Reduction of Oxamyl and Related Pesticides by FeII: Influence

Reduction of Oxamyl and Related Pesticides by Fe: Influence of Organic Ligands .... Role of Fe(III), Phosphate, Dissolved Organic Matter, and Nitrate ...
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Environ. Sci. Technol. 2002, 36, 5172-5183

Reduction of Oxamyl and Related Pesticides by FeII: Influence of Organic Ligands and Natural Organic Matter TIMOTHY J. STRATHMANN* AND ALAN T. STONE Department of Geography and Environmental Engineering, Johns Hopkins University, Baltimore, Maryland 21218

The reduction of oxamyl and related oxime carbamate pesticides (OCPs; methomyl and aldicarb) by FeII is an important pathway for the degradation of these compounds in soil and groundwater. A series of batch kinetic experiments was carried out to assess the effects that selected carboxylate and aminocarboxylate ligands have on these reactions. In the absence of FeII, no OCP reduction by the ligands is observed. In the presence of FeII, the rate of OCP reduction varies by several orders of magnitude and can be described by the expression kred ) [FeII]∑ikiRi, where kred is the observed pseudo-first-order rate constant for OCP reduction, [FeII] is the total FeII concentration, Ri is the fraction of each FeII species in solution, and ki is the second-order rate constant for OCP reduction by each FeII species. The reactivity of individual FeII species is dependent upon the standard one-electron reduction potential of the corresponding FeIII/FeII redox couple (EH°) and the availability of inner-sphere FeII coordination sites for bonding with Lewis base donor groups within the OCP structure. A linear free energy relationship is proposed. Kinetic measurements demonstrate that natural organic matter from the Great Dismal Swamp facilitates OCP reduction by FeII in the same manner as the individual organic ligands. Results from this study improve our understanding of the pathways and rates of pesticide degradation in reducing subsurface environments, especially those rich in organic matter.

Introduction Oxamyl and related oxime carbamate pesticides (OCPs) are heavily used in agriculture for crop protection purposes (1). They are especially effective at controlling plant-parasitic nematodes (2). OCPs are extremely toxic to both insects and mammals (3). Their toxicity and efficacy both result from their ability to inhibit the activity of the enzyme acetylcholinesterase (3); inhibition of the enzyme results in buildup of acetylcholine, a neurotransmitter, to toxic levels. Groundwater contamination by pesticides has been of increased concern in recent years (4). A large number of pesticides, including OCPs, have been detected in groundwaters that serve as sources of drinking water (4). Reported OCP concentrations have often been found to exceed recommended health guidelines (4-6). Concern about groundwater contamination by OCPs is especially high because their properties (high aqueous solubility and low * Corresponding author e-mail: [email protected] 5172



sorption to most soils) suggest a high degree of mobility in soils and groundwater (7). Little is known about the molecular mechanisms that control pesticide fate in subsurface environments. Particularly lacking is information on pesticide transformation processes that occur under suboxic/anoxic conditions that prevail in many soil and groundwater environments, including anoxic microenvironments that exist within oxic settings. Several different reducing agents are ubiquitous in anoxic/suboxic settings (8, 9). Reactions with these compounds may play a significant role in determining the environmental fate of many pesticides. Among the most abundant reducing agents are various dissolved and solid-phase FeII species (8, 9). FeII has been shown to react with a variety of pollutants, including nitrate and nitrite (10, 11), nitroaromatics (12, 13), halogenated organics (14-16), CrVI (17-20), UVI (21), and SeVI (22). FeII may also react with widely used pesticides, many of which possess more involved chemical structures containing multiple functional groups. The complex nature of pesticide structures often makes it difficult to predict a priori whether they will react with FeII or what products will be produced in such a reaction. The reduction of oxamyl and related OCPs by FeII has previously been reported (23-25). In the presence of FeII, oxamyl undergoes a net two-electron reduction that is coupled with the one-electron oxidation of two FeII ions (24); the observed reduction products are N,N-dimethyl-1-cyanoformamide (DMCF), methanethiol, methylamine, and CO2:

Results indicate that reaction with FeII is a major degradation pathway for oxamyl and related pesticides in reducing environments. This is especially true considering the lack of reactivity observed between oxamyl and other common reductants (e.g., bisulfide, cysteine, dihydroxybenzenes) found in these settings (24). We reported in an earlier study that selected inorganic ligands have a marked effect on the rate of oxamyl reduction by FeII (25). Reaction kinetics can be explained in terms of changing FeII speciation. The overall pseudo-first-order rate constant for oxamyl reduction (kred) is equal to the weighted sum of the rate constants for individual FeII species reacting with the pesticide:

∑k R

kred ) [FeII]

i i



where [FeII] is the total FeII concentration, Ri is the fractional concentration of FeII species i, and ki is the second-order rate constant for species i reacting with oxamyl. The logarithmic ki values correlate with the standard one-electron reduction potentials (EH°) of the corresponding FeIII/FeII redox 10.1021/es0205939 CCC: $22.00

 2002 American Chemical Society Published on Web 10/31/2002

TABLE 1. Organic Ligands Included in This Study

couples in a linear free energy relationship (LFER). Similar LFERs have been reported for the reactivity of FeII with O2 (26, 27), H2O2 (27), and CrVI (19). In addition to common inorganic ligands, agricultural subsurface environments often contain considerable concentrations of organic ligands that can alter the speciation and reactivity of FeII with OCPs. Microbial decomposition of plant and animal biomass releases a variety of low molecular weight organic ligands and polymeric humic substances into soil and groundwater (28). Microorganisms and plants also synthesize and excrete specialized ligands for a variety of purposes, including nutrient acquisition and toxic metal sequestration (29-31). In addition, many synthetic organic ligands that are recalcitrant to conventional wastewater treatment processes (32) migrate into soil and groundwater. Furthermore, metal chelates are often added directly to fertilizers to facilitate plant uptake of micronutrient metal ions (2). In this study, we examine the effects that representative low molecular weight organic ligands (Table 1) as well a natural organic matter sample collected from the Great Dismal Swamp, in Virginia, have on the reactivity of FeII with oxamyl and two related OCPs (methomyl and aldicarb). In keeping with the theme of our previous work with inorganic ligands (25), emphasis is placed on quantifying the reactivity of OCPs with individual FeII-organic ligand complexes and assessing the molecular-scale factors that control FeII reactivity.

Materials and Methods Chemical Reagents. All chemical reagents were of the highest purity available. Oxamyl, methomyl, oxamyl oxime, methomyl oxime, and N,N-dimethyl-1-cyanoformamide (DMCF) were provided by DuPont Crop Protection (Wilmington, DE). Aldicarb was provided by Rhoˆne-Poulenc Agriculture Limited (Essex, England). FeCl2‚4H2O, NaCl, NaOH, sodium acetate, citric acid, malonic acid, oxalic acid, disodium ethylenediaminetetraacetate (EDTA), disodium nitrilotriacetate (NTA), 2-(N-morpholino)ethanesulfonic acid monohydrate (MES buffer), and 3-(2-pyridyl)-5,6-bis(4-phenylsulfonic acid)1,2,4-triazine monosodium salt monohydrate (ferrozine colorimetric reagent) were purchased from Aldrich (Milwaukee, WI). Acetic acid, iminodiacetic acid (IDA), 3-(Nmorpholino)propanesulfonic acid (MOPS buffer), and N-tris(hydroxymethyl)methyl-3-aminopropanesulfonic acid (TAPS buffer) were purchased from Sigma (St. Louis, MO). Trimethylenediamine-N,N,N′,N′-tetraacetic acid (TMDTA) was obtained from Fluka (Buchs, Switzerland). Acetonitrile, disodium oxalate, HNO3, and HCl were purchased from J. T. Baker (Phillipsburg, NJ). Water collected from the Great Dismal Swamp (Suffolk, VA) was filtered through glass wool in the field, stored under

darkness at 4 °C, and then freeze-dried within 20 days. A stock solution of reconstituted Great Dismal Swamp water (RGDSW) was then prepared by dissolving 200 mg L-1 of the freeze-dried solid in deoxygenated aqueous solution and used without further modification (i.e., stock solution was not subsequently filtered or treated to remove metal ions). The RGDSW stock solution was determined to contain 150 mg of C L-1 total organic carbon (TOC; Phoenix 8000 UVPersulfate TOC analyzer; Dohrmann, Cincinnati, OH). Flame atomic absorption spectrophotometric analysis (AAS; Aanalyst 100, Perkin-Elmer, Norwalk, CT) indicated the presence of 55.2 µM iron, 270 µM sodium, 55.0 µM calcium, 27.7 µM magnesium, and 5.9 µM zinc in the RGDSW stock solution. Dissolved iron was determined to be 56% FeII (30.8 µM), using a ferrozine colorimetric method described previously (24). AAS measurements of manganese, copper, and silicon in the RGDSW stock solution were unchanged from Milli-Q water blanks. Kinetic Experiments. Details of the experimental setup were described previously (24). Because strict oxygen exclusion was required, all experiments were conducted inside a controlled-atmosphere glovebox (95% N2, 5% H2; Pd catalyst; Coy Laboratory Products, Grass Lake, MI). Kinetic experiments were conducted in batch reactors (250-mL polypropylene bottles mixed by magnetic stir plate) that were maintained under darkness in a constant temperature circulating water bath (25.0 ( 0.1 °C) contained within the glovebox. Batch reactions were prepared by initially mixing the appropriate deoxygenated aqueous stock solutions of pH buffer, electrolyte, ligand of interest, and FeCl2. After the mixture was equilibrated overnight, reactions were initiated by adding the appropriate OCP from a deoxygenated aqueous stock solution. Aliquots of solution were then periodically collected from the batch reactors for immediate HPLC analysis of OCPs and their reaction products. Samples collected from very fast reactions (t1/2 < 0.5 h) were immediately quenched by addition of a large excess of EDTA (early results demonstrated that FeII complexed by EDTA reacts very slowly with OCPs). The pH stability of each reaction was verified by periodic measurement (Fisher Accumet 825MP meter with Orion combination semimicroprobe; NIST standard buffers). Unless otherwise indicated, the initial concentration of FeII was 0.5 mM, and parent OCP was 25 µM. The specific pH and ligand concentrations examined varied from one ligand to another in an effort to obtain the widest possible variation in FeII speciation. A complete description of the solution conditions for all batch reactions is provided in the Supporting Information section (Table S1). Analytical. The extent of oxamyl and methomyl disappearance and reaction product formation was monitored by reverse-phase HPLC using the setup and method previously described (24). The extent of aldicarb disappearance was monitored using the same HPLC setup, with the exceptions that detection was at 210 nm and a different isocratic eluent was used (3 mM acetic acid in 30% acetonitrile/70% water). Data Analysis. Kinetic data from each reaction were analyzed using the software package Scientist (33). Pseudofirst-order rate constants for the reduction (kred, h-1) and E1cb elimination (kelim, h-1) of oxamyl and methomyl were determined for each batch reaction using data for both parent compound disappearance and reaction product appearance according to a procedure outlined previously (25) (details provided in the Supporting Information). Bromilow et al. (23) also reported FeII-facilitated redox breakdown of aldicarb, a related OCP. Although the mechanism proposed for aldicarb differs slightly from that proposed for oxamyl and methomyl, the rate-limiting step (transfer of the first electron from FeII to the OCP) is believed to be the same (23). For experiments conducted with aldicarb, only parent compound disappearVOL. 36, NO. 23, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY



FIGURE 1. Influence of carboxylate (A and B) and aminocarboxylate (C and D) ligands on the rate of oxamyl reduction by FeII. Error bars (smaller than symbol if not visible) indicate one standard deviation. Downward arrow for TMDTA indicates an upper limit estimate. Reaction conditions: 0.5 mM FeII, 25 µM oxamyl, 25 °C, pH and organic ligand concentration provided in figure, various pH buffers (complete list of conditions provided in Table S1 in Supporting Information). ance was monitored; no product formation data were obtained. Consequently, kred for aldicarb was calculated as the difference between kobs (pseudo-first-order rate constant for parent compound disappearance; h-1) measured in the presence of FeII and kobs measured in FeII-free solution at the same pH and temperature. Equilibrium FeII speciation was calculated using the software package HYDRAQL (34). Equilibrium expressions and the corresponding stability constants used for model input are provided in the Supporting Information (Table S2) along with appropriate references (35-38). Stability constants were also used to calculate the standard one-electron reduction potentials (EH°) for selected FeIII/FeII redox couples (FeIIIL + e- ) FeIIL) as previously described (25). Ionic strength corrections were made using the Davies equation (36). Scientist was also used to determine second-order rate constants for oxamyl reduction by individual FeII-organic ligand complexes (ki; M-1 h-1). Rate constants were determined by least-squares fit of eq 2 to measured values of kred and HYDRAQL calculations of FeII speciation. Details of the data fitting procedure are provided in the Supporting Information Section.

Results and Discussion A complete list of the results from 198 kinetic experiments is provided in the Supporting Information section (Table S1). Control Reactions. Results from a series of control experiments demonstrate that the organic ligands examined in this study are unreactive with the OCPs in the absence of FeII. In addition, the organic ligands have no effect on the rates of a parallel E1cb elimination reaction (discussed in detail in ref 24). Therefore, we conclude that the ligands 5174



examined in this study affect OCP degradation solely by altering FeII speciation. Oxamyl Reduction by FeII. Rates of oxamyl reduction by 0.5 mM FeII were measured in the presence of a series of structurally related carboxylate and aminocarboxylate ligands (see Table 1 for structures). Individual ligands were selected, and reaction conditions (e.g., ligand concentration, pH) were designed in an effort to obtain a wide variation in FeII species properties (e.g., EH° of FeIII/FeII redox couple, coordination geometry, FeII-ligand complex stability, species charge). Figure 1 is a group of log-log plots that illustrate the results of these experiments. Measured values of kred vary by nearly 6 orders of magnitude, depending upon (i) the identity of the ligand added to solution, (ii) the concentration of added ligand, and (iii) the solution pH. Figure 1A,B shows the effects of the carboxylate ligands acetate, malonate, oxalate, and citrate on oxamyl reduction kinetics. The overall effect on kinetics varies considerably from one carboxylate ligand to another. For example, in solutions containing 10 mM of each ligand, kred does not change significantly for acetate, increases 200-fold for citrate, increases 400-fold for malonate, and increases 6000-fold for oxalate. Two different kinetic trends are observed when ligand concentrations are varied at a constant pH. For acetate and malonate, the rate of oxamyl reduction increases continually with increasing ligand concentration. For oxalate and citrate, however, kred increases with increasing ligand concentration until a maximum value is reached and then decreases when ligand concentrations are increased further. In organic ligandfree solution, the rate of reduction is relatively low and constant for pH < 7 but increases sharply between pH 7 and pH 8. In the presence of the three polycarboxylates, kred increases from the ligand-free value at very low pH until a

plateau value is reached that depends on the polycarboxylate present. As shown in Figure 1C, increasing concentrations of both EDTA and TMDTA have virtually no effect on the rate of oxamyl reduction at pH 5.4. kred measured in the presence of these ligands is similar to that measured in organic ligandfree solution at the same pH. In contrast, IDA or NTA addition results in dramatic increases in rates of oxamyl reduction. The kinetic trends for increasing IDA and NTA concentration are similar to those observed for oxalate and citrate, but the maximum value of kred occurs at much lower ligand concentrations. For NTA, the maximum kred value is already reached at 0.5 mM NTA. The effect of pH is somewhat different than that observed for the carboxylate ligands (Figure 1D). For IDA and NTA, kred increases from the ligandfree value at low pH until a maximum value is reached but begins decreasing again as pH is raised further. For EDTA at pH > 4, kred is relatively low and constant, but at pH < 4, kred increases with decreasing pH. For TMDTA, kred is at a maximum at pH 4 and decreases with either increasing or decreasing pH. It is also notable that kred measured in the presence of TMDTA drops significantly below the values observed in ligand-free solution at pH > 5.5 (i.e., TMDTA inhibits oxamyl reduction by FeII). FeII Speciation and Reactivity. The effect that organic ligands have on the kinetics of oxamyl reduction can be explained by examining the effects that ligands have on FeII speciation. Various FeII-ligand complexes form in aqueous solution (35). For example, the following are equilibrium expressions for the formation of FeII-malonate complexes:

Fe2+ + malonate2- T Fe(malonate)0 Fe2+ + 2malonate2- T Fe(malonate)22-

KFeMal (3a)



where KFeMal and KFe(Mal)2 represent stability constants for formation of 1:1 and 1:2 FeII-malonate complexes, respectively. A complete list of the equilibrium expressions and stability constants used to model FeII speciation are provided in the Supporting Information (Table S2). As discussed previously (25), the reduction of oxamyl by FeII actually occurs simultaneously along several parallel paths involving the different FeII species present in solution. For example, the following two reactions need to be considered (in addition to reactions involving Fe2+ and FeII-hydroxo complexes) in solutions containing malonate:

Fe(malonate)0 + oxamyl f products Fe(malonate)22- + oxamyl f products

kFeMal (4a)



where kFeMal and kFe(Mal)2 represent the second-order rate constants for oxamyl reduction by the 1:1 and 1:2 FeIImalonate complexes, respectively. If FeII is present in considerable excess of oxamyl and if we assume that equilibrium FeII speciation is maintained (i.e., metal complexation and protonation reactions are fast relative to electron transfer), then the overall pseudo-first-order rate constant for oxamyl reduction can be described by eq 2, which accounts for the weighted contribution from each pathway. In solutions containing malonate, eq 2 can be expanded to

kred ) [FeII](kFe(2+)RFe(2+) + kFe(OH)2RFe(OH)2 + kFeMalRFeMal + kFe(Mal)2RFe(Mal)2) (5) Note that the contribution made by each FeII species is equal to the product of ki times Ri. Therefore, for any given species to significantly influence the rate of oxamyl reduction, it must

TABLE 2. Model-Derived Second-Order Rate Constants ki for Selected FeII Species Reacting with Oxamyl at 25 °C, One-Electron Reduction Potentials for the Corresponding FeIII/FeII Redox Couples, and Number of FeII Coordination Positions Occupied by Ligands Other Than Water FeII Fe2+


FeOH+ Fe(OH)20 FeCl+ FeF+ FeHCO3+ FeCO30 Fe(CO3)22FeHPO40 FeH2PO4+ Fe(CN)64Fe(acetate)+ Fe(malonate)0 Fe(malonate)22Fe(oxalate)0 Fe(oxalate)22Fe(oxalate)34Fe(citrate)FeH(citrate)0 FeH2(citrate)+ FeH(citrate)23Fe(IDA)0 Fe(IDA)22Fe(NTA)Fe(NTA)24Fe(OH)(NTA)2Fe(EDTA)2FeH(EDTA)Fe(TMDTA)2FeH(TMDTA)-