WATER CHLORINATION
`Chemistry,
`Environmental Impact
`and Health Effects
`Volume 5
`
`Edited by
`Robert L. Jolley
`Richard J. Bull
`William P. Davis
`Sidney Katz
`Morris H. Roberts, Jr.
`Vivian A. Jacobs
`
`Proceedings of the .Fifth Conference on
`Water Chlorination: Environmental Impact and Health Effects
`Williamsburg, Virginia
`June 3-8, 1984
`
`Sponsored by
`Electric Power Research Institute, National Cancer Institute,
`Oak Ridge National Laboratory, Tennessee Valley Authority,
`U.S. Department of Energy, U.S. Environmental Protection Agency,
`Virginia Institute of Marine Science- School of Marine Science, and
`College of William and Mary
`
`- - LEWIS PUBLISHERS, INC.
`
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`

`

`7
`
`\_:>l\ ~ .t
`. (
`"'
`.
`'-\
`Library of Congress Cataloging-in-Publication Data
`Main entry under title:
`
`\
`
`Water chlorination. Volume 5.
`
`Papers presented at the fifth Conference on Water
`Chlorination: Environmental Impact and Health Effects,
`held at the College of William and Mary, Williamsburg,
`Virginia, June 3-8, 1984.
`Includes bibliographies and index.
`1. Water- Purification - Chlorination - Environmental
`aspects. 2. Water-Purification-Chlorination-Hygienic
`aspects. 3. Water chemistry.
`I. Jolley, Robert L.
`II. Conference on Water Chlorination: Environmental
`Impact and Health Effects (5th : 1984 : College of
`William and Mary)
`
`1985
`TD462.W38
`ISBN 0-87371-005-3
`
`363.6'1
`
`85-18122
`
`COPYRIGHT © 1985 by LEWIS PUBLISHERS, INC.
`ALL RIGHTS RESERVED
`
`Neither this book nor any part may be reproduced or transmitted in
`any form or by any means, electronic or mechanical, including
`photocopying, microfilming, and recording, or by any information
`storage and retrieval system, without permission in writing from the
`publisher.
`
`LEWIS PUBLISHERS, INC.
`121 South Main Street, P.O. Drawer 519, Chelsea, Michigan 48118
`
`PRINTED IN THE UNITED STATES OF AMERICA
`
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`

`

`This material may be protected by Copyright law (Title 17 U.S. Code)
`
`CHAPTER 62
`
`The Chemistry of Oxo-Chlorine Compounds
`Relevant to Chlorine Dioxide Generation
`
`E. Marco Aieta and Paul V. Roberts
`
`Within the last decade, halogenated organic compound formed in drinking
`water by disinfection with chlorine have caused considerable concern through(cid:173)
`out the drinking water industry. 1 The U.S. Environmental Protection Agency
`(EP ) has regulated the concentration of THM to a maximum of l 00 µ,g/L
`total THM in drinking water a a fir t tep in an effort to control the chlori(cid:173)
`nated organic compounds formed during the di infection process. Three strat(cid:173)
`egie
`for meeting this regulation are removal of the organic precursor ;
`removal of the THM that are formed; or substitution of an alternative disin(cid:173)
`fectant that doe not form THM .
`One of the alternative di infectants that has received considerable attention
`is chlorine dioxide (Cl02), which offers several attributes that make it an
`acceptable ub titute for chlorine. 2 It does not form THMs. It does, however,
`produce ome chlorinated organic compounds in aqueous solution, but under
`condition encountered in drinking water treatment the concentration of
`the e compound are IO to 100 time Jes
`than the concentration produced by
`chlorine under the same condition .3 Chlorine dioxide is an effective disinfec(cid:173)
`tant over a broad pH range and in ome cases ignificantly more effective than
`chlorine. 4
`hlorine dioxide can be mea ured easily at the point of u e of the
`disinfected water, permitting re idual disinfectant measurements to be u ed a
`an indicator of the microbiological afety of the delivered water.
`Chlorine dioxide is an un table explosive ga at concentrations greater than
`about 10% in air. 5 The gas, either pure or in mixture , cannot be compressed
`tored; hence, Cl02 must be generated on- ite for immediate u e. For
`and
`drinking water treatment, CI02 is most commonly generated by the oxidation
`of odium chlorite by chlorine.
`Thi chapter reviews the chemistry of oxo-chlorine compounds o that the
`reactions u ed to generate Cl02 for water treatment may be more fully under-
`tood. Thi chapter also includes results from our re earch on the kinetic of
`the oxidation of . odium chlorite by chlorine. This reaction i the preferred
`Cl02 generation reaction for water treatment, becau e of the high yields that
`can be achie ed and because there are no waste streams that po e dispo al
`problem .
`
`783
`
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`

`

`784
`
`WATER CHLORINATION
`
`AQUEOUS CHEMISTRY OF CHLORINE AND
`OXO-CHLORINE COMPOUNDS
`
`Table I lists the major chlorine compounds of the oxo-chlorine family and
`their oxidation states.
`
`Chlorine Hydrolysis Reaction
`
`Chlorine gas is moderately soluble in water. In addition to solvated halogen
`molecules, hypochlorous acid and chloride ion are present in aqueous solution
`due to the hydrolysis/ disproportionation reaction of the solvated molecular
`chlorine. These reactions are
`
`Cli(g) = Cli(aq)
`Clz(aq) + H20 = HOC! + CJ- + H +
`
`The chlorine hydrolysis constant is given by
`
`[HOCI] [H +] [CI-]
`[Cl2(aq)]
`
`(1)
`
`(2)
`
`(3)
`
`The rate of chlorine hydrolysis has been widely studied. The interpretation
`of the results from early studies was complicated because of confusion about
`the mechanism of the chlorine hydrolysis reaction. Shilov and So1odushenkov6
`assumed the reaction to be as shown by Equation (2). Their results, however,
`indicated that the rate constant decreased as the reaction proceeded. Morris7
`recalculated the results of Shilov and Solodushenkov, 6 assuming the reaction
`involved the hydroxide ion
`
`HOCl + Cl-
`
`(4)
`
`Table I. Chlorine and Oxo-Chlorine Species
`
`Oxidation
`State
`
`Oxidation
`No.
`
`Compounds
`
`Name
`
`Chlorine ( -1)
`Chlorine (0)
`Chlorine (I)
`Chlorine (Ill)
`Chlorine (IV)
`Chlorine (V)
`Chlorine (Vil)
`
`- 1
`0
`1
`3
`4
`5
`7
`
`HCI, c1 -
`Cl2
`HOCI, oc1-
`HClO2, c1O - 2
`ClO2
`HCIO3, C1O - 3
`HClO4 , CIO - 4
`
`Hydrochloric acid, chloride
`Molecular chlorine
`Hypochlorous acid, hypochlorite ion
`Chlorous acid,chlorite ion
`Chlorine dioxide
`Chloric acid, chlorate ion
`Perchloric acid, perchlorate ion
`
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`

`REACTION DYNAMICS
`
`785
`
`and found much better constancy in the calculated second-order rate constant.
`The value of the second-order rate constant determined by Morris was 5 x
`10 14 M -1 s- 1, an extremely high value. The Morris7 calculations also indicated
`zero activation energy for the chlorine hydrolysis reaction. A reinvestigation
`of the chlorine hydrolysis reaction by Shilov and Solodushenkov8 did not
`reproduce the decreasing rate constant observed in their first study, so that
`Morris' argument did not seem to apply. Lifshitz and Perlmutter-Hayman9
`tried to establish which of Equations (2) or (4) represented the hydrolysis in
`pure water. Their findings supported the original assumption of Shilov and
`Solodushenkov6 that Equation (2) correctly represented the chlorine hydrolysis
`reaction in pure water. Furthermore, Lifshitz and Perlmutter-Hayman9
`pointed out that the maximum second-order rate constant to be expected in
`solution corresponded to the diffusion-controlled limit of l x 1010 M I s- 1, a
`much lower value than that calculated by Morris,7 and that the rate of forma(cid:173)
`tion of hydroxide ions was not rapid enough in acid solution for Equation (4)
`to contribute significantly to the rate of chlorine hydrolysis.
`Morris'7 ideas relating to the mechanism of chlorine hydrolysis were not
`totally in error, however, as later demonstrated by Spalding. 10 The equilibrium
`pH of a chlorine solution in pure, unbuffered water at a given temperature is
`controlled by the total chlorine concentration of the system, as given by the
`chlorine hydrolysis equilibrium constant, Equation (3). In the studies of Shilov
`and Solodushenkov6,8 and of Lifshitz and Perlmutter-Hayman,9 chlorine gas
`was dissolved in pure, unbuffered water so that the pH of the reaction solution
`was near pH 2 and may have been somewhat lower at the lower temperatures
`of these studies. Eigen and Kustin II proposed a general mechanism for the
`hydrolysis of halogens, including chlorine. Both Equations (2) and (4) are
`included in the general mechanisms. Eigen and Kustin's 11 results indicated that
`at pH 2.2, the highest pH they investigated, the hydrolysis mechanism repre(cid:173)
`sented by Equation (4) did not contribute to the rate of chlorine hydrolysis.
`Spalding, 10 however, absorbed chlorine gas into water at initial pH values
`between 3 and 10.2. His results indicated that below an initial pH of 10.2, the
`primary reaction mechanism for the hydrolysis of chlorine was that given in
`Equation (2); however, above pH 12.5, the reaction mechanism was that pro(cid:173)
`posed by Morris,7 which is given in Equation (4). It should be emphasized that
`the actual pH value reported by Spalding, 10 at which Equation (2) became
`dominant, was the initial pH of the absorbing fluid and that the contact time
`was very short. Under different experimental conditions this pH value might
`be shifted. Spalding10 estimated the first-order chlorine forward rate constant
`for Equation (2) to be 20.9 s I at 25°C. His estimate for the second-order rate
`constant of Equation (4) was 1 x 106 M I s 1 at 25 °C. More rec~nt work by
`Sandal et al., 12 who studied the absorption of chlorine gas in strong sodium
`hydroxide solution, gave a value for the second-order rate constant of Equa(cid:173)
`tion (4) as 2.7 x 107 M 1 s-1 at 0°c.
`Figure I shows a comparison of the first-order forward rate constants for
`the chlorine hydrolysis reaction determined in our laboratory with the results
`
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`

`786
`
`WATER CHLORINATION
`
`-...
`I • -
`
`50
`40
`
`30
`
`20
`
`19
`8
`7
`6
`5
`4
`
`3
`
`+
`
`◊
`
`+
`+
`
`D
`
`+ Spaldlnc (1982)
`◊ Brillll et al. (1988)
`
`0 Eicen lllld Kuatln (1982)
`+ Perlmutter-Haymllll et al. (1973)
`X W■hits lllld Perlmutter-Hayman (1980)
`
`):( Praent Study
`
`X
`
`+
`
`+
`2 ~__._____,__.........._..._~___.____,__........__..._~___.____.___._......._~----....___.___._......._.......__..___._~
`3.2
`3.3
`3.4
`3.5
`3.1
`3.6
`1000 / T
`
`Figure 1. Chlorine hydrolysis rate constant as a function of temperature.
`
`from the work of others. A rather wide range of values for the chlorine
`hydrolysis rate constant is reported at 20°C. The kinetic methods used include
`temperature jump relaxation 11 •13 and stopped flow techniques, 9• 13 as well as
`gas-liquid mass transfer techniques. 10 • 14•15 If the data in Figure 1 are analyzed
`according to the Arrhenius rate law by the regression of ln(k) vs l / T, the
`activation energy is found to be 60.3 kJ/ mol, with a 95% confidence interval
`of 49.4 to 71.2 kJ/ mol. The preexponential factor is 6.76 x 10 11 and the
`coefficient of determination, r2 , is 0.94. The value of the chlorine hydrolysis
`first-order forward rate constant predicted from all the data is 12.2 s- 1 at
`20°C, with a 95% confidence interval of 10.7 to 14.0 s- 1•
`
`Aqueous Reactions of Hypochlorite,
`Chlorate, and Chlorite
`
`Hypochlorite ion tends to disproportionate to form chlorate ion and chlo(cid:173)
`ride ion,
`
`3 OCl = ClO3 + 2 CJ-
`
`(5)
`
`This reaction is slow at or below room temperature, but in hot (e.g., 80°C)
`solutions, the reaction is rapid and produces high yields of chlorate ion. Equa-
`
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`

`REACTION DYNAMICS
`
`787
`
`tion (5) forms the basis of the industrial production of chlorate. Chlorine
`dioxide is produced for the pulp and paper industry by the reduction of chlo(cid:173)
`rate in strong acid medium. Either hydrochloric or sulfuric acid may be used.
`The most common reducing agents are chloride, sulfur dioxide, and methanol.
`The reduction of chlorate by chloride in acid medium is given by
`
`There is a competing parallel reaction that produces only chlorine,
`
`(7)
`
`Commercially, chlorites are produced from the selective reduction of ClO2 .
`The ClO2 used in the manufacture of chlorites is produced as described above
`and is reduced in alkaline solution in the presence of a weak reducing agent to
`form the metal chlorite.
`Chlorites are used commercially as bleaching agents and for small-scale
`CIO2 production for water treatment. In alkaline solution, chlorite ion is very
`stable, even at 100°C. In acid solution, chlorous acid rapidly disproportion(cid:173)
`ates to form CIO2, chlorate ion, and chloride ion,
`
`In the presence of appreciable amounts of chloride ion, only small amounts of
`chlorate are formed and the stoichiometry is approximated by
`
`(9)
`
`Equation (9) forms the basis for some commercial, small-scale ClO2 genera(cid:173)
`tion systems. These systems find only limited application in water treatment
`systems because of the availability of more efficient ClO2 generation systems
`that use the oxidation of chlorites by chlorine. 16
`
`Chlorine - Chlorite Reaction and Mechanism
`
`The pioneering work of Taube and Dodgen 17 using radioactive chlorine
`provided significant insight into the reaction mechanisms of the oxidation of
`chlorite by chlorine (or hypochlorous acid), the reduction of chlorate by chlo(cid:173)
`ride ion (and the reverse reaction), and the disproportionation ·of chlorous
`acid. The observations of Taube and Dodgen 17 led them to postulate an
`unsymmetrical activated complex that was common to all three reaction mech(cid:173)
`anisms. Subsequent work by Emmenegger and Gordon 18 has amplified and
`refined the mechanism proposed by Taube and Dodgen, 17 but their original
`
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`

`

`788
`
`WATER CHLORINATION
`
`findings are still the foundation for many of the subsequent studies of the
`kinetics and mechanisms of aqueous chlorine chemistry.
`In the oxidation of aqueous chlorite by chlorine or hypochlorous acid, both
`Cl02 and chlorate appear as products. In acid solution, where the chlorine is
`present mainly as dissolved molecular chlorine, the stoichiometries of the two
`reactions are
`
`and
`
`2CI02 + 2CI-
`
`(10)
`
`(11)
`
`In solutions near neutral pH, where chlorine is present largely as hypochlorous
`acid, the stoichiometries are
`
`HOCI + 2CI0-2
`
`2Cl02 + CI- + OH-
`
`and
`
`(12)
`
`(13)
`
`In alkaline solutions in which the chlorine is present as hypochlorite ion, the
`reaction is very slow and the only product formed is chlorate ion:
`
`oc1- + c10-2 = c10-3 + c1-
`
`(14)
`
`The product ratio of CI02 to chlorate has been observed to vary as a func(cid:173)
`tion of the experimental conditions such that:
`
`I . The reaction between chlorite ion (or chlorous acid) and chlorine (in acid or neutral
`solution) is second order overall, being first order in both chlorine and chlorite.
`2. Acidic pH values favor the formation of CIO2, whereas at neutral and alkaline pH,
`chlorate is the primary reaction product.
`3. For a given pH value, an increase in chlorite concentration results in relatively more ClO2
`production.
`4. Higher chloride ion concentrations favor the formation of CIO2 relative to chlorate,
`especially at acidic pH.
`5. Proportional increases in all reactant concentrations favor the formation of CIO2 relative
`to chlorate.
`
`The mechanism for the chlorine-chlorite reaction proposed by Taube and
`Dodgen 17 as consistent with experimental observations is presented below. The
`starred (*) chlorine atoms trace the fate of the chlorine originally present as
`molecular chlorine or hypochlorous acid. Hypothetical, metastable, interme(cid:173)
`diate species are enclosed in brackets { } . For the reaction of chlorite with
`molecular chlorine,
`
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`

`, REACTION DYNAMICS
`
`789
`
`(15)
`
`(16)
`
`(17)
`
`and with hypochlorous acid
`
`For the production of CIO2,
`
`For the production of chlorate,
`
`This mechanism satisfies the experimental observations discussed above
`according to the following rationale. The metastable intermediate, { Cl2O2},
`can decompose by either a first-order process, Equation (18), to give chlorate
`ion or by a second-order process, Equation (17), to give CIO2 • The formation
`of the intermediate, { Cl2O2}, is, in either case, the rate-limiting step in the
`mechanism as determined from the observed rate law . 18 Reaction conditions
`that promote the formation of higher concentrations of the intermediate will
`favor the product ClO2 via the second-order route over the product chlorate
`ion via the first-order route. Emmenegger and Gordon 18 have shown that the
`rate of oxidation of chlorite by molecular chlorine, Equation (15), is consider(cid:173)
`ably faster than the oxidation by hypochlorous acid, Equation (16). Therefore,
`reaction conditions that favor molecular chlorine as opposed to hypochlorous
`acid will result in higher intermediate concentrations and relatively more C1O2•
`In aqueous systems, higher hydrogen ion concentrations and chloride ion
`concentrations shift the chlorine hydrolysis equilibrium, Equation (3), toward
`increased molecular chlorine concentration. These reaction conditions are, in
`fact, those that are observed to produce relatively more ClO2 both experimen(cid:173)
`tally and in practice. An excess of chlorite also promotes higher concentrations
`of the intermediate, as does increasing the absolute concentrations of all reac(cid:173)
`tants, while maintaining the same reactant ratios, which results in proportion(cid:173)
`ally more ClO2•
`In our research, chlorine gas was contacted with aqueous sodium chlorite
`solution in a gas-liquid system. As shown by Emmenegger and Gordon, 18 the
`rate of reaction of the dissolved molecular chlorine gas with the chlorite ion is
`much faster than the rate of chlorine hydrolysis. Hence, the chlorine molecule
`reacted preferentially with the chlorite ion, and the reaction solution pH did
`not influence the rate of reaction in this gas-liquid system as it had been shown
`to do in a liquid-liquid system.
`The details of this kinetic apparatus and the methodology have been pre(cid:173)
`sented elsewhere. 15 The results of these experiments are presented in Table II.
`
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`

`

`0 z
`~
`:!! z
`0
`r-
`:I:
`n
`lJ
`~ m
`=:
`
`0
`co
`....,
`
`END-H/1308
`
`dCV = Coefficient of variation = standard deviation/mean.
`<Cl = Confidence interval.
`bN = Number of runs.
`aExperiment = experiment identification number.
`
`14.9
`24.3
`26.6
`21.5
`20.2
`16.5
`
`(%)
`CVd
`
`1.55 -1.79
`0.98 -1.33
`0.49 -0.61
`1.50 -1.88
`1.19 -1.79
`1.23 -1.40
`
`(M-1 s-1)
`X 10-4
`95% Cle
`
`0.06
`0.08
`0.03
`0.09
`0.12
`0.04
`
`0.25
`0.28
`0.15
`0.36
`0.30
`0.22
`
`1.67
`1.16
`0.55
`1.69
`1.49
`1.31
`
`Error, x 10-4
`
`(M-1 s-1)
`
`Standard
`
`Deviation, x 10-4
`
`(M-1 s-1)
`
`Standard
`
`Constant, x 10-4
`
`(M-1 s-1)
`
`Rate
`
`30
`20
`10
`20
`20
`20
`
`(°C)
`
`T
`
`1.42
`4.26
`1.39
`0.17
`0.12
`1.42
`
`(M)
`I
`
`19
`12
`25
`16
`6
`28
`
`Nb
`
`6
`5
`4
`3
`2
`1
`
`Experiment•
`
`Table II. Summary of Experimentally Determined Second-Order Rate Constants of the Chlorine -Chlorite Reaction
`
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`

`REACTION DYNAMICS
`
`791
`
`Effect of Ionic Strength
`
`As discussed by Moore and Pearson 19 for ionic strengths >0.01 M, the
`logarithm of the reaction rate constant should be a linear function of the first
`power of ionic strength for reactions involving a neutral molecule and an ion.
`The chlorine-chlorite reaction falls into this category.
`Experiments 1, 3, and 5 reported in Table II were all performed at 20°C, but
`the absorbing liquids were of different ionic strengths. A regression analysis of
`the natural logarithms of the second-order rate constants, k2, of the chlorine(cid:173)
`chlorite reaction from these experiments vs ionic strength yielded the relation(cid:173)
`ship:
`
`In ( ~ ) = -0.0851
`k2,o
`
`(19)
`
`where k2,0 is the second-order rate constant determined by extrapolation to
`zero ionic strength and is 1.62 x 104 M-1s-1• The 95% confidence interval for
`the zero ionic strength second-order rate constant is 1.49 x I 04 to 1. 76 x 104
`M- 1s- 1• The 95% confidence interval for the slope in Equation (19) is -0.046 to
`-0.124. This relationship is shown in Figure 2 together with the means of the
`experimental data and the respective standard deviations.
`
`2.0
`
`1.0
`0.9
`
`0 .8
`
`0 .7
`
`0.6
`
`0.5
`
`0
`
`Ill
`~
`
`3
`2
`1
`Ionic Strength ( mol / 1 )
`
`4
`
`Figure 2. The effect of ionic strength on the second-order rate constant of the chlorine(cid:173)
`chlorite reaction.
`
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`

`

`792
`
`WATER CHLORINATION
`
`Effect of Temperature
`
`The rate of the chlorine-chlorite reaction increases with increasing tempera(cid:173)
`ture. The temperature dependence can be described by the Arrhenius relation(cid:173)
`ship. 19 The data from Experiments 1, 4, and 6 shown in Table II for the
`second-order rate constant of the chlorine-chlorite reaction at 20, 10, and
`30°C, respectively, were used to determine the parameters of the Arrhenius
`expression. Correction of the second-order rate constant determined from
`Experiment 4 at ionic strength of 1. 39 to the ionic strength of Experiments 1
`and 6 resulted in a decrease in the value given in Table II of only 0.03%. A
`regression analysis of In (rate) vs 1/T yielded
`
`In k2 = 25.6 - 4766 (1/T)
`
`or
`
`k2 = 1.31 x 1011
`
`• exp
`
`l -39.9 kJ · moI-1 l
`
`RT
`
`(20)
`
`(21)
`
`Equation (21) is plotted in Figure 3 along with the means of the experimental
`data points and their respective standard deviations. The 95 % confidence
`interval for the activation energy is 35.6 to 45.2 kJ mol.
`
`20000
`
`10000
`9000
`8000
`7000
`
`6000
`
`5000
`
`4000
`
`--
`-I
`Ill') .
`
`1
`
`)I
`..._,
`DI
`.!lit
`
`3000
`3.25
`
`3.3
`
`3.35
`
`3.4
`1000 / T
`Figure 3. The effect of temperature on the second-order rate constant of the chlorine-chlorite
`reaction.
`
`3.45
`
`3.5
`
`3.55
`
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`

`REACTION DYNAMICS
`
`793
`
`SUMMARY
`
`A review of oxo-chlorine chemistry and results from recent experimental
`investigations have been presented. Relevant to ClO2 generation for water
`treatment is the observation that ClO2 yield can be optimized by selecting
`reaction conditions such that high concentrations of molecular chlorine are in
`contact with high concentrations of sodium chlorite.
`
`ACKNOWLEDGMENT
`
`This work was funded by the U.S. Environmental Protection Agency under
`Research Grant R-808686. This paper has not been subjected to the Agency's
`required administrative review, and therefore does not necessarily reflect the
`views of the Agency, nor should any official endorsement be inferred.
`
`REFERENCES
`
`1. Bellar, T. A., J. J. Lichtenberg, and R. D. Kroner. "The Occurrence of Organoha(cid:173)
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`WATER CHLORINATION
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`14. Brian, P. L. T., J.E. Vivian, and C. Piazza. "The Effect of Temperature on the
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