solution. These models were selected on the basis of results from oxidative degradation of aquatic humic material. Their data (Table III-3) suggest that aromatic substitution patterns that are typical of lignin and soil humic acid (1,3,5- and 1,3,4,5-) produce less chloroform than the basic m-dihydroxy model (resorcinol). The decrease is particularly marked when hydroxy is absent or methylated. The yield data from the substituted cinnamic acid (3-phenyl-2-propenoic acid, C₉H₈O₂) indicate that the propyl side chain may be involved in the production of chloroform.

TABLE III-3

Comparative Chloroform Yields for Some Humic Model Compounds.

The yield of chloroform from resorcinol (Table III-3) compares favorably with the data of Rook (1977). In addition, the configuration of the cyclopentene intermediate suggests that Rook's hypothesis concerning the location of the chloroform-producing carbon is correct.

Reactions Producing Compounds Other than THM's

The previous section emphasized model compound studies that were designed primarily to elucidate the mechanisms of THM formation. This section contains a review of published data from investigators who were principally interested in chlorination products other than THM's.

Carlson et al. (1975) and Carlson and Caple (1978) addressed the problems of predicting the extent of chlorine incorporation into aromatic substances. They chlorinated aqueous solutions of a variety of compounds whose parent structure was the benzene ring but which contained a variety of substituent groups. Included in their studies were phenol (C₆H₅OH), anisole (C₆H₅OCH₃), acetanilide (CH₃CONHC₆H₅), toluene (C₆H₅CH₃), benzyl alcohol (C₆H₅CH₂OH), benzonitrile (C₆H₅CN), nitrobenzene (C₆H₅NO₂), chlorobenzene (C₆H₅Cl), methyl benzoate (C₆H₅COOCH₃), and benzene (C₆H₆) in concentrations of 9.5 ± 0.6 x 10^-4 M. Chlorine (7.0 × 10^-4 M) was added and allowed to react at 25°C for 20 min. Experiments were conducted at pH 3, 7, and 10. They also studied biphenyl (C₆H₅C₆H₅), the parent compound of polychlorinated biphenyls (PCB's).

As expected their results showed that phenol was the only aromatic compound with one substituent that reacted readily at pH's that are common to water treatment (e.g., pH 7 and 10). Others, like anisole and acetanilide, reacted to some degree, but only at pH 3. The substituent groups in anisole and acetanilide (—O—CH₃ and —NH—CO—CH₃, respectively) are not highly reactive with hydrochlorous acid at near-neutral pH's that are encountered in water treatment. Biphenyl reactions with chlorine under conditions that are normally found at water treatment plants were not significant. Either extremely long reacting times (days), high dosages of chlorine (> 100 mg/liter), or low pH's were required to obtain relatively small yields of the chlorinated biphenyls, but one or more of these conditions might prevail where superchlorination is practiced.

Chloramines

Combining ammonia (NH₃) with chlorine (Cl₂) to form chloramines for the treatment of drinking water has been called combined residual chlorination, chloramination, or the chloramine process. Objectives of this water treatment are to provide disinfecting residual that is more persistent than free chlorine in distribution systems and to reduce the unpleasant tastes and odors that are associated with the formation of chlorophenolic compounds (Symons et al., 1977). Thus, this process utilizes the formation of monochloramine (NH₂Cl) as indicated in the following reaction:

(4)

The production of monochloramine is optimized by a pH range of 7 to 8 and a chlorine-to-ammonia ratio of 5:1 (by weight) or less (Symons et al., 1977). White (1972) states that the preferred ratio is 3:1 (by weight).

At higher chlorine-to-ammonia ratios or at lower pH values, dichloramine (NHCl₂) and trichloramine (NCl₃), also called nitrogen trichloride, are formed according to the following reactions:

(5)

(6)

These and organic chloramines are produced during the chlorination of water containing ammonia or organic amines. Their presence may contribute to taste and odor problems in the finished water (Symons et al., 1977).

Ammonia may be added to the water before (preammoniation) or after (postammoniation) addition of chlorine. Preammoniation can prevent the formation of tastes and odors that are caused by reaction of chlorine with phenols and other substances. According to White (1972), postammoniation is the most often used ammonia-chlorine water treatment process.

The Inventory of Municipal Water Supplies (U.S. Public Health Service, 1963) indicated that only 308 of the 11,590 water treatment plants surveyed used an ammonia-chlorine process (Symons et al., 1977).

Properties of Chloramines

Hypochlorous acid (HOCl) reacts rapidly with ammonia to form monochloramine, dichloramine, or nitrogen trichloride as shown in Reactions 4 through 6, respectively. The reaction products are dependent upon the pH, the relative concentrations of hypochlorous acid and ammonia, the reaction time, and the temperature (Morris, 1978).

Usually monochloramine is the only chloramine that is observed when pH values are greater than 8 and when the molar ratio of hypochlorous acid to ammonia is 1:1 or less. At pH values less than 3, only nitrogen trichloride is ordinarily detected (Morris, 1978).

Drago (1957) described the general chemistry of monochloramine, and Theilacker and Wegner (1964) and Kovacic et al. (1970) reported its organic reactions in some detail. Most of these organic reactions occurred in nonaqueous media. Although extremely useful for organic syntheses, these experiments have only limited value for predicting effects from drinking water treatment.

Monochloramine has been the principal subject of several reviews. Its properties and chemistry have been described by Metcalf (1942), Colton and Jones (1955), Jander (1955a), Drago (1957), Theilacker and Wegner (1964), Czech et al. (1961), Gmelin (1969), and Kovacic et al. (1970). Monochloramine is a colorless, water-soluble liquid with a freezing point of -66C. It may decompose violently above that temperature (Kirk-Othmer, 1964, p. 914).

Relatively little is known about dichloramine. Chapin (1929) determined that its odor, volatility from aqueous solution, and relative solubility in various solvents are intermediate between those of monochloramine and nitrogen trichloride. He also found that dichloramine liberates iodine from acidified potassium iodide (KI) solution as do the other chloramines. Dichloramine solutions are reported to be unstable (Corbett et al., 1953; Palin, 1950). Recently, Gray and Margerum (1978) reported that aqueous dichloramine solutions are more stable than previously thought and, thus, may be more significant in the water treatment process.

Nitrogen trichloride is a bright yellow liquid with a strong irritating odor and lachrymatory fumes. Its melting point is below -40°C; its boiling point is 70°C. It is extremely explosive and, therefore, dangerous, except at very low concentrations. Its solubility in water is limited (Kirk-Othmer, 1964, p. 916). In aqueous solutions it decomposes slowly to ammonia and hypochlorous acid (Remick, 1942). Corbett et al. (1953) observed that aqueous solutions of nitrogen trichloride are stabilized by small amounts of acid. The compound is an effective chlorinating agent, particularly in nonaqueous media (Dowell and Bray, 1917; Houben and Weyl, 1962; Jander, 1955b; Kovacic et al., 1970).

Monochloramine is the principal chloramine that is encountered under the usual conditions of water treatment. The rate of its formation, shown in Reaction 4, is extremely rapid at the concentrations and conditions of water treatment. At the pH range of most water supplies, the reaction is usually 90% complete in approximately 1 min. As shown in Figure III-2, the reaction rate is maximum at pH 8.5 (Morris, 1978; Weil and Morris, 1949).

Figure III-2

Variation in rate of chloramine formation with pH. Calculated with log K Z 8.48. From Morris, 1978.

The specific rate of formation of dichloramine is much slower than that for monochloramine except at pH values less than 5.5 (Morris, 1978) (see Table III-4). Because dichloramine forms much more slowly than monochloramine at near-neutral pH values, dichloramine does not constitute a large percentage of the available chlorine unless the waters are quite acid or when the molar ratio of chlorine to ammonia is greater than 1. The relative proportion of dichloramine and monochloramine for equimolar chlorine and ammonia and for 25% excess ammonia from pH 4-9 are shown in Figure III-3 (Morris, 1978). Dichloramine is much less stable than monochloramine or nitrogen trichloride. The decomposition of dichloramine is simplified in Reaction 7. The actual reaction is more complicated: more chlorine is reduced and some nitrate is formed (Morris, 1978).

TABLE III-4

Specific Rates for Chloramine Formation at 20°C.

Figure III-3

Proportions of monochloramine (NH₂Cl) and dichloramine (NHCl₂) formed in water chlorination with equimolar concentrations of chlorine and ammonia (Morris, personal communication).

(7)

In acid solutions (pH 4 or less) or in solutions where chlorine concentrations far exceed those of ammonia, nitrogen trichloride is formed. Because nitrogen trichloride is formed from dichloramine, the reaction occurs only under conditions in which the dichloramine is reasonably stable. Thus, nitrogen trichloride is the only chloramine at pH values less than 3. At chlorine-to-ammonia molar ratios greater than 2, nitrogen trichloride occurs in diminishing proportions up to pH values of 7.5. Above pH 7.5, no nitrogen trichloride is found regardless of the ratio of chlorine to ammonia (Morris, 1978).

When chlorine is added to waters containing ammonia, the ''breakpoint" phenomenon becomes significant in the pH range of 6 to 9. At chlorine-to-ammonia m olar ratios of 0 up to 1, monochloramine is formed, thereby creating the "peak" shown in Figure II-3 (Chap. II). At values greater than 1, dichloramine is formed. Being unstable, it decomposes, usually as indicated in Reaction 7. Thus, with the addition of chlorine the apparent chlorine residual decreases from a chlorine-to-ammonia molar ratio of I up to approximately 1.65, at which the breakpoint occurs, i.e., after the ammonia has been converted principally to nitrogen (N₂) and some nitrate (see Figure II-3 in Chap. II). Chlorine that is added after the breakpoint exists as free chlorine, i.e., hypochlorous acid and the hypochlorite ion (OCl^-) (Morris, 1978; Wei, 1972; Wei and Morris, 1974).

Chloramine By-Products Found in Drinking Water and Selected Nonpotable Waters

Very few chemical studies of drinking water have been designed to identify the products resulting from the reaction of chloramines with organic or inorganic constituents of the water supplies.

Monochloramine is less effective as a chlorinating agent than hypochlorous acid by a factor of approximately 104 (Morris, 1967). Presumably, many of the reaction products of chlorination of water will be formed from chloramination because of the hydrolysis of chloramines to hypochlorous acid; however, the products should occur in lower concentrations because of the low equilibrium concentrations of the acid. According to Margerum and Gray (1978) the hydrolysis of monochloramine,

(8)

has a reaction half-time of 10 hr. The equilibrium constant is K = 6.7 × 10^-12.

Margerum and Gray (1978) indicate that the formation of hydroxylamine (NH₂OH),

(9)

at pH 8 has a reaction half-time of 350 yr. Therefore, it probably does not occur in water treatment.

Stevens et al. (1978) determined that trihalomethane (THM) formation was minimized when chloramines (mostly monochloramine) were used to treat raw water. Chlorine was added at 5.5 mg/liter to raw water and to raw water spiked with 20 mg/liter ammonium chloride (NH₄Cl) (ammonia nitrogen, 5.2 mg/liter) (see Figure III-4). Thus, during chlorination of water where the ammonia breakpoint is not achieved, THM production may not be great (Stevens et al., 1978).

Figure III-4

Comparison of THM formation with free and combined chlorine in Ohio River water, at pH 7.0 and 25°C. (Chlorine dosage 5.5 mg/liter: free chlorine values represent raw water, combined chlorine values represent raw water spiked with 20 mg/liter (more...)

Rickabaugh and Kinman (1978) determined that chloramination of Ohio River water with monochloramine at 10 mg/liter, pH 7 to 9, and 25°C resulted in 90.7% to 99.9% less production of THM as compared with THM production from chlorination with 10 mg/liter chlorine as hypochlorous and hypochlorite ion.

In the 1975 National Organics Reconnaissance Survey (NORS), 10 of the 80 water supplies sampled had been disinfected with chloramines. The concentration of THM's in the finished water of these utilities ranged from 1 to 81 µg/liter with an average of 19 µg/liter. The THM concentrations in treatment plants using breakpoint chlorination ranged from 1 to 472 µg/liter, averaging 72 µg/liter (Symons et al., 1977).

Chloramination of Model Compounds

N-Chloroorganic compounds, such as N-chloroglycine [H(Cl)NCH₂COOH] and chlorophenols, may be by-products of the chloramine water treatment process.

According to Margerum and Gray (1978), monochloramine is a chlorinating agent for N-compounds in aqueous solutions. For example, with 10^-4 M glycine the following reaction takes place:

(10)

The rate constant (K) for this reaction is 1.5 M ^-1s^-1 at pH 5 to 9, making it or similar reactions probable in aqueous systems. This report also corroborates the observation of Ellis and Soper (1954) that monochloramine undergoes chlorine-exchange reactions with primary and secondary aliphatic amines.

Although initially no chlorophenols are formed when low concentrations (milligram-per-liter range) of monochloramine and phenol are mixed, they appear after a reaction time of several days (Burttschell et al., 1959).

Inferences of Possible Chloramine Reaction By-Products

Reaction products from aqueous chloramine reactions are presented in Table III-5. The types of reaction are quite varied. They include addition, chlorine substitution, oxidation, amination, and free radical reactions. The majority of these reactions involve either highly acidic or highly alkaline solutions. Thus, deducing or inferring the presence of such reaction by-products in the chloramine water treatment process is a highly speculative activity. This is particularly true if the substrate is present in the water supply only in trace or very low concentrations. Much of the considerable reported research concerning chloramine reactions used organic solvents, such as ether. Because reactions in organic solvents have limited value in predicting by-products of the chloramine water treatment process, they were not included in Table III-5.

TABLE III-5

Aqueous Chloramine Reactions with Organic Compounds: Summary of Reaction Products.

Reactions of chloramides, such as chloramine-T (p-CH₃C₆H₄SO₂NCl^-Na⁺) have been studied extensively. Selected recent papers have been referenced in Table III-6. If present at all, such chloramides would occur in the water treatment process only at very low concentrations. However, similar reaction products would be anticipated from chloramine reactions. The reaction summaries in Table III-6 include several substrates, e.g., carbohydrates that might be present in water supplies. They indicate that reaction by-products such as aldonic acids could occur in the chloramine water treatment process.

TABLE III-6

Aqueous Chloramide Reactions with Organic Compounds: Summary of Reaction Products.

Bromine and Iodine

The use of chlorine for disinfection of public water supplies has proved so successful that there has been little impetus to test alternative halogens. In the late 1930's, bromine was used briefly at Irvington, California, but was abandoned because it was difficult to handle, it produced undesirable tastes, and a residual could not be maintained in the distribution system (White, 1972, p. 712). During the 1920's, iodine was added to the water supplies at Rochester, New York, and Sault Sainte Marie, Michigan. One reason for its use in these cases was goiter prevention, but this motivation disappeared with the advent of iodized table salt (Laubusch, 1971). A long-term study of physiological effects of iodinated drinking water was conducted in Florida prisons (Black et al., 1968; Freund et al., 1966).

Both bromine and iodine have been used in swimming pool disinfection. In addition, iodine that is administered in Globaline tablets, whose active ingredient is tetraglycine hydroperiodide, 2[(NH₂CH₂COOH)₄HI] 2½ I₂, has been adopted by the military to disinfect individual water supplies. Unfortunately, none of these applications has provided data on trace by-products. Therefore, an assessment of the potential of bromine and iodine to produce hazardous substances at the trace level must be based on less direct evidence.

Properties of Bromine and Iodine

Bromine and iodine may be introduced into water in several ways: in pure elemental form (Br₂ liquid or I₂ solid), as interhalogen compounds such as bromine chloride (BrCl) (Mills, 1975), or they may be released from various organic substrates (Goodenough et al., 1969; Morris et al., 1953). They may also be generated within the water by oxidizing bromide or iodide with an agent such as chlorine. Once the halogen has been introduced into water, it reacts rapidly. Consequently, its ultimate chemical form is controlled mainly by the composition and temperature of the water rather than by the nature of the halogen source.

Various aspects of bromine chemistry have been discussed recently by Johnson and Overby (1971), LaPointe et al. (1975), and Sugam (1977). In pure water containing no amino-nitrogen, the predominant form of bromine at near-neutral pH is hypobromous acid (HOBr). Above pH 8.6 at 25°C, this gives way to the hypobromite ion (OBr^-). At pH values below 6, hypobromous acid will convert to Br₂, Br₃^-, bromine chloride, or other halide complexes. The exact nature of the predominant species and the exact pH of the transition is controlled by the halide composition of the water (Sugam, 1977). However, at the pH and halide levels of drinking waters, halide complexes are not abundant. In the presence of ammonia or organic amines, bromamines are formed. Unlike the analogous chloramines, the inorganic bromamines are labile and interconvert rapidly so that the species distribution is controlled by chemical equilibrium (Johnson and Overby, 1971). The bromamines decompose readily into primarily nitrogen gas and bromide (LaPointe et al., 1975). At temperatures above approximately 70°C, alkaline hypobromite solutions decompose rapidly to yield the thermodynamically more stable anions bromate (BrO₃^-) and bromide (Br^-). At ordinary temperatures and at near-neutral pH, this reaction is very slow (Engel et al., 1954); however, Macalady et al. (1977) have shown that the rate is appreciable in seawater under bright sunlight.

Iodine chemistry in drinking water has been discussed in detail by Chang (1958). Iodine differs from chlorine and bromine in several notable ways. Iodoamines are unstable and do not form to any appreciable extent in aqueous solutions. Hypoiodite (OI^-), which becomes a predominant species only at very high pH, is not of practical importance because of the tendency of hypoiodous acid (HOI) to decompose rapidly to iodate (IO₃^-) plus iodide (I^-) above pH 9.

(11)

This decomposition is much faster than the analogous bromine reaction and results in a loss of most of the disinfecting capacity of the iodine. Unlike oxyanions of chlorine and bromine, iodate is thermodynamically stable in oxygenated water. Once formed, it is likely to be persistent. Also in contrast to both bromine and chlorine, the diatomic form of iodine (I₂) is fairly stable in aqueous solutions and can be the predominant species in neutral-to-acidic solutions. The exact pH at which hypoiodous acid predominance gives way to iodine depends upon the iodide concentration. If iodide is greater than approximately 10^-3 M, then iodine is replaced by I₃^- as the predominant species in neutral-to-acidic solutions.

Although the predominant chemical form of a halogen under specified conditions of dose, pH, ionic strength of the halide concentration, and temperature can usually be predicted precisely, it may not be the active chemical species in disinfection or in the formation of organic byproducts. The long-standing debate over the role of transient species, e.g., Cl⁺, Br⁺, H₂OBr⁺, H₂OI⁺, etc., in aromatic substitution reactions (Berliner, 1966; Gilow and Ridd, 1973; Swain and Crist, 1972) illustrates the difficulties, and, commonly, the ambiguities, that are associated with attempts to identify the active species.

With respect to their persistence in distribution systems, alkyl bromides and iodides are somewhat more labile than alkyl chlorides. The data in Table III-7 indicate that carbon-bromine and carbon-iodine bonds are somewhat weaker than carbon-chlorine bonds and are considerably weaker than the carbon-fluoride bond. The relative hydrolysis rates for methyl bromide (CH₃Br) and methyl iodide (CH₃I) are correspondingly greater than for methyl chloride (CH₃Cl) and methyl fluoride (CH₃F). Actual hydrolysis rates for trihalomethanes (THM's) and other disinfection by-products in distribution systems have not yet been determined. Although the bromo and iodo compounds hydrolyze relatively faster than the chloro compounds, absolute rates may be too slow for these reactions to have practical significance.

TABLE III-7

Comparison of Bond Energies and Relative Hydrolysis Rates for Methyl Halides.

Bromine and Iodine By-Products Found in Drinking Water or Selected Nonpotable Waters

Experience with Drinking Water Supplies

Since neither bromine nor iodine are used to disinfect major public water supplies, it is only possible to establish the tendency of these disinfectants to produce hazardous by-products in the laboratory. However, it is possible to gain some insight into the problem by examining data for chlorinated water supplies. In these, traces of bromine and iodine are commonly produced by oxidation of very small amounts of bromide and iodide in the raw water. This phenomenon was first reported by Rook (1974), who discovered chloroform (CHCl₃), bromodichloromethane (CHCl₂Br), dibromochloromethane (CHClBr₂), and bromoform (CHBr₃) in chlorinated drinking water in Rotterdam. He showed that the ratio of ΣBr to ΣCl in the THM's was in excess of the Br(I) to Cl(I) ratio in the water, even if quantitative oxidation of the trace bromide was assumed. This indicated that the THM formation kinetics favored bromine selectively.

Rook's observations have been confirmed in many water treatment systems throughout the world. Abundant documentation substantiates that THM's are produced mostly by the chlorination process and are quite rare in the raw water (e.g., see Bellar et al., 1974; Environmental Health Directorate, Canada, 1977; Henderson et al., 1976). In addition to the four bromo-chloro haloforms, dichloroiodomethane (CHCl₂I) is sometimes observed. In the National Organic Monitoring Survey (NOMS) of the Environmental Protection Agency, this compound was found in 85 of 111 supplies during Phase II of the survey and in 50 of 105 supplies during Phase III, but it was not quantified. Iodine-containing THM's are probably underrepresented in the data on water supplies because the isolation techniques that are used in analysis are usually not very suitable for these relatively nonvolatile compounds. Nevertheless, the total organically bound iodine that can be produced by chlorinating drinking waters will be severely limited by the iodide in the raw water. Typical iodide concentrations in river waters are below 10 µg/liter (Livingstone, 1963; Turekian, 1971).

Shackelford and Keith (1976) listed 20 organic compounds containing bromine and iodine and the number of reports of their presence in water supplies (Table III-8). These compounds were not necessarily produced by disinfection. Indeed, this would be unlikely in many cases. However, the list is very instructive because of the clear analytical bias that it displays. All of the compounds are nonpolar, and most have relatively low molecular weight and high volatility. Undoubtedly, many compounds are overlooked because of limitations in the analytical methods that are used most frequently. For example, in view of the ease with which they form, it is surprising that there have been no reports of bromophenols in drinking water. The NOMS study produced evidence that 2,4-dichlorophenol [C₆H₃(Cl)₂OH] is produced by chlorination of drinking waters. It seems likely that some bromophenols would also have been produced, but this remains to be established.

TABLE III-8

Organobromides and Iodides Reported to be Found in Finished Drinking Water.

These studies of chlorinated water supplies show that large amounts of THM's would probably result if bromine or iodine replaced chlorine as the principal disinfectant, since THM's containing bromine and iodine are produced from traces of oxidized bromine and iodine.

Laboratory Studies with Potable Waters

The conclusion in the previous paragraph is also supported by several laboratory studies of river waters. Bunn et al. (1975) treated Missouri River water both with chlorine, in the form of calcium hypochlorite [Ca(OCl)₂], and with mixtures of chlorine and halide salts. This is analogous to the treatment of some swimming pools to which bromide or iodide salts are added and then oxidized with chlorine. The results of Bunn et al. reveal that the addition of potassium fluoride (KF) or potassium chloride (KCl) has little effect on the composition of the THM products (Table III-9). But in the presence of potassium bromide (KBr) or potassium iodide (KI), appreciable amounts of these halogens are incorporated into the THM products.

TABLE III-9

Effect of Added Potassium Halide on THM Formation in Missouri River Water.

Similar experiments were performed by Rickabaugh (1977) and Rickabaugh and Kinman (1978). They treated Ohio River waters with hypochlorous acid (HOCl), hypochlorous acid plus iodide, monochloramine (NH₂Cl), monochloramine plus iodide, and iodine at pH's between 7.5 and 9.0. In addition to the usual four chlorine- and bromine- containing THM's, they found dichloroiodomethane and bromochloroiodomethane (CHBrClI) when iodine was present. On a molar basis, the total yield of THM's increased with pH. Ten mg/liter of iodine produced 5 to 10 times less total THM than 10 mg/liter of hypochlorous acid. Treatment with monochloramine or monochloramine plus iodide produced even smaller amounts. It would be unwise to generalize too much from these data, since the doses were not comparable on a molar basis. Moreover, little is known about the relative halogen demands under the different treatment conditions. Nevertheless, they suggest that iodine may be a less potent generator of THM's than chlorine.

Experience with Nonpotable Waters

Kuehl et al. (1978) studied brominated organic compounds in the tissues of fathead minnows that had been exposed to sewage effluent treated with bromine chloride. The compounds that they found and their estimated concentrations are given in Table III-10. These compounds were identified by mass spectrometry. None of the compounds were found in control fish, which were exposed to effluent that had not been treated with bromine chloride. An obvious interpretation is that these compounds were synthesized by reaction of the bromine chloride with organic compounds in the effluent. The authors caution that they cannot rule out the possibility that the enhanced bromide levels in the exposed fish caused biosynthesis of brominated compounds. Unfortunately, the effluent water was not analyzed.

TABLE III-10

Brominated Compounds in Fish Exposed to Secondary Effluent Treated with Bromine Chloride at Grandville, Michigan.

Bean et al. (1978) examined the compounds that were recovered from XAD-2 resin through which chlorinated seawater had been passed. The bromide content of untreated seawater is about 65 mg/liter, well in excess of chlorine doses that are typically used for disinfection. Thus, in chlorinated seawater, bromide oxidizes rapidly, forming mainly brominated organic by-products (Helz and Hsu, 1978). Using gas chromatography and mass spectrometry, Bean et al. found several halogenated compounds including bromoform, bromoacetal [BrCH₂CH(OC₂H₅)₂], and various brominated aromatic compounds. Since the seawater was not dechlorinated prior to its passage through the resin, it is not clear which compounds were formed in the seawater and which were formed in the resin by reaction with bromine in the seawater. The authors concluded that, except for bromoform, the concentration of nonpolar halogenated compounds that were generated by chlorination of relatively pristine seawater appears to be very low, i.e., in the nanogram per liter range.

Chlorine Dioxide

Chlorine dioxide (ClO₂) is an orange-yellow gas with a liquefaction temperature of 9.7°C at atmospheric pressure. It is explosive in either its gaseous or pure liquid state (Robson, 1964). Accordingly, it is usually prepared in solution at the place and time of use.

Properties of Aqueous Chlorine Dioxide

According to Henry's law, the coefficient for aqueous solubility of chlorine dioxide is 1.0 mol/liter/atm at 25°C and 3.1 mol at 0°C (Gordon et al., 1972). Dilute gas mixtures with less than 0.1 atm partial pressure of chlorine dioxide and dilute aqueous solutions are not explosive; they may be handled with ordinary precautions. However, even dilute aqueous solutions are somewhat unstable, particularly in light, tending to decompose through reactions that will be described later. Fresh solutions must be prepared daily and must be protected from light to be suitable as reproducible stock or test solutions.

Two principal methods are used for the preparation of aqueous solutions of chlorine dioxide for water treatment (Gall, 1978; Masschelein, 1967, 1969). The first, which is used in the United States almost exclusively, is the reaction of sodium chlorite (NaClO₂) and chlorine (Cl₂) in acidic aqueous solution, primarily by the following equation:

(12)

Reaction 12 is not entirely stoichiometric. Side reactions producing chlorate (ClO₃^-) consume some extra chlorine, so that an excess over the proportion that is indicated by the equation is required to utilize the chlorite (ClO₂^-) fully. For stoichiometric reaction, according to Reaction 12, 0.39 part of chlorine as Cl₂ is required for each part of sodium chlorite. In practice, 0.5 to 1.0 part of chlorine per part of technical sodium chlorite (80%) is commonly used (Dowling, 1974; Gall, 1978; Symons et al., 1977). The solution resulting after completion of Reaction 12 then contains excess aqueous chlorine and some chlorate along with the chlorine dioxide.

According to Masschelein (1969), optimal yields of chlorine dioxide (90% to 95% of the theoretical yield based on sodium chlorite) are obtained by mixing a strong sodium chlorite solution (300 g/liter) with obtained by mixing a strong sodium chlorite solution (300 g/liter) with aqueous chlorine solution containing 2 to 3 g/liter of Cl₂. The mixed solution is then usually passed through a reactor that is packed with inert turbulence-promoting material to give a contact time of 1 to several minutes for reaction. The resulting solution, containing 3 to 5 g/liter of chlorine dioxide, may be dosed directly into the water being treated or may be diluted to about 1 g/liter for short-term intermediate storage before dosing. The other method of preparing chlorine dioxide for use in water treatment is acidification of strong sodium chlorite solution, usually with hydrochloric acid (HCl). The major reaction is:

(13)

Side reactions leading to the formation of chlorate and even elemental chlorine also occur in this system, as in the other (Gordon et al., 1972; Taylor, et al., 1940). The side reactions are minimized when hydrochloric acid is used at a 1 to 2 molar concentration and mixed in excess with a strong sodium chlorite solution (Beuerman, 1965; Kieffer and Gordon, 1968; Masschelein, 1969; Toussaint, 1972).

According to Reaction 13, 0.322 part of hydrochloric acid is required per part of sodium chlorite for stoichiometric reaction. In practice, approximately equal parts, representing a threefold excess of hydrochloric acid, have been used. After a few minutes of contact to allow reaction to proceed to completion, the concentrated mixture may be dosed directly or given intermediate storage, as in the previous method.

The latter method, sometimes with an acid other than hydrochloric acid, has often been preferred for laboratory preparation of chlorine dioxide solutions as a means of avoiding the presence of excess aqueous chlorine (Dowling, 1974; Granstrom and Lee, 1957, 1958). Since chlorine or hypochlorous acid (HOCl) may be a product of decomposition of chlorous acid (HClO₂) under some conditions, however, the method does not assure freedom from contamination by aqueous chlorine unless special precautionary techniques are followed (Feuss, 1964; Miltner, 1977). Failure to observe such precautions casts doubts on a number of reported findings of chlorinated products from the reaction of aqueous chlorine dioxide with organic matter.

For industrial use, especially in the pulp and paper industry, where much greater amounts are used than in water treatment, chlorine dioxide is prepared by reduction of sodium chlorate (NaClO₃) with agents such as sulfur dioxide (SO₂), methanol (CH₃OH), or chloride (Cl^-), generally in several molar sulfuric acid (H₂SO₄) solutions (Gall, 1978; Sussman and Rauh, 1978). These processes are more economical than those based on sodium chlorite for the manufacture of large quantities of chlorine dioxide. However, they are more complex in operation, the formed chlorine dioxide must be swept from the reaction mixture with a gas stream and redissolved in water, contamination with chlorine usually occurs, and disposal of by-product wastes can be a problem. Nonetheless, adaptation of some of these processes to conditions of water treatment is being attempted (Gall, 1978), and practical methods may develop if demand for chlorine dioxide disinfection increases sufficiently.

The principal side reactions occurring during the preparation of chlorine dioxide solutions are:

(14)

and

(15)

These lead to the formation of chlorate as a significant by-product Accordingly, the health effects of chlorate must be considered in connection with the treatment of water with chlorine dioxide. Gordon et al. (1972) have discussed these and other reactions that may occur in solutions containing chlorine dioxide, chlorite, and aqueous chlorine They pointed out that many of the observed reactions appear to proceed through a transitory, common intermediate, Cl₂O₂, with a suggested structure:

This intermediate may undergo several reactions, including the following:

(16)

This reaction is invoked to account for the appearance of aqueous chlorine in the acid solutions of chlorine dioxide that are used to bleach pulp (Lindgren and Ericsson, 1969; Lindgren and Nilsson, 1972). It may also account for the reported presence of hypochlorous acid or chlorine in reaction mixtures that are used for the preparation of supposedly ''chlorine-free" chlorine dioxide (Kieffer and Gordon, 1968; Rosenblatt, 1975, 1978). On the other hand, many of these reports may simply reflect the inadequacy of most analytical methods for discriminating accurately between chlorine dioxide and free aqueous chlorine. Gordon et al. (1972) concluded that the reactions of chlorite plus acid do give chlorine dioxide solutions that are free of chlorine. Chlorine dioxide is reasonably stable in neutral or mildly acid aqueous solutions at concentrations of several milligrams per liter or less, provided the solutions are shielded from light (Bowen and Cheung, 1932; Gordon et al., 1972). However, in alkaline solutions, it disproportionates according to the reaction:

(17)

whose rate has been measured extensively (Gordon, 1964). Another decomposition reaction:

(18)

occurs with measurable speed only in quite acid solution (Bray, 1906). As noted earlier, photochemical breakdown takes place readily. The initial reaction appears to be:

(19)

but the complete mechanism of the photochemical reaction in solution is unknown. Hydrogen peroxide (H₂O₂), chlorite, chlorate, Cl₂O₃, oxygen (02), and chlorine have all been reported as intermediates or products (Gordon et al., 1972). Presumably, the chlorine dioxide will not persist in open basins or reservoirs, although it can remain for days in clean distribution systems. Residual persistence is discussed more fully in Chapter II.

Chlorine dioxide is a strong oxidizing agent. The normal pathway of action in mildly acid, neutral, and alkaline solutions is:

(20)

for which the standard potential is E° = 0.954 V at 25°C. Reduction to chloride also occurs, but only in acid solutions, so that the reaction is of little interest for the usual conditions of water treatment.

Inorganic Reactants.

Iodide (I^-), arsenite (AsO₃^-), and bisulfite (HSO₃^-) are some of the inorganic reductants that are oxidized by chlorine dioxide between pH 5 and 9. The chlorine dioxide is reduced to chlorite (Reaction 19). Manganous ions are also oxidized readily to manganic hydroxides, a reaction that is often used to precipitate manganese as hydroxide in water treatment (Sussman and Rauh, 1978; Symons et al., 1977). Oxidation of ammonia (NH₃) in aqueous solution is not known to occur, but may not have been looked for adequately.

Chlorine Dioxide Reactions with Model Organic Compounds

There have been extensive investigations of reactions of chlorine dioxide with organic compounds, particularly phenols and amines. Results of these studies have been reviewed in detail in several articles and monographs (Gordon et al., 1972; Masschelein, 1969; Miller et al., 1978; Stevens et al., 1978). Unfortunately, most of the studies were conducted in acidic solutions in which concentrations of chlorine dioxide and organic substrate were much greater than those encountered during water treatment. Consequently, the extent to which the findings are applicable to the pH values and concentrations of organic compounds that are found in drinking water sources is not clear. In addition, most reviewers did not specify the ranges of pH and concentrations to which their statements apply, thereby making it difficult to distinguish the conclusions that pertain to highly dilute, nearly neutral aqueous solutions.

Since chlorine dioxide reduces to chlorite when it acts as an oxidant in the nearly neutral pH range, it functions as a one-electron acceptor. The immediate oxidation product must have an odd number of electrons to conform to the general definition of a free radical. Except for the studies of Lindgren with lignin simulators (Lindgren and Ericsson, 1969; Lindgren and Nilsson, 1972) and of Rosenblatt on reactions with tertiary amines (Gordon et al., 1972; Rosenblatt, 1978), this aspect of the reaction of chlorine dioxide with organic materials has seldom been considered in defining reaction pathways and products. Yet, the reactivity of chlorine dioxide with organic material often differs greatly from that of other oxidants because of this one-electron transfer. Despite this, some investigators have stated that products resulting from treatment of organic material in natural waters with chlorine dioxide are generally similar to those found after ozonization (Buydens, 1970; Miller et al., 1978).

In a broad sense, chlorine dioxide is most reactive with tertiary amines and phenols (Masschelein, 1969; Rosenblatt, 1978); it is moderately reactive with olefins (Gordon et al., 1972); and its reactivity with alcohols and aldehydes, yielding the corresponding carboxylic acids, appears greater than that of ozone and aqueous chlorine (Masschelein, 1969; White et al., 1942). It is less reactive with secondary amines than with tertiary ones, and it is very unreactive toward primary amines (Gordon et al., 1972). As noted earlier, its reaction with ammonia is unknown (Miller et al., 1978). Chlorine dioxide is unreactive toward saturated aliphatic hydrocarbons and aliphatic side chains, but the latter may be split from aromatic rings or other functional groups. Saturated carboxylic acids, carbonyl compounds, and amino acids are usually inert with chlorine dioxide (Kennaugh, 1957; Sarkar, 1935).

Chlorine dioxide has sometimes been called a "pure" oxidant, which refers to its failure to form chlorinated derivatives. This is certainly incorrect, for numerous chlorinated aromatic compounds (Masschelein, 1969; Miller et al., 1978; Paluch et al., 1965) and some chlorinated aliphatic substances (Lindgren et al., 1965) have been obtained from aqueous oxidations by chlorine dioxide. However, no THM's have been detected as reaction products of chlorine dioxide with organic materials, although investigators have looked for them carefully (Miller et al., 1978; Vilagenes et al., 1977).

One complication is encountered when considering the reactions of chlorine dioxide with organic materials: the normal reaction product of chlorine dioxide, chlorite, may also be a reactive agent in some circumstances. A nucleophilic reagent (Rosenblatt, 1975), chlorite has been proposed as a possible source of some of the chlorinated products that are found after reaction of aromatic compounds with chlorine dioxide (Lindgren and Nilsson, 1972). Chlorite also reacts with aldehydes in neutral or mildly acidic aqueous solution yielding chlorine dioxide in accord with the following overall equation (White et al., 1942):

(21)

This reaction has been used, for example, in the colorimetric estimation of sugars (Jeanes and Isbell, 1941; Spinks and Porter, 1934). According to Lindgren and co-workers (Lindgren and Ericsson, 1969; Lindgren and Nilsson, 1972), Reaction 21 may involve formation of hypochlorous acid as an initial product, which then reacts on other chlorite to regenerate chlorine dioxide in a chain reaction or reacts to chlorinate or oxidize other organic substrates. Formation of hypochlorous acid as a product in the oxidation of phenols by chlorine dioxide has also been postulated by Lindgren (Lindgren and Ericsson, 1969). His ideas are based on studies at pH 1 to 5 and millimolar or greater concentrations. The extent of these types of reactions for conditions of water treatment is not known.

Reactions of chlorine dioxide with phenol and substituted phenols have been studied by many investigators. An extensive list of references is given by Gordon et al. (1972). Probably the most detailed investigations that are applicable to conditions of water treatment have been those of Glabisz (1968) and Paluch (1964). See Gordon et al. (1972) for other citations. The major products of the reactions are quinones and chlorinated quinones, in which the initial ring structure is retained, and carboxylic acids such as oxalic (HOOCCOOH), fumaric (trans-HOOKAH = CHCOOH), and maleic acids (cis-HOOCCH = CHCOOH), which result from rupture of the aromatic ring.

Phenol has been found to give 1,4-benzoquinone, 2-chloro-1,4-benzoquinone, 2,5-dichloro-1,4-benzoquinone, and the previously listed products of ring rupture (Masschelein, 1969). The chlorine-substituted products have been found to increase with decreasing ratio of chlorine dioxide to phenol, while ring rupture increases with increasing excess of chlorine dioxide over phenol.

Substituted phenols are also oxidized to p-quinones. Para substituents are often stripped away in the process, as is the case with p-nitrophenol, p-chlorophenol, and p-hydroxybenzaldehyde (Miller et al., 1978; Symons et al., 1977). According to Glabisz (1968), p-alkylphenols and o- or m-dihydric phenols are more subject to ring cleavage than are other types of phenols.

There is some evidence that formation of chlorophenols precedes oxidation to chlorinated quinones (Glabisz, 1967). In his work with a limited concentration of chlorine dioxide that was mixed with 0.67 millimolar phenol, Spanggord found 2-chlorophenol, 4-chlorophenol, 2,6-dichlorophenol, 2,4-dichlorophenol, 2,4,6-trichlorophenol, 2-chloro-2,4-dihydroxybenzene, resorcinol [C₆H₄^-1,3^-(OH)₂], and fumaric acid as products (Miller et al., 1978). Also, at the Water Supply Research Laboratory of the Environmental Protection Agency in Cincinnati, chlorophenols were observed as reaction products at a chlorine dioxide to phenol molar ratio of 0.8, but not at 3:1 or 14:1 (Miller et al., 1978). Interestingly, hydroquinone [C₆H₄^-1,4^-(OH)₂] was observed at all molar ratios. Some formation of polyquinones was also indicated.

Some other types of aromatic organic compounds are much less reactive than the phenols. For example, it has been reported that benzoic acid (C₆H₅COOH), benzenesulfonic acid (C₆H₅SO₃H) (Schmidt and Braunsdorf, 1922), and cinnamic acid (C₆H₅CH = CHCOOH) (Sarkar, 1935) do not react with chlorine dioxide. Benzylic acid [(C₆H₅)₂C(OH)COOH] reacts only slowly (Paluch et al., 1965). Dinitrophenols and trinitrophenols have also been found to be inert (Paluch, 1964). According to Gordon et al. (1972), anilines should react in much the same manner as do phenols. Reactions with amines have been studied extensively, particularly by Rosenblatt and co-workers (Gordon et al., 1972; Rosenblatt, 1978). Tertiary aliphatic amines are very reactive, undergoing a free-radical, oxidative dealkylation yielding aldehydes and secondary amines as products. The typical reaction with triethylamine [(C₂H₅)₃N] has the overall equation:

(22)

Secondary amines react through a similar pathway, but more slowly. Most primary amines react only slightly (Rosenblatt, 1978). Benzylamines undergo a hydrogen-abstraction reaction, and β-substituted amines may exhibit oxidative fragmentation as a result of α,β-scission.

Most amino acids are not reactive under conditions that prevail during water treatment (Kennaugh, 1957). In mildly acid solutions, tryptophan (2-amino-3-indoleproprionic acid) is oxidized to indoxyl, isatin, and indigo red, the acetic acid side chain being concurrently removed (Fujii and Ukita, 1957; Schmidt and Braunsdorf, 1922). Tyrosine [p-HOC₆H₄CH₂CH(NH₂)COOH] gives dopaquinone and dopachrome at pH 4.5 (Hodgden and Ingols, 1954). Histidine [2-amino-4(5)-imidazole proprionic acid] is also somewhat reactive (Schmidt and Braunsdorf, 1922).

Additional reactions of nitrogenous materials are described in considerable detail by Gordon et al. (1972). Most of the compounds that are covered are not known as common pollutants in water supplies. Consequently, their findings are not clearly relevant to water treatment with chlorine dioxide.

Studies with other materials have been sparse and have been confined mostly to acid conditions, pH 1 to 4, and to the concentrations between 10^-3 and 10^-1 M, which are characteristic of pulp and paper bleaching. However, some of these studies have shown that a number of organic structures are quite unreactive with chlorine dioxide even under these conditions. Observed reactions are often attributed to hypochlorous acid or Cl₂ that are formed in other processes.

In contrast to ozone (O₃), chlorine dioxide is not highly reactive toward the olefinic double bond. The oleic chain of methyl oleate [CH₃(CH₂)₇CH = CH(CH₂)₇COOCH₃] (Leopold and Mutton, 1959; Lindgren and Svahn, 1966) or triolein (glyceryl trioleate) (Leopold and Mutton, 1959) reacts, even in acidic solution over several days, primarily by oxidation at a position α or β to the double bond or by scission at the α-β link. Cyclohexene (1,2,3,4-tetrahydrobenzene, C₆H₁₀) does undergo ring opening at the double bond, but also gives cyclohex-l-ene-3-one and 3-chlorocyclohexene as major initial products (Lindgren et al., 1965). With excess ClO₂, ring splitting occurs when dicarboxylic acids form. Moreover, Schmidt and Braunsdorf (1922) found fumaric, maleic, and crotonic (CH₃CH = CHCOOH) acids to be unreactive with chlorine dioxide.

Similarly, although primary alcohols may be oxidized to carboxylic acids or ketones in relatively concentrated acidic solutions, much less reactivity is apparent in dilute, nearly neutral aqueous media. For example, glucose at pH 2 oxidizes at the —CH₂OH group, without opening of the furananose or pyranose ring (Flis et al., 1955), but Somsen (1960) found no reaction with ethanol or 2,3-butandiol [CH₃CH(OH)CH(OH)CH₃] at pH 7, even at 80°C.

Aldehyde groups are more reactive. For example, Somsen (1960) found that acetaldehyde (CH₃CHO), butyraldehyde (CH₃CH_2-CH₂CHO), and acetoin [CH₃CH(OH)COCH₃] are oxidized to the corresponding carboxylic acids at pH 7. Otto and Paluch (1965) reported that an aqueous dispersion of benzaldehyde (C₆H₅CHO) reacts violently with chlorine dioxide. On the other hand, Stevens et al. (1978) reported that several C₂ to C₈ aliphatic aldehydes were products of the treatment of Ohio River water with chlorine dioxide.

The situation is complicated further by the reaction of chlorite ion with aldehydes to produce chlorine dioxide with concurrent oxidation of the aldehyde to carboxylic acid or carbon dioxide (Lindgren and Nilsson, 1972; White et al., 1942). Reaction takes place in neutral or slightly acidic solution:

(23)

Formaldehyde (HCHO), in particular, reacts very readily, but acetaldehyde, benzaldehyde, and reducing sugars are also known to react.

Hydrocarbons and simple carboxylic acids are generally inert towards dilute aqueous chlorine dioxide (Stevens et al., 1978). However, Reichert (1968a,b) has shown that the polynuclear aromatic hydrocarbon, 3,4-benzpyrene (C₂₀H₁₂), reacts to give quinones and chlorinated derivatives. Reactions of polycyclic aromatic hydrocarbons have been described by Thielmann (1972).

Humates and fulvates react with chlorine dioxide (Buydens, 1970; Fuchs and Leopold, 1927), as its use for the bleaching of colored waters attests, but the nature of the reactions is not known. Phenolic materials are reported to be released from the humic matrix under certain conditions (Buydens, 1970). Some investigators have conducted studies with model building units for lignin and humic acid; however, they have used acidic, high-concentration solutions. For example, vanillin [4(HO)C₆H₃—3—(OCH₃)CHO] at pH 4 or less suffers ring rupture to give β-formylmuconic acid methyl ester (Sarkanen et al., 1962), while vanillyl alcohol [4—(HO)C₆H₃-3—(OCH₃)OH] at a final pH of 1 forms chlorinated benzoquinones (Dence et al., 1962). Under similar conditions, veratryl alcohol [3,4-(OCH₃)₂C₆H₃0H] produces 4,5-dichloroveratrole [4,5-Cl₂C₆H₄(OH₃)₂] (Dence and Sarkanen, 1960). However, the pertinence of these findings to water treatment conditions is questionable, in view of the pH at which these reactions were observed.

Ozone

Although ozone (O₃) has been used as a disinfectant of water for over eight decades (Lawrence and Cappelli, 1977), relatively little is known about its chemistry as an oxidizing agent of organic and inorganic solutes in aqueous solution. On the other hand, ozone reactions in nonaqueous systems have been studied thoroughly, mainly because of the widespread use of ozone as a degradative tool in the elucidation of organic structure (Murray, 1968). While it is probably true that much of this knowledge may be transferred to aqueous systems, particularly at low pH values (Hoigne and Bader, 1975), few definitive mechanistic studies have been conducted to confirm this view. Thus, a discussion of the chemistry of ozone by-products that are formed during water treatment is severely limited by a lack of specifically designed studies and by the poor definition of the organic materials involved.

This unfortunate situation probably will not prevail very long. Several studies in progress have been stimulated by the prospect that ozone will be used increasingly as a water treatment chemical and by the need for more information on by-products that may be formed from ozone oxidations of water contaminants. Some of these studies have been reported in preliminary form at scientific meetings (Rice and Cotruvo, 1978). Combined with the earlier literature, they allow one to see an emerging picture of ozone by-product chemistry. Clearly, much is left to be accomplished in this area.

This part of the chapter is divided into four sections. In the first, some of the physical and chemical properties of ozone as they pertain to its use as a water disinfectant are examined. The second lists those few cases where reaction by-products from actual in-plant use of ozone have been isolated and identified. The third section examines some studies that have been conducted with model compounds. The final section discusses reactions that are thought to be possible on the basis of the known chemistry of ozone and the various substrates that occur in municipal drinking waters.

Ozonization of Model Organic and Inorganic Compounds

Rice (1977) has reviewed several studies on the reaction products from the ozonization of various organic and inorganic compounds in aqueous solution. A recent monograph by Rice and Cotruvo (1978) contains papers from a 1976 conference on the products arising from organic materials treated with ozone or chlorine dioxide (ClO₂). Together, these two references contain the vast majority of papers on this subject. No critical examinations of the reliability of the data contained in the papers have been attempted in this report.

Tables III-11 through III-18 list products of ozonization of various classes of organic compounds. In most of the studies listed, if ozone was measured at all, it was simply reported that ozone was bubbled, at a measured gas phase level, into water plus substrate. Exhausting ozone gas or decomposition by virtue of contact with the bubbler first, etc., were generally not measured. Similarly, pH was seldom measured, despite its importance in such studies.

TABLE III-11

Summary of Ozonization Byproducts of Organic Compounds—Aromatic Compounds: Benzene and Its Homologs.

TABLE III-18

Summary of Ozonization Byproducts of Organic Compounds— Miscellaneous Compounds.

Reactions with Inorganic Compounds

In general, ozone oxidizes inorganic elements to a stable high oxidation state. Among the common metallic elements, iron is converted to the ferric state (Mallevialle et al., 1978) and manganese to either the IV or VII oxidation state, manganate (MnO₄2^-) or manganese dioxide (MnO₂), and permanganate (MnO₄^-), respectively. Ozone followed by filtration has been used to reduce iron and manganese to levels below those required by public health standards (Kjos et al., 1975). Ozone effectively removes a variety of trace metals as insoluble oxides at pH values from 7 to 9 (Netzer and Bowers, 1975). Sulfur in organic compounds often is converted to sulfate (SO₄2^-) (Gunther et al., 1970; Mallevialle et al., 1978), although in this and other cases the lower oxidation states of the element will result if small amounts of ozone are used.

Bromine is oxidized first to hypobromite (OBr^-) and then to bromate (BrO₃^-) (Ingols, 1978). Although hypobromite may be unstable in comparison to bromate under usual conditions, its transient existence is indicated by the formation of bromoform (CHBr₃) from the ozonization of seawater (Helz et al., 1978). Presumably one would also find iodinated organics under the same conditions.

The ammonium ion (NH₄⁺) is resistant to ozone attack, but above pH 7 ammonia is oxidized to nitrate (NO₃^-), presumably with several forms of intermediaries, including nitrite (NO₂^-) (Hoigne and Bader, 1978b; Huibers et al., 1969; Singer and Zilli, 1975; Wynn et al., 1973). Organic amines have yielded both nitrate and nitrite (Rogozhkin et al., 1970). Again, the dose of ozone could determine the relative yield of the oxidized forms in many cases.

Ozone has been used extensively for the oxidation of free and complexed cyanide (CN^-) to cyanate (OCN^-) (Bollyky, 1975; Garrison et al., 1975; Mathieu, 1975; Selm, 1959).

TABLE III-12Summary of Ozonization Byproducts of Organic Compounds— Aromatic Compounds: Polynuclear Aromatics

View in own window

Inferences of Possible Ozone Reaction By-Products

As indicated above, attempts to identify new by-products from the reaction of ozone with organic materials in water have been limited to rather conventional separation and identification methods. As expected, rather conventional compounds have been identified as reaction byproducts; however, some of them have important characteristics. For example, paroxon (C₉H₁₀O₂) from parathion (C₁₀H₁₄NO₅PS) is reported to be more toxic than its precursor. Nonetheless, the most important byproducts from many organic substrates may have been overlooked because of shortcomings of the analytical method that was used. There is indirect evidence of the presence of these species. Carlson and Caple (1977) observed residual oxidizing species in ozonized water long after ozone would have decomposed. Moreover, analysis of ozone in water by a purge method yields only 81% of the value that was obtained by the direct iodide/thiosulfate titration procedure (Brody, 1975).

It appears that studies have not shown that many of the labile intermediates, such as hydroxyhydroperoxides, peroxides, etc., are absent in ozonized organic solutions. Investigators simply have not looked for them carefully. These by-products may be of considerable significance to human health and need evaluating in this respect.

The Criegee mechanism for the reaction of ozone with carbon-carbon double bonds (Reaction 24) predicts that a hydroxyhydroperoxide will be formed as an intermediate when water is used as the solvent (Criegee, 1959).

Moreover, Oehlschaeger (1978) reported that various polymeric peroxides form in nonaqueous solvents, but there have been no known studies to determine if this occurs in water.

(24)

TABLE III-13Summary of Ozonization Byproducts or Organic Compounds— Aromatic Compounds: Phenolics

View in own window

TABLE III-14Summary of Ozonization Byproducts of Organic Compounds— Aromatic Compounds: Miscellaneous

View in own window

TABLE III-15Summary of Ozonization Byproducts of Organic Compounds—Humic Acids and Similar Natural Products

View in own window

Substrate	Products	Reference
Humic acid (Aldrich Chemical)	Malonic acida	Lawrence, 1977
	Hexanoic acida
	Succinic acida,b
	Heptanoic acid
	Benzoic acida,b
	Octanoic acid
	Glutaric acida,b
	Adipic acida,b
	Pimelic acida,b
	Suberic acida,b
	Azelaic acida,b
	Sebacic acida,b
	1,2,3,-Propane tricarboxylic acida,b
	Di-, tri-, tetra-, and pentacarboxylic Benzenes
	1,2,3,4-Butane tetracarboxylic acid
Humic acid (Peat)	Crenic and Apocrenic acids	Shevchenko and Taran, 1966
	Oxalic acid
Humic acid (Coal Derived)	Oxalic acid (as above) plus ''ozone-resistant acids"	Ahmed and Kinney, 1950

a

Also identified from ozonized lignosulfonic acids.

b

Also identified from ozonized tannic acid.

TABLE III-16Summary of Ozonization Byproducts of Organic Compounds—Pesticides

View in own window

TABLE III-17Summary of Ozonization Byproducts of Organic Compounds— Aliphatic Compounds

View in own window

Epoxide Formation

The possibility that ozone will produce potentially dangerous epoxides has been the subject of much informal speculation, but few studies on the subject have been reported. Carlson and Caple (1977) found oleic acid oxide from the aqueous phase ozonization of oleic acid, but subsequent experiments indicated that the compound results from the oxidation of the acid by peracid intermediates in the reaction mixture arising from the autooxidation of aldehydic products. No examples of epoxide formation from polynuclear aromatic compounds with aqueous ozone have been found, but the possibility that these compounds exist as minor products cannot be dismissed. Further study in this area is needed.

Conclusions

The preceding sections of this chapter summarize the results of an extensive survey of the chemistry of disinfectants that are or may be used in water treatment. Although it has been recognized since 1974 that chlorine produces potentially harmful by-products, such as the ubiquitous chloroform, little attention has been given to the other prospective disinfectants and oxidants such as ozone, chlorine dioxide, chloramines, and the other halogens, bromine and iodine. Quite apart from the question of the efficacy of these substances as alternative disinfectants, there remains the question, "Will the substitution of a disinfectant for chlorine in water treatment merely produce a different set of by-products whose effects on human health may be as significant, or more so, than those by-products known to be produced from chlorine?"

Unfortunately, a definitive answer cannot yet be given to this question. The surveys that are reported in the foregoing sections illustrate the lack of information on this subject even for the two common disinfectant chemicals, chlorine and ozone. There have been few studies in which the water of a treatment plant has been characterized before and after treatment, and, in most of these cases, chlorine was used as the disinfectant chemical. There are only two studies of water that had been treated with ozone, which is a common treatment in Western Europe, and no in-plant studies of water for which chlorine dioxide or the other halogens were used.

The authors were forced to resort to a survey of laboratory studies in which the various disinfectants were used or to works where a secondary wastewater was treated with one of the alternative disinfectants. As expected, these surveys showed that each disinfectant produces a set of by-product compounds that reflect the nature of the starting materials and the disinfectants that were used. Unfortunately, many of the laboratory studies were conducted under conditions that were not comparable to those in water treatment. For example, the initial concentrations of starting materials or disinfectant doses were too high, pH was too low, etc. Moreover, in most cases, the characterization of the by-products was incomplete or in other ways equivocated, so that a direct parallel could not be drawn between the laboratory results and those that could be expected in an actual water treatment plant.

Nonetheless, it is clear that each disinfectant chemical examined in this survey produces by-products that may occur in actual water treatment applications. Of particular concern are the following substances that result from the use of the various disinfectants.

• From chlorine: the trihalomethanes (THM's), trichloroacetone (CCl₃COCH₃), and other largely uncharacterized chlorinated and oxidized intermediates that are formed from the complex set of precursors in natural waters; chloramines; chlorophenols; and the largely unknown products of dechlorination.
• From ozone: epoxides that may in principle result from unsaturated substrates such as oleic acid, although none have yet been found in drinking water; peroxides and other highly oxidized intermediates such as glyoxal (OHCCHO) and methylglyoxal (CH₃COCHO) from aromatic precursors.
• From bromine and iodine: THM's and other bromine and iodine analogs of chlorinated species; bromophenols, bromoindoles, and bromoanisoles; plus the halogens themselves, which may remain in drinking water as residual.
• From chlorine dioxide: chlorinated aromatic compounds; chlorate (ClO₃^-) and chlorite (ClO₂^-), which are often present as by-product or unreacted starting material from production of chloride dioxide; and chlorine dioxide itself.

This list, incomplete as it is, is compelling in that it shows that each disinfectant produces chemical side effects that should be examined in more detail before the disinfectant is widely adopted for water treatment. It is clear that each of these disinfectants, being highly reactive chemical agents, will have inevitable side effects. Even in the case of chlorine, all of these side effects are not fully understood. This situation is due partly to the lack of systematic studies designed to evaluate by-products from each alternative disinfectant under actual water treatment conditions. But of equal importance is the need for these studies to utilize the complete spectrum of analytical instrumentation for the detection and identification of by-product compounds, not just for compounds that are amenable to extraction and GC/MS analysis. It appears that a preoccupation with these techniques has led chemists to omit from their characterizations those very polar, water-soluble components that might elude extraction and those high-molecular-weight or labile compounds that might not be detectable by GC/MS.

The subcommittee hopes that this report will stimulate future research that will avoid these deficiencies and provide more complete information on the by-products that are associated with the use of each alternative disinfectant under realistic conditions.

References

Precursor Compounds and the Haloform Reaction

• Bell, R.P., and O.M. Lidwell. 1940. The base catalyzed prototropy of substituted acetones. Proc. R. Soc. (Lond.) A176:88-113.
• Black, A.P., and R.F. Christman. 1963. Chemical characteristics of fulvic acids. J. Am. Water Works Assoc. 55:897-912.
• Burges, N.A., H.M. Hurst, and B. Walkden. 1964. The phenolic constituents of humic acid and their relation to the lignin of the plant cover. Geochim. Cosmochim. Acta 28:1547-1554.
• Christman, R.F. 1978. a. Progress Report, EPA Research Grant No. R 804430, for the period October 15, 1977, through January 15, 1978. University of North Carolina, Raleigh.
• Christman, R.F. 1978. b. Progress Report, EPA Research Grant No. R 804430, for the period July 15, 1978, through October 15, 1978. University of North Carolina, Raleigh.
• Christman, R.F., and M. Ghassemi. 1966. Chemical nature of organic color in water. J. Am. Water Works Assoc. 58:723-741.
• Christman, R.R., and R.T. Oglesby. 1971. Microbiological degradation and the formation of humus. Pp. 769-795 in K.V. Sarkanen, editor; and C.H. Ludwig, editor. , eds. Lignins: Occurrence, Formation, Structure, and Reactions. Wiley Interscience Publishers, New York.
• Dubach, P., N.C. Mehta, T. Jakab, F. Martin, and N. Roulet. 1964. Chemical investigations on soil humic substances. Geochim. Cosmochim. Acta 28:1567-1578.
• Fuson, R.C., and R.A. Bull. 1934. The haloform reaction. Chem. Rev. 15:275-309.
• Gagosian, R.B., and D.H. Stuermer. 1977. The cycling of biogenic compounds and their diagenetically transformed products in seawater. Mar. Chem. 5:605-632.
• Gjessing, E.T. 1976. Physical and Chemical Characteristics of Aquatic Humus. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 120 pp.
• Green, G., and G. Steelink. 1962. Structure of soil humic acid. II. Some copper oxide oxidation products. J. Org. Chem. 27:170-174.
• Khan, S.U., and F.J. Sowden. 1972. Distribution of nitrogen in fulvic acid fraction extracted from the black solonetzic and black chernozemic soils of Alberta. Can. J. Soil Sci. 52:116-118.
• Morris, J.C. 1975. Formation of Halogenated Organics by Chlorination (A Review). EPA-600/1-75-002, Office of Research and Development, U.S. Environmental Protection Agency, Washington, D.C. 154 pp.
• Morris, J.C, and B. Baum. 1978. Precursors and mechanisms of haloform formation in the chlorination of water supplies. Pp. 29-48 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Neyroud, J.A., and M. Schnitzer. 1975. Alkaline hydrolysis of humic substances. Geoderma 13:171-188.
• Nissenbaum, A., and I.R. Kaplan. 1972. Chemical and isotopic evidence for the in situ origin of marine humic substances. Limnol. Oceanogr. 17:570-582.
• Schnitzer, M., and S.U. Kahn. 1972. Humic Substances in the Environment. Marcel Dekker, New York. 327 pp.
• Shapiro, J. 1957. Chemical and biological studies on the yellow organic acids of lake water. Limnol. Oceanogr. 2:161-179.
• Steelink, C. 1977. Humates and other natural organic substances in the aquatic environment J. Chem. Ed. 54:599-603.
• Stevens, A.A., C.J. Solcum, D.R. Seeger, and G.G. Robeck. 1978. Chlorination of organics in drinking water. Pp. 77-104 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Stuermer, D.H., and G.R. Harvey. 1978. Structural studies on marine humus: a new reduction sequence for carbon skeleton determination. Mar. Chem. 6:55-70.
• Whittaker, R.H., and G.E. Likens. 1973. Carbon in the biota. In G.M. Woodwell, editor; and E.V. Pecan, editor. , eds. Carbon and the Biospere. AEC Symposium Series 30:281-302. Technical Information Center, U.S. Atomic Energy Commission, Washington, D.C.

Chlorine

• Andelman, J.B., and J.E. Snodgrass. 1974. Incidence and significance of polynuclear aromatic hydrocarbons in the water environment. Crit. Rev. Environ. Control 4:69-83.
• Babcock, D.P., and P.C. Singer. 1977. Chlorination and coagulation of humic and fulvic acids. Proceedings of the 97th Annual American Water Works Association Conference held in Anaheim, California, May 8-13, 1977. American Water Works Association, Denver, Colo.
• Barnes, D.B. 1978. Trihalomethane-forming potential of algal extracellular products and biomass. MS. thesis. Virginia Polytechnic Institute and State University, Blacksburg.
• Bell, R.P., and O.M. Lidwell. 1940. The base catalyzed prototropy of substituted acetones. 1940. Proc. R. Soc. (Lond.) A176:88-113.
• Bellar, T.A., J.J. Kichtenberg, and R.C. Kroner. 1974. The occurrence of organohalides in chlorinated drinking water. J. Am. Water Works Assoc. 66(12):703-706.
• Blakenship, W.M. 1978. Personal communication with R.C. Hoehn. U.S. Environmental Protection Agency, Region III, Water Supply Branch, Philadelphia, Pa.
• Blumer, M., and W.W. Youngblood. 1975. Polycyclic aromatic hydrocarbons in soils and recent sediments. Science 188:53-55. [PubMed: 17760164]
• Brass, H.J., M.A. Feige, T. Halloran, J.W. Mello, D. Munch, and R.F. Thomas. 1977. The national organic monitoring survey: samplings and analyses for purgeable organic compounds. Pp. 393-416 in R.B. Pojasek, editor. , ed. Drinking Water Quality Enhancement Through Source Protection. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Burttschell, R.H., A.A. Rosen, F.M. Middleton, and M.B. Ettinger. 1959. Chlorine derivatives of phenol causing taste and odor. J. Am. Water Works Assoc. 51:205-214.
• Carlson, R.M., and R Caple. 1978. Organochemical implications of water chlorination. Pp. 65-75 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Carlson, R.M., R.E. Carlson, H.L. Kopperman, and R. Caple. 1975. Facile incorporation of chlorine into aromatic systems during aqueous chlorination processes. Environ. Sci. Technol. 9:674-675.
• Christman, R.F., and M. Ghassemi. 1966. Chemical nature of organic color in water. J. Am. Water Works Assoc. 58:723-741.
• Christman, R.F., J.D. Johnson, J.R. Haas, F.K. Pfaender, W.T. Liao, D.L Norwood, and H.J. Alexander. 1978. a. Natural and model aquatic humics: reactions with chlorine. Pp. 15-28 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Christman, R.F., J.D. Johnson, and D.L. Norwood. 1978. b. Progress Reports, EPA Research Grant No. R 804430 for Oct. 15, 1977, through Oct. 15, 1978.
• Coleman, W.E, R.L. Lingg, R.G. Melton, and F.C. Kopfler. 1976. The occurrence of volatile organics in five drinking water supplies using gas chromatography/mass spectrometry. Pp 305-327 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 718 pp.
• Fuchs, F., and W. Kuhn. 1976. The use of activated carbon to analyze natural waters with regard to their behaviour in waterworks filters. In H. Sontheimer, editor. , ed. Translation of Reports on Special Problems of Water Technology, Vol. 9. Adsorption. Conference held in Karlsruhe, West Germany, 1975. EPA-600/9-76-030. U.S. Environmental Protection Agency, Water Supply Research Division, Cincinnati, Ohio.
• Garrison, A.W. 1978. Personal communication with R.C. Hoehn, U.S. Environmental Protection Agency, Southeastern Regional Laboratory, Athens, Ga.
• Garrison, A.W., J.D. Pope, and F.R. Allen. 1976. GC/MS analysis of organic compounds in domestic wastewaters. Pp. 517-556 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Gehrs, C.W., and G.R. Southworth. 1978. Investigating the effect of chlorinated organics. Pp. 329-342 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Giger, W., M. Reinhard, C. Schaffner, and F. Zurcher. 1976. Analyses of organic constituents in water by high-resolution gas chromatography in combination with specific detection and computer-assisted mass spectrometry. Pp. 433-452 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Glaze, W.H., and J.E. Henderson, IV. 1975. Formation of organochlorine compounds from the chlorination of a municipal secondary effluent. J. Water Pollut. Control Fed. 47:2511-2515.
• Glaze, W.H., J.E. Henderson, IV., and G. Smith. 1976. Analysis of new chlorinated organic compounds in municipal wastewaters after terminal chlorination. Pp. 247-254 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Glaze, W.H., J.E. Henderson, IV., and G. Smith. 1978. Analysis of new chlorinated organic compounds formed by chlorination of municipal wastewater. Pp. 139-159 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc, Ann Arbor, Mich. 439 pp.
• Harrison, R.M., R. Perry, and R.A. Wellings. 1975. Review paper: polynuclear aromatic hydrocarbons in raw, potable, and waste waters. Water Res. 9:331-346.
• Hase, A., and R.A. Hites. 1976. On the origin of polycyclic aromatic hydrocarbons in the aqueous environment Pp. 205-214 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Hoehn, R.C., and C.W. Randall. 1977. Drinking water disinfection and chlorinated organics formation. Proceedings AWWA Disinfection Seminar. Presented at the 97th Annual Conference of the American Water Works Association in Anaheim, California, May 8, 1977. Paper No. 11, pp.1-18.
• Hoehn, R.C., C.W. Randall, F.A. Bell, Jr., and P.T.B. Shaffer. 1977. Triahlomethanes and viruses in a water supply. J. Environ. Eng. Div., ASCE 103:803-814.
• Hoehn, R.C., R.P. Goode, C.W. Randall, and P.T.B. Shaffer. 1978. Chlorination and treatment for minimizing trihalomethanes in drinking water. Pp. 519-535 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 PP.
• Jolley, R.L. 1975. Chlorine-containing organic constituents in chlorinated effluents. J. Water Pollut. Control Fed. 47:601-618.
• Jolley, R.L., G. Jones, W.W. Pitt, Jr., and J.E. Thompson. 1976. Determination of chlorination effects on organic constituents in natural and process waters using high-pressure liquid chromatography. Pp. 233-246 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Jolley, R.L., G. Jones, W.W. Pitt, and J.E. Thompson. 1978. Chlorination of organics in cooling waters and process effluents. Pp. 105-138 in R.L Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Keith, L.H., A.W. Garrison, F.R. Allen, M.H. Carter, T.L. Floyd, J.D. Pope, and A.D. Thruston, Jr. 1976. Identification of organic compounds in drinking water from thirteen U.S. cities. Pp. 329-373 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Kleopfer, R.D. 1976. Analysis of drinking water for organic compounds. Pp. 399-416 in L.H. Keith, editor. , ed. Identification and Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Laubusch, E.J. 1959. Standards of purity for liquid chlorine. J. Am. Water Works Assoc. 51:742-748.
• Lee, G.F. 1967. Kinetics of reactions between chlorine and phenolic compounds. Pp. 54-74 in S.D. Faust, editor; and J.V. Hunter, editor. , eds. Principles and Applications of Water Chemistry. Proceedings of the Fourth Rudolfs Research Conference held at Rutgers State University, New Brunswick, N.J., 1965. John Wiley & Sons, Inc., New York.
• Lee, G.F., and J.C. Morris. 1962. Kinetics and chlorination of phenol-chlorophenolic tastes and odors. Int. J. Air Water Pollut. 6:419-431. [PubMed: 13929076]
• Miller, G.W., R.G. Rice, C.M. Robson, W. Kuhn, and H. Wolf. 1978. An assessment of ozone and chlorine dioxide technologies for treatment of municipal water supplies, Part 2, Section 9. Draft Report. EPA Grant No. R 80435-01, Municipal Environmental Research Laboratory, Office of Water Supply, US. Environmental Protection Agency, Cincinnati, Ohio. 96 pp.
• Morris, J.C. 1967. Kinetics of reactions between aqueous chlorine and nitrogen compounds. Pp. 23-53 in S.D. Faust, editor; and J.V. Hunter, editor. , ed. Principles and Applications of Water Chemistry. Proceedings of the Fourth Rudolfs Research Conference, held at Rutgers State University, New Brunswick, NJ., 1965. John Wiley & Sons, Inc., New York.
• Morris, J.C. 1975. Formation of Halogenated Organics by Chlorination of Water Supplies (A Review). EPA-600/l-75-002, Office of Research and Development. U.S. Environmental Protection Agency, Washington, D.C. 154 pp.
• Morris, J.C. 1978. The chemistry of aqueous chlorine in relation to water chlorination. Pp. 21-35 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Morris, J.C., and B. Baum. 1978. Precursors and mechanisms of haloformation in the chlorination of water supplies. Pp. 29-48 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Munch, D.J., M.A. Feige, and H.J. Brass. 1977. The analyses of purgeable compounds in the national organic monitoring survey by gas chromatography/mass spectrometry. In Water Quality in the Distribution System. American Water Works Association's Fifth Annual Water Quality Technology Conference held in Kansas City, Dec. 4-7, 1977. Paper 3A-6, 5 pp. American Water Works Association, Denver, Colo.
• Oyler, A.R., D.L. Bodenner, K.J. Welch, R.J. Liukkonen, R.M. Carlson, H.L. Kopperman, and R. Caple. 1978. Determination of aqueous chlorination reaction products of polynuclear aromatic hydrocarbons by reversed phase high performance liquid chromatography-gas chromatography. Anal. Chem. 50:837-842.
• Pfaender, F.K., R.B. Jonas, A.A. Stevens, L. Moore, and J.R. Haas. 1978. Evaluation of direct aqueous injection method for analysis of chloroform in drinking water. Environ. Sci. Technol. 12:438-441.
• Rook, J.J. 1974. Formation of haloforms during chlorination of natural waters. Water Treat. Exam. 23:234-243.
• Rook, J.J. 1976. Haloforms in drinking water. J. Am. Water Works Assoc. 68:168-172.
• Rook, J.J. 1977. Chlorination reactions of fulvic acids in natural waters. Environ. Sci. Technol. 11:478-482.
• Rosenblatt, D.H. 1975. Chlorine and oxychlorine species reactivity with organic substances. Pp. 249-276 in J.D. Johnson, editor. , ed. Disinfection: Water and Wastewater. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 425 pp.
• Shackelford, W.M., and L.H. Keith. 1976. Frequency of Organic Compounds Identified in Water. EPA-600/4-76-062. Environmental Research Laboratory, Office of Research and Development, U.S. Environmental Protection Agency, Athens, Ga. 629 pp.
• Sievers, R.E., R.M. Barkley, G.A. Eiceman, L.P. Haack, R.H. Shapiro, and H.F. Walton. 1978. Generation of volatile organic compounds from non-volatile precursors by treatment with chlorine or ozone. Pp. 615-624 in R.L Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Stevens, A.A., C.J. Slocum, D.R. Seeger, and G.G. Robeck. 1976. Chlorination of organics in drinking water. J. Am. Water Works Assoc. 68:615-620.
• Stevens, A.A., C.J. Slocum, D.R. Seeger, and G.G. Robeck. 1978. Chlorination of organics in drinking water. Pp. 77-104 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. 1. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Suffet, I.H., L. Brenner, and B. Silver. 1976. Identification of 1,1,1-trichloroacetone (1,1,1-trichloropropanone) in two drinking waters: a known precursor in the haloform reaction. Environ. Sci. Technol. 10:1273-1275. [PubMed: 22175753]
• Symons, J.M., T.A. Bellar, J.K. Carswell, J. Demarco, K.L. Kropp, G.G. Robeck, D.R Seeger, C.J. Slocum, B.L. Smith, and A.A. Stevens. 1975. National organic reconnaissance survey for halogenated organics. J. Am. Water Works Assoc. 67:634-647.
• Thompson, B.C. 1978. Trihalomethane formation potential of algal extracellular products and biomass. M.S. thesis. Virginia Polytechnic Institute and State University, Blacksburg. 135 pp.
• Vallentyne, J.R. 1957. The molecular nature of organic matter in lakes and oceans, with lesser reference to sewage and terrestrial soils. J. Fish. Res. Bd. Can. 14:33-82.

Chloramines

• Agrawal, M.C., and S.P. Mushran. 1973. Mechanism of oxidation of aldoses by chloramine-T. J. Chem. Soc. Perkin Trans. 2:762-765.
• Antelo, J.M., J.M. Cachaza, J. Casado, and M.A. Herraez. 1974. Influence of pH on the reaction of cresols with chloramine-T. An. Quim. 70:555-558; Chem. Abstr. 82:15941z (1975).
• Banerji, K.K. 1977. Kinetics and mechanism of the oxidation of substituted benzyl alcohols by chloramine-T in acid solution. Bull. Chem. Soc. Jpn. 50(6): 1616-1618.
• Bauer, R.C., and V.L. Snoeyink. 1973. Reactions of chloramines with active carbon. J. Water Pollut. Control Fed. 45:2290-2301.
• Burttschell, R.H., A.A. Rosen, F.M. Middleton, and M.B. Ettinger. 1959. Chlorine derivatives of phenol causing taste and odor. J. Am. Water Works Assoc. 51:205-214.
• Chapin, R.M. 1929. Dichloroamine. J. Am. Chem. Soc. 51:2117-2122.
• Colton, E., and M.M. Jones. 1955. Monochloramine. J. Chem. Educ. 32:488-489.
• Corbett, R.E., W.S. Metcalf, and F.G. Soper. 1953. Studies of N-halogeno-compounds. Part IV. The reaction between ammonia and chlorine in aqueous solution, and the hydrolysis constants of chloroamines. J. Chem. Soc. 1953:1927-1929.
• Crochet, R.A., and P. Kovacic. 1973. Conversion of o-hydroxyaldehydes and ketones into o-hydroxyanilides by monochloramine. J. Chem. Soc. Chem. Commun. 197(19):716-717.
• Cross, C.F., E.J. Bevan, and W. Bacon. 1910. Chloramine reactions. Methylenechloroamine. J. Chem. Soc. Trans. 97:2404-2406.
• Czech, F.W., R.J. Fuchs, and H.F. Antczak. 1961. Determination of mono-, di-, and trichloramine by ultraviolet absorption spectrophotometry. Anal. Chem. 33:705-707.
• Dowell, C.T., and W.C. Bray. 1917. Experiments with nitrogen trichloride. J. Am. Chem. Soc. 39:896-905.
• Drago, R.S. 1957. Chloramine. J. Chem. Educ. 34:541-545.
• Ellis, A.J., and F.G. Soper. 1954. Studies of N-halogeno-compounds. Part VI. The kinetics of chlorination of tertiary amines. J. Chem. Soc. 1954:1750-1755.
• Gmelin's Handbuch der Anorganischen Chemie. 1969. Chlor und Stickstoff. Pp. 483-500 in Gmelin's Handbuch der Anorganischen Chemie, 8. Auflage. Chlor Erganzungsband, Teil B-Lieferung 2, System-Nr. 6. Verlag Chemie, G.m.b.H., Weinheim/Bergstr, W. Germany.
• Gowda, N.M.M., A.S.A. Murthy, and D.S. Mahadevappa. 1975. Oxidation of cysteine with chlorimine-T and dichloramine-T. Curr. Sci. 44:5-6.
• Gray, E.T., Jr., and D.W. Margerum. 1978. Kinetics and equilibria of chloramine species. In Abstracts of Papers of the 175th American Chemical Society National Meeting, Anaheim, Calif., March 13-17. Abstr. No. INOR 257.
• Harwood, J.E., and A.L. Kuhn. 1970. A colorimetric method for ammonia in natural waters. Water Res. 4:805-811.
• Hauser, C.R., and M.L. Hauser. 1930. Researches on chloramines. I. Orthochlorobenzalchlorimine and anisalchlorimine. J. Am. Chem. Soc. 52:2050-2054.
• Houben, J., and T. Weyl. 1962. Halogen-Verbindungen. Pp. 760-811 in J. Houben, editor; and T. Weyl, editor. , eds. Methoden der Organischen Chemie (Houben-Weyl), 4. Auflage, Band 5, Teil 3. Geroge Thieme Verlag, Stuttgart, W. Germany.
• Ingols, R.S., H.A. Wyckoff, T.W. Kethley, H.W. Hodgden, E.L. Fincher, J.C. Hildebrand, and J.E. Mandel. 1953. Bactericidal studies of chlorine. Ind. Eng. Chem. 45:996-1000.
• Jander, J. 1955. a. Ein Beitrag zur Kenntnis des Monochloramins. Naturwissenschaften 42:178-179.
• Jander, J. 1955. b. Zum Verstandnis der Chemie der Chlor-Stickstoff-und Chlor-Sauerstoff-Verbindungen. Z. Anorg. Allg. Chem. 280:276-283.
• Kinman, R.N., and R.F. Layton. 1976. New method for water disinfection. J. Am. Water Works Assoc. 68(6):298-302.
• Kirk-Othmer Encyclopedia of Chemical Technology. 1964. Vol. 4, 2nd ed. Wiley Interscience Publishers, New York.
• Kovacic, P., M.K. Lowery, and K.W. Field. 1970. Chemistry of N-bromamines and N-chloramines. Chem. Rev. 70:639-665.
• Kumar, A., A.K. Bose, and S.P. Mushran. 1975. Kinetic study of oxidation of acetophenone by chloramine-T. Ann. Soc. Sci. Bruxelles, Ser. 1, 89:567-574.
• Kumar, A., R.M. Mehrostra, and S.P. Mushran. 1976. a. Kinetics and mechanism of oxidation of cyclohexanol by chloramine-T. Bull. Acad. Pol. Sci., Ser. Sci. Chim. XXIV: 181-185.
• Kumar, A., A.K. Bose, and S.P. Mushran. 1976. b. Kinetics and mechanism of oxidation of phenyalanine and serine by chloramine-T. J. Indian Chem. Soc. LIII:755-758.
• Lindsay, M., and F.G. Soper. 1946. Methylenechloroamine. J. Chem. Soc. 1946:791-792.
• Mahadevappa, D.S., and N.H.M. Gowda. 1975. Estimation of glutathione with chloramine-T and dichloramine-T. Talanta 22:771-773. [PubMed: 18961731]
• Mahadevappa, D.S., and N.H.M. Gowda. 1975. Estimation of glutathione with chloramine-T and dichloramine-T. Talanta 22:771-773. [PubMed: 18961731]
• Mahadevappa, D.S., and H.M.K. Naidu. 1976. Oxidation of valine, leucine, and phenyl alanine by chloramine-T. Curr. Sci. 45:652-653.
• Mani, U.V., and A.N. Radhakrishnan. 1976. The oxidation of hydroxyproline by chloramine-T. Evidence discounting pryrrole-2-carboxylate as an intermediate in the reaction. Indian J. Biochem. Biophys. 13:185-186. [PubMed: 1010556]
• Margerum, W., and E.T. Gray, Jr. 1978. Chlorination and the formation of N-chloro compounds in water treatment. In Abstracts of Papers of the 175th American Chemical Society National Meeting, Anaheim, Calif., March 13-17. Abstr. No. INOR 158.
• Metcalf, W.S. 1942. The absorption spectra of mono-, di-, and trichloramines and some aliphatic derivatives. J. Chem. Soc. 1942:148-150.
• Morris, J.C. 1967. Kinetics of reactions between aqueous chlorine and nitrogen compounds. Pp. 23-53 in S.D. Faust, editor; and J.V. Hunter, editor. , eds. Principles and Applications of Water Chemistry. Proceedings of the Fourth Rudolfs Research Conference held at Rutgers State University, New Brunswick, NJ., 1965. John Wiley & Sons, Inc., New York.
• Morris, J.C. 1978. Modern chemical methods in water and waste treatment, Vol. I. International Course in Sanitary Engineering, International Institute for Hydraulic and Environmental Engineering, Delft, The Netherlands.
• Mushran, S.P., R.M. Mehrotra, and R. Sanehi. 1974. a. Kinetics and mechanism of chloraminometric reactions involving some primary alcohols. J. Indian Chem. Soc. LI:594-596.
• Mushran, S.P., K.C. Gupta, and R. Sanehi. 1974. b. Kinetics and mechanism of oxidation of D-(-)-ribose by chloramine-T. J. Indian Chem. Soc. LI:145-148.
• Naidu, H.M.K., and D.S. Mahadevappa. 1976. Oxidation of crotoyl alcohol with chloramine-T. Curr. Sci. 45:216-218.
• Nair, C.G.R., and V.R. Nair. 1973. Dichloramine-T as a new oxidimetric titrant in nonaqueous and partially aqueous media. II. Potentiometric determination of hydroquinone, hydrazine, oxine, cinnamic acid, tin (II), antimony (III), thallium (I), and ferrocyanide. Talanta 20:696-699. [PubMed: 18961334]
• Natarajan, N.M., and V. Thiagarajan. 1975. Kinetics of oxidation of secondary alcohols by chloramine-T. J. Chem. Soc. Perkin Trans. 2:1590-1594.
• Neale, R. 1964. The chemistry of ion radicals: The free-radical addition of N-chloramines to olefinic and acetylenic hydrocarbons. J. Am. Chem. 86:5340-5342.
• Palin, A.T. 1950. A study of the chloro derivatives of ammonia and related compounds, with special reference to their formation in the chlorination of natural and polluted waters. Water Water Eng. 54:151-159.
• Petrov, A.A., G.A. Galaev, and D.B. Ioffe. 1953. Halogenation and thiocyanation of aromatic amines with liquid chloramines, chloramides, alkyl hypochlorites, and organic hydroperoxide. J. Gen. Chem. USSR 23:689-692.
• Ramanujam, V.M.S., and N.M. Trieff. 1977. Kinetic and mechanistic studies of reactions of aniline and substituted anilines with chloramine-T. J. Chem. Soc. Perkin Trans. 2:1275-1280.
• Remick, A.E. 1942. A semiquantitative extension of the electronic theory of the English school J. Org. Chem. 7:534-545.
• Rickabaugh, J., and R.N. Kinman. 1978. Trihalomethane formation from iodine and chlorine disinfection of Ohio River water. Pp. 583-591 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Robson, H.L. 1964. Chloramines and chloroamines. Pp. 908-929 in Kirk-Othmer Encyclopedia of Chemical Technology, Vol. 4, 2nd ed. Interscience Publishers, New York.
• Saguinsin, J.L.S., and J.C. Morris. 1975. The chemistry of aqueous nitrogen trichloride. Pp. 277-299 in J.D. Johnson, editor. , ed. Disinfection Water and Wastewater. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 425 pp.
• Sanehi, R., K.C. Gupta, R.M. Mehrotra, and S.P. Mushran. 1975. Kinetics and mechanism of oxidation of D-(+)-sorbose by chloramine-T. Bull. Chem. Soc. Jpn. 48:330-332.
• Shih, K.L., and J. Lederberg. 1976. a. Chloramine mutagenesis in Bacillus subtilis . Science 192:1141-1143. [PubMed: 818709]
• Shih, K.L., and J. Lederberg. 1976. b. Effects of chloramine on Bacillus subtilis deoxyribonucleic acid. J. Bacteriol. 125:934-945. [PMC free article: PMC236169] [PubMed: 815252]
• Sisler, H.H., N.K. Kotia, and R.E. Highsmith. 1970. The formation of sulfur-sulfur bonds by the chloramination of thiols. J. Org. Chem. 35:1742-1745.
• Srivastava, A., and S. Bose. 1975. a. Determination of mercaptans and xanthates with chloramine-T and chloramine-B. J. Indian Chem. Soc. LII:214-216.
• Srivastava, A., and S. Bose. 1975. b. Analytical applications of N-haloamines N-haloamides for the determination of some sulfur containing functional groups. J. Indian Chem. Soc. LII:217-220.
• Standard Methods for the Examination of Water and Wastewater, 14th ed. 1976. American Public Health Association, Washington, D.C. 1193 pp.
• Stasiuk, W.N., Jr. 1974. Reactions of chloramines with activated carbon. Ph.D. thesis. Rensselaer Polytechnic University, Troy, N.Y. 136 pp.
• Stevens, A.A., C.J. Slocum, D.R. Seeger, and G.G. Pobeck. 1978. Chlorination of organics in drinking water. Pp. 77-104 in R.L. Jolley, editor. , ed. Water Chlorination: Environmental Impact and Health Effects, Vol. I. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 439 pp.
• Symons, J.M., J.K. Carswell, R.M. Clark, P. Dorsey, E.E. Geldreich, W.P. Heffernam, J.C. Hoff, O.T. Love, L.J. McCabe, and A.A. Stevens. 1977. Ozone, Chlorine Dioxide, and Chloramines as Alternatives to Chlorine for Disinfection of Drinking Water: State of the Art. Water Supply Research Division, U.S. Environmental Protection Agency, Cincinnati, Ohio. 84 pp.
• Theilacker, W., and E. Wegner. 1964. Organic syntheses using chloramine. Pp. 303-317 in W. Foerst, editor. , ed. Newer Methods of Preparative Organic Chemistry. Vol. III. (Trans. by H. Birnbaum). Academic Press, Inc., New York.
• Trieff, N.M., and V.M.S. Ramanujam. 1977. Removal of odorous aromatic amine environmental pollutants by chloramine-T. Bull. Environ. Contam. Toxicol. 18(1):2628. [PubMed: 884333]
• U.S. Public Health Service. 1963. Inventory of Municipal Water Supplies. PHS Publication No. 1039. U.S. Public Health Service, Washington, D.C.
• Wei, I.W. 1972. Chlorine-ammonia breakpoint reactions: kinetics and mechanism. Ph.D. dissertation. Harvard University, Cambridge, Mass.
• Wei, I.W., and J.C. Morris. 1974. Dynamics of breakpoint chlorination. Pp. 297-332 in A.J. Rubin, editor. , ed. Chemistry of Water Supply, Treatment, and Distribution. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Weil, I., and J.C. Morris. 1949. Kinetic studies on the chloramines. I. The rates of formation of monochloramine, N-chlormethylamine and N-chlorodimethylamine. J. Am. Chem. Soc. 71:1664-1671.
• White, G.C. 1972. Handbook of Chlorination. Van Nostrand Reinhold, New York. 744 pp.

Bromine and Iodine

• Bean, R.M., R.G. Riley, and P.W. Ryan. 1978. Investigation of halogenated components formed from chlorination of marine water. Pp. 223-233 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds., Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Bellar, T.A., J.J. Lichtenberg, and R.C. Kroner. 1974. The occurrence of organohalides in chlorinated drinking waters. J. Am. Water Works Assoc. 66:703-706.
• Berliner, E. 1966. The current state of positive halogenating agents. J. Chem. Educ. 43:124-133.
• Black, A.P., W.C. Thomas, Jr., R.N. Kinman, W.P. Bonner, M.A. Keirn, J.J. Smith, Jr., and A.A. Jabero. 1968. Iodine for the disinfection of water. J. Am. Water Works Assoc. 60:69-83.
• Bunn, W.W., B.B. Haas, E.R. Deane, and R.D. Kleopfer. 1975. Formation of trihalomethanes by chlorination of surface water. Environ. Lett. 10:205-213. [PubMed: 1213009]
• Chang, S.L. 1958. The use of active iodine as a water disinfectant J. Am. Pharm. Assoc. 47:417-423. [PubMed: 13549285]
• Engel, P., A. Oplatka, and B. Perlmutter-Hayman. 1954. The decomposition of hypobromite and bromite solutions. J. Am. Chem. Soc. 76:2010-2015.
• Environmental Health Directorate. 1977. National Survey for Halomethanes in Drinking Water. Department of National Health and Welfare, Health Protective Branch, Ottawa, Canada. Publ. No. 77-EHD. 119 pp.
• Freund, G., W.C. Thomas, Jr., E.D. Bird, R.N. Kinman, and A.P. Black. 1966. Effect of iodinated water supplies on thyroid function. J. Clin. Endocrinol. Metab. 26:619-624. [PubMed: 4161810]
• Gilow, H.M., and J.H. Ridd. 1973. Mechanism of aromatic bromination by hypobromous acid in aqueous perchloric acid. Kinetic evidence against the prior formation of 'positive bromine.' J. Chem. Soc. Lond. Perkin Trans. 2:1321-1327.
• Goodenough, R.D., J.F. Mills, and J. Place. 1969. Anion exchange resin (polybromide form) as a source of active bromine for water disinfection. Environ. Sci. Technol. 3:854-856.
• Helz, G.R., and R.Y. Hsu. 1978. Volatile chloro- and bromocarbons in coastal waters. Limnol. Oceanogr. 23:858-869.
• Henderson, J.R., G.R. Peyton, and W.H. Glaze. 1976. A convenient liquid-liquid extraction method for the determination of halomethanes in water at the parts-per-billion level. Pp. 105-111 in L.H. Keith, editor. , ed. Identification & Analysis of Organic Pollutants in Water. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich.
• Kuehl, D.W., G.D. Veith, and E.N. Leonard. 1978. Brominated compounds found in waste-treatment effluents and their capacity to bioaccumulate. Pp. 175-192 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 909 pp.
• Johnson, J.D., and R. Overby. 1971. Bromine and bromamine disinfection chemistry. J. Sanit. Eng. Div., Am. Soc. Civ. Eng. 97:617-628.
• LaPointe, T.F., G. Inman, and J.D. Johnson. 1975. Kinetics of tribromamine decomposition. Pp. 301-338 in J.D. Johnson, editor. , ed. Disinfection: Water and Wastewater. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 425 pp.
• Laubusch, E.J. 1971. Chlorination and other disinfection processes. Pp. 158-224 in American Water Works Association. Water Quality and Treatment. A Handbook of Public Water Supplies, 3rd ed. McGraw-Hill, New York.
• Livingstone, D.A. 1963. Chemical Composition of Rivers and Lakes. U.S. Geological Survey Professional Paper 440-G. Washington, D.C. 64 pp.
• Macalady, D.L., J.H. Carpenter, and C.A. Moore. 1977. Sunlight-induced bromate formation in chlorinated seawater. Science 195:1335-1337. [PubMed: 17738417]
• Mills, J.F. 1975. Interhalogens and halogen mixtures as disinfectants. Pp. 113-143 in J.D. Johnson, editor. , ed. Disinfection: Water and Wastewater. Ann Arbor Science Publishers, Ann Arbor, Mich. 425 pp.
• Moelwyn-Hughes, E.A. 1971. The Chemical Statics and Kinetics of Solutions. Academic Press, New York. 530 pp.
• Morris, J.C., S.L. Chang, G.M. Fair, and G.H. Conant, Jr. 1953. Disinfection of drinking water under field conditions. Ind. Eng. Chem. 45:1013-1015.
• Pauling, L. 1960. The Nature of the Chemical Bond, and the Structure of Molecules and Crystals: An Introduction to Modern Structural Chemistry, 3rd ed. Cornell University Press, Ithaca, N.Y. 644 pp.
• Rickabaugh, J.F., and R.N. Kinman, 1978. Trihalomethane formation from iodine and chlorine disinfection of Ohio River water. Pp. 583-591 in R.L. Jolley, editor; , H. Gorchev, editor; , and D.H. Hamilton, Jr., editor. , eds. Water Chlorination: Environmental Impact and Health Effects, Vol. 2. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 905 pp.
• Rickabaugh, J.F. 1977. The study of trihalomethane formation when iodine is used for disinfection of drinking water. M.S. thesis. University of Cincinnati, Cincinnati, Ohio. 110 pp.
• Rook, J.J. 1974. Formation of haloforms during chlorination of natural waters. Water Treat. Exam. 23:234-243.
• Rook, J.J., A.A. Gras, B.G. van der Heijden, and J. de Wee. 1978. Bromide oxidation and organic substitution in water treatment J. Environ. Sci. Health A13:91-116.
• Shackleford, W.M., and L.H. Keith. 1976. Frequency of Organic Compounds Identified in Water. EPA-600/4-76-062. U.S. Environmental Protection Agency, Environmental Research Laboratory, Athens, Ga. 629 pp.
• Sugam, R.J. 1977. Chlorine degradation in estuarine waters. Ph.D. dissertation. University of Maryland, College Park, Md. 221 pp.
• Swain, C.G., and D.R. Crist. 1972. Mechanisms of chlorination by hypochlorous acid. The last of chlorinium ion, Cl+. J. Am. Chem. Soc. 94:3195-3200.
• Turekian, K.K. 1971. Rivers, tributaries, and estuaries. Pp. 9-73 in D.W. Hood, editor. , ed. Impingement of Man on the Oceans. Wiley Interscience Publishers, New York.
• White, G.C. 1972. Handbook of Chlorination. Van Nostrand Reinhold, New York. 744 pp.

Chlorine Dioxide

• Beuermann, L. 1965. Preparation of chlorine dioxide from sodium chlorite and hydrochloric acid. Gas-Wasserfach. 106:783-788.
• Bowen, E.J., and W.M. Cheung. 1932. The photodecomposition of chlorine dioxide solutions. J. Chem. Soc. Lond. 1932(Part I): 1200-1208.
• Bray, W. 1906. Beitrage zur Kenntnis der Halogensauerstoffverbindungen. Abhandlung III. Zur Kenntnis des Chlordioxyds. Z. Physik. Chem. 54:569-608.
• Buydens, R. 1970. Ozonization and its effects on the mode of purification of river waters. (In French) Trib. CEBEDEAU 23:319-320, 286-291.
• Dence, C.W., and K.V. Sarkanen. 1960. A proposed mechanism for the acidic chlorination of softwood lignin. Tappi 43:87-96.
• Dence, C.W., M.K. Gupta, and K.V. Sarkanen. 1962. Studies on oxidative delignification mechanisms. Part II. Reactions of vanillyl alcohol with chlorine dioxide and sodium chlorite. Tappi 45:29-38.
• Dowling, L.T. 1974. Chlorine dioxide in potable water treatment Water Treat Exam. 23(2): 190-204.
• Feuss, J.V. 1964. Problems in the determination of chlorine dioxide residuals. J. Am. Water Works Assoc. 56:607-615.
• Flis, I.E. et al. 1955. Trans. Leningr. Techn. Inst. Tsell, Bumazhu Prom. 16:62-67.
• Fuchs, W., and H. Leopold. 1927. Humic acids. II. The action of bromine, thionyl chloride and chlorine dioxide on artificial humic acids. Brennstoff Chem. 8:101-103.
• Fujii, M., and M. Ukita. 1957. Mechanism of wheat protein coloring by chlorine dioxide. Nippon Nogei Kagaku Kaishi 31:101-109.
• Gall, R.J. 1978. Chlorine dioxide—an overview of its preparation, properties and uses. Pp. 356-382 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute, Inc., and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Glabisz, U. 1968. The Reaction of Chlorine Dioxide with Components of Phenolic Wastewaters—A Summary. Monograph 44. Wyd. Uczln. Politech., Szczecin, Poland. 127 pp.
• Glabisz, U. 1967. Action of chlorine dioxide on monohydric phenols. Chem. Tech. (Berlin) 19:352-355.
• Gordon, G., and F. Feldman. 1964. Stoichiometry of the reaction between uranium (IV) and chlorite. J. Inorg. Chem. 3:1728-1733.
• Gordon, G., R.G. Kieffer, and D.H. Rosenblatt. 1972. The chemistry of chlorine dioxide. Pp. 201-286 in S.J. Kippard, editor. , ed. Progress in Inorganic Chemistry, Vol. 15. John Wiley & Sons, Inc., New York.
• Granstrom, M.L., and G.F. Lee. 1957. Rates and mechanisms of reactions involving oxychloro compounds. Public Works 88:90-92.
• Granstrom, M.L., and G.F. Lee. 1958. Generation and use of chlorine dioxide in water treatment. J. Am. Water Works Assoc. 50:1453-1466.
• Hodgen, H.W., and R.S. Ingols. 1954. Direct colorimetric method for the determination of chlorine dioxide in water. Anal. Chem. 26:1224-1226.
• Jeanes, A., and H.S. Isbell. 1941. Chemical reactions of the chlorites with carbohydrates. J. Res. Natl. Bur. Stand. 27:125-142.
• Kennaugh, J. 1957. Action of diaphanol on arthropod cuticles. Nature 180:238.
• Kieffer, R.G., and G. Gordon. 1968. Inorg. Chem. 7:235-238, 239-244.
• Leopold, B., and D.B. Mutton. 1959. The effect of chlorinating and oxidizing agents on derivatives of oleic acid. Tappi 42:218-225.
• Lindgren, B.O., and B. Ericsson. 1969. Reaction of chlorine dioxide with phenols: formation of α,b-epoxy ketones from mesitol and 2,6-xylenol. Acta Chem. Scand. 23:3451-3460.
• Lindgren, B.O., and T. Nilsson. 1972. Lignin reactions during chlorine dioxide bleaching of pulp. Oxidation by chlorite. Sven. Papperstidn. 75:161-168.
• Lindgren, B.O., and C.M. Svahn. 1966. Reactions of chlorine dioxide with unsaturated compounds. II. Methyl oleate. Acta Chem. Scand. 20:211-218.
• Lindgren, B.O., C.M. Svahn, and G. Widmark. 1965. Chlorine dioxide oxidation of cyclohexene. Acta Chem. Scand. 19:7-13.
• Love, O.T., Jr., J.K. Carswell, R.J. Miltner, and J.M. Symons. 1976. Treatment for the prevention of removal of trihalomethanes in drinking water. In J.M. Symons. Interim Treatment Guide for the Control of Chloroform and Other Trihalomethanes. Water Supply Research Division, Municipal Environmental Research Laboratory, U.S. Environmental Protection Agency, Cincinnati, Ohio: (App. 3).
• Mallevialle, J. 1976. Ozonation des substances de type humique dans les eaux. Pp. 262-270 in R.G. Rice, editor; , P. Pichet, editor; , and M.A. Vincent, editor. , eds. Proceedings of the 2nd International Conference on Ozone Technology. International Ozone Institute, Cleveland, Ohio.
• Masschelein, W. 1967. Development in the chemistry of chlorine dioxide and its applications. Chim. Ind. Genie Chim. 97:49-61.
• Masschelein, W. 1969. Les oxydes de chlore et le chlorite de sodium. Monogr. Dunod 74:16-57.
• Miller, G.W., R.G. Rice, C.M. Robson, W. Kuhn, and H. Wolf. 1978. An assessment of ozone and chlorine dioxide technologies for treatment of municipal water supplies. Pp. 9-57 to 9-89 in Report of EPA Grant R804385-01. Municipal Environmental Research Laboratory, Office of Water Supply, U.S. Environmental Protection Agency, Cincinnati, Ohio.
• Miltner, R.J. 1977. Measurement of chlorine dioxide and related products. In Proceedings, American Water Works Association Water Quality Technology Conference, San Diego, Calif., Dec. 6-7, 1976. Paper No. 2A-5. American Water Works Association, Denver, Colo.
• Otto, J., and K. Paluch. 1965. Reactions of chlorine dioxide with some organic compounds. V. Reaction of benzaldehyde with chlorine dioxide. Roczniki Chem. 39:1711-1712.
• Paluch, K. 1964. The reaction of chlorine dioxide with phenols. I. Phenol and chlorophenols. II. Hydroquinone, chloro derivatives of hydroquinone, and nitrophenols. Rocznicki Chem. 38:35-42, 43-46.
• Paluch, K., J. Otto, and K. Kozlowski. 1965. Reaction of chlorine dioxide with some organic compounds. VI. Reaction of benzyl alcohol with chlorine dioxide and with acidified sodium chlorite solution. Rocznicki Chem. 39:1603-1608.
• Reichert, J.K. 1968. a. Kanzerogene Substanzen in Wasser und Boden. XXI. Die Entfernung polyzyklischer Aromaten bei der Trinkwasser-Aufbereitung durch Chlordioxid: Quantitative Befunde. Arch. Hyg. 152:37-44. [PubMed: 5707367]
• Reichert, J.K. 1968. b. Kanzerogene Substanzen in Wasser und Boden. XXIII. Die Entfernung polyzyklischer Aromaten bei der Trinkwasseraufbereitung durch Isolierung und Identifizierung der 3,4-Benzpyrenfolgeprodukte. Arch. Hyg. 152:265-276. [PubMed: 5711174]
• Robson, H.L. 1964. Pp. 35-50 in H.F. Mark, editor; , J.J. McKetta, Jr., editor; , and D.F. Othmer, editor. , eds. Kirk-Othmer Encyclopedia of Chemical Technology. Vol. 5, 2nd ed. Interscience Publishers, New York.
• Rosenblatt, D.H. 1975. Chlorine and oxychlorine species reactivity with organic substances. Pp. 249-276 in J.D. Johnson, editor. , ed. Disinfection: Water and Wastewater. Ann Arbor Science Publishers, Inc., Ann Arbor, Mich. 425 pp.
• Rosenblatt, D.H. 1978. Chlorine dioxide: chemical and physical properties. Pp 332-343 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Sarkanen, K.V., K. Kakehi, R.A. Murphy, and H. White. 1962. Studies on oxidative delignification mechanisms. Part I. Oxidation of vanillin with chlorine dioxide. Tappi 45:24-29.
• Sarkar, P.B. 1935. Chemistry of jute lignin. VII. Behaviour of organic compounds towards chlorine dioxide and its significance on the constitution of lignin. J. Indian Chem. Soc. 12:470-482.
• Schmidt, E., and K. Braunsdorf. 1922. Natural proteins. I. Behavior of chlorine dioxide towards organic compounds. Ber. 55B: 1529-1534.
• Somsen, R.A. 1960. Oxidation of some simple organic molecules with aqueous chlorine dioxide solutions. I. Kinetics. II. Reaction products. Tappi 43:154-156, 157-160.
• Spinks, J.W.T., and J.M. Porter. 1934. Photodecomposition of chlorine dioxide. J. Am. Chem. Soc. 56:264-270.
• Stevens, A.A., D.R. Seeger, and C.J. Slocum. 1978. Products of chlorine dioxide treatment of organic materials in water in ozone/chlorine oxidation products of organic materials. Pp. 383-399 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Sussman, S., and J.S. Rauh. 1978. Use of chlorine dioxide in water and wastewater treatment. Pp. 344-355 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the US. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Symons, J.F., J.K. Carswell, R.M. Clark, P. Dorsey, E.E. Geldreich, W.P. Heffernam, J.C. Hoff, O.T. Love, L.J. McCabe, and A.A. Stevens. 1977. Ozone, Chlorine Dioxide, and Chloramines as Alternatives to Chlorine for Disinfection of Drinking Water: State of the Art. U.S. Environmental Protection Agency, Water Supply Research Division, Cincinnati, Ohio. 84 pp.
• Taylor, M.C, J.F. White, G.P. Vincent, and G.L. Cunningham 1940. Sodium chlorite, properties and reactions. Ind. Eng. Chem. 32:899-903.
• Thielemann, H. 1972. Uber die Einwirkung von Chlordioxid auf einige polycyklische aromatische Kohlenwasserstoffe. Mikrochim. Acta 575-577.
• Toussaint, M. 1972. Chlorine dioxide in drinking water treatment. Trib. CEBEDEAU 25(342):260-266.
• Vilagenes, R., A. Monteil, A. Derremaux, and M. Lambert. 1977. A comparative study of halomethane formation during drinking water treatment by chlorine or its derivatives in a slow and sand filtration plant and in wastewater treatment plants. Paper presented at 96th Annual Conference, American Water Works Association, Anaheim, Calif., May 8, 1977.
• White, J.F., M.C. Taylor, and G.P. Vincent. 1942. The chemistry of chlorites. Ind. Eng. Chem. 34:782-792.

Ozone Reactions and Products

• Ahmed, M., and C.R. Kinney. 1950. Ozonation of humic acids prepared from oxidized bituminous coal. J. Am. Chem. Soc. 72:559-561.
• Bailey, P.S. 1975. Reactivity of ozone with various organic functional groups important to water purification. Pp. 101-119 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Bauch, H., and H. Burchard. 1970. Untersuchungen uber die Einwirkung von Ozon auf Wasser mit geringen Verunreinigungen. Wasser Luft Beitr. 14:270-273.
• Bauch, H., H. Burchard, and H.M. Arsovic. 1970. Ozone as an oxidant for phenol degradation in aqueous solutions. Gesund. Ing. 91(9):258-262.
• Bollyky, L.J. 1975. Ozone treatment of cyanide and plating wastes. Pp. 522-532 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Briner, E. 1959. Photochemical production of ozone. Pp. 1-6 in Ozone Chemistry and Technology. Advances in Chemistry Series No. 21. American Chemical Society, Washington, D.C.
• Brody, S.S. 1975. A proposed new analysis for ozone in water using a field portable chemiluminescent ozone analyzer. Pp. 84-92 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Carlson, R.M., and R. Caple. 1977. Chemical/Biological Implications of Using Chlorine and Ozone for Disinfection. U.S. Environmental Protection Agency, Environmental Research Laboratory, Duluth, Minn. Report No. EPA/600/3-77/066. 99 pp.
• Cerkinsky, S.N., and N. Trahtman. 1972. The present status of research on the disinfection of drinking water in the USSR. Bull. WHO 46:277-283. [PMC free article: PMC2480699] [PubMed: 4537489]
• Chian, E.S.K., and P.P.K. Kuo. 1976. Fundamental study on the post-treatment of RO permeates from army wastewater. Second Annual Summary Report, Report No. UILU-ENG-76-2019. U.S. Army Medical R & D Command, Washington, D.C.
• Criegee, R. 1959. Products of ozonization of some olefins. Pp. 133-135 in Ozone Chemistry and Technology. Advances in Chemistry Series No. 21. American Chemical Society, Washington, D.C.
• Dobinson, F. 1959. Ozonization of malonic acid in aqueous solution. Chem. Ind. Lond. 26:853-854.
• Eisenhauer, H.R. 1968. The ozonization of phenolic wastes. J. Water Pollut. Control Fed. 40:1887-1899.
• Falk, H.L., and J.E. Moyer. 1978. Ozone as a disinfectant of water. Pp. 38-58 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Press International, Cleveland, Ohio. 487 pp.
• Garrison, R.L., C.E. Mauk, and H.W. Prengle, Jr. 1975. Advanced ozone-oxidation system for complexed cyanides. Pp. 551-577 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Gilbert, E. 1976. Ozonolysis of chlorophenols and maleic acid in aqueous solution. Pp. 253-261 in R.G. Rice, editor; , P. Pichet, editor; , and M.A. Vincent, editor. , eds. Proceedings of the Second International Symposium on Ozone Technology, Montreal, Canada, May 11-14, 1975. Ozone Press International, Jamesville, N.Y. 725 pp.
• Gilbert, E. 1978. Reactions of ozone with organic compounds in dilute aqueous solution: identification of their oxidation products. Pp. 227-242 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Gould, J.P., and W.J. Weber, Jr. 1976. Oxidation of phenols by ozone. J. Water Pollut. Control Fed. 48:47-60.
• Gunther, F.A., D.E. Ott, and M. Ittig. 1970. The oxidation of parathion to paraoxon. II. By use of ozone. Bull. Environ. Contam. Toxicol. 5:87-94. [PubMed: 24185733]
• Helz, G.R., R.Y. Hsu, and R.M. Block. 1978. Bromoform production by oxidative biocides in marine waters. Pp. 68-76 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Hoigne, J., and H. Bader. 1975. Ozonation of water: role of hydroxyl radicals as oxidizing intermediates. Science 190:782-784.
• Hoigne, J., and H. Bader. 1976. Identification and kinetic properties of the oxidizing decomposition products of ozone in water and its impact on water purification. Pp. 271-282 in R.G. Rice, editor; , P. Pichet, editor; , and M.A. Vincent, editor. , eds. Proceedings of the Second International Symposium on Ozone Technology, Montreal, Canada, May 11-16, 1975. Ozone Press International, Jamesville, N.Y. 725 pp.
• Hoigne, J., and H. Bader, 1977. Rate constants for the reactions of ozone and organic pollutants and ammonia in water. Symposium on Advanced Ozone Technology. International Ozone Institute, Toronto, Ontario, Canada.
• Hoigne, J., and H. Bader. 1978. a. Ozone initiated oxidations of solutes in wastewater. A reaction kinetic approach. Paper presented at International Conference on Water Pollution, Stockholm, Sweden.
• Hoigne, J., and H. Bader. 1978. b. Ozonation of water: kinetics of oxidation of ammonia by ozone and hydroxyl radicals. Environ. Sci. Technol. 12:79-84.
• Huibers, D.T.A., R. McNabney, and A. Halfon. 1969. Ozone treatment of secondary effluents from wastewater treatment plants. Federal Water Pollution Control Administration, U.S. Department of the Interior, Cincinnati, Ohio. Robert A. Taft Water Research Center Report No. TWRC-4. 62 pp.
• Ingols, R.S. 1978. Ozonation of seawater. Pp. 77-81 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Jurs, R.H. 1966. Die Wirkung des Ozons auf in Wasser geloste Stoffe. Fortschr. Wasserchem. Ihrer Grenzgeb., Heft 4:40-64.
• Kilpatrick, M.L., C.C. Herrick, and M. Kilpatrick. 1956. The decomposition of ozone in aqueous solution. J. Am. Chem. Soc. 78:1784-1789.
• Kinman, R.N., J. Rickabaugh, V. Elia, K. McGinnis, T. Cody, S. Clark, and R. Christian. 1978. Effect of ozone on hospital wastewater cytotoxicity. Pp. 97-114 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Kinney, C.R., and L.D. Friedman. 1952. Ozonization studies of coal constitution. J. Am. Chem. Soc. 74:57-61.
• Kjos, D.J., R.R. Furgason, and L.L. Edwards. 1975. Ozone treatment of potable water to remove iron and manganese: preliminary pilot plant results and economic evaluation. Pp. 194-203 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Klein, M.J., R.I. Brabets, and L.C. Kinney. 1975. Generation of ozone. Pp. 1-9 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Kuo, P.P.K., E.S.K. Chian, and B.J. Chang. 1977. Identification of end products resulting from ozonization and chlorination of organic compounds commonly found in water. Environ. Sci. Technol. 11:1177-1181.
• Lawrence, J. 1977. Identification of ozonization products in natural waters. Presented at the Symposium on Advanced Ozone Technology. International Ozone Institute, Toronto, Ontario, Canada.
• Lawrence, J., and F.P. Cappelli. 1977. Ozone in drinking water treatment: a review. Sci. Total Environ. 7:99-108. [PubMed: 841303]
• Maggiolo, A. 1978. Ozone's radical and ionic mechanisms of reaction with organic compounds in water. Pp. 59-67 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Mallevialle, J. 1975. Action de l'ozone dans la degradation des composes phenoliques simples et polymerises: Application aux matieres humiques contenues dans les eaux. Tech. Sci. Munic. Revue l'Eau 70:107-113.
• Mallevialle, J., Y. Laval, M. Lefebvre, and C. Rousseau. 1978. The degradation of humic substances in water by various oxidation agents (ozone, chlorine, chlorine dioxide). Pp. 189-199 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held at Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Manley, T.C., and S.J. Niegowski. 1967. Ozone. Pp. 410-432 in Kirk-Othmer Encyclopedia of Chemical Technology, Vol. 14, 2nd ed. Interscience Publishers, New York.
• Mathieu, G.I. 1975. Application of film layer purifying chamber process to cyanide destruction—a progress report. Pp. 533-550 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Miller, G.W., and R.G. Rice. 1978. Testimony before House Science and Technology Committee, Subcommittee on the Environment and the Atmosphere (transcript by private communication).
• Murray, R.W. 1968. Ozone chemistry of organic compounds. Pp. 1-4 in F.R. Mayo, editor. , ed. Oxidation of Organic Compounds. III. Ozone Chemistry Photo and Singlet Oxygen and Biochemical Oxidation. Advances in Chemistry Series No. 77. American Chemical Society, Washington, D.C.
• Netzer, A., and A. Bowers. 1975. Removal of trace metals from wastewater by lime and ozonation. Pp. 731-747 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Oehlschlaeger, H.F. 1978. Reactions of ozone with organic compounds. Pp. 20-37 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Rice, R.G. 1977. Reaction products of organic materials with ozone and with chlorine dioxide in water. Presented at International Ozone Institute Symposium on Advanced Ozone Technology, Toronto, Ontario, Canada.
• Rice, R.G., and J.A. Cotruvo. 1978. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Rice, R.G., C. Gomella, and G.W. Miller. 1978. Rouen, France water treatment plant: Good organics and ammonia removal with no need to regenerate carbon beds. Civil Eng. ASCE (N.Y.) 48(5):76-82.
• Richard, Y., and L. Brener. 1978. Organic materials produced upon ozonization of water. Pp. 169-188 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 17-19, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Rogozhkin, G.I., T. Vsesayyzn, and I. Nauchno. 1970. Inta vodosnabzheniya kanalizatsii. Pp. 27-45 in Gidrotekhnicheskikh Sooruzheniy i Gidrogeologii.
• Schalekamp, M. 1978. Experiences in Switzerland with ozone particularly in connection with the charge of undesirable elements present in water. Ozone Technology Symposium and Exposition. International Ozone Institute, Los Angeles, Calif.
• Selm, R.P. 1959. Ozone oxidation of aqueous cyanide waste solutions in stirred batch reactors and packed towers. Pp. 66-77 in Ozone Chemistry and Technology. Advances in Chemistry Series No. 21. American Chemical Society, Washington, D.C.
• Shevchenko, M.A., and P.N. Taran. 1966. Investigation of the ozonolysis products of humus materials. Sov. Prog. Chem. 32:408-410. [Trans. of Ukr. Khim Zn. 32(5):532536.]
• Sievers, R.R., R.M. Barkley, G.Z. Eiceman, R.H. Shapiro, H.F. Walton, K.J. Kolonko, and L.R. Field. 1977. a. Environmental trace analysis of organics in water by glass capillary column chromatography and ancillary techniques. J. Chromatogr. 142:745-754. [PubMed: 914947]
• Sievers, R.E., R.M. Barkley, G.A. Eiceman, L.P. Haack, R.H. Shapiro, and H.F. Walton. 1977. b. Generation of volatile organic compounds from non-volatile precursors in water by treatment with chlorine or ozone. Presented at Conference on Water Chlorination, Gatlinsburg, Tenn.
• Sievers, R.E., R.H. Shapiro, H.F. Walton, G.A. Eiceman, and R.M. Barkley. 1977. c. High resolution gas chromatographic determination of organic compounds in ozonized wastewater. Presented at ACS National Meeting, Chicago, Ill.
• Singer, P.C., and W.B. Zilli. 1975. Ozonation of ammonia: application to wastewater treatment. Pp. 269-287 in R.G. Rice, editor; and M.E. Browning, editor. , eds. First International Symposium on Ozone for Water and Wastewater Treatment. International Ozone Institute, Waterbury, Conn.
• Spanggord, R.J., and V.J. McClurg. 1978. Ozone methods and ozone chemistry of selected organics in water. 1. Basic chemistry. Pp. 115-125 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 15-17, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio. 487 pp.
• Stumm, W. 1954. Der Zerfall von Ozon in wassriger Losung. Helv. Chim. Acta 37:773-778.
• Sturrock, M.G., E.L. Cline, and K.R. Robinson. 1963. The ozonation of phenanthrene with water as participating solvent. J. Org. Chem. 28:2340-2343.
• Symons, J., J.K. Carswell, R.M. Clark, P. Dorsey, E.E. Geldreich, W.P. Heffernam, J.C. Hoff, O.T. Love, L.J. McCabe, and A.A. Stevens. 1977. Ozone, Chlorine Dioxide and Chloramine as Alternatives to Chlorine for Disinfection of Water: State of the Art. Water Supply Research Division, U.S. Environmental Protection Agency, Cincinnati, Ohio. 84 pp.
• Weil, L., B. Sttruif, and K.E. Quentin. 1977. Reaction mechanisms upon reaction of organic substances in water with ozone. Presented at the International Symposium on Ozone and Water, Berlin. International Ozone Institute, Cleveland, Ohio.
• Westgate Research Corp. 1978. The continued investigation into the chemistry of the UV-ozone water purification process. Westgate Research Corporation.
• Wynn, C.S., B.S. Kirk, and R. McNabney. 1973. Pilot plant for tertiary treatment of wastewater with ozone. U.S. Environmental Protection Agency, Washington, D.C. No. EPA-R2-73-146. 229 pp.
• Yocum, F.H. 1978. Oxidation of styrene with ozone in aqueous solution. Pp. 243-263 in R.G. Rice, editor; and J.A. Cotruvo, editor. , eds. Ozone/Chlorine Dioxide Oxidation Products of Organic Materials. Proceedings of a Conference held in Cincinnati, Ohio, November 15-17, 1976. Sponsored by the International Ozone Institute and the U.S. Environmental Protection Agency. Ozone Press International, Cleveland, Ohio.
• Yokoyama, K., S. Sato, I. Yoshiyasu, and T. Imamura. 1974. Degradation of organic substances in water, by ozone. Mitsubishi Denki Giho Tech. Rev. 48:1233-1238.