Environ. Sci. Technol. 1991, 25, 1774-1779
Degradation of Aldicarb, Aldicarb Sulfoxide, and Aldicarb Sulfone in Chlorinated Water Carl J. Miles
Environmental Biochemistry Department, University of Hawaii, 1800 East-West Road, Honolulu, Hawaii 96822 The three insecticidal species of Temik, aldicarb (I), aldicarb sulfoxide (II),and aldicarb sulfone (III),degrade rapidly in chlorinated water. I is sulfoxidized to 11, while I1 forms 111, (chloromethy1)sulfonyl species, and Nchloro-11. I11 forms N-chloro-111, which decomposes to I11 acid and dichloromethylamine. Over the pH range of 6-9, the degradation rates of 1-111 in chlorinated water are several orders of magnitude faster than base hydrolysis, the primary detoxification mechanism in water. Chlorination of aldicarb species during drinking water treatment can reduce the initial concentrations but may produce compounds of concern to public health. W
Introduction
Human exposure to pesticide residues in drinking water is a potential health problem. Although many rural homes and communities consume raw groundwater, the majority of the U.S. population drinks treated water, which usually includes chlorination. The fate of pesticide residues in chlorinated water needs to be determined. While controlling many pathogenic bacteria, chlorination of natural organic matter can produce undesirable compounds such as the trihalomethanes and the haloacetonitriles in drinking water. On the other hand, chlorination can degrade undesirable compounds such as pesticides and rates vary depending upon the pesticide (1-5). Carbamate and carbamoyloxime pesticides are susceptible to degradation in chlorinated water (3-5). In many cases, degradation products have not been identified and this data gap must be filled to help determine potential health effects. Field studies have shown that typical water treatment processes that include chlorination do not significantly degrade many of pesticides present in the treatment plant influent (6, 7). Aldicarb [2-methyl-2-(methylthio)propionaldehyde 0(methylcarbamoyl)oxime],the active ingredient in Temik insecticide/nematicide, is used on a wide variety of crops in several states. In soils, it is rapidly oxidized to the sulfoxide and the sulfone, which are the primary species found in groundwater contamination episodes (8). In a study of 43000 potable well samples collected in 34 states, aldicarb residues above the EPA Health Advisory level of 10 pg/L were found in the groundwater of 8 states (8). On the basis of the above evidence, reaction of chlorine and aldicarb sulfoxide and sulfone during drinking water treatment is possible. This investigation examined the fate of aldicarb (I) and its metabolites, aldicarb sulfoxide (11)and aldicarb sulfone (III),in chlorinated water. Degradation rates and products for these compounds were determined and an overall reaction scheme is proposed. Experimental Section
I, I oxime, I nitrile, 11, I1 nitrile, 111, I11 oxime, and I11 nitrile were obtained from Rhone-Poulenc (Research Triangle Park, NC). I11 acid was prepared by refluxing I11 in 3 M HC1 and then 3 M NaOH for 24 h. The alkaline/ethyl ether extract was discarded while the acidic/ ethyl ether extract contained I11 acid. Hypochlorite solutions were prepared by bubbling chlorine gas into a so1774
Environ. Sci. Technol., Vol. 25, No. 10, 1991
lution of sodium hydroxide (or NaOD in DzO). Hypochlorite concentrations were determined by iodometric titrations. Chloramine solutions were prepared by 1:1 mixtures of hypochlorite and ammonia (pH 9). Water was obtained from a Milli-Q (Millipore, Bedford, MA) water system. All other chemicals and reagents were obtained from Aldrich Chemical Co. (Milwaukee, WI) and used without further purification. Pesticide concentrations were measured by liquid chromatography (LC) using a Perkin-Elmer (Norwalk, CN) Series 10 pump, a Rheodyne (Cotati, CA) 7125 injector (100-pL loop), Perkin-Elmer 3 X 0.46 cm C18column (3 pm; reduced activity), and a Linear Instruments (Reno, NV) Model 200 UVIS variable-wavelength detector (200 nm). Mobile phases with 5 4 0 % acetonitrile in water were used after being degassed with helium. Solutions of the test pesticide (0.1 mM in 0.1 M phosphate or deionized water) were treated with 0.1-10 mM hypochlorite over the pH range 6-9, and pesticide concentration was measured periodically. All kinetic experiments used a chlorine/pesticide ratio of 240, and rates were assumed to be pseudo first order. Solutions were stored in darkness (24 "C) or simulated sunlight (30 "C; see ref 9 for photoreactor design), and the pH was measured after analyses were completed. Proton (300 MHz) and 13C (76 MHz) nuclear magnetic resonance (NMR) spectra were obtained on a General Electric (Fremont, CA) QE 300 or a Nicolet (Madison, WI) NT 300 spectrometer. Spectra were determined in D 2 0 or a deuterated phosphate buffer (0.1 M; pH 7) with 2,2dimethyl-2-silapentane-5-sulfonate (DSS) and dioxane as proton and 13C internal standards, respectively. Initial pesticide concentrations were -20 mM. In certain experiments, deuterated hypochlorite solutions were added in molar equivalents or excess amounts and spectra recorded over time. Deuterium-hydrogen exchange was measured by proton NMR on solutions of I1 oxime or I11 oxime (0.008 M) in D 2 0 and 0.42 M NaOD. The carbamoyloximes were not used in these experiments because of rapid hydrolysis to the oximes. Mass spectrometry was performed using a VG Trio 2A quadrupole instrument (VG BioTech, Altrincham, U.K.). Gas chromatography (Model 5890A; Hewlett-Packard, Palo Alto, CA) employed a 30 m X 0.25 mm i.d. (0.25-pm film) DB-5 capillary column (J&W Scientific, Cordova, CA). One-microliter injections were made on a splitless injector at 250 "C with the MS interface and source at 200 "C. After 1 min a t 40 "C, the column was temperature programmed to 200 "C at 10 "C/min. Positive ion electron ionization (EI+) spectra were obtained at 70 eV with a mass range of 45-650 amu (1-s scan rate). Chemical ionization (CI) spectra were obtained with methane a t a source manifold pressure of 4 x mbar and a source temperature of 200 "C. Positive ion CI (CI+) spectra were recorded over the 45-650 amu range (1 s) while the negative ion CI (CI-) spectral range was 100-650 amu (1s). For all modes of operation, the MS was tuned with Heptacosa (perfluorotributylamine). Thermospray (TSP) spectra were obtained with a VG Thermospray/Plasmaspray LC/MS probe a t a source
0013-936X/91/0925-1774$02.50/0
0 1991 American Chemical Society
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o = s\
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Figure 1. Proposed reaction pathways for 1-111 in chlorinated water. Reactions other than chlorination are noted.
manifold pressure of 5 X lo4 mbar and source temperature of 250 OC. A vaporizer temperature of 190 OC was optimized for maximum M + H+ (m/z 207) formation with aldicarb sulfoxide. The discharge electrode was operated at 250-400 p A and was necessary to stabilize ion transmission. Positive and negative ion spectra were acquired over the 90-650 amu range (1-s scan rate). Particle beam (PB) spectra were obtained with a VG LINC LC/MS interface with a Hildebrand grid nebulizer (Leeman Labs, Lowell, MA). The helium pressure into the nebulizer was maintained a t 40 psi while the temperature of the desolvation chamber gas flow was held at 30 "C. mbar Typical operating pressures were 8,0.9, and 3 X a t the first-stage momentum separator pump, the second-stage momentum separator pump, and the ion source, respectively. Positive ion electron ionization (EI+) spectra were obtained a t 70 eV with a source temperature of 200 "C and a mass range of 50-650 amu (1-s scan rate). CI spectra were obtained with 1%ammonia in methane a t mbar and a source a source manifold pressure of 4 X temperature of 150 "C. CI+ spectra were recorded over the 45-650 amu range (1-s scan rate) while negative ion CI spectral range was 60-650 amu (1-s scan rate). For all modes of operation, the MS was tuned and optimized by using Heptacosa. LC/MS was performed with the instrument described above and a variable-wavelength UV detector (220 nm) in-line to monitor products. Flow rates of 0.5 mL/min were used for all separations. Mobile phases consisted of 25% or 50% acetonitrile (Fisher Optima, Fairlawn, NJ) in Milli-Q water (PB/LC/MS) or 0.1 M ammonium acetate (TSP/LC/MS; Fisher HPLC grade) and were de-
gassed with helium purging. Neutral and acidic mixtures of pesticide and hypochlorite were either analyzed directly or extracted with ethyl ether at various stages of reaction. Some samples were treated with diazomethane prior to analysis by GC/MS. Retention times and spectra of unknowns were compared to standard materials when possible. Chloroform was determined by headspace analysis GC with electron capture detection using a 25 m X 0.53 mm, 5.0-pm BP-1 (SGE Scientific, Austin, TX) column a t 70 "C.
Results and Discussion I was rapidly oxidized to I1 in chlorinated water (Figure 1). The reaction appeared to be complete and there was no evidence of N-chloro-I formation by observations with proton NMR or LC. With 40 or more molar equivalents of hypochlorite, the kinetics were too fast to measure by NMR or LC (i.e., t l / z < 1 min). The disappearance of I1 in chlorinated water ([chlorine]/[II] 2 40) followed a two-step rate process ( 4 ) in which approximately half of the initial I1 rapidly disappeared and then the rate slowed considerably. This suggests a change in reaction mechanism. Competition between sulfoxidation and chlorination of the methylsulfinyl group (see below) could be the cause of the complex kinetics. The rates of I1 loss determined from the second phase of degradation increased slightly with an increase in pH and were several orders of magnitude faster than rates for base hydrolysis (10) (Figure 2). In simulated sunlight at pH 9, the reaction rate was similar until total chlorine concentrations were not detectable (ca. 30 min), upon which the reaction rate slowed significantly. Environ. Sci. Technol., Vol. 25, No. 10, 1991
1775
.-'c
:I
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.-v
Table I. Spectral Data of Aldicarb Derivatives and Degradation Products
without chlorine
0-0 1
0
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.'
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I I1 111
V 04 5.5
6.0
6.5
7.0
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Flgure 2. Degradation rate of I1 as a function of pH with and without chlorine.
VI1
I1 reacted rapidly in chlorinated water ([chlorine]/ [11] 2 10) forming several products including I11 (ca. 25%) and IV, which were determined by LC/MS, GC/MS, or NMR. The formation of IV was confirmed by the regeneration of -25% I1 after treatment of chlorine/II solutions with excess thiosulfate after 1 half-life of degradation. N chloro compounds are usually unstable and the parent compound is regenerated with reductants such as ascorbic acid (11)and sulfite (12). Several other products (V-VIII) have been tentatively identified by GC/MS and NMR (see Table I) and accounted for the remaining I1 identified (Figure 1). However, the actual products (V-VII) could be the respective carbamoyloximes since the carbamoyloxime thermally degrades to the nitrile under GC conditions. Chloroform also was measured in these mixtures. I11 degradation products also were observed in chlorine/II solutions and are described later. When the reaction was carried out in simulated sunlight, significantly lower amounts of the chlorinated products were observed. MS analysis of I1 solutions treated with 10 or more molar equivalents of chlorine showed the presence of (chloromethy1)sulfonyl derivatives (V-VIII) whereas I11 solutions treated similarly did not. Disappearance of the methylsulfinyl resonance in the proton NMR after addition of 3 molar equiv of chlorine suggests that this group is preferentially chlorinated. Analysis of solutions with various chlorine/II ratios (pH 7; 24 h) showed that VI1 (or the respective carbamoyloxime) was detectable only when the chlorine/II ratio was 23 and concentrations increased rapidly with increasing proportions of chlorine. Identification of chloroform in chlorine/II solutions suggests hydrolysis of the trichloromethyl species. In a similar experiment, chloroform concentrations increased rapidly only after a chlorine/II ratio of 3 was reached. For the highest chlorine/II ratio tested (7), the chloroform produced represents only a 0.4% yield. Much lower chloroform concentrations were observed in chlorine/III solutions and, as opposed to 11, there was no sudden increase in chloroform concentration at high chlorine/III ratios. These results suggest that I1 and I11 react with chlorine by different mechanisms and suggest preferential chlorination of the methylsulfinyl group. Preferential chlorination of the methylsulfonyl group in 111, as opposed to the methylsulfinyl group in 11, is predicted by acidity differences. The base-catalyzed haloform-type reaction should be favored when the negative charge is a to a sulfonyl group as opposed to a sulfinyl group. Formation of trichloromethyl sulfones from alkyl methyl sulfoxides was reported to occur by this mechanism following oxidation of the sulfoxide to the sulfone (13). Subsequent hydrolysis of the trichloromethyl sulfones
-
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Environ. Sci. Technoi., Vol. 25, No. 10, 1991
VI11
lH NMR 6 1.45 (s, 6 H), 1.99 (s, 3 H), 2.80 ( 3 , 3 H), 7.73 (s, 1 H) 'H NMR 6 1.51 (d, J = 24.9 Hz,6 H), 2.64 (s, 3 H), 2.81 (9, 3 H), 7.94 (s, 1 H) NMR 6 18.2, 26.9, 31.3, 57.7, 156.6, 157.6 lH NMR 6 1.64 (9, 6 H), 3.08 (s, 3 H),2.81 (9, 3 H), 8.02 ( 3 , 1 H) EI(+)MS m / z 82 (251, 68 (100) CI(+)MS m / z 184 (7), 182 (19), 155 (7), 117 (111, 115 (32), 97 (19), 82 (16), 68 (100) CI(-)MS m / z 115 (70), 113 (100) EI(+)MS m / z 85 (45), 83 (85), 68 (100) CI(+)MS m/z 218 (2), 216 (4), 181 (2), 179 (5), 154 (Z), 152 (5), 97 (22), 83 (31), 68 (100) CI(-)MS m / z 149 (65), 147 (loo), 132 (22), 113 (31) EI(+)MS m / z 121 (33), 119 (93), 117 (100). 84 (18). 82 (201, 68 (88) CI(+)MS miz 254 (8). 252 (23). 250 (25). 227 (26). 225 '(62), 223'(60), 218'(28), 216.(40), 68 (100) CI(-)MS m / z 185 (21), 183 (63), 181 (69), 151 (9), 149 (35), 147 (45), 132 (100) EI(+)MS m / z 183 (3), 101 (85), 73 (loo), 59 (65), 49 (11)
IX X
XI
CI(+)MS m/z 217 (3), 215 (8), 185 (4), 183 (9), 101 (loo), 86 (181, 73 (431, 59 (12) CI(-)MS m / z 115 (33), 113 (100) 'H NMR 6 1.66 (s, 6 H), 3.11 (8, 3 H), 3.40 (9, 3 H), 8.15 (s, 1 H) 'H NMR 6 3.66 (5, 3 H) 13C NMR 6 64.7 EI(+)MS m / z 103 (12), 102 (14), 101 (51), 100 (86), 99 (981, 98 (loo), 65 (14), 63 (27), 62 (51) 'H NMR 8 1.55 (s, 6 H), 3.16 (s, 3 H) I3C NMR 6 18.7, 36.3, 68.3, 172.1 EI(+)MS m / z 149 (8), 1 2 1 (3), 101 (IOO), 73 (85), 59 (72) CI(+)MS m / z 181 (51), 149 (34), 121 (12), 101 (loo), 73 (44), 59 (21)
yielded chloroform. Conversely, dimethyl sulfoxide (DMSO) reacted rapidly with tert-butyl hypochlorite to form a trichloromethyl sulfoxide (14),and formation of an HC1 adduct appeared to block further oxidation of DMSO. Similar to the reaction kinetics observed in this study with 11, DMSO reacted rapidly with tert-butyl hypochlorite until 3 equiv was added and then reacted more slowly with additional hypochlorite (14). Other evidence indicates that chlorine attacks the sulfur forming a chlorosulfoxonium ion, which rearranges to form the a-chloro sulfoxide (15, 16). The absence of a-carbanion formation during chlorination of I1 oxime was suggested by the lack of deuterium-hydrogen exchange. Also, the minor influence of base on the chlorine/II reaction rate does not support a-carbanion formation. This suggests that I1 is chlorinated via the chlorosulfoxonium ion. However, rapid S-oxidation of the a-chloro sulfoxide must occur since only a-chloro sulfone products are observed. The lack of evidence of a (chloromethy1)sulfinyl species is difficult to explain. The inductive effect of the chlorine atoms should make oxidation of the (chloromethy1)sulfinyl more difficult. One possible explanation is that the high polarity of sulfoxides lowers extraction efficiency, making them undetectable by GC/MS analysis. Also, unreacted I1 gives a low GC response compared to 111despite the use of lower injection port temperatures reported as successful by others (17). LC/MS circumvents these problems, but all efforts to detect (chloromethy1)sulfinyl derivatives were unsuccessful. I11 reacted completely and rapidly in chlorinated water to form N-chloro-111 (IX), which slowly degraded to 111acid
water are much slower than rates observed for reaction of primary amines and chlorine [ca. lo9 M-lnrnin-' (20)].The slower rates are consistent with the decreased amine basicity, resulting in decreased rates of N-chloroamine formation (20). For example, amides are N-chlorinated much less rapidly at neutral and alkaline pH than amines (20). The reaction rate of hypochlorous acid and amine will be optimal a t pH values where the reactants are at maximum concentration (pH < pK~oc1/0~1(7.5) and pH > pKAof the amine). The pK's for I1 and 111 are unknown I 1 but should be much lower than for methylamine (pKA= 10.7). Thus, the maximal reaction rate for I1 or I11 and chlorine should be at slightly basic pH values and decrease a t higher or lower values. Rate decreases were observed at higher pH values for both compounds; however, the rate decreased only for I1 and only at slightly acidic pH values (Figures 2 and 4). The reaction rate of I11 with hypochlorite (pH 7) in2 8 7 6 5 4 3 creased with increasing temperature over the range of PPM 15-35 "C, and an activation energy of 12.4 kca1.M-l was Flgure 3. Proton NMR spectra of I11 (A) before addition of DOCI, (8) calculated from the slope of a plot of log k vs 1/T. The 5 min after 7 equiv of DOC1 added, and (C) 16 h after 7 equiv of DOC1 activation energy and enthalpy (11.8 kcal) calculated for added. (pD ca. 8.) this reaction are similar to values for base hydrolysis of 111 and other carbamates (19,21,22).However, the cal7 1 1 culated entropy (-24 eu) is much lower than that determined for base hydrolysis of I and other N-methyl carbamates (-5.8 eu) (21-23). Lower entropy is associated with a tetrahedral activated complex and a SN2 mechanism as opposed to the deprotonated amine in the elimination mechanism proposed for base hydrolysis of I (23)and other carbamates (22). Reaction rates of 111 with 100 equiv of chloramine (10 with chlorine mM) were significantly slower (ca. 500-fold) than with free chlorine and varied little over the pH range of 7-9. The average reaction rate of I11 and chloramine was 1.1f 0.5 X 10" mi&, which corresponds to a half-life of 105 h. The 5.5 6.0 6.5 7.0 7.5 8.0 8.5 9.0 9.5 reaction rate of methylamine and hypochlorous acid was PH 200-fold faster than the reaction of methylamine and Flgure 4. Degradation rate of I11 as a functlon of pH with and without chloramine (24) and was relatively insensitive to pH chlorine. changes in the neutral to slightly alkaline range. Chlorine/II mixtures produced the most products ob(XI) and dichloromethylamine (X) (Figure 1). The served by GC/MS after a 24-h reaction time. Since I1 and products were determined by LC/MS, GC/MS, or NMR. 111 and their chlorine reaction products fragment readily, The formation of IX was confirmed by quantitative rechemical ionization was helpful to determine structures. generation of I11 by addition of thiosulfate to chlorine/III Products V-VI11 and XI all formed M + H+ ions but most solutions. 'H NMR analyses showed instantaneous Nwere in low abundance (Table I). Major ions in the CI+ chlorination followed by slower formation of X and XI (see and EI+ spectra were the isopropyl nitrile fragment (m/z Figure 3 and Table I). The chemical shift for the N-methyl 68) for nitrile (or carbamoyloxime) derivatives and the group in IX (+0.59 ppm) is consistent with the 0.57 ppm isopropyl acid/ester fragment (m/z 101) for the acid/ester downfield shift observed for N-chloromethomyl (18). derivatives. Loss of the cyano group was also observed in Disappearance of I11 in chlorinated water followed the CI+ spectra of V and VI1 (Figure 5). pseudo-first-order kinetics and the rate increased slightly The CI- spectra contained significant information for with pH (Figure 4). As observed with 11, degradation of structure elucidation as well as selective detection of 111 in chlorinated water was several orders of magnitude chlorinated compounds in the GC runs. Mono-, di-, and faster than base hydrolysis. Others observed an increased trichlorosulfonyl derivatives fragmented a t the sulfone/ reaction rate for 111in chlorinated water and attributed isopropyl bond forming the monochloro ( m / z 113/115), the enhanced rate to pH increase from the hypochlorite dichloro ( m / z 147/149/151), and trichloro ( m / z 181/ addition (19). However, their rates slowed after -3 h, 183/ 185) sulfone ions with characteristic chlorine isotopic probably because of chlorine depletion. As observed here patterns. Within these spectra, loss of chlorine was also with 11, reaction rates in simulated sunlight were similar observed as was the sulfone isopropyl nitrile fragment (m/z until chlorine was depleted, after which the reaction rate 132) (Figure 5). slowed significantly. The similar reaction rates of I1 and 111in chlorinated water suggest that decomposition of I11 The positive ion thermospray (TSP+) spectrum of I11 controls the rate of I1 degradation. consisted of [M + HI+ and [M + NH4]+ adducts (both The second-order rate constants (ca. lo4 M-lsmin-l) es-100%). The major ions for I1 were the protonated nitrile timated for the reaction of I1 or I11 with chlorine are ( m / z 132) and the nitrile-ammonium adduct ( m / z 149) considerably higher than literature values for the reaction while [M + HI+ was -10% and [M + NH4]+was 220 in CI+.
1
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Flgure 8. TSP+ spectra of I V (A) and I X (B).
temperature, or the relatively high source temperature used. The N-chloro derivatives behaved similar to the parent compounds with IX forming a large (90%) [M + NH4]+ ( m / z 274) while the [M NH4]+ ( m / z 258) for IV was only -0.8% of the base peak (Figure 6). TSP/ LC/MS was useful for confirmation of these N-chloro derivatives and offers a sensitive and selective route to detection of these compounds in complex matrices. Negative ion thermospray (TSP-) spectrum of I11 consisted only of m/z 139, which is probably the (methylthio)sulfonyl-acetate adduct. The TSP- spectrum of I1 was very weak and was mostly m / z 132. The pesticide/ chlorine reaction mixtures showed several compounds with a monochloro species [ m / z 95 (100%) and 97 (35%)], which was probably the chloroacetate ion formed by reaction of chlorine with acetate in the thermospray probe
+
1778
Environ. Sci. Technol., Voi. 25, No. 10, 1991
or source. Compound VI1 had a TSP- spectrum very similar to that observed by GC/MS in CI- (Figure 5). LC/MS analysis of chlorine/II reaction mixtures over several days showed that the maximum concentration of VI1 was reached in -1 day and was stable for a t least 1 week. VI1 eluted in 3.6 min with 50% acetonitrile in 0.1 M ammonium acetate. Several other products with the mono- (m/z 113/115), di- (m/z 147/149/151), and trichloro ( m / z 181/183/185) sulfone ions with characteristic chlorine isotopic patterns were observed. Absence of complimentary TSP+ or particle beam (EI+ or CI+) spectra hindered structural determination of these minor products. Dichloromethylamine (X) was observed by GC/MS in neutral extracts of both pesticides and chlorine. The peak eluted shortly after the solvent front and was identified by a search of the NBS mass spectral library. This degradation product was also observed by LC (retention time of 3.5 min with 25% acetonitrile/water). Negative ion thermospray was the only LC/MS mode that produced a significant spectrum and consisted of the chloroacetate ion. Particle beam LC/MS of pesticide/chlorine mixtures showed several spectra with characteristic fragments observed by GC/MS such as m/z 68,79,132, etc. Compound IX produced mostly fragment ions but the [M - 79]+ ion [ m / z 177 (1%) and 179 (0.3%)] was observed. Particle beam CI+ and CI- spectra did not yield any significant new structural information for the major UV peaks. In summary, I rapidly and completely oxidizes to 11. I11 quickly and quantitatively forms the N-chloro derivative (IX), which decomposes to the acid (XI) and dichloromethylamine (X). Reaction of I1 is more complex with sulfoxidation, N-chlorination, and chlorination of the methylsulfinyl group all occurring rapidly. Over the pH range of natural waters, the degradation rates of 1-111 in chlorinated water are several orders of magnitude faster than base hydrolysis, the primary detoxification mechanism. Reaction of I11 with chloramines results in significantly lower reaction rates than reaction with free chlorine. Reaction rates of chlorine and pesticide are not significantly affected by simulated sunlight until chlorine is depleted by photoreaction after which the reaction rate decreases. Sunlight also decomposes many of the chlorinated products to unidentified products. These results suggest that chlorine treatment of drinking water contaminated with these pesticide residues will result in significantly lower residues but also can contain different species than the untreated water. However, the human health effects of most of these species is unknown.
Acknowledgments
I thank Wendy Oshiro for technical assistance, Daniel Doerge for helpful discussions, Susie Scherer for graphics, and Rhone-Poulenc for the aldicarb derivatives. This is Journal Series No. 3577 from the Hawaii Institute of Tropical Agrriculture and Human Resources. Registry No. I, 116-06-3; 11, 1646-87-3; 111, 1646-88-4; IV, 135043-57-1;V, 135043-58-2;VI, 135043-59-3;VII, 135043-60-6; VIII, 135043-61-7;IX,135043-62-8;X, 7651-91-4; XI, 25841-43-4; CHC13, 67-66-3.
Literature Cited (1) El-Dib, M. A.; Aly, 0. A. Water Res. 1977, 11, 611. (2) Dennis, W. H.; Meier, E. P.; Randall, W. F.; Rosencrance, A. B.; Rosenblatt, D. H. Environ. Sci. Technol. 1979,13, 594. (3) Lemley, A. T.; Zhong, W. Z.; Janauer, G. E.; Rossi, R. In Treatment and Disposal of Pesticide Wastes;Krueger, R. F., Seiber, J. N., Eds.; ACS Symposium Series 259; Americal Chemical Society: Washington, DC, 1984; pp 245-260.
(4) Miles, C. J.; Trehy, M. L.; Yost, R. A. Bull. Environ. Contam. Toxicol. 1988, 41, 838. (5) Miles, C. J.; Oshiro, W. C. Environ. Toxicol. Chem. 1990, 9, 535. (6) Lykins, B. W.; Koffskey, W. E.; Miller, R. G. J.-Am. Water Works Assoc. 1986, 78, 66. (7) Miltner, R. J.; Baker, D. B.; Speth, T. F.; Fronk, C. A. J.-Am. Water Works Assoc. 1989, 81, 43. (8) Moye, H. A,; Miles, C. J. Rev. Enuiron. Contam. Toxicol. 1988, 105, 99. (9) Crosby, D. G.; Tang, C. S. J. Agric. Food Chem. 1969,17, 1041. (10) Lightfoot, E. N.; Thorne, P. S.; Jones, R. L.; Hansen, J. L.; Romine, R. R. Enuiron. Toxicol. Chem. 1987, 6 , 377. (11) Uetrecht, J. P.; Zahid, N. Chem. Res. Toxicol. 1991,4, 218. (12) Nickelsen, M. G.; Nweke, A,; Scully, F. E., Jr.; Ringhand, H.P. Chem. Res. Toxicol. 1991, 4 , 94. (13) Haszeldine, R. N.; Rigby, R. B.; Tipping, A. E. J. Chem. SOC.,Perkin Trans. 1 1973, 676. (14) Walling, C.; Mintz, M. J. J. Org. Chem. 1967, 32, 1286. (15) Cinquini, M.; Colonna, S.;Landini, D. J. Chem. SOC.,Perkin Trans. 2 1972, 296.
(16) Klein, J.; Stollar, H. J. Am. Chem. SOC.1973, 95, 7437. (17) Zhong, W. Z.; Lemley, A. T.; Spalik, J. J. Chromutogr. 1984, 299, 269. (18) Miles, C. J.; Oshiro, W. C. Environ. Toxicol. Chem. 1990, 9, 535. (19) Lemley, A. T.; Zhong, W. Z. J. Agric. Food Chem. 1984,32, 714. (20) Morris, J. C. In Principles and Applications of Water Chemistry; Faust, S.D., Hunter, J. V., Eds.; John Wiley & Sons, Inc.: New York, 1967; Chapter 2. (21) Christenson, I. Acta Chem. Scand. 1964, 18, 904. (22) Fukuto, T. R.; Fahmy, M. A. H.; Metcalf, R. L. J. Agric. Food Chem. 1967, 15,273. (23) Bank, S.; Tyrrell, R. J. J.Agric. Food Chem. 1984,32,1223. (24) Isaac, R. A.; Morris, J. C. Environ. Sci. Technol. 1983, 17, 738.
Received f o r review February 19, 1991. Revised manuscript received June 7, 1991. Accepted June 10, 1991. The officeof Research Administration, University of Hawaii, provided financial support.
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