Environ. Sci. Technol. 1989, 23, 1412- 1419
Registry No. SO2, 7446-09-5; HN03, 7697-37-2; HN02, 7782-77-6; NHS, 7664-41-7; NH4+, 14798-03-9;H+, 12586-59-3. L i t e r a t u r e Cited Schwartz, S. E. Science 1989, 243, 753-163. Schindler, D. W. Science 1988, 239, 149-157. Lippmann, M. EHP, Environ. Health Perspect. 1985,63, 63-70. Lippmann, M. EHP Environ. Health Perspect. 1989, 79, 3-6. Bates, D. V.; Sizto, R. In Aerosols: Research, Risk Assessment and Control Strategies; Lee, Z., Schneider, Z., Grant, Z., Eds.; Lewis Publishers: Chelsea, MI, 1987; pp 761-771. Bates, D. V.; Sizto, R. EHP, Environ. Health Perspect. 1989, 79, 69-72. Speizer, F. E. EHP, Environ. Health Perspect. 1989, 79, 61-68. Spengler, J. D.; Sexton, K. Science 1983, 221, 9-16. Spengler, J. D.; Soczek, M. Enuiron. Sci. Technol. 1984, 18, 268A-280A. Dockery, D. W.; Spengler, J. D. Atmos. Environ. 1981,15, 335-343. Sinclair, J. D.; Psota-Kelty, L. A,; Weschler, C. J. Atmos. Environ. 1988, 22, 461-469. Koutrakis, P. et al., Atmos. Enuiron. (in preparation).
Koutrakis, P. et al. Environ. Sci. Technol. 1988, 22, 1463-1468. Slater, J. L., Brauer, M., Koutrakis, P., and Keeler, G. J. In Proc. of the 1988 EPAIAPCA Symposium on Measurement of Toxic and Related Air Pollutants. pp. 176-181. Brauer, M.; Koutrakis, P.; Wolfson, J. M.; Spengler, J. D. Atmos. Enuiron. 1989,23, 1981-1986. Koutrakis, P.; Wolfson, J. M.; Spengler, J. D. Atmos. Environ. 1988, 22, 157-162. Allegrini, I. et al. Sci. Total Environ. 1987, 67, 1-16. Pitts, J. N.; Wallington, T. J.; Biermann, H. W.; Winer, A. M. Atmos. Environ. 1985, 19, 763-767. Biermann, H. W.; Pitts, J. N., Jr.; Winer, A. M. In Advances in Air Sampling; ACGIH, Lewis Publishers: Chelsea, MI, 1988; pp 265-289. Brauer, M., unpublished observations. Nishimura, H.; Hayamizu, T.; Yanagisawa, Y. Environ. Sci. Technol. 1986,20, 413-416. Jenkin, M. E.; Cox, R. A.; Williams, D. J. Atmos. Environ. 1988,22, 487-498. Larson, T. V.; Covert, D. S.; Frank, R.; Charlson, R. J. Science 1977, 197, 161-163. Received for review March 20, 1989. Accepted July 17, 1989. Supported by EPA Cooperative Agreement CR-812667-02-2 and NIEHS Training Grant ES07155.
Destruction of Chlorination Byproducts with Sulfite Jean-Phlllppe Crouet and David A. Reckhow Environmental Engineering Program, Department of Civil Engineering, University of Massachusetts, Amherst, Massachusetts 0 1003
The observation that sulfite can destroy mutagenic activity in chlorinated waters has important implications with respect to the use of S(IV) in water treatment and sample preservation. The purpose of this study was to evaluate the reaction of sodium sulfite with specific organohalides formed during the chlorination of drinking water. In the first phase, chlorinated fulvic acid solutions were analyzed by closed-loop stripping and GC/MS. Compounds susceptible to decomposition were identified by treating some solutions with sulfite and some without. Phase two included a series of kinetic experiments using pure solutions of chloropicrin, trichloroacetonitrile, dichloroacetonitrile, dibromoacetonitrile, l,l,l-trichloropropanone, chloral, 1,l-dichloropropanone, 2,3,6-trichloroanisole, and 3-chloro-4-(dichloromethyl)-5hydroxy-2(5H)-furanone(MX). The reactions were found to be first order in the compound and first order in sulfite (specifically, [Sot-]). Each of the reactive compounds gave reduced products (loss of halogen) at rates that suggest the use of sulfite is a feasible means of controlling selected chlorination byproducts in drinking water treatment. Introduction
It is well-known that the chlorination of humic substances leads to the formation of volatile and nonvolatile halogenated compounds. In general, these compounds are the same as those observed after chlorination of natural surface water or drinking water. These compounds are also believed to be responsible for much of the mutagenic act Present address: University of Poitiers, France; at the time of this work, Dr. Croue was an employee of the Compagnie General des Eaux.
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Environ. Sci. Technol., Vol. 23, No. 11, 1989
tivity found in finished drinking waters. Recent work suggests that many chlorination byproducts are destroyed by reaction with sodium sulfite, a reducing agent added to remove residual chlorine before analysis (1-3). Knowledge of the reactions of commonly used reducing agents with disinfection byproducts is critical for the proper development of analytical procedures. Because the U S . Environmental Protection Agency (EPA) is considering the adoption of more restrictive chlorination byproduct standards, there is an urgent need for new or improved control technologies. Most approaches to chlorination byproduct control fall under one of three categories: (A) removal of organic precursors, (B) minimization of chlorine contact, and (C)removal of byproducts. The first two have been responsible for most of the progress to date in reducing byproduct levels. The direct removal of chlorination byproducts has been less widely practiced, because of poor performance using current technologies (e.g., coagulation of low molecular weight byproducts) or high costs associated with newer technologies (e.g., activated carbon adsorption). The research presented here has implications with respect to a new and inexpensive process for byproduct removal, chemical reduction with S(IV) species. Since sulfur dioxide and its aqueous forms are used for dechlorination in drinking water systems, the application of these compounds represents a currently available and acceptable technology. Background
Hazardous Chlorination Byproducts. Most of the chlorination byproducts that are being considered for regulation by the U.S.EPA and state agencies are among the group of haloorganic mutagens or suspected carcinogens [e.g., halonitriles, haloketones ( 4 , 5 ) ] . Researchers at the US.EPA Health Effects Laboratory in Cincinnati
0013-936X/89/0923-1412$01.50/0
0 1989 American Chemical Society
have been actively pursuing the identity of these compounds. One very useful approach has been to use the Ames mutagenicity test in conjunction with GC/MS analysis to determine specific compounds responsible for the mutagenic activity in chlorinated natural waters. Despite extensive studies of this type, these researchers have until recently only been able to account for 7 4 % of the overall mutagenicity in chlorinated humic acid solutions (6). The discovery, in 1986, of low concentrations of an extremely potent mutagen, 3-chloro-4-(dichloromethyl)-5-hydroxy-2(5H)-furanone(MX), in chlorinated humic acid solutions and drinking waters has significantly advanced this effort (7). In a survey of six chlorinated drinking waters, MX alone was found to account for 22-44% of the overall mutagenicity (8). Nevertheless, more than 50% of the mutagenicity remains unaccounted for, and the importance of strong direct-acting mutagens in human carcinogenesis is still uncertain. Nucleophilic Dehalogenation and Mutagenicity. It is now well established that the mutagenic activity in chlorinated natural waters (9,10) and humic acid solutions (11) can be rapidly destroyed by alkaline treatment. Significant decomposition is seen as low as pH 4 (IO),and the decomposition increases with increasing pH so that at pH 11.5 complete elimination of mutagenic activity may occur within 7 days (11). Specifically,the compounds MX, 2-chloropropenal, 1,3-dichloropropanone, and 1,1,3,3tetrachloropropanone have all been observed to undergo base-catalyzed decomposition with half-lives on the order of minutes to days depending on the pH (5,12,13). The latter two chloroketones undergo loss of chlorine and loss of mutagenic activity (14 ) upon decomposition, whereas pentachloropropanone decomposes without the breaking of any C-X bonds [chloroform and dichloroacetic acid are produced (15)]. Since most ultimate carcinogens are electrophiles (16), one might expect a variety of nucleophiles, not just hydroxide, to be effective at destroying these compounds. Cheh and co-workers (1)found that a small excess of sodium sulfite over that needed to quench the residual chlorine could lead to a 4040% reduction in the mutagenic activity (Ames test) of chlorinated natural waters (pH 7.5-8.7). Wilcox and Denny ( I 7) qualitatively confirmed this observation. However, in experiments using actual raw waters, they noted that partial dechlorination was only marginally successful at destroying mutagens. This indicates that the reaction between sulfite and residual chlorine is fast and that chlorine may out-compete the organic mutagens for sulfite. Wilcox and Denny (17)also rechlorinated water samples that were treated with sulfite, and retested them for Ames mutagenicity. They found that only about 0-50% of the original mutagenicity was regenerated upon rechlorination. As for the effects of pH, Donnini (18)found that S(IV) (i.e., sulfite, bisulfite, and aqueous sulfur dioxide) was most effective at destroying mutagenicity of pulp effluents under alkaline conditions. He concluded that SOptreatment was the most economical method for reducing environmental impacts of chlorinated pulp wastewaters. Although there is only a small amount of information on sulfite's effects on drinking water mutagenicity, there was even fewer data on the specific compounds responsible for these effects. Perhaps the most significant studies were those of Trehy and Bieber (2) and Fam (3),which showed that dihaloacetonitriles and a variety of unsaturated alkyl halides can readily undergo dehalogenation reactions with sodium sulfite. Chemical Basis for Nucleophilic Dehalogenation. Due to the high electronegativities of the halogens, car-
bon-halogen bonds tend to be susceptible to nucleophilic attack at the carbon atom, often giving simple substitution products. These reactions may be strongly affected by steric hindrance. As the a- and 8-carbons become substituted with bulky groups, the reaction rate may drop off sharply. Also, the reactions are fastest for iodo compounds and slowest for chloro compounds, since the heavier halogens are better leaving groups. For many compounds, these types of reactions occur under environmental conditions at significant rates (19). One can make two very important generalizations regarding nucleophilicity. First, for a given row on the periodic table, the stronger the base, the stronger the nucleophile. Second, bases formed from elements in the first row are invariably less nucleophilic than second-row compounds of the same basicity (20). Thus, one would expect that moderately basic sulfur and phosphorus compounds (e.g., sulfite, triethylphosphine) to be as nucleophilic as highly basic compounds containing only oxygen, nitrogen, or carbon (e.g., hydroxide, diethylamine). Indeed, sulfite readily displaces halides in what is called the Strecker reaction (21). This SN2reaction occurs for primary and secondary alkyl halides, benzyl halides, and a wide range of chloro acids and chloro ketones.
Materials and Methods The objective of this study was to evaluate the reaction of sodium sulfite with organohalides formed during the chlorination of dilute solutions of aquatic fulvic acid. These experiments were performed at neutral pHs and in the presence or absence of bromide. The first phase of experiments was designed to investigate the overall effect of sulfite on chlorination byproducts and to identify those volatile and semivolatile organohalogen products that were especially reactive with sodium sulfite. The second phase was designed to determine the kinetics of these reactions for selected chlorination byproducts. Analytical Techniques. The isolation and extraction techniques used included closed-loop stripping and liquid/liquid extraction. The compounds were analyzed by capillary gas chromatography (DB-5,30M) with electron capture detection and a low thermal mass on-column injector. Compound identifications were made with a Hewlett-Packard 5985 quadrupole GC/MS using both electron impact and chemical ionization (with methane). All model compounds quantifications were made from pentane extractions of the reaction solution with an internal standard. In addition, analysis of MX required further concentration and derivitization (22). MX samples were first extracted with ethyl acetate and concentrated by rotary evaporation to dryness. Then, they were reconstituted in methanol with 2% sulfuric acid and methylated by heating the liquid to 70 "C for 1h. Finally, this solution was extracted with hexane following neutralization with sodium bicarbonate. Chlorination of Humic Materials. For the phase I experiments, buffered solutions of a commercial humic acid and an aquatic fulvic acid were prepared both with and without added bromide (0.4 mg of Br-/mg of DOC). The aquatic fulvic acid was extracted from Thousand Acre Reservoir, an emergency water supply for the town of Athol, MA. A commonly used hydrophobic resin extraction procedure was employed (23). By inclusion of commercial humic acid in the experimental design, high concentrations of DOC could be used so that high-quality maw spectra could be obtained. Other researchers have shown that commercial (e.g., Aldrich, Fulka) humic acid and natural aquatic humic material give similar chlorination products (I3). Environ. Scl. Technol., Vol. 23, No. 11, 1989
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The fulvic solutions (7.5 mg/L DOC) were chlorinated at a chlorine to carbon ratio of 0.1 mol/mol. This low dose resulted in the complete loss of residual halogen within the 2-h contact period. As a result, it was not necessary to add a reducing agent prior to closed-loop stripping analysis. The humic solutions were chlorinated at a chlorine to carbon ratio of 0.4 mol/mol. This dose resulted in the complete loss of residual halogen within 48 h. After this point, both sets of solutions were either extraded directly or extracted after an addition of sulfite (235 pM for 1h, in the case of fulvic acid). The solutions were subjected to closed-loop stripping and on-column capillary GC with electron capture detection. The humic acid solutions were also analyzed by electron impact and chemical ionization GC/MS. Model Compound Studies. The phase 2 experiments included carefully controlled laboratory studies for the purpose of determining the effect of reaction time, pH, and sulfite concentration on the extent of reaction of sulfite with selected chlorination byproducts. The compounds chosen for this work were chloropicrin, trichloroacetonitrile, dichloroacetonitrile, dibromoacetonitrile, l,l,ltrichloropropanone, chloral, 1,l-dichloropropanone, 2,3,6trichloroanisole, and MX. All are mutagens that are produced during drinking water chlorination. Most are currently being considered (either as a single compound or as part of a group) for regulation by the U.S. EPA. Dilute aqueous solutions were prepared for each from the pure compound and buffered with phosphate (15 mM). Experimental conditions for all but the MX studies were as follows: pH 6.1-8.5, 5-100 pg/L model compound starting concentration, 25 pM sulfite concentration, 0-4-h reaction time, 20 "C in the absence of light. Experiments with MX were conducted with starting concentrations of 5 pg/L, a sulfite dose of 100 pM, and reaction times of up to 48 h.
Results and Discussion Experiments with Humic Materials. Figure 1shows two chromatograms for chlorinated Thousand Acre fulvic acid without added bromide. Figure l a represents the sample that was not treated with sulfite. Figure l b presents the chromatogram from the same treated with sulfite. Many of the compounds identified are shown in these figures. Those that have been confirmed (with an actual standard) are underlined. Others are considered tentative, although in many cases, their spectra have been matched to those published by Coleman and co-workers (24).Certain peaks are significantly reduced in size by the sulfite treatment and those are marked by an asterisk. Note that trichloroacetonitrile, chloropicrin, 3,3-dichloropropenal, dichloropropenenitrile, and trichloropropenenitrile are among the compounds that appear to be lost as result of the sulfite treatment. However, due to the nature of the concentration technique, the phase I studies must be viewed as being only semiquantitative. In the presence of bromide (0.4 mg of Br-/mg of C) many new compounds are formed (Figure 2a). Figure 2b shows a chromatogram representing the sample treated with sulfite. Again, those compounds that are lost to a significant extent by reaction with sulfite are indicated by an asterisk (Figure 2a). Among these compounds are bromochloroacetonitrile, dibromoacetonitrile, bromochloropropanone, bromodichloropropanone,and bromodichloropropenenitrile. Note also the formation of tetrahalomethanes in the presence of bromide. This was a surprise, because these compounds are not commonly reported as chlorination byproduds. However, Snoeyink and co-workers (25) observed the formation of carbon tet,rs1414
Environ. Sci. Technol., Vol. 23, No. 11, 1989
-"
m
I " h
L
I
10
5
20
30
40
50
60
TIME (Mlfi)
5
10
20
30
40
53
60
TIME (MIN) Flgure 1. GC/ECD chromatogam of chkrkrated Thousand Acre Mvic acid without added bromide (0.1 mol of Cl,/mol of DOC, 2-h reaction time). Analytical conditions: on-column injection: 30-m DE5 capillary column; oven, 30-40 OC at 20 OC/min, hold for 10 min, 40-90 OC at 2 OC/min, 90-230 OC at 5 OC/min. (a) Without sulfite; (b) 235 pM Na,S03, 1-h contact time.
chloride from the reaction of free chlorine with activated carbon. In the present study, it was hypothesized that a small bromine residual may have existed at the time of stripping and this residual could have reacted with the carbon adsorbant to produce tetrahalomethanes. Finally, it should be pointed out that since closed-loop stripping was used to extract these compounds, the chlorination products in Figures 1 and 2 are all volatile. Many halo acids were certainly produced in these experiments; however, the techniques used did not permit their detection, nor could their susceptibility to degradation by sulfite be assessed. Experiments with Model Organic Compounds. A. Chloropicrin. Solutions of chloropicrin were prepared (100 pg/L) in a phosphate buffer (15 mM) and treated with small concentrations of sodium sulfite (0-25 pM). The concentration of chloropicrin and the relative peak height of a major degradation product were followed over time. This product was later identified as dichloronitromethane on the basis of a comparison of its mass spectrum with that published by Coleman and co-workers (26). Therefore, it is presumed that sultite leads to a reduction of chloropicrin as shown in reaction 1. nso3-/so8~ CCl3N02 CHClzNOz
Table I. Pseudo-First-OrderReaction Rate Constants for the Decomposition of Six Chlorination Byproducts at Three Different pHs
compound CClaN02 CClaCN CHClzCN CHBrzCN CClaCOCHa MXd
kht(l lo4 s-l pH 7.2 12.8 f 3 17.4 f 0.9 3.07 i 0.15 -9 i 4
pH 6.1 10 f 3 6 f 1.4 C
-6i4 C C
C
pH 8.5 9.5 174 f 14 4.21 f 0.18 -2 f 5 73 i 18
0.9 f 0.5
C
k,,b lo4 s-l pH 7.2
pH 6.1 269 i 7 111 f 4
C
pH 8.5 1950 f 125 950 i 66 0.5 i 0.3 1310 f 90 0 f 20
22 f 3
C
1410 f 40 620 f 40 0 i 0.2 1081 f 7
C
127 i 5 C C
Rate constants for hydrolysis. *Rate constants for reaction with 25 pM [S(IV)]. e Constanta not determined. Neutral MX values were determined at pH 7.0 rather than at pH 7.2,sulfite concentration waa 100 NM. 1.5
-.f
-
4
1.0
X
2
0.5
3 a3 3
A! 0.1 10
15
20
IO
40
50
60
20
40
60
TIME
(MINI
80
100
120
FlQure 9. Reaction of chloropicrinwith sulfite: loss of chloropialn and formation of dichloronitromethane.
-1 .o
0 PH = 0 pH = V PH =
4
3
8,s 7,2
6,l
-3.0 5
10
15
20
TIM
30
40
50
60
(HIN)
Figure 2. QC/ECD chromatogams of chbrinated thousand Acre fuMc acM In the presence of 0.4 mg of W/mg of DOC (0.1 mol of Cl,/mol of DOC, 241 reactlon Ume). Analytical conditkns: orrcolumn Injection; 30-m D&5 caplllary column; oven, 30-40 OC at 20 OC/mIn, hoM for 10 mln, 40-90 OC at 2 OC/min, 90-230 OC at 5 Oc/mIn. (a) Without sulfite; (b) 235 pM Na,SO,, 1-h contact time.
Figure 3 shows the concentration of chloropicrin as a function of time at three different pHs. In addition, the formation curve for dichloronitromethane is shown in units of relative peak area. The exact concentration of dichloronitromethane could not be determined, because a standard was not available. Figure 3 indicates that the rate of decomposition of chloropicrin increases with increasing pH. Also note that the pH 7.2 curve for dichloronitromethane reaches a maximum peak area ratio much below that of pH 6.1. This suggests that the dichloronitromethane is also subject to decomposition under alkaline conditions. The concentration of dichloronitromethane at pH 8.5 was near the detection limit (not shown), which supports this hypothesis. In the absence of sulfite, only a slight loss of chloropicrin was observed (less than 10% after 3 h, regardless of pH). The hydrolysis rate is, therefore, insignificant under these
10
50
90
TIME (MINI
130
Flgure 4. Reactlon of chloropicrin with sulfite: semllogarlthmlc plot at three pHs.
conditions. Figure 4 is a semilogarithmic plot of the chloropicrin data in Figure 3. The fact that these data form a straight line suggests a first order reaction of the following type. - (d[Cl/dt) = kob[Cl (2) The slopes of the lines represent the pseudo-first-order reaction rate constants, kob Table I presents a summary of the pseudo-first-order reaction rate constants as determined for solutions without sulfite (hydrolysis) and solutions with 25 p M sulfite (reduction). The hydrolysis constants, kh, are simply equal to the slopes of the semilogarithmic plots of concentration vs time, kob, for experiments conducted in the absence of sulfite. The sulfite/bisulfite reaction constants, k,,are equal to the slopes of the semilogarithmic plots of concentration vs time, k+, in the presence of 25 p M sulfite minus the hydrolysis constants determined at the same pH. (3) k, = kob - k h Environ. Sci. Technol., Vol. 23, No. 11, 1989
1415
-t
-7 10
'
1
w m v
-
I
-
-3.0
10
50
30
70
90
0
TIME (tm) Flgure 5. Reaction of chloropicrin with sulfite: semilogarithmic plot at three sulfite doses.
9 3 12
1
2
4
p] 7'01
a
6
(lm~L)
Figure 7. Reaction of chloropicrin with sulfite: determination of k , and k,.
V
,
3
4 x
"
"
,
9 0 PH = 0 pH =
2o 10 5
10
15 (UI,WLJL)
[Sd
20
25
Flgure 6. Reactlon of chloropicrin with sulfite: reaction of sulfite dose with pseudo-firstorder reaction rate constant (pH 7.2).
To test for the reaction rate dependence on total sulfite, identical solutions of chloropicrin were buffered at pH 7.2 and treated with different concentrations of sulfite. The decay curves are shown in Figure 5 on a semilogarithmic scale. As the concentration of total sulfite increases, the rate of decomposition increases. Figure 6 shows the slopes plotted as a function of total sulfite concentration, [S(IV)]. The fact that a straight line that passes through the origin was obtained indicates that the decomposition of chloropicrin is first order in total sulfite. The increase in k, for chloropicrin (Table I) with increasing pH may reflect changes in the speciation of S(IV), or some other type of base catalysis. Since sulfite is known to be a more powerful nucleophile than bisulfite, it is proposed that this pH effect is due to the shift in the sulfite species with pH. Thus, k, might be expressed as the s u m of the rate of reaction with sulfite and the rate of reaction with bisulfite (eq 4).
8,5 7,2
I I
I
I
10
30
60
I
TIME(MINI
90
I
120
'
1
300
Figure 8. Hydrolysis of trichloroacetonitrile: effect of pH.
ref 28). The uncertainties are calculated based on the standard deviations for k, in Table I, a standard deviation of 0.03 for pKa and 0.05 for pH. Note that the lack of a significant slope indicates that kl is very small compared to kz. This is expected since sulfite is a much stronger nucleophile than bisulfite. For example, the ratio of k2 to klfor nucleophilic attack on aldehydes (carbonyl)has been observed to be in the range of 1W105 (29). The intercept from Figure 7 indicates that k2 is -85 M-l si.As a result of the insignificant kl,a plot of k, as a function of [SOS2-] was prepared (not shown), and it gave a straight line (r2 = 0.984) that passed near the origin. When forced through the origin it had a slope of -85 M-' s-l. Thus, one can represent the decomposition of chloropicrin in the presence of sulfite at 20 "C by eq 7, where k2 is -85 M-' s-l.
- (d[C]/dt)
= k2[S032-][C]
(7)
B. Trichloroacetonitrile. In contrast to chloropicrin, trichloroacetonitrile shows a significant rate of hydrolysis (4) in the absence of sulfite. The final degradation produch were not identified; however, research by Exner and coEquation 4 can be rewritten as workers (30) suggests that the alkaline degradation of k, = (k2a2 + k,a,)[S(IV)I (5) haloacetonitriles gives haloacetamides (and ultimately haloacetic acids). Figure 8 indicates that this reaction is where [S(IV)] is the molar sum of the sulfite species, base catalyzed. The pseudo-first-order reaction rate conprimarily [S032-]and [HS03-]. Substituting for the a's stants are presented in Table I. In the presence of sulfite, in eq 5 and rearranging, one obtains an accelerated decomposition is observed with the forks(Ka + [H+l)/[S(IV)l = kl[H+I + k2Ka (6) mation of dichloroacetonitrile (DCAN). Equation 8 summarizes these pathways. Figure 9 shows the raw data for Figure 7 shows a plot of k,(K, + [H+])/[S(IV)Ivs [H+]. By use of the extended DebyeHuckel correction (27) for HSOa-fS08" OHCC13CN CHC12CN CC13CONH2 ionic strength, the apparent pKa for bisulfite deprotonation CC1,COOH (8) is calculated to be 7.03 (pKa = 7.18 f 0.03 at I = 0, 25 "C;
k, = k2[S032-] + kl[HS03-]
-
1416
Environ. Sci. Technol., Vol. 23, No. 11, 1989
-
b ffl = 7,2
0
0
A W I W SULFITE AWITH SULFITE
h
A
50-
I
I
0
n
TW ACCOUNTED FQR
120 200 TIME(MIN) Flguro 9. Reaction of trichloroacetonitrile with sulfite at pH 6.1. 30
10
60
0
0 W I W SULFITE 0 WITH SULFITE
-
1
I
I
10
20
I
50
1
60
TIME(HOURS) Flgure 11. Decomposition of dichloroacetonitrile with and without sulfite. 0.5
-1
1
40
1
30
0
ffl = 8,s
0 ffl = 7,2
v PH = 6,l 0 BROWACETONITRILE
0.4
ffl = 8,5
0.3 0.2 \
~n
10
50
100
1
150
I
200
0.1
TIME(MINI Fi(pm 10. Reactkn of tr-ichkroacetonitrila with sulfite: semilogarithmic plot at three pHs.
trichloroacetonitrile (TCAN) and dichloroacetonitrile (DCAN) at pH 6.1. The squares represent the amount of starting compound accounted for by the pathways in eq 8. This is calculated from the sum of TCAN remaining, DCAN formed, and cumulative TCAN hydrolyzed, which is based on kh and the TCAN decay curve. Note that this analysis allows one to account for nearly all of the initial TCAN throughout the reaction period. The same is true for the data collected at pH 7.2 and 8.5 (not shown). The semilogarithmic plots of the TCAN data are combined in Figure 10. As with chloropicrin, the reaction rate constants calculated from Figures 8 and 10 (see Table I) increase with increasing pH in the same way that sulfite concentration increases with pH. A plot of k, vs [SOB2-] gives a straight line (r2 = 0.998) that passes through the origin and has a slope of 41 M-I (not shown). Thus, eq 7 is the suggested rate law with k2 = 41 M-I s-l. In order to verify that TCAN, once decomposed, would not re-form from its primary degradation product, an additional experiment was performed. Pure solutions of dichloroacetonitrile (100 pg/L) were chlorinated and analyzed over a period of 8 h. The absence of TCAN in the quenched solutions indicated that TCAN could not be formed by this route under conditions typical of drinking water treatment. C. Dichloroacetonitrile. Dichloroacetonitrile is known to undergo a base-catalyzed decomposition (2). Although hydrolysis rates noted in this study were small, the value at pH 8.5 was significntly higher than the one determined for pH 7.2 (see Table I). These constants are similar to the constant one would calculate from the data of Bieber and Trehy (31)at pH 8.32 and 25 "C (i-e.,kh = 6.4 X lo4 5-l). No significant additional decomposition could be detected in the presence of 25 pM sulfite (Figure 11). This is in agreement with data from Trehy and Bieber (2).
10
60
TIME(MINI
I
I
120
180
Figure 12. Reaction of dibromoacetonitrile with sulfite at three pHs.
D. Dibromoacetonitrile. Dibromoacetonitrile undergoes a very slow base-catalyzed hydrolysis in water. From the data of Trehy and Bieber (2) one calculates a pseudo-first-order reaction rate of 2.3 X lV sw1at pH 8.32 and 25 "C. Data for Exner and co-workers (30) give 0.4 X lo4 s-I at pH 7.4 and 10.3 X lo4 s-l at pH 9.0, all at 20 "C. Trehy and Bieber (2) found that dibromoacetonitrile was rapidly destroyed by sulfite; however, their data do not permit the calculation of a reaction rate. In this study it was found that dibromoacetonitrile gives bromoacetonitrile quantitatively in the presence of sulfite (Figure 12). Figure 12 shows the dibromoacetonitrile concentrations at three pHs as a function of time, along with one bromoacetonitrile formation curve as an example. Note that the reaction rate constants increase with increasing pH in the same fashion as the divalent sulfite concentration (see Table I). Thus, the reaction may be characterized by eq 7 with k2 = 54 M-l s-l. E. l,l,l-Trichloropropanone.This chlorination byproduct shows a significant decomposition rate at pH 8.5, however, no additional loss could be detected at this pH in the presence of sulfite. The second-order basecatalyzed hydrolysis constant calculated from this rate is 23 f 6 M-' s-l. This is similar to the constant determined by Reckhow and Singer (15)of 18 f 3 M-' s-' (20"C, 14 mM phosphate buffer), and that of Gurol and co-workers (32)of 36 f 6 (25 "C, 200 mM phosphate buffer). It is possible that a significant amount of the a-hydroxy sulfonate derivative of trichloropropanone was formed at the start of the reaction and remained at equilibrium throughout. One might expect this to occur because of the electron-withdrawing nature of the CC1, group (33). This might not have affected the trichloroactone quantification, if a rapid Environ. Sci. Technoi., Vol. 23, No. 11, 1989
1417
h
B
loo
k-4
19
WITHObT SULFITE A THIS STUDY . A KRONBERG (1987)
I
I
I
10
20
TIME
u
WITH SULFITE
gLiBz1987,
I
30
I
40
I
50
(HCURS)
Flgure 13. Decomposition of MX wRh and without sulfffe at pH 7.0.
loss of bisulfite had occurred either during extraction or in the GC injector. Experiments run at pH 6.1 and 7.2 in the presence of sulfite showed only a very slight loss of trichloropropanone. It was presumed that this was due to the slow hydrolysis expected at these pHs. F. MX. Solutions of MX (3-chloro-4-(dichloromethyl)-B-hydroxy-2(5H)-furanone) were prepared from the pure compound at concentrations of 5 pg/L. The mass spectra of this compound were matched with published spectra (22). The decomposition of this compound at pH 7 in the presence [lo0 pM initial S(1V) concentration] and absence of sulfite (Figure 13) was found to be similar to that reported by Holmbom and Kronberg (34). This is true despite the fact that the starting MX concentrations were 5 pg/L for this study as opposed to 20 OOO pg/L for that of Holmbom and Kronberg (34). A semilogarithmic plot of these data shows a great deal of curvature at long reaction times. This may have been due to the slow loss of sulfite with time. At the end of 48 h the analytical sulfite concentration was determined to be -40 pM by iodometry (35).Assuming a pseudo-first-order reaction, one can calculate an initial rate with the first four data pointa (Table I). Note that this constant is much smaller than those determined for chloropicrin, trichloroacetonitrile, and dibromoacetonitrile. Similar data collected at pH 6 show little or no hydrolysis and a reduced rate of loss due to sulfite (not shown). The low reactivity of MX is surprising in light of the work of Cheh and co-workers ( I ) , and the importance of MX to the overall mutagenicity of chlorinated water. Recall that these authors found reductions of mutagenicity of 40440% from dechlorination with an excess of a few micromoles of sulfite at pH 7.5-8.7. This apparent anomaly suggests that compounds other than MX are responsible for the drop in mutagenicity as water is dechlorinated. G. Other Compounds. In addition to those discussed above, chloral, 2,3,6-trichloroanisole, and 1,l-dichloropropanone were tested for decomposition in the presence of sulfite. None showed an effect of sulfite within the conditions used with this research. The fact that chloral was unaffected by sulfite is especially interesting, since this compound is expected to react very rapidly with S(1V) to form an a-hydroxy sulfonate derivative. If this reaction occurred to the extent proposed by Betterton and coworkers (33),the loss of sulfite or bisulfite must have had predominance over ail other subsequent reactions. The fact that all of the chloral could be recovered after reaction with sulfite may be explained by a rapid equilibrium between the aldehyde and the a-hydroxy sulfonate. Chemical Interpretations. In determining the rate of reaction of a gem polyhalide with a given nucleophile 1418
such as sulfite, three characteristics of the parent molecule may be important: the number of halogens, the type of halogen, and the electron-withdrawing tendency of the neighboring groups. Regarding the type of halogen, Table I shows a much greater rate of attack of sulfite with dibromoacetonitrile than with dichloroacetonitrile. This is expected for bimolecular nucleophilic substitution reactions (SN2)since bromide is a much better leaving group than chloride. Bromide has a greater polarizability and it is more easily solvated. In contrast, alkaline hydrolysis is fastest with the dichloroacetonitrile. This would be predicted based on the greater electronegativity of chlorine, as hydrolysis presumably involves the nucleophilic attack of hydroxide on the adjacent nitrile group with no loss of halide. Based on steric hindrance of the nucleophile, one would expect compounds with greater numbers of halogens on one carbon to be more stable. On the other hand, chlorine is strongly electronegative, and the presence of an additional chlorine atom will render the carbon more electrophilic. Its seems the latter effect is more important with the chloroacetonitriles (i.e., DCAN and TCAN). One also sees the effect of electron-withdrawing substituents in comparing the rates of chloropicrin, trichloroacetonitrile, and trichloropropanone. The more the substituent is electron withdrawing, the greater the reaction rate (Le., NO2 > CN > COCHJ. Implications with Respect to Drinking Water Treatment. The results presented here allow one to calculate removal of chloropicrin, trichloroacetonitrile,and dibromoacetonitrile as a function of sulfite dose and contact time. For example, consider the destruction of chloropicrin in a treated water with 0.5 mg/L residual chlorine. If one adds 2 mg/L sulfur dioxide to this water, the residual chlorine will be reduced to zero and a total sulfite residual of 22 pM will exist. Since chloropicrin's hydrolysis rate is very small, the overall removal rate is simply equal to the rate of reaction with sulfite species.
Environ. Sci. Technol., Vol. 23, No. 11, 1989
hobs
- ks
k!2a2[S(IV)]
(9)
One can then define a pH-dependent rate constant, k':
k' = k2a2
(10)
And finally the concentration of chloropicrin as a function of time can be calculated from C = Co exp(-kTS(IV)]t)
(11)
where
k'=
k2Ka 5'6 Ka + [H+] 6.6 x
lo+
+ 10-pH
M-1 s-l
('2)
If the water in question has a pH of 8, then k'is calculated to be 74 M-' s-l. Now, if the water is held for 10 min prior to rechlorination and discharge into the distribution system, the fiial concentration is calculated to be 38% of that which was originally formed (eq 11). A similar analysis for trichloroacetonitrile and dibromoacetonitrile indicates that these compounds would be reduced to 62% and 54% of their respective original concentrations under the same conditions (ignoring hydrolysis). Conclusions 1. Chloropicrin, trichloroacetonitrile, and dibromoacetonitrile can be rapidly destroyed in chlorinated drinking waters with small doses of sulfite. This may represent a viable treatment technology for the control of these chlorination byproducts. This also means that sulfite
should not be used as a quench when analyzing for these compounds. 2. The major degradation products of the reaction of sulfite with chloropicrin, trichloroacetonitrile, and dibromoacetonitrile are the partially dehalogenated species, dichloronitromethane, dichloroacetonitrile, and bromoacetonitrile. 3. Either chloral hydrate, l,l,l-trichloropropanone, dichloroacetonitrile, 2,3,6-trichloroanisole, and 1,l-dichloropropanone do not react irreversibly with sulfite or they react very slowly. 4. The potent mutagen, MX, is removed over several days of contact with sulfite at [S(IV)]concentrations on the order of 100 pM. 5. Other major mutagens, besides MX, may be highly susceptible to decomposition by sulfite. These unknown compounds may be primarily responsible for the rapid loss of mutagenicity noted in chlorinated waters that have been treated with sulfite.
Acknowledgments We express our gratitude to Thomas Potter for his help with the GC/MS analyses. Thanks also go to Dr. Leif Kronberg and Dr. Emile Coleman for providing synthesized MX, and Dr. James N. Jensen for his helpful discussions. Registry No. MX, 77439-76-0; CClsN02, 76-06-2; CCISCN, 54506-2; CHC12CN,3018-12-0; CHBr2CN,3252-43-5; CClSCOCHS,
918-00-3; CHC12COCHs, 513-88-2; chloral, 75-87-6; 2,3,6-trichloroanisole, 50375-10-5; sodium sulfite, 7757-83-7.
Literature Cited Cheh, A. M.; Skochdopole, J.; Koski, P.; Cole, L. Science 1980, 207, 90. Trehy, M. L.; Bieber, T. I. In Advances in the Identification and Analysis of Organic Pollutants in Water; Keith, L. H., Ed.; Ann Arbor Science: Ann Arbor, MI, 1981; Vol. 2, pp 941-975. F m , S. A. Ph.D. Dissertation, University of California at Los Angeles, CA, 1986. Bull, R. J.; Robinson, M.; Meier, J. R.; Sorber, J. EHP, Environ. Health Perspect. 1982, 46, 215. Kringstad, K. P.; Ljungquist, P.0.;de Sousa, J.; Stromberg, L. M. Environ. Sei. Technol. 1983, 17, 468. Meier, J. R.; Bull, R. J. In Water Chlorination: Chemistry, Environmental Impact and Health Effects; Jolley et al., Eds., Lewis: Chelsea, MI, 1985; Vol. 5, p 207. Coleman, W. E.; Munch, J. W.; Streicher, R. P.; Hodakievic, P.A. Prepr. Pap. Natl. Meet.-Am. Chem. SOC., Diu. Environ. Chem. 1986, 332. Holmbom, B.; Kronberg, L.; Backlund, P.;Langvik, V.-A,; Hemming, J.; Reunanen, M.; Smeds, A.; Tikkanen, L. Presentation at the Sixth Conference on Water Chlorination, Oak Ridge, TN, 1987. Loper, J. C. In Water Chlorination: Environmental Impact and Health Effects; Jolley et al., Eds.; Ann Arbor Science: Ann Arbor, MI, 1980; Vol. 3, p 937. Kronberg, L., Holmbom, B.; Tikkanen, L. In Aquatic Micropollutants in the Environment; Bjorseth, Z., Angeletti, Z., Eds.; D. Reidel: Boston, MA, 1986. Meier, J. R.; Lingg, R. D., Bull. R. J. Mutat. Res. 1983,118, 25.
(12) Holmbom, B.; Voss, R. H.; Mortimer, R. D.; Wong, A. Environ. Sei. Technol. 1984, 18, 333. (13) Kopfler, F. C.; Ringhand, H. P.; Coleman, W. E.; Meier, J. R. In Water Chlorination: Chemistry, Environmental Impact and Health Effects; Jolley et al., Eds., Lewis: Chelsea, MI, 1985; Vol. 5, p 161. (14) Nazar, M. A.; Rapson, W. H. Environ. Mutagen. 1982,4, 435. (15) Reckhow, D. A.; Singer, P. C. In Water Chlorination: Chemistry, Environmental Impact and Health Effects; Jolley et al., Eds.; Lewis: Chelsea, MI, 1985; Vol. 5, pp 1129-1257. (16) Miller, E. C.; Miller, J. A. Chemical Carcinogens; American Chemical Society: Washington, DC, 1976. (17) Wilcox, P.; Denny, S. In Water Chlorination: Chemistry, Environmental Impact and Health Effects; Jolley et al., Eds.; Lewis: Chelsea, MI, 1985; Vol. 5, p 1341. (18) Donnini, G. P. Pulp Pap. Can. 1983,84(4), T93-T98. (19) Mabey, W.; Mill, T. J. Phys. Chem. Ref. Data 1978,7,383. (20) Streitwieser, A.; Heathcock, C. H. Introduction to Organic Chemistry; Macmillan: ’New York, 1976. (21) Gilbert, E. E. Sulfonation and Related Reactions; Interscience: New York, 1965. (22) Kronberg, L. Ph.D. Dissertation, Abo Akademi, Finland, 1987. (23) Thurman, E. M.; Malcolm, R. L. Enuiron. Sei. Technol. 1981,15,463-466. (24) Coleman, W. E.; Munch, J. W.; Kaylor, W. H.; Streicher, R. P.; Ringhand, H. P.;Meier, J. R. Environ. Sci. Technol. 1984,18,674-681. (25) Snoeyink, V. L.; Clark, R. R.; McCreary, J. J.; McHie, W. F. Environ. Sci. Technol. 1981,15, 188-192. (26) Coleman, W. E.; Lingg, R. D.; Melton, R. G.; Kopfler, F. C. In Identification and Analysis of Organic Pollutants in Water; Keith, L. H., Ed.; Ann Arbor Science: Ann Arbor, MI, 1976; pp 305-327. (27) Stumm,W.; Morgan, J. J. Aquatic Chemistry, 2nd ed.,John Wiley & Sons: New York, 1981. (28) Wedzicha, B. L. Chemistry of Sulfur Dioxide in Foods; Elsevier Applied Science: Essex, UK, 1984. (29) Betterton, E. A.; Hoffmann, M. R. J. Phys. Chem. 1987, 91, 3011-3020. (30) Exner, J. H.; Burk, G. A.; Kyriacou, D. J.Agric. Food Chem. 1973,21,838-842. (31) Bieber, T. I.; Trehy, M. L. In Water Chlorination: Environmental Impact and Health Effects; Jolley, R. L. et al. Us.; Ann Arbor Science: Ann Arbor, MI, 1983; Vol. 4, Book 2, pp 85-96. (32) Gurol, M. D.; Wowk, A.; Myers, S.; Suffet, I. H. et al. In Water Chlorination: Environmental Impact and Health Effects; Jolley et al., Eds., Ann Arbor Science: Ann Arbor, MI, 1983; Vol. 4, Book 1, pp 269-284. (33) Betterton, E. A.; Erel, Y.; Hoffmann, M. R. Environ. Sei. Technol. 1988,22, 92-99. (34) Holmbom, B.; Kronberg, L. Presentation at the Fifth European Symposium on Organic Micropollutants in the Aquatic Environment, Rome, October 1987. (35) APHA, AWWA; WPCF Standard Methods for the Examination of Water and Wastewater; APHA Washington, DC, 1985.
Received for review August 29,1988. Accepted March 20,1989. This work was supported by Anjou Recherche of the Compagnie Generale des Eaux (Paris). Much of the material in this paper was presented at the 1988 American Water Works Association Annual Conference, Orlando, FL.
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