Environ. Sci. Technol. 2002, 36, 653-661
Reduction of the Pesticides Oxamyl and Methomyl by FeII: Effect of pH and Inorganic Ligands TIMOTHY J. STRATHMANN* AND ALAN T. STONE Department of Geography and Environmental Engineering, Johns Hopkins University, Baltimore, Maryland 21218
This work examines the effect that pH and selected inorganic ligands have on the kinetics of reactions between FeII and two structurally related oxime carbamate pesticides, oxamyl and methomyl. In anoxic solutions containing FeII, these compounds degrade by parallel elimination and reduction pathways. Rates of FeII-independent carbamate elimination (Elcb mechanism) are proportional to [OH-], increasing 10-fold for each unit increase in pH. In homogeneous solution, rates of carbamate reduction by 0.5 mM FeII are relatively constant at pH 8.3, Fe(OH)2(s) precipitation occurs, and carbamates react with both solution-phase and solid-phase FeII. Carbamate reduction by FeII is not significantly affected by the presence of chloride, bromide, nitrate, perchlorate, and sulfate. In contrast, increased rates of carbamate reduction are observed in solutions containing fluoride, carbonate, and phosphate. Kinetic measurements are interpreted in terms of changing FeII speciation according to the expression kred ) [FeII]∑ikiRi, where kred is the pseudo-first-order rate constant for carbamate reduction, [FeII] is the total FeII concentration, and ki and Ri are the second-order rate constant and fractional concentration of each FeII species, respectively. It follows that the overall kinetics of carbamate reduction is a function of the identity and concentration of individual FeII species present in solution as well as the inherent reactivity of each species with carbamates. The magnitude of ki is related to the standard one-electron reduction potential (EH°) of the corresponding FeIII/FeII redox couple.
Introduction Contamination of soil, groundwater, and surface water by pesticides is a significant concern in the United States (1-3) and throughout the world (4-6). Many of these compounds are a threat to human health and the environment. In recent years, there has been increased interest in elucidating the mechanisms by which pesticides degrade in a variety of environmental settings (7-9). Considerable attention is also being given to improving our understanding of the environmental factors that control the pathways and rates of pesticide degradation. These factors include temperature and * Corresponding author e-mail:
[email protected]; fax: (609)258-1274; phone: (609)258-3827. Present address: Department of Geosciences, Guyot Hall, Princeton University, Princeton, NJ 08544. 10.1021/es011029l CCC: $22.00 Published on Web 01/15/2002
2002 American Chemical Society
FIGURE 1. Schematic representation of the abiotic degradation pathways for oxamyl and methomyl observed in the presence of FeII (15, 26). pH (10-12) as well as the presence of chemical oxidants or reductants (13-16), catalytic mineral phases (17-19), and favorable electron acceptors or donors for microbial-mediated degradation processes (20). Recent studies have shown that several classes of organic contaminants (e.g., organohalides, nitroaromatics, sulfoxides) undergo reductive transformation reactions in anoxic/ suboxic subsurface environments (9). These compounds are reduced through either microbial enzymatic processes (21, 22) or abiotic reactions with reductants present in these settings (e.g., reduced sulfur compounds, FeII, reduced organic matter) (23-25). Surprisingly, little is known about the reductive transformation reactions of many widely used classes of pesticides. In a previous paper (26), we examined the abiotic degradation of two structurally related oxime carbamate pesticides, oxamyl and methomyl, in anoxic solutions containing various reductants (e.g., bisulfide, cysteine, hydroquinone) and metal ions (e.g., CaII, MnII, NiII). Of the reagents surveyed, only FeII, CuI, and CuII accelerate oxime carbamate (hereafter referred to as carbamate) degradation relative to reagent-free blank reactions. FeII is particularly important because of its abundance in suboxic and anoxic subsurface environments (27, 28). Figure 1 illustrates the parallel pathways for carbamate degradation observed in anoxic solution containing FeII (15, 26). An FeII-independent, oxygen-independent elimination reaction (E1cb mechanism) results in the formation of an oxime product, methylamine and CO2. This reaction is well-documented for compounds containing the N-methyl carbamate moiety (10, 29). In parallel with the elimination reaction, a net two-electron reduction of the carbamate is coupled with the one-electron oxidation of two FeII ions. Carbamate reduction products include a nitrile, methanethiol, methylamine, and CO2. Overall carbamate degradation kinetics conforms to the following rate expression:
-
d[carbamate] ) (kelim + k2[FeII])[carbamate] (1) dt
where [carbamate] and [FeII] are the molar concentrations of the parent carbamate and total ferrous iron, respectively; kelim (h-1) is the pseudo-first-order rate constant for carbamate elimination and k2 (M-1 h-1) is a generalized second-order rate coefficient for carbamate reduction by FeII. Equation 1 does not account for the pH dependence of elimination kinetics, nor does it account for the influence of FeII speciation, which can vary considerably in aqueous environments. Even in “simple” homogeneous electrolyte VOL. 36, NO. 4, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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solutions, several distinct FeII species are present (e.g., Fe2+, FeOH+, Fe(OH)20, FeCl+, FeHCO3+), the result of metal ion coordination by ligands present in solution (e.g., H2O, OH-, Cl-, HCO3-) (30). Several studies on FeII oxidation by O2 have reported that reaction rates vary dramatically with changing FeII speciation (31-37). The influence of speciation is explained by considering the apparent FeII oxidation rate (kapp) to be a result of several individual FeII species reacting in parallel:
-
d[FeII] dt
∑k R
) kapp[O2][FeII] ) [O2][FeII]
i i
(2)
i
where [O2] is the concentration of dissolved O2, Ri is the fractional concentration of each FeII species, and ki is the second-order rate constant for oxidation of each species by O2. Selected FeII complexes with hydroxide, fluoride, carbonate, and phosphate are significantly more reactive (several orders of magnitude in some cases) with O2 than is Fe2+ (hexaquo FeII) (32, 37). As a result, even if only a small fraction of [FeII] is present in the form of these reactive species, it can have a marked impact on the apparent rate of FeII oxidation. FeII speciation may exert a similar influence on rates of carbamate reduction. Therefore, a speciation-dependent analysis of reaction kinetics is of interest. This study examines the effects of solution pH and the presence of selected inorganic ligands on the reactivity of FeII with two carbamate pesticides, oxamyl and methomyl. Emphasis is placed on quantitatively correlating observed kinetic trends with changes in FeII speciation. Determining the reactivity of individual FeII species with carbamate pesticides helps us better understand the molecular scale reaction mechanism. In addition, results from this study dramatically improve our ability to predict rates of carbamate degradation in diverse environmental settings. This information is important for evaluating the risks associated with using these chemicals.
Materials and Methods Experimental Procedures. A detailed description of the experimental setup, analytical instrumentation, and listing of most chemical reagents is provided elsewhere (26). NaNO3 (J. T. Baker, Phillipsburg, NJ), NaClO4 (Aldrich Chemical, Milwaukee, WI), NaBr (Aldrich), NaF (MCB Reagents, Darmstadt, Germany), Na2SO4 (Baker), NaHCO3 (Baker), NaH2PO4 (Baker), N-tris(hydroxymethyl)methyl-3aminopropanesulfonic acid (TAPS buffer, pKa 8.4; Sigma Chemical, St. Louis, MO), and 3-[(1,1-dimethyl-2-hydroxyethyl)amino]-2-hydroxy propane sulfonic acid (AMPSO buffer, pKa 9.0; Sigma) were also used in this work. Because strict oxygen exclusion was required, all experiments were conducted within a controlled-atmosphere glovebox (95% N2, 5% H2; Pd catalyst; Coy Laboratory Products, Grass Lake, MI). All stock solutions and reactions were prepared in the glovebox with autoclaved, deoxygenated water from a Milli-Q water purification system (18 MΩ‚cm resistivity; Millipore Corp., Milford, MA). Carbamate degradation was monitored in batch reactors (250 mL polypropylene containers) that were continuously mixed under darkness at 25.0 ( 0.1 °C in a circulating water bath within the glovebox. Reaction solutions were prepared by mixing together appropriate pH buffer, electrolyte, ligand of interest, and FeCl2 from aqueous stock solutions. After the mixtures were equilibrated overnight, an appropriate amount of carbamate was added from an aqueous stock solution to initiate each reaction. Reactions carried out in carbonatecontaining solutions were conducted in zero-headspace syringes to prevent CO2 outgassing. Unless otherwise indi654
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cated, the initial concentration of FeII was 0.5 mM, and parent carbamate was 25 µM. Reactions were monitored for 3 halflives or 3 months, whichever was shorter. At least five aliquots (usually several more) were collected during this time period and were analyzed for the concentration of parent carbamate and selected degradation products using HPLC-UV. Detailed descriptions of the HPLC equipment and operating conditions used are provided elsewhere (26). Solution pH was measured at least three times during the course of each reaction (Fisher Accumet 825MP meter with Orion combination semi-microprobe; NIST standard buffers); these measurements confirmed that pH varied by less than 0.05 units in each reaction. The concentration of dissolved FeII (out of 0.5 mM total FeII) was measured in a series of anoxic carbamate-free solutions buffered at different pH. After being equilibrated for 1 day, the solutions were filtered (0.02 µm Anatop 25-Plus inorganic membrane; Whatman Scientific, Maidstone, England), and the filtrate was analyzed using a ferrozine colorimetric method adapted from Stookey (38). FeII Precipitate Characterization. An FeII precipitate formed in anoxic solution under alkaline pH conditions. The color of the precipitate varied from light gray to dark green, depending upon solution pH, FeII concentration, and even the buffer present. Freshly prepared “pure” Fe(OH)2(s) is reported to be white but turns green when traces of FeIII are present (39). Any FeIII impurities in our system are presumed to result from trace FeIII initially present in FeII stock solutions. Tests verified that dissolved FeII is stable at pH 7 for extended periods of time (at least weeks), demonstrating that oxygen contamination of the experimental setup is negligible. The FeII precipitate was characterized by powder X-ray diffraction (XRD; Philips XRG 3100 X-ray generator, Cu KR radiation source). A precipitate was prepared under conditions similar to those used in carbamate degradation experiments (25 °C, pH 8.75, 5 mM FeII, 0.1 M NaCl, 25 mM TAPS buffer, 24 h equilibration time). The precipitate was then collected by filtration (0.2 µm Nucleopore polycarbonate), mixed with glycerol to prevent FeII oxidation (40) and analyzed immediately. The resulting diffraction pattern (scan 2θ from 5 to 110°) exhibits three peaks that correspond to crystallographic d-spacings of 4.5, 6.7, and 13.6 Å. These do not correspond to d-spacings reported for Fe(OH)2(s) or FeO(s), nor do they match those reported for mixed FeII/FeIII solid phases that could result from partial oxidation of the precipitate (e.g., magnetite, green rusts) (25, 40, 41). Chemical Speciation and Redox Potential Calculations. Equilibrium FeII speciation was calculated using the software package HYDRAQL (42). Table 1 lists the equilibrium expressions and stability constants used for model input along with appropriate references (41-45). Ionic strength corrections were made using the Davies equation (46). Stability constants for the formation of FeII- and FeIIIligand complexes (41-44, 47, 48) were also used to calculate the standard one-electron reduction potentials (EH°; V vs NHE) for selected FeIII/FeII redox couples (FeIII stability constants provided in Supporting Information). The following half-reaction and thermodynamic description are widely reported for reduction of uncomplexed Fe3+ (49):
Fe3+ + e- ) Fe2 +
E°H,Fe(3+)/Fe(2+) ) - ∆G°Fe(3+)/Fe(2+)/F ) + 0.77 V (3)
where ∆G° is Gibbs free energy and F is the Faraday constant. We can derive the half-reaction and thermodynamic description for reduction of an FeIII-ligand complex by combining eq 3 with equilibrium expressions for the formation of equivalent FeII- and FeIII-ligand complexes (i.e.,
TABLE 1. Equilibrium Expressions and Stability Constants for Calculating FeII Speciationa log K
ref
-13.997 -9.397 -20.494 -28.991 -45.988 -12.844
43 43 43 43 43 41
Reactions Involving Chloride Fe2+ + Cl- ) FeCl+ -0.200 Na+ + Cl- ) NaCl0 -0.500
43 43
Reactions Involving Fluoride Fe2+ + F- ) FeF+ 1.208 Na+ + F- ) NaF0 -0.200 H+ + F- ) HF0 3.170 2H+ + F- ) H2F+ 3.750
43 43 43 43
Reactions Involving Carbonate Fe2+ + CO32- + H+ ) FeHCO3+ 11.799 Fe2+ + CO32- ) FeCO30 5.450 Fe2+ + 2CO32- ) Fe(CO3)227.160 Na+ + CO32- + H+ ) NaHCO30 10.079 Na+ + CO32- ) NaCO31.270 H+ + CO32- ) HCO310.329 2H+ + CO32- ) H2CO30 16.681
44 44 44 43 43 43 43
Reactions Involving Phosphate Fe2+ + PO43- + H+ ) FeHPO40 15.975 Fe2+ + PO43- + 2H+ ) FeH2PO4+ 22.273 3Fe2+ + 2PO43- ) Fe3(PO4)2(s) 33.30 32+ H + PO4 ) HPO4 12.375 2H+ + PO43- ) H2PO419.573 3H+ + PO43- ) H3PO40 21.721 Na+ + PO43- ) NaPO421.43 3+ + Na + H + PO4 ) NaHPO4 13.445 Na+ + 2H+ + PO43- ) NaH2PO40 19.873 2Na+ + PO43- ) Na2PO42.590 2Na+ + H+ + PO43- ) Na2HPO40 13.320
43 43 42 43 43 43 43 43 43 43 43
Hydrolysis Reactions H2O - H+ ) OH2+ + Fe + H2O - H ) FeOH+ Fe2+ + 2H2O - 2H+ ) Fe(OH)20 Fe2+ + 3H2O - 3H+ ) Fe(OH)3Fe2+ + 4H2O - 4H+ ) Fe(OH)42Fe2+ + 2H2O - 2H+ ) Fe(OH)2(s)
a Stability constants corrected with Davies equation (46) to I ) 0 and T ) 25 °C.
complexes that differ only in their oxidation state):
Fe2+ + xH+ + yL ) FeIIHxLy ∆G°Fe(II)L ) - RT ln KFe(II)L (4) Fe3+ + xH+ + yL ) FeIIIHxLy ∆G°Fe(III)L ) - RT ln KFe(III)L (5) where x and y represent stoichiometric coefficients for protons and ligand molecules (L) in the metal complex, respectively; R is the gas constant; T is absolute temperature; and K is the stability constant. Combining eqs 3-5, we arrive at the following half-reaction and Gibbs free energy description for reduction of the FeIII-ligand complex:
FeIIIHxLy + e- ) FeIIHxLy
∆G°Fe(III)L/Fe(II)L )
∆G°Fe(3+)/Fe(2+) + ∆G°Fe(II)L - ∆G°Fe(III)L (6) By plugging the expressions for ∆G° from eqs 3-5 into eq 6, we arrive at an expression for EH° of the FeIII-ligand/ FeII-ligand redox couple in terms of the stability constants provided in Tables 1 and S1 (Supporting Information):
( )
RT KFe(III)L EH° ) + 0.77 ln F KFe(II)L
Inspection of eq 7 reveals that ligands that form more stable complexes with FeIII than FeII (i.e., KFe(III)L/KFe(II)L > 1) will lead to EH° values less than 0.77 V and vice versa. Several studies have reported that the reactivity of FeII correlates with EH° (35, 50, 51). Kinetic Modeling. Plots of ln[carbamate] vs time are linear for all reactions monitored (FeII present in large excess to carbamate), indicating that overall degradation follows pseudo-first-order kinetics (kobs, h-1):
-
d[carbamate] ) kobs[carbamate] dt
(8)
According to Figure 1, carbamate degradation occurs by two parallel reaction pathways. Equation 8 can be expanded to account for degradation due to each pathway:
-
d[carbamate] ) (kelim + kred)[carbamate] dt
(9)
where kelim (h-1) and kred (h-1) are defined as pseudo-firstorder rate constants for carbamate elimination and reduction, respectively. Note that eq 9 is equivalent to eq 1 because kred ) k2[FeII] when FeII is present in considerable excess of carbamate. Pseudo-first-order rate expressions can also be written for the formation of the oxime elimination and nitrile reduction products that were also monitored using HPLC:
d[oxime] ) kelim[carbamate] dt
(10)
d[nitrile] ) kred[carbamate] dt
(11)
Pseudo-first-order rate constants were calculated for each batch reaction using the software package Scientist for Windows (52). Scientist calculates these parameters (along with the initial concentration of the parent compound) by fitting (method of least squares) numerically integrated solutions of the system of differential rate expressions (eqs 9-11) to experimental data for parent compound loss and reaction product appearance. By fitting product formation data simultaneously with parent compound loss data, we are better able to assess the importance of each reaction pathway to overall carbamate degradation. We demonstrated in a prior study that kelim is unaffected by the presence of FeII (26). As a result, kelim values were constrained when modeling carbamate degradation in solutions containing FeII (fixed at values determined in FeII-free reactions at the same pH), thereby reducing the number of adjustable fitting parameters from 3 to 2 (kred and [carbamate]initial). Only rate constants greater than 7.0 × 10-5 h-1 could be reliably calculated using our experimental approach and data analysis methods. Scientist for Windows was also used to estimate secondorder rate constants for oxamyl reduction by selected FeII species. This was done by least-squares fitting eq 17 (see Speciation Kinetics Model section) to experimentally measured rate constants and calculated FeII speciation. Model fitting was carried out in a stepwise fashion. Systems with the least number of FeII species were fit first (e.g., only Fe2+, FeOH+, and Fe(OH)20). Rate constants derived from fitting data in these systems were then fixed when applying the model to more complex systems that included additional FeII species (e.g., FeII-carbonate species). Upper limit estimates for the reactivity of selected FeII species were made in cases where kinetic trends with respect to a given species were extremely small or nonexistent.
Results and Discussion (7)
For all batch reactions conducted in this study, the observed carbamate transformation products and product mass balVOL. 36, NO. 4, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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program SPARC (53, 54). A complete list of SPARC results are provided in the Supporting Information. The estimated pKa values are all less than pH 1.0 (conjugate acids of the oxime nitrogen, amide nitrogen, and carbonyl oxygen moieties) or greater than pH 10.5 (amide nitrogen). According to these results, the nonionic carbamate species predominates throughout the pH range investigated here. FeII speciation was calculated for solution conditions used in the kinetic experiments (pH 2-10, 0.5 mM FeII, 0.1 M NaCl). Results of the calculations are shown in Figure 2A. Fe2+ (hexaquo FeII) is the predominant dissolved FeII species for pH 2-8.3; Fe2+ is orders of magnitude higher in concentration than all other species. FeCl+ is the next most concentrated species. As the pH increases, the concentration of various FeII-hydroxo species increase, and eventually saturation with respect to the (hydr)oxide solid, Fe(OH)2(s), is reached. For experimental conditions used in this study, Fe(OH)2(s) is predicted to begin precipitating at pH 8.3. Measurements of dissolved FeII (passes through a 0.02 µm filter), also shown in Figure 2A, agree closely with the calculated speciation; FeII solubility begins to drop off considerably only when the pH is raised above 8.3. Figure 2, panels B and C, shows the rate constants for carbamate elimination and reduction, kelim and kred, respectively, measured as a function of pH. It was reported in an earlier study that rates of carbamate elimination are independent of FeII concentration (26). From pH 2 to pH 10, elimination kinetics observed in the presence of 0.5 mM FeII are adequately described using kelim values measured in FeIIfree solutions at the same pH (values shown were measured in FeII-free solution). Carbamate elimination reactions are exceedingly slow at subneutral pH. Below pH 5, the downward arrows on kelim values for oxamyl data indicate that the rate constant for elimination is too low (
