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Environ. Sci. Technol. 2001, 35, 2461-2469

Reduction of the Carbamate Pesticides Oxamyl and Methomyl by Dissolved FeII and CuI TIMOTHY J. STRATHMANN* AND ALAN T. STONE Department of Geography and Environmental Engineering, Johns Hopkins University, Baltimore, Maryland 21218

FIGURE 1. Structures of oxamyl and methomyl.

The degradation of two oxime carbamate pesticides, oxamyl and methomyl, was investigated in anoxic solutions containing various metal ions and reducing agents. In reagent-free solutions, these carbamates degrade slowly via base-catalyzed elimination. Rates of carbamate degradation are accelerated by FeII, CuI, and CuII, but not by several other metal ions and reducing agents. In the presence of FeII, carbamates undergo a net two-electron reduction that is coupled to the sequential one-electron oxidation of two FeII ions. The observed products are a substituted nitrile, methanethiol, and methylamine. A radical intermediate is inferred by polymerization of the radical scavenger acrylonitrile. Redox kinetics (i) vary with carbamate identity, (ii) exhibit first-order dependence on both FeII and carbamate concentration, (iii) are relatively independent of pH, (iv) follow Arrhenius temperature dependence, and (v) are only indirectly influenced by the presence of O2. Coordinatively saturated FeII complexes (FeIIEDTA2- and FeII(CN)64-) react with oxamyl at rates equal to and greater than hexaquo FeII, respectively, indicating that an inner-sphere FeII-carbamate coordination complex is not required for electron transfer. Experimental results indicate that CuI reduces the carbamates by the same mechanism as FeII but at much higher rates. In contrast, CuII acts as a catalyst for both elimination and reduction reactions.

In the absence of other reactive species, oxime carbamates (I; hereafter referred to as carbamates) degrade in aqueous solution via a base-catalyzed elimination reaction (E1cb mechanism) (7-9):

E1cb elimination is well documented for compounds containing the N-methyl-carbamate moiety (10, 11). According to this mechanism, the parent compound is transformed into a substituted oxime (II), methylamine (III) and CO2 (10). Degradation of oxamyl and methomyl by elimination can be extremely slow, especially at subneutral pH where reaction half-lives can range from months to years (10). In contrast, rapid degradation of these carbamates is observed in anoxic soil suspensions; it has been attributed to a redox pathway involving FeII (12, 13). To date, the work of Bromilow et al. (13) represents the only detailed study of the abiotic degradation of oxamyl and methomyl in the presence of FeII. In addition to the oxime and methylamine elimination products, a substituted nitrile (IV) and methanethiol (V) are observed as degradation products in homogeneous solutions containing FeII. The authors hypothesize that these products form as a result of a net two-electron carbamate reduction that is coupled with the one-electron oxidations of two FeII ions:

Introduction Oxamyl and methomyl (Figure 1) belong to a class of compounds known as oxime carbamates (or carbamoyloximes). They are widely used for the control of insect and nematode pests. Nearly three million pounds are applied annually to U.S. agricultural fields (1). They are systemic pesticides, taken up through root or leaf surfaces and distributed throughout the plant (2). Like other closely related carbamates, oxamyl and methomyl act by inhibiting the enzyme acetylcholinesterase (2). The fate of oxamyl and methomyl in the environment is of interest because their acute toxicity ranks among the highest of all pesticides in use today (oral LDrat 50 ) 5 mg/kg for oxamyl and 21 mg/kg for methomyl) (2). Concern also arises from their extremely high water solubility (280 g/L for oxamyl and 58 g/L for methomyl) (3) and low sorption affinity to soils (4, 5); both properties suggest a high degree of mobility in soil and aquatic environments. In fact, oxamyl and methomyl have been detected by groundwater surveys conducted throughout the United States (6). * Corresponding author phone: (410) 516-5527; fax: (410) 5168996; e-mail: [email protected]. 10.1021/es001824j CCC: $20.00 Published on Web 05/01/2001

 2001 American Chemical Society

Results from a limited number of kinetic experiments suggest that the rate of carbamate reduction is proportional to FeII concentration, but relatively independent of pH over the small range investigated (5.65-7.65) (13). A more thorough kinetic study, covering a wider range of conditions, is of interest. Several studies indicate that FeII is an important reductant for a variety of organic and inorganic pollutants found in natural waters (13-19). FeII is often abundant in anoxic and suboxic soil and sediment environments (20, 21). In these settings, FeII arises primarily from the chemical and microbial reduction of FeIII-bearing (hydr)oxide and silicate minerals (22-24). FeII is also released by weathering of FeII-bearing minerals (25, 26). Additional FeII is produced in surface waters via photochemical FeIII reduction (27, 28). This study examines the abiotic degradation of two carbamates, oxamyl and methomyl, in the presence of a VOL. 35, NO. 12, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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variety of chemicals that are found in anoxic and suboxic aqueous environments. These include organic reducing agents (e.g., hydroquinone, cysteine), inorganic reducing agents (e.g., bisulfide), metal ion hydrolysis catalysts (e.g., CuII, ZnII), and metal ion reducing agents (e.g., FeII, CuI). Carbamate reduction by FeII and CuI is examined in detail, and two alternative mechanisms are proposed to account for the observed products, intermediates, kinetics, and stoichiometry. Because of its abundance in subsurface environments, FeII may play an important role in the degradation of these and related agrochemicals. Determining the pathways and rates through which carbamates react with FeII will improve our ability to predict the fate of these compounds in diverse environmental settings.

Materials and Methods All chemical reagents were of the highest purity available. Unless otherwise stated, all solutions were prepared inside a controlled-atmosphere glovebox (95% N2, 5% H2, Pd catalyst; Coy Laboratory Products, Grass Lake, MI) using distilled, deionized water (DDW) with a resistivity of 18 MΩcm (Millipore Corp., Milford, MA). Prior to use, DDW was autoclaved and was then purged (>3 h/L) with ultrahigh purity nitrogen (BOC gases, Baltimore, MD); this was done prior to, as well as immediately after, bringing DDW into the glovebox. Metal ion salt and organic buffer stock solutions were filtered (0.22 µm Millipore Millex-GS) prior to use. All glassware was soaked in 4 N HNO3 (J. T. Baker, Phillipsburg, NJ) and was rinsed several times with DDW prior to use. Glassware having prior contact with iron solutions was first soaked in 0.1 M oxalic acid (Aldrich Chemical Co., Milwaukee, WI) and was rinsed several times with DDW prior to soaking in HNO3. Chemicals. Oxamyl, methomyl, oxamyl oxime, methomyl oxime, and N,N-dimethyl-1-cyanoformamide (DMCF) were provided by DuPont Crop Protection (Wilmington, DE). Additional quantities of oxamyl and methomyl were purchased from Chem Service (West Chester, PA). Sodium methanethiolate and methylamine hydrochloride were obtained from Fluka Chemie AG (Buchs, Switzerland). Dimethyl disulfide was purchased from Aldrich. 3-(2-Pyridyl)-5,6-bis(4-phenylsulfonic acid)-1,2,4-triazine monosodium salt monohydrate (ferrozine colorimetric reagent, Aldrich) was used for determination of dissolved FeII. Acetonitrile (J. T. Baker), acetic acid (J. T. Baker), methanol (J. T. Baker), and difluoronitrobenzene (Aldrich) were used for HPLC analysis of carbamates and their degradation products. Acrylonitrile (Aldrich) was used to test for the presence of a radical intermediate. Ionic strength was adjusted using NaCl (Aldrich). HCl (J. T. Baker), 2-(N-morpholino)ethanesulfonic acid monohydrate (MES; Aldrich), and 3-(N-morpholino)propanesulfonic acid (MOPS; Sigma) were used to buffer pH. These buffers were selected because they complex metal ions weakly, and therefore do not significantly influence FeII speciation. Control experiments were conducted to verify that carbamate reaction rates are not significantly influenced by the buffer concentrations used in this study. Several metal ions and reductants were surveyed to evaluate their potential role as stoichiometric reagents or catalysts for the degradation of carbamates. FeCl2‚4H2O, NiCl2‚6H2O, CuCl2‚2H2O, CuCl, ZnCl2, PbCl2, K4Fe(CN)6, ethylenediaminetetraacetate (EDTA), catechol, citrate, malonate, hydroquinone, hydroxylamine hydrochloride, lactic acid, and sodium oxalate were obtained from Aldrich. CaCl2‚ 2H2O, MgCl2‚6H2O, MnCl2‚4H2O, CoCl2‚6H2O, and Na2S were obtained from J. T. Baker. L-Cysteine hydrochloride and 3-mercaptopropionic acid were obtained from Sigma Chemical (St. Louis, MO). FeII stock solutions were prepared by filtering 25 mL of a ∼1.1 M FeCl2 solution through a 0.02 µm filter (Anatop 2462

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25-Plus inorganic membrane; Whatman Scientific, Maidstone, England) into 2 mL of 0.1 M HCl. Filtrate concentration was determined using flame atomic absorption spectrophotometry (AAS; Aanalyst 100, Perkin-Elmer, Norwalk, CT); absorbance of the stock solution was compared against an iron AAS authentic standard (Aldrich). 10 mM CuI stock solutions were prepared daily in 1 M NaCl. Experimental Setup. Unless otherwise stated, experiments were conducted in an oxygen-free glovebox (95% N2, 5% H2). Batch reactors were mixed continuously under darkness in a constant temperature circulating water bath contained within the glovebox. Except for the temperature variation experiments, all reactions were conducted at 25 ( 0.1 °C. Most reactions were performed in 250-mL polypropylene containers to prevent silica contamination and adsorption of iron to reactor surfaces. Selected reactions were conducted in zero-headspace syringes to prevent volatilization of sulfur compounds (H2S, methanethiol, and dimethyl disulfide). The reaction solutions were prepared by initially mixing together pH buffer, electrolyte, and FeCl2 or other metal ion/ reductant from aqueous stock solutions. After equilibrating overnight, an appropriate amount of carbamate was added from an aqueous stock solution to initiate the reaction. The initial pesticide concentration in most kinetic experiments was 25 µM, and FeII was present in considerable excess (g0.5 mM). Aliquots of solution were then periodically collected for analysis as the reaction proceeded. To quench reactions, aliquots were filtered (0.22 µm Millipore Millex-GS) to remove colloidal FeIII solids, were passed through a cation-exchange resin (OnGuard-H; Dionex Corp., Sunnyvale, CA) to remove soluble FeII, and were acidified with HCl. In cases where reactions could not be quenched, analysis was carried out immediately after collecting each aliquot. Reactions were monitored for three months or through approximately three half-lives, whichever was shorter. At least five aliquots were collected during this time period; several more were collected for most reactions. Metal ion-free control experiments were conducted to account for non-reductive loss processes, most notably base-catalyzed elimination. pH stability was verified by periodic measurement (Fisher Accumet 825MP meter with Orion combination semi-micro probe; NIST-traceable standards) throughout the course of each reaction. The loss of parent carbamate and the formation of selected reaction products was monitored by reverse-phase HPLC with UV detection at 233 nm. An HPLC system (Waters Corp., Milford, MA) was used with a Spherisorb S5ODS-2 column (4.6 × 150 mm column and 4.6 × 10 mm guard column containing 5 µm C8 packing material). The flow rate was set at 1 mL/min and an isocratic eluent (5 mM acetic acid in 15% acetonitrile/85% H2O v/v) was used. The injected volume was 200 µL. Chromatographic peaks for oxamyl, methomyl, oxamyl oxime, methomyl oxime, and DMCF were identified by retention time comparison with authentic standards at three different eluent ratios (7, 15, and 25% acetonitrile). Methylamine was identified in selected reaction solutions using HPLC with precolumn derivatization (adapted from Gui et al.; ref 29). Briefly, 2-mL aliquots were added to vials containing a large molar excess (>50×) of difluoronitrobenzene dissolved in 0.8 mL acetonitrile and 0.9 mL of 0.1 M NaOH. After vortexing and equilibrating for at least five minutes, samples were analyzed by HPLC with UV detection at 380 nm. Separation of 20 µL samples was achieved using a µBondapak column (3.9 × 300 mm column and Guard Pak precolumn insert containing 10 µm C18 packing material) and an isocratic eluent (17 mM acetic acid in 50% methanol/ 50% H2O v/v) at a flow rate of 1 mL/min. Methanethiol and dimethyl disulfide were identified in selected reactions by GC-MS headspace analysis (HewlettPackard 6890 gas chromatograph equipped with a 5973 mass

FIGURE 2. Oxamyl degradation in the presence of various dissolved metal ions. Reaction conditions: 1 mM metal ion, 25 µM oxamyl, pH 5.5 (50 mM MES buffer), 25 °C, anoxic unless otherwise indicated. 0.1 M NaCl added to the CuI reaction to increase metal ion solubility. selective detector; Palo Alto, CA). Reactor headspace gas samples (100 µL) were injected in split mode (5:1 ratio) onto a Phenomenex (Torrance, CA) ZB-5 capillary column (30 m × 0.25 mm i.d. with 0.25 µm thick 5% phenylpolysiloxane film). Conditions were helium gas flow of 0.6 mL min-1; injection port temperature set at 270 °C; constant oven temperature of 50 °C; interface temperature set at 285 °C; ion source temperature set at 230 °C; electron multiplier voltage set at 1800 V; and electron energy set at 70 eV. Mass spectra were scanned every 0.6 s from mass 43 to 100. FeII oxidation kinetics were monitored in selected reactions using a colorimetric method adapted from Stookey (30). Samples (0.3 mL) were mixed with 10 mL of a solution containing 1 mM ferrozine and 0.5 M MOPS buffer (pH 7.1). The mixture was then filtered (0.22 µm Millipore Millex-GS) to remove colloidal FeIII, and absorbance of the filtrate was measured at 562 nm (UV-160 UV/VIS spectrophotometer; Shimadzu, Kyoto, Japan). The presence of a radical intermediate was evaluated using the radical scavenger acrylonitrile. Aqueous solutions containing 5 mM FeII (or CuI) + 2.5 mM carbamate + 5 mM acrylonitrile were prepared. These solutions, as well as metalfree and carbamate-free controls, were then visually examined for evidence of monomer polymerization after 24 h.

Results and Discussion Survey of Reductants and Metal Ions. Initially, several inorganic and organic compounds were surveyed in single reagent experiments to evaluate their ability to promote oxamyl degradation in anoxic solution. These chemicals are representative of naturally occurring inorganic reductants (1 mM bisulfide, hydroxylamine), aromatic organic reductants (1 mM hydroquinone, catechol), carboxylic acids (1-5 mM acetate, oxalate, malonate, citrate, lactate), and organosulfur reductants (1 mM L-cysteine, 3-mercaptopropionate). However, results from these experiments indicate that the rate of oxamyl degradation is unchanged, relative to the reagent-free blank, in the presence of these compounds. Natural waters, both anoxic and oxic, also contain a variety of metal ions that may facilitate carbamate degradation. Figure 2 shows the results of an experiment in which oxamyl concentration was monitored for 9 days in the presence of the following metal ions: 1 mM CaII, MgII, MnII, FeII, CoII, NiII, CuII, CuI, ZnII, and PbII. Of the metal ions tested, only

three (CuI, FeII, and CuII) accelerate oxamyl degradation relative to the metal ion-free blank solution; the most rapid degradation is observed for CuI. It is also notable that oxamyl degradation in the presence of CuII is much faster in anoxic solution than in an air-saturated solution of the same composition. Reaction Products. With the exception of solutions containing CuII, CuI, and FeII, carbamate degradation observed in this study results in stoichiometric production of the corresponding oxime and methylamine, as indicated in eq 1. Hence, we conclude that base-catalyzed elimination is the only reaction by which oxamyl and methomyl degrade in iron-free and copper-free anoxic solutions. In addition to the elimination products mentioned above, a nitrile and two sulfur-containing organic products are detected in anoxic solutions containing FeII. Stoichiometric production of a nitrile, N,N-dimethyl-1-cyanoformamide (DMCF), is observed for oxamyl, and stoichiometric production of methylamine is observed for both carbamates using HPLC-UV. The equivalent nitrile for methomyl, acetonitrile, cannot be detected with our HPLC setup. Reactor headspace gases were analyzed after allowing sufficient time for complete degradation of the carbamates. GC-MS analysis identified methanethiol (H-S-CH3) and dimethyl disulfide (H3C-S-S-CH3) as products in reactors containing FeII. Total sulfur equivalents detected in these reactions ([H-S-CH3] + 2[H3C-S-S-CH3]) are roughly equivalent to the initial concentration of parent carbamate, indicating that these compounds provide a good mass balance for sulfur (each parent carbamate molecule contains one sulfur equivalent). However, relative concentrations of the two sulfur-containing products vary with pH. At pH 5.5, carbamate degradation results in stoichiometric production of methanethiol and only traces of dimethyl disulfide. In contrast, at pH 2.0, dimethyl disulfide is higher in concentration than methanethiol. We presume methanethiol formation results directly from carbamate reduction, according to eq 2, while dimethyl disulfide forms by a subsequent redox reaction between methanethiol and FeIII:

2 H-S-CH3 + 2 FeIII f H3C-S-S-CH3 + 2 FeII + 2 H+ (3) FeIII has been shown to oxidize a variety of aromatic and aliphatic thiol compounds to their corresponding disulfides according to the stoichiometry depicted in eq 3 (31). FeII production consistent with eq 3 is observed when methanethiol and FeIII are mixed together, and the rate of FeII production is significantly faster at pH 2.0 than pH 5.5 (see Figure S1 in Supporting Information). The higher rates of FeIII reduction and dimethyl disulfide formation observed at pH 2.0 is likely a reflection of the relative reactivity of dissolved FeIII versus precipitated FeIII with methanethiol; FeIII solubility increases markedly with decreasing pH (32). The same carbamate degradation products are observed for solutions containing CuI as for FeII. Therefore, we conclude that carbamate reduction can be coupled with either FeII or CuI oxidation. This is the first report, to our knowledge, of oxime carbamate reduction by CuI. For oxamyl degradation in aerobic solutions containing CuII, only E1cb elimination products are observed (eq 1). However, the rate of degradation is significantly faster than in CuII-free solutions of the same composition (a significant difference is apparent when reactions are monitored for three months). These findings are consistent with earlier studies that attribute CuII-facilitated degradation of oxamyl and related carbamates to metal ion catalysis of the E1cb reaction (9, 33). For oxamyl degradation in anoxic solutions containing CuII, oxamyl oxime, DMCF and dimethyl disulfide are VOL. 35, NO. 12, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 3. Representative time courses for oxamyl degradation in the presence of FeII and CuI. Reaction conditions: (A) 0.5 mM FeII, 25 µM oxamyl, pH 5.5 (25 mM MES buffer), 25 °C, 0.1 M NaCl; (B) 0.5 mM FeII, 25 µM oxamyl, pH 6.1 (25 mM MES), 25 °C, 0.1 M NaCl; (C) 0.5 mM FeII, 25 µM oxamyl, pH 7.1 (25 mM MOPS), 25 °C, 0.1 M NaCl; (D) 0.5 mM CuI, 25 µM oxamyl, pH 5.5 (25 mM MES), 25 °C, 0.1 M NaCl. Calculated rate constants: (A) kred ) 4.60 × 10-3 h-1, kelim , kred (elimination too slow to observe on the time scale of reduction); (B) kred ) 4.76 × 10-3 h-1, kelim ) 4.45 × 10-4 h-1; (C) kred ) 6.47 × 10-3 h-1, kelim ) 4.47 × 10-3 h-1; (D) kred ) 1.21 × 10-1 h-1, kelim , kred. detected. It is unclear how CuII facilitates production of oxamyl reduction products. Preliminary evidence, including autocatalytic DMCF production, suggests that CuII is somehow reduced to CuI (possibly by one of the intermediates or products of the elimination reaction), and subsequent reoxidation of CuI is coupled with reduction of oxamyl. Dimethyl disulfide is then formed by a reaction analogous to eq 3, in which CuII acts as the one-electron oxidant. No CuII-catalyzed elimination or reduction is observed for methomyl under both aerobic and anoxic conditions. The remainder of this study focuses on characterizing the reaction between carbamates and reduced metal ions (FeII and CuI). The reaction between carbamates and FeII is of particular importance because FeII is one of the most abundant reducing agents present in soils and groundwater. CuII-facilitated carbamate degradation will be the subject of a future investigation. FeII: Reaction Time Courses. Figure 3A-C illustrates typical reaction time courses for oxamyl in the presence of excess FeII. Using the HPLC method described above, we are able to simultaneously follow oxamyl disappearance and the formation of two reaction products: (a) oxamyl oxime, resulting from the E1cb elimination reaction shown in eq 1, and (b) DMCF, a nitrile resulting from the redox reaction shown in eq 2. For methomyl (time courses provided in Figure S2 of Supporting Information), the analogous oxime and nitrile products are methomyl oxime and acetonitrile, respectively. Only oxime elimination products are observed in FeII-free solution (FeII-free oxamyl time courses provided in Figure S3 of Supporting Information). For oxamyl, a mass balance can be calculated for a portion of the molecule, shown within the dotted lines below, by summation of the concentrations of HPLC-monitored compounds ([oxamyl] + [oxamyl oxime] + [DMCF]):

Nearly 100% mass balance on this portion of the oxamyl molecule is maintained throughout all reactions monitored, as illustrated in Figure 3A-C. As mentioned above, aceto2464

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FIGURE 4. Representative time courses that illustrate the stoichiometry for FeII reacting with (A) oxamyl and (B) methomyl. Initial reaction conditions: 1 mM FeII, 0.5 mM carbamate, pH 5.5 (25 mM MES buffer), 25 °C, 0.1 M NaCl. nitrile cannot be detected by HPLC-UV, so a similar mass balance cannot be verified for methomyl. Reduction and elimination products form by independent processes, so it is not surprising that their rates of formation (kred and kelim, respectively; see reaction kinetics below) exhibit differing dependencies on solution conditions. The rate of elimination increases with increasing pH and is unchanged by the presence of FeII. In contrast, the rate of reduction is relatively independent over the pH range 2.0-7.1 (pH 2.0 time course similar to Figure 3A) and is proportional to the FeII concentration (rate dependence on [FeII] and [carbamate] discussed below). FeII: Reaction Stoichiometry. Redox stoichiometry was evaluated by simultaneously monitoring the disappearance of 1 mM FeII and 0.5 mM carbamate, as well as the formation of DMCF. Figure 4 shows the time courses observed at pH 5.5; two moles of FeII are oxidized for each mole of carbamate reduced. This net stoichiometry, which is also observed at pH 7.1, agrees with the hypothesized reaction pathway, a two-electron-transfer process. At pH 3.0, however, a different net stoichiometry is observed. Under these conditions, only 0.6 mol of FeII are oxidized for each mole of carbamate reduced. This deviation can be explained, in part, by the rereduction of FeIII by methanethiol according to eq 3. Deviation from the ideal stoichiometry is not observed at pH 5.5 and 7.1, though, because the rate of FeIII reduction by methanethiol decreases with increasing pH. FeII: Radical Intermediate. The reduction of organic compounds by one-electron reductants, such as FeII, often results in unstable radicals as reaction intermediates (34). The presence of radicals can be inferred by observing polymerization of selected radical-scavenging monomers added to a reaction solution (35-37). Acrylonitrile polymerization is induced in solutions containing FeII + carbamate,

but not in either FeII- or carbamate-free control reactions. Furthermore, polymerization is not observed in solutions containing FeIII + methanethiol, demonstrating that the reaction depicted in eq 3 is not the source of the polymerinducing radical. Together, these observations suggest that carbamate reduction by FeII involves a free radical mechanism. FeII: Reaction Kinetics. It has been shown above that carbamate degradation in the presence of FeII occurs by parallel elimination and reduction pathways. Carbamate loss can be described by a rate expression that accounts for both reactions:

-

d[carbamate] ) dt kelim[carbamate] + k2[FeII]R[carbamate]β (4)

where [carbamate] and [FeII] represent the molar concentrations of parent carbamate and FeII, respectively, kelim (h-1) is the pseudo first-order rate constant for carbamate elimination, k2 is the rate constant for carbamate reaction with FeII, and R and β are the reaction orders with respect to FeII and carbamate concentration, respectively. Because FeII oxidation is coupled with carbamate reduction, it follows that

-

d[FeII] ) 2k2[FeII]R[carbamate]β dt

(5)

under conditions where FeIII reduction by methanethiol is not significant. The factor 2 reflects the observed stoichiometry of FeII oxidation by carbamates when the first oxidation step is rate limiting. If the redox reaction is first order in FeII and carbamate concentration (i.e., R ) β ) 1), then k2 is a second-order rate constant with units of M-1 h-1. Depending on the solution conditions, eqs 4-5 can be further simplified. If FeII is present in considerable excess of carbamate, and if carbamate loss is assumed to be a pseudo first-order process (β ) 1), then the following expression for carbamate loss is obtained:

-

d[carbamate] ) kobs[carbamate] ) dt (kelim + kred)[carbamate] (6)

where kobs (h-1) is the overall pseudo first-order rate constant for carbamate loss and kred (h-1) is the pseudo first-order rate constant for carbamate reduction. Under these conditions, expressions for the formation of oxime elimination and nitrile reduction products can also be defined:

d[oxime] ) kelim[carbamate] dt

(7)

d[nitrile] ) kred[carbamate] dt

(8)

By fitting product formation data along with parent compound loss data, we are better able to assess the importance of each reaction to overall carbamate degradation. Likewise, if carbamate is present in considerable excess of FeII and if FeII oxidation is assumed to be a pseudo firstorder process (R ) 1), then eq 5 can be simplified:

-

d[FeII] ) kox[FeII] (9) dt

(9)

where kox (h-1) is the pseudo first-order rate constant for FeII oxidation.

kelim, kred, and kox were calculated for individual batch reactions using the software package Scientist for Windows (v. 2.01; Micromath, Salt Lake City, UT). Scientist calculates these parameters by fitting (method of least squares) numerically integrated solutions of the system of differential rate expressions (eqs 6-8 for oxamyl degradation; eqs 6-7 for methomyl degradation; eq 9 for FeII oxidation) to experimental data for an individual reaction. Given our experimental setup and the limitations of our analytical methods, only rate constants greater than 7.0 × 10-5 h-1 can be reliably calculated. Model fits for carbamate degradation (Figure 3 and Figures S2 and S3 in Supporting Information) indicate that parent compound loss and reaction product formation are adequately fit by a kinetic model that accounts for parallel elimination and reduction pathways. Likewise, experimental data for FeII oxidation is adequately fit by pseudo first-order kinetics (Figure S4 in Supporting Information). It is also worth noting that kelim values determined in the presence of FeII agree quite well with values determined in FeII-free solutions of the same pH. Values of R, β, and k2 were obtained by measuring kred values for oxamyl and methomyl as a function of excess FeII concentration, and by measuring kox for FeII as a function of excess carbamate concentration. Slopes of unity are observed in log-log plots of kred vs [FeII] and kox vs [carbamate] (Figure 5A,B, respectively), demonstrating that the reactions are firstorder with respect to both FeII and carbamate concentration (i.e., R ) β ) 1). Analysis of this same data in terms of eqs 4 and 5 results in k2 values of 9.29 (( 0.06) M-1 h-1 for oxamyl and 27.1 (( 0.2) M-1 h-1 for methomyl (uncertainties represent one standard deviation). These values agree quite well with half-life measurements reported in Bromilow et al. (13). After correcting for temperature, their half-life measurements correspond to k2 values of 11 M-1 h-1 for oxamyl and 24 M-1 h-1 for methomyl. It is not entirely clear why FeII is more reactive with methomyl than oxamyl (k2oxamyl ≈ 1/3 k2methomyl). The structures of the two carbamates differ only in the substituent represented by R in eqs 1-2 (Roxamyl ) -(CdO)N(CH3)2; Rmethomyl ) -CH3). Electron withdrawing groups adjacent to the reaction center should stabilize electron-rich radical intermediates and reduced products; kinetics are related to the stability of these species if the transition state is intermediate-like or product-like in nature. Using a Taft σ* coefficient approach, Roxamyl is more electron-withdrawing than Rmethomyl (38). Furthermore, a conjugated π bonding arrangement created by Roxamyl (OdC-CdN) is expected to provide additional resonance stability to a radical intermediate and the reduced nitrile (34). Therefore, electronic substituent effects predict relative rates of carbamate reduction that are opposite of our experimental observations. This suggests that other factors play a role in determining the relative reactivity of the two carbamates. For example, if bond cleavage occurs in a concerted step with electron transfer, then the influence of electronic substituent properties on the leaving group of interest needs to be considered; electrondonating substituents may enhance leaving group departure. Comparing a larger set of structural analogues would greatly improve our understanding of the relationship between carbamate chemical structure and their reactivity with FeII. FeII: Effect of Temperature and O2. The temperature dependence of kred follows Arrhenius behavior for both carbamates. kred increases by a factor of 13 when the temperature is raised from 16 to 45 °C. The apparent activation parameters were determined according to transition-state theory (34). The activation enthalpies (∆H‡), calculated from the slopes of the plots shown in Figure 5C, are 69.1 (( 1.1) kJ mol-1 for oxamyl and 66.0 (( 1.4) kJ mol-1 for methomyl. The activation entropies (∆S‡), calculated at VOL. 35, NO. 12, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 6. Effect of FeII coordination to EDTA4- and to CN- on the rate of oxamyl reduction. Conditions are such that >99% of FeII is complexed and all six of the FeII coordination positions are occupied by the complexing ligands. Reaction conditions: 25 µM oxamyl, pH 5.5 (25 mM MES buffer), 25 °C, FeII and ligand concentrations provided in figure legend.

FIGURE 5. (A) Relationship between kred and [FeII] when FeII is present in considerable excess, (B) relationship between kox and [carbamate] when carbamate is present in considerable excess, and (C) plot for temperature dependence of kred. Reaction conditions: (A) 0.5-10 mM FeII, 25 µM carbamate, 25 °C, pH 5.5 (25 mM MES buffer), 0.1 M NaCl; (B) 0.5-8 mM carbamate, 100 µM FeII, 25 °C, pH 5.5 (25 mM MES), 0.1 M NaCl; (C) 16-45 °C, 4 mM FeII, 100 µM carbamate, pH 5.8 (50 mM MES), 0.1 M NaCl. 25 °C from the slopes and intercepts of Figure 5C, are 28.2 (( 6.3) J mol-1 deg-1 for oxamyl and 25.1 (( 7.8) J mol-1 deg-1 for methomyl. At 25 °C, the enthalpic contribution to the activation free energy (∆G‡) is much larger than the entropic contribution (∆G‡ ) ∆H‡ - T∆S‡). The activation parameters for the two carbamates are not significantly different from one another, supporting our hypothesis that both compounds are reduced by the same mechanism. At the same time, the similarity in activation parameters provides no information to help us rationalize the relative reactivities of the two compounds. The presence of dissolved oxygen has no direct effect on the rate of oxamyl reduction by FeII. At pH 2, the rates of oxamyl disappearance and DMCF formation measured under air-saturated conditions (0.21 atm O2 in reactor headspace) are unchanged from the rates measured in anoxic solution of the same composition. The competing reaction between 2466

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FeII and O2 is extremely slow under acidic conditions; it can be ignored on the time scale of the reaction with oxamyl (39, 40). However, the rate of FeII oxygenation increases dramatically with increasing pH (39, 40). As a result, the presence of O2 indirectly inhibits carbamate reduction by FeII under neutral and alkaline pH conditions. O2 does influence the production of volatile sulfurcontaining products. Almost no methanethiol and little dimethyl disulfide are detected in the headspace of O2containing reaction solutions. This indicates that methanethiol reacts with O2 to yield final product(s) other than dimethyl disulfide, or that dimethyl disulfide undergoes further oxidation (possibly to less volatile sulfoxides and sulfones) (31). Is FeII-Carbamate Coordination Required? Some of the organic and inorganic reductants surveyed in this work have a lower standard reduction potential (E°H) than dissolved FeII and CuI. It is interesting that none of these reagents are able to reduce oxamyl; kinetic factors, not thermodynamics, must be limiting the reaction with non-metal reductants. One possible kinetic factor is the formation of inner-sphere bonds between the reduced metal ion and ligand donor groups present on the carbamate molecule. These bonds would provide a unique route for electron transfer from the metal ion to the carbamate molecule; this pathway would not be available for non-metal ion reducing agents. Oximes and oxime esters are well known for their ability to form coordination complexes with transition metals (41). In addition, metal ion coordination by oxime esters has often been considered an important factor in metal ion-catalyzed hydrolysis of these compounds (9, 33, 42). No direct evidence for the role of metal ion-carbamate coordination was obtained in this study. We expect that any bonding between FeII and Lewis base groups present on the carbamates would be weak and difficult to detect by spectroscopic methods (33). Indirect evidence, however, indicates that metal ion coordination is not an essential component of the reaction mechanism. If it were essential, we would expect inhibition of the reaction when all six FeII inner-sphere coordination positions are occupied by competing ligands (thus not allowing FeII-carbamate bonding) (43). Instead, results shown in Figure 6 demonstrate that coordinatively-saturated FeII is capable of reducing oxamyl. In fact, oxamyl reduction by FeII(CN)64-, a substitution-inert complex (44) which cannot form coordination bonds with carbamate Lewis base groups, is significantly faster than

reduction by uncomplexed FeII (predominantly FeII(H2O)62+, a substitution-labile species at pH 5.5). Furthermore, the rate of reduction by FeII is unchanged in the presence of a 10-fold excess of EDTA, a strongly complexing ligand [>99.99% of FeII will be bound as a hexadentate complex, FeII(EDTA)2- under conditions shown; ref (45)]. These observations do not rule out the formation of inner-sphere FeII-carbamate coordination bonds for other FeII species. They do indicate, though, that inner-sphere coordination bonds are not required for electron transfer from FeII to the carbamate. FeII: Reaction Mechanism. Figure 7 illustrates two alternative mechanisms for carbamate reduction by FeII. Both mechanisms are consistent with reaction products, intermediates, stoichiometry, and kinetics observed in this study as well as Bromilow et al. (13). The first two steps are the same for both mechanisms. The first step is the formation of an FeII-carbamate precursor complex. No direct evidence has been obtained regarding the nature (i.e., inner- vs outersphere) or structure of the precursor complex. Indirect evidence mentioned above, however, demonstrates that inner-sphere complex formation between FeII and the carbamate molecule is not necessary for the reaction to occur. Following precursor complex formation, the second step in both mechanisms is the initial electron transfer from FeII to the carbamate, resulting in an FeIII-carbamate successor complex. The two mechanisms differ in the steps following the initial electron-transfer step. In “mechanism A”, the third step involves dissociation of the successor complex and oxime ester bond cleavage to form an imino radical, FeIII, and N-methylcarbamic acid. N-methylcarbamic acid then rapidly degrades to form methylamine and carbon dioxide (10). In the fourth step, the imino radical rapidly reacts with a second FeII ion resulting in the corresponding imino anion and FeIII. To be complete, this step must also involve the formation of a precursor complex and the dissociation of a successor complex. The final step is then elimination of the imino anion, via thio-methyl cleavage, to form a substituted nitrile and methanethiol. “Mechanism B” differs from “mechanism A” only in that thio-methyl cleavage occurs following the first electrontransfer step (resulting in a carbon-centered radical), and oxime ester bond cleavage occurs following the second electron-transfer step. Bromilow et al. (13) presumed that the reaction occurs according to mechanism A. Unambiguous assignment of one mechanism in favor of the other, however, will require more experimental information on the structure of the radical intermediate(s). Experimental observations are consistent with either of the first two steps being rate determining. Both are bimolecular elementary reactions involving a single FeII ion and a carbamate molecule, consistent with the observed rate dependence on reactant concentration. Furthermore, the time course for FeII disappearance parallels carbamate disappearance and reduction product formation (see Figure 4); this is consistent with a mechanism in which all electron transfer steps occur either during or following the ratedetermining step. Formation of a precursor complex would only be rate limiting if an inner-sphere complex forms and if ligand exchange rates for FeII are slow. We demonstrated above that an inner-sphere complex is not required for electron transfer, and with the exception of ligands such as CN-, which forms a low-spin complex that is substitution inert, FeII undergoes very rapid ligand exchange (46). Altogether, this information leads us to believe that the initial electron transfer is the rate-determining step in the reaction mechanism. CuI: A Comparison. A limited number of carbamate degradation experiments were also conducted in the presence

of CuI. Results from these experiments indicate that the carbamate reduction mechanism described above applies for CuI as well as FeII. It was already mentioned that the same degradation products are observed for CuI as FeII. In addition, just as with FeII, a free radical intermediate is inferred from acrylonitrile polymerization, and pseudo first-order oxamyl disappearance in the presence of excess CuI results in stoichiometric production of DMCF (Figure 3D). Likewise, kred is relatively independent of pH and follows a first-order dependence on CuI concentration. These results suggest that the kinetic model described for FeII (eqs 4-9) is valid for CuI as well. If we assume this is true, k2 for CuI is estimated to equal 243 (( 7) M-1 h-1 for oxamyl and 117 (( 3) M-1 h-1 for methomyl. The k2 values measured for CuI are considerably larger than those measured for FeII under equivalent solution conditions. The difference can be explained if the relative reactivity of carbamates with different metal ion reductants is related in some manner to the thermodynamic redox activity of the metal ions. Under the conditions examined (pH 5.5, 0.1 M NaCl), CuI speciation is dominated by the formation of various chloro complexes (principally CuCl2and CuCl0), while hexaquo Fe2+ is the predominant FeII species. Reduction potentials for the half reactions involving the chloro complexes of CuI (E°H ) 0.33 and 0.46 V for CuCl0/ CuCl+ and CuCl2-/CuCl20, respectively) are lower than the reduction potential for the half reaction involving hexaquo FeII (E°H ) 0.77 V for Fe2+/Fe3+) (45, 47). Therefore, based on thermodynamics, the predominant CuI species are stronger reductants than Fe2+. It is also worth noting that k2 for CuI reacting with oxamyl is approximately double the value for the reaction with methomyl. This is reversed from the trend observed for FeII, where k2 measured for oxamyl is only one-third the value for methomyl. The trend observed for CuI is consistent with inductive and resonance electronic contributions of the R group substituents discussed earlier. Nonetheless, the reversal of trends from FeII to CuI suggests that other factors might play a role in determining the relative reactivity of the two carbamates. Possible factors might include varying degrees of steric inhibition or orbital overlap for each metalcarbamate combination. Environmental Significance. The environmental fate of agrochemicals is controlled by a number of abiotic and biological processes. Results from this investigation demonstrate that reduction of oxamyl and methomyl can be coupled with the oxidation of either FeII or CuI. FeII is abundant in many anoxic and suboxic environments. As a result, this reaction likely plays a significant role in the environmental degradation of these and related agrochemicals. In fact, reduction is the primary route of oxamyl degradation observed in anaerobic saturated subsoils collected from field settings (13, 48). Other reducing agents that may be found in these settings, such as bisulfide and dihydroxybenzenes, fail to react with these pesticides. Consequently, reaction with FeII may be the dominant abiotic sink for these compounds in reducing environments. Determining the factors that control reaction rates is of great importance. FeII speciation is a factor that is likely to be of major importance in determining carbamate reduction kinetics. It should be stressed that the kinetic parameters reported in this study and in Bromilow et al. (13) are only valid for a narrow range of conditions (pH 2.0-7.1, Fe2+ species predominates). Several studies have reported that FeII speciation has a dramatic influence on its reactivity with a variety of oxidants, including O2, H2O2, and CrVI (e.g., see refs 49-52). Thus, assessing the influence of FeII speciation on carbamate reduction kinetics is of great interest; future papers in this series will address the topic in detail. VOL. 35, NO. 12, 2001 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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FIGURE 7. Proposed mechanisms for carbamate reduction by FeII (adapted from Bromilow et al.; ref 13).

Acknowledgments Financial support for T.J.S. was provided by the U.S. Environmental Protection Agency, through a STAR Fellowship. Additional support was provided by Grant R82-6376, U.S. Environmental Protection Agency. DuPont Crop Protection (Wilmington, DE) provided samples of oxamyl, methomyl, oxamyl oxime, methomyl oxime, and DMCF. GCMS was conducted with the help of E. Silveira (DuPont) and D. Carlson (Johns Hopkins). The manuscript benefited greatly from comments and discussion provided by K. Karlin, A. L. Roberts, P. Vikesland, and C. Whitehead (Johns Hopkins), B. Nowack (E.T.H.), R. Warren, A. Barefoot, and A. Taylor (DuPont), and four anonymous reviewers.

Supporting Information Available A complete list of rate constants (Table S1), as well as time courses for FeIII reduction by methanethiol (Figure S1), methomyl degradation in solutions containing FeII (Figure S2), oxamyl degradation in reagent-free solutions (Figure S3), and FeII oxidation by oxamyl and methomyl (Figure S4) are provided. This material is available free of charge via the Internet at http://pubs.acs.org.

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Received for review October 30, 2000. Revised manuscript received March 16, 2001. Accepted March 19, 2001. ES001824J

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