Fenton-Mediated Oxidation in the Presence and Absence of Oxygen

School of Civil and Environmental Engineering, The University of New South Wales, Sydney, New South Wales 2052, Australia, and Department of Chemistry...
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Environ. Sci. Technol. 2005, 39, 5052-5058

Fenton-Mediated Oxidation in the Presence and Absence of Oxygen CHRISTOPHER K. DUESTERBERG,† WILLIAM J. COOPER,‡ AND T . D A V I D W A I T E * ,† School of Civil and Environmental Engineering, The University of New South Wales, Sydney, New South Wales 2052, Australia, and Department of Chemistry and Biochemistry and Center for Marine Science, University of North Carolina Wilmington, 5600 Marvin K. Moss Lane, Wilmington, North Carolina 28409

The increased use of Fenton systems for the treatment of contaminated waters and wastewaters necessitates the development of kinetic models capable of accurately simulating key species concentrations in order to optimize system performance and efficiency. In this work a reaction mechanism in which the hydroxyl radical is nominated to be the active oxidant in Fenton systems is used to describe the oxidation of formic acid (HCOOH) under a variety of experimental conditions. A kinetic model based on this reaction mechanism is shown to adequately describe results of experiments in which starting concentrations of H2O2 and HCOOH varied over 1 and 4 orders of magnitude, respectively, under both air-saturated and deaerated conditions. The intermediate generated during HCOOH oxidation was observed to increase oxidation efficiency, especially at high initial organic concentrations [relative to Fe(II)], by assisting in the redox cycling of iron. In the presence of oxygen, however, such improvement was attenuated through competition for the organic intermediates. While mechanistic analysis and associated kinetic modeling is invaluable in optimization of Fenton systems, a clear understanding of reaction byproducts and their reactivity toward other species in the system is critical for accurate simulations.

Introduction Aqueous mixtures of ferrous iron [Fe(II)] and hydrogen peroxide (H2O2), otherwise known as Fenton’s reagent, represent one of the more common treatment methods for contaminated surface and groundwaters and industrial wastewater effluents. First discovered over a century ago (1), the Fenton [Fe(II)/Fe(III)/H2O2] and photo-Fenton [Fe(II)/ Fe(III)/H2O2/UV light] systems have been the subject of numerous investigations attempting to identify optimal process parameters and the reaction mechanism involved (2-4 and references therein). While the qualitative impact of some parameters (such as pH) on process performance has generally been well-established (4), some variation has been observed in the optimal choice of other parameters, such as the amount and ratio of iron and hydrogen peroxide. Tang and Huang (3) observed the optimal ratio of H2O2 to * Corresponding author phone: +61-2-9385 5060; fax: +61-29385 6139; e-mail: [email protected]. † The University of New South Wales. ‡ University of North Carolina Wilmington. 5052

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Fe(II) to be around 5-11 for the initial rate of oxidation of trichloroethylene (TCE), in agreement with a theoretical optimum of 11. Teel et al. (5) demonstrated optimum conditions for TCE degradation using a H2O2 to Fe(II) ratio of 1:1, whereas Gallard and De Laat (6) observed maximum rates for the oxidation of atrazine at H2O2:Fe(II) ratios approaching 500. While such differences are likely attributable to the varying measures used to define the optimum ratio or the differing target organic species, these and other works (7, 8) highlight the importance of an appropriate understanding of the reactions governing the Fenton system for determining optimal process parameters. The generally accepted mechanism for the Fenton process identifies the hydroxyl radical (•OH) as the active oxidizing intermediate in the system (7, 9, 10). According to this mechanism (see Table 1), the combination of ferrous iron and hydrogen peroxide induces a series of chain reactions initiated by the degradation of peroxide to the hydroxyl radical and the hydroxide ion (reaction 1). The hydroxyl radical serves as a chain carrier that may react with Fe(II), H2O2, or any organic species present. These reactions may either propagate the chain cycle, through the production of additional radicals (superoxide and its conjugate acid, reaction 3) that can reduce Fe(III) back to Fe(II) (reaction 4), or terminate the chain by oxidizing Fe(II) (reactions 5 and 6). When reaction 2 is taken into account, Fe(III) may also be considered a chain carrier, producing Fe(II) and superoxide, although this cycling occurs at a much slower rate (k2 , k3, k4). Additional chain termination reactions include the more minor radical-radical recombination pathways (reactions 6-8). Depending on the type of organic species present, reactions with •OH may either propagate the chain by producing HO2•/O2• - or organic radicals capable of reducing Fe(III) directly, or terminate the cycle by scavenging •OH (7). Despite its long history and wide acceptance, the role of the hydroxyl radical as the active oxidizing intermediate in the Fenton system continues to be debated, perhaps more so now than ever. This is partly due to the difficulty in directly measuring the hydroxyl radical itself, whose high reactivity limits its lifetime in solution. Over the past decade, however, an increasing number of theoretical and experimental studies have presented results inconsistent with the proposed hydroxyl radical mechanism, instead presenting evidence for a high-valent iron-oxo complex as the active intermediate responsible for oxidation in the Fenton system (11-15). In fact, Kremer (13) recently proposed an alternative reaction mechanism in which the ferryl ion (FeO2+) acts as the key oxidant in order to explain his measurements of the amount of O2 produced during the Fenton reaction. Regardless of whether the hydroxyl radical or a high-valent iron-oxo complex such as FeO2+ is the actual oxidizing intermediate involved, a kinetic model that accurately predicts species concentrations over a range of initial conditions provides a valuable tool for the application of Fenton’s technology to water and wastewater treatment. Recent work by a number of groups has demonstrated considerable success in predicting Fe(II), H2O2, and organic species concentrations over time during laboratory-controlled experiments using reaction sets similar to that described in Table 1 (6, 16-20). Despite the number and variety of studies undertaken, no one publication has comprehensively applied the basic Fenton’s model to a simple target organic compound for a wide range of initial reagent and substrate concentrations. In many cases the target organic chosen has been a complex aromatic compound whose oxidative degradation pathway and/or products 10.1021/es048378a CCC: $30.25

 2005 American Chemical Society Published on Web 05/18/2005

TABLE 1. Mechanism of Fe(II)-Initiated Chain Reaction

reaction 1 2 3 4 5 6 7 8 9

•OH

OH-

Fe(II) + H2O2 f Fe(III) + + Fe(III) + H2O2 f Fe(II) + HO2•/O2• - + H+ H2O2 + •OH f HO2•/O2• - + H2O Fe(III) + HO2•/O2• - f Fe(II) + O2 + H+ Fe(II) + •OH f Fe(III) + OHFe(II) + HO2•/O2• - f Fe(III) + H2O2 HO2•/O2• - + HO2•/O2• - f H2O2 + O2 •OH + HO •/O • - f H O + O 2 2 2 2 •OH + •OH f H O 2 2

rate constant (M-1 s-1)

ref

51 2.0 × 10-3 3.3 × 107 7.8 × 105 3.2 × 108 1.3 × 106 2.3 × 106 7.1 × 109 5.2 × 109

7 17 17 17 17 17 17 17 17

are ill-defined (18-20). At high initial organic concentrations ([org]0 > [Fe(II)]0, [H2O2]0), knowledge of such degradation products and their reactivity toward other species in the system, especially Fe, becomes important for the accuracy of the model. Other studies in which the degradation of simpler compounds has been investigated (17) were limited to low initial substrate concentrations ([org]0 , [H2O2]0). The primary objective of this work was to investigate the validity of the reaction mechanism presented in previous studies by applying a kinetic model based on the proposed mechanism over a wide range of experimental conditions. Initial molar ratios of the Fenton reagents and organic compound ([H2O2]0:[Fe(II)]0:[org]0) were varied from 1:1:0.001 to 22:1:20 under both oxygenated and anoxic conditions at pH 3 and 0.01 M ionic strength. Formic acid was chosen as the target organic compound as its reaction with •OH is fast (108-109 M-1 s-1) and its degradation pathway well-defined (21-23).

Materials and Methods Reagents. All chemicals were reagent-grade and were used as received, and all solutions were prepared with 18 MΩ‚cm Milli-Q water. H2O2 (30%) was from APS Chemicals Ltd., and stock solutions were calibrated by the iodometric method (24). Prior to calibration of H2O2 stock solutions, the sodium thiosulfate solution was standardized by the dichromate method (24). 14C-Labeled formic acid (Na14COOH, from Sigma) was used for formic acid analysis. The H2O2 stock, 14C-labeled formic acid stock, peroxidase (type II from horseradish), and N,N-diethyl-p-phenylenediamine (DPD) solutions were stored in the dark at 4 °C. The peroxidase and DPD solutions were kept for no more than 2 weeks. A 10.0 mM Fe(II) stock solution was prepared from ferrous perchlorate [Fe(ClO4)2, Sigma] in 0.1 M HClO4. Standard calibrations of Fe(II) solutions were conducted regularly to ensure oxidation of Fe(II) did not occur over time. All other solutions were prepared by dissolution in Milli-Q water to specified concentrations. The experimental apparatus was flushed both before and after experiments with 0.1 M HCl prepared from 30% HCl (Sigma). All glassware and containers used were soaked in 5% HNO3 for 10 h and rinsed thoroughly with Milli-Q water prior to use. pH measurements were made on a Metrohm 692 pH/ion meter calibrated against standard buffers. Analytical Techniques. Ferrous iron concentrations were measured spectrophotometrically in a flow-injection apparatus by the 1,10-phenanthroline method (25). The sample solution was drawn continuously from the reaction vessel during the experiment and combined in a mixing T with a solution of 1,10-phenanthroline in an acetate buffer prior to passing through a 1 cm Teflon flow cell. A Mikropack DH2000 UV-vis-NIR lamp was used as the light source for the flow cell, and absorbance measurements were obtained on an Ocean Optics SD2000 spectrometer and recorded on a

personal computer. Correlation between absorbance measurements and Fe(II) concentrations was determined through calibration with standardized Fe(II) solutions. Samples for H2O2 and HCOOH analysis were taken at discrete intervals during the experiments according to the following procedure. At specified times, 1 mL of sample solution was added to 9 mL of a 1,10-phenanthroline solution in a 0.05 M phosphate buffer. The 1,10-phenanthroline served to quench the Fenton reaction by complexing Fe(II), and the presence of the phosphate buffer established the pH required for H2O2 analysis by the DPD method (26). Control experiments showed negligible degradation of H2O2 in the presence of the Fe(II)-phenanthroline complex over a period of several hours. An appropriate volume (300 µL-2 mL) of the quenched sample was then diluted with 0.5 M EDTA in a 1 cm quartz cuvette in order to minimize interference by Fe(III) (27). The absorbance was recorded at 551 nm on a Cary 1E UV-vis spectrometer after addition of DPD and peroxidase. As ferric iron was still present, the samples were analyzed after each experiment in order to minimize the H2O2 degradation by reaction with Fe(III). H2O2 concentrations were calculated from calibration curves by use of H2O2 standards. Following H2O2 analysis, 1 mL of 1.4 M hydroxylamine hydrochloride (NH2OH-HCl) was added to the remaining sample volume to reduce the Fe(III) present to ferrous iron. This allowed all of the iron present to be complexed by 1,10-phenanthroline and prevented further oxidation of HCOOH during its analysis. The samples were then vigorously sparged with air for at least 30 s to drive out any 14CO2 before 1 mL of the sample solution was added to 10 mL of Beckman liquid scintillation fluid for analysis in a Packard Tri-Carb 2100TR liquid scintillation analyzer. Kinetic Experiments. All experiments were performed in a 250 mL water-jacketed glass reaction vessel at 25 °C and pH 3 under constant stirring with a magnetic stirring bar. Appropriate volumes of stock solutions of Fe(II) and HCOOH were diluted in water containing 10 mM NaClO4 (for ionic strength uniformity) to yield the target initial concentrations and a final volume of 200 mL. For the deaerated experiments, solutions were continuously sparged with argon for at least 5 min prior to addition of H2O2 and for the duration of the reaction. Kinetic Modeling. Modeling results were generated by use of the kinetic modeling programs Acuchem (28) and Kintecus (29). Table 1 shows the reaction set for the Fenton system used as input for both programs. The reaction set was based on the kinetic model proposed by De Laat and Gallard (16) with minor modifications as described by Kwan and Voelker (17). These modifications included the use of composite rate constants (reactions 6, 7, 8, and 10) in order to account for the effect of pH on superoxide/hydroperoxyl radical speciation. Fe(III)-hydroperoxy complexes were omitted from the model as their concentrations were negligible under the experimental conditions used. Rate constants for the reactions listed in Table 1 are relatively well established in the literature, though a slight discrepancy exists over the exact value for the initiation reaction, k1, with published values ranging from 51 to 76 M-1 s-1 (3, 7, 16). A value of 51 M-1 s-1 for k1 (3) was used in our model as it provided the best fit to experimental data across the entire range of conditions examined. To accurately simulate species profiles measured during HCOOH oxidation, some reactions for which rate constants were unknown were deemed necessary for the model (see Results below). Rate constants for these reactions were determined by use of numerical fitting routines built into the Kintecus program. As the optimized k values varied for each set of data, the rate constants used were chosen on the basis of visual comparison of the model and experimental results across the entire range of conditions investigated. VOL. 39, NO. 13, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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Results and Discussion Model Verification at Low Organic Concentrations. In an effort to verify the applicability of the reaction mechanism proposed by Kwan and Voelker (17), initial experiments were conducted with similar starting concentrations of the target organic, H14COOH (200 nM). Experiments were run for a variety of H2O2 concentrations such that initial mole ratios of [H2O2]0:[Fe(II)]0 varied from 1:1 to 22:1. In all cases [Fe(II)]0 ) 200 µM and [H14COOH]0 ) 200 nM except for the runs with the 22:1 initial mole ratio, where [Fe(II)]0 ) 100 µM. Concentration profiles of the measured species are shown in Figure 1. The reaction between formic acid/formate and •OH is fast and well-characterized, with a rate constant on the order of 108-109 M-1 s-1 (21, 23, 30). To account for the oxidation of H14COOH, reaction 10 was included in the reaction set used for computer simulations of these experiments (Figure 1). The rate constant for reaction 10 is a composite

reaction •OH

10

+ HCOOH f CO2 + HO2•/O2• -

rate constant (M-1 s-1)

ref

6.5 × 108

17

value calculated using the rate constants for the reactions of HCOOH and HCOO- with •OH and HCOO- and a pKa of 3.75 for HCOOH/HCOO- (17). While the results presented in Figure 1 indicate that the reaction set described in Table 1 and reaction 10 effectively predicts species concentration profiles under certain conditions, it should be noted that reaction 10 represents a vastly simplified version of the oxidation pathway of formic acid. Previous studies have shown that •OH initially attacks HCOOH/HCOO- via hydrogen abstraction, resulting in generation of the carboxyl radical, HCOO•/COO• -, and water (21-23, 31). In the presence of oxygen, the carboxyl radical then reacts at diffusion-limited rates to yield CO2 and superoxide (30, 32). Under the experimental conditions described above, any carboxyl radicals generated by the reaction of •OH with HCOOH are immediately oxidized to CO2 by the oxygen present in the system. Considering the low starting concentration of HCOOH relative to the dissolved oxygen concentration (at 25 °C, DOsat ) 2.7 × 10-4 M; 33) it seems reasonable that the presence of the organic radical species can be neglected under these conditions. However, as the initial concentration of formic acid increases and/or the system is deoxygenated, this simplification is no longer valid and the full oxidation pathway for HCOOH must be considered. Initial Conditions and the Role of Oxygen. After validating of the reaction mechanism using low target organic concentrations in air-saturated systems, additional experiments were conducted to examine species profiles at higher initial HCOOH concentrations under both air-saturated and deaerated conditions. Initial [H2O2]0:[Fe(II)]0 mole ratios were consistent with previous experiments (1:1 up to 22:1) while initial HCOOH concentrations varied from 20 µM to 2 mM. Concentration profiles of the measured species for airsaturated and deaerated experiments in which [HCOOH]0 ) 2 mM and 200 µM are shown in Figures 2-5. As expected, examination of the Fe(II) concentration profiles in Figures 2-5 show the rate of Fe(II) oxidation increases in proportion to the initial H2O2 concentration under all the experimental conditions studied. Comparison of species concentration profiles obtained for air-saturated and deaerated systems confirms that oxygen has a significant effect on these systems. The reaction set described earlier, comprised of the reactions in Table 1 and reaction 10, would obviously be 5054

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FIGURE 1. Experimental (symbols) and model (lines) data for oxidation of HCOOH. [HCOOH]0 ) 200 nM; [H2O2]0 ) 200 µM (a), 1.1 mM (b), or 2.2 mM (c, d); [Fe(II)]0 ) 200 µM (a-c) or 100 µM (d). Simulated data are based on mechanism in Table 1 and reaction 10. Some profiles have been omitted for clarity. inadequate in predicting the species behavior observed as it makes no allowance for the presence or absence of oxygen in the system. The large difference in species profiles between air-saturated and deaerated experiments demonstrates that oxygen reactions must be included in order for the model to accurately simulate system behavior. To this end, reactions 11-21, shown in Table 2, were implemented in place of reaction 10 in order to better describe the pathway involved in the oxidation of HCOOH. The simulated results shown in Figures 2-5 were generated with the reaction mechanism described in Tables 1 and 2 (reactions 1-9 and 11-21). Examination of Table 2 reveals that reactions 11 and 12 contain the same reactants but different product species in the protonated and deprotonated forms of the carboxyl radical. Where we have previously dealt with such equilibrium speciation in the form of composite rate constants, the carboxyl radical species is treated differently due to the differing products generated by disproportionation of the carboxyl radical anion (reactions 15-17) and the carboxyl radical (reactions 18 and 19). Considering the pKa value (2.3) for the CO2• -/HCOO• radical (21) and the pH 3 used in our experiments, [CO2• -]/[HCOO•] ≈ 5 at equilibrium, so it seemed relevant to include both recombination pathways.

FIGURE 2. Experimental (symbols) and model (lines) data for oxidation of HCOOH in solution open to the atmosphere. [HCOOH]0 ) 2 mM; all other concentrations are as in Figure 1. Simulated data are based on mechanism in Tables 1 and 2. Some profiles have been omitted for clarity.

FIGURE 3. Experimental (symbols) and model (lines) data for oxidation of HCOOH in argon-saturated solution. [HCOOH]0 ) 2 mM; all other concentrations are as in Figure 1. Simulated data are based on mechanism in Tables 1 and 2. Some profiles have been omitted for clarity.

TABLE 2. Mechanism for HCOOH/HCOO- Oxidation

in light of its affinity for Fe(III) and its effect on the iron redox cycle (34). Using the composite rate constant for the •OH + HCOOH/HCOO- reaction, the rate constants k and 11 k12 were calculated such that [CO2• -]/[HCOO•] ≈ 5 (i.e., k12/ k11 ≈ 5) throughout the simulation. In addition to disproportionation reactions between carboxyl radicals (reactions 15-19), Table 2 includes reactions that account for the ability of oxygen (reactions 13 and 14) and ferric iron (reactions 20 and 21) to participate in HCOOH oxidation. Comparison of the respective species profiles between Figures 2 and 3 and between Figures 4 and 5 highlights the significant role such species play in the oxidative decomposition of HCOOH. As an oxidant in plentiful supply in air-saturated systems, oxygen serves as the primary scavenger of the carboxyl radicals produced in solution, yielding CO2 and a superoxide radical that can both propagate and terminate the Fenton chain cycle through reduction of Fe(III) (reaction 4) or oxidation of Fe(II) (reaction 6). Even at high initial HCOOH and H2O2 concentrations, when production of carboxyl radicals is greatest, the near diffusion-limited rate constant for the oxidation of carboxyl radicals by O2 (30, 32) continues to dominate the oxidation pathway of HCOOH. In the absence of oxygen, however, the

reaction

rate constant (M-1 s-1)

+ HCOOH f + H2O 1.3 × + HCOOH f CO2• - + H2O + H+ 6.5 × 108 HCOO• + O2 f CO2 + HO2•/O2• 4.2 × 109 CO2• - + O2 + H+ f CO2 + HO2•/O2• - 4.2 × 109 CO2• - + CO2• - f int 1.4 × 109 int + H+ f CO2 + HCOOH 1.0 × 1010 int f oxalate 1.6 × 106 HCOO• + HCOO• f CO2 + HCOOH 1.7 × 109 HCOO• + CO2• - + H+ f 1.7 × 109 CO2 + HCOOH 20 CO2• - + Fe(III) f CO2 + Fe(II) 9.83 × 106 a 2.77 × 106 b 21 HCOO• + Fe(III) f CO2 + Fe(II) + H+ 3.55 × 106 a 1.0 × 105 b 11 12 13 14 15 16 17 18 19

•OH

HCOO•

108

•OH

a Rate constants used when [HCOOH] ) 2 mM. 0 used when [HCOOH]0 ) 200 µM.

b

ref this study 17 21 21 21 21 21 21 21 this study this study this study this study

Rate constants

In addition, while model calculations indicate the concentration of oxalate is always less than 5 µM under the experimental conditions examined in this work, the presence of oxalate would need to be monitored under other conditions

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FIGURE 4. Experimental (symbols) and model (lines) data for oxidation of HCOOH in solution open to the atmosphere. [HCOOH]0 ) 200 µM; all other concentrations are as in Figure 1. Simulated data are based on mechanism in Tables 1 and 2. Some profiles have been omitted for clarity.

FIGURE 5. Experimental (symbols) and model (lines) data for oxidation of HCOOH in argon-saturated solution. [HCOOH]0 ) 200 µM; all other concentrations are as in Figure 1. Simulated data are based on mechanism in Tables 1 and 2. Some profiles have been omitted for clarity.

ferric iron available in the system becomes more significant in oxidizing the carboxyl radicals to CO2, at the same time assisting the Fenton chain cycle by regenerating Fe(II). The use of oxidants, especially oxygen and ferric iron, to assist in the hydroxylation or oxidation of organic compounds has been recognized in previous studies (31, 35-37). Variations in degradation efficiency with the addition or removal of such species are generally attributed to the ability of the species to oxidize or reduce the radical intermediates produced by the reaction of hydroxyl with the target organic. As a result, reactions 20 and 21 were included to account for the Fe(II) behavior observed in the argon-saturated experiments, where Fe(II) concentration profiles were consistently higher than their oxygenated counterparts. Rate constants for reactions 20 and 21 were determined through numerical fitting routines and visual comparison of experimental and simulated data. Without reactions 20 and 21, computer simulations of the deaerated experiments were unable to replicate the temporal changes in species concentrations observed. Closer inspection of Figure 3 shows that the simulated profiles for Fe(II) do not exhibit the regenerative behavior observed experimentally, especially at lower initial H2O2

concentrations (200 µM and 1 mM), even though the experimental and model data for H2O2 and HCOOH are similar. This difference can most likely be attributed to the equilibrium speciation of Fe(III). For the HCOOH concentrations used in these experiments, equilibrium calculations reveal that approximately 70% of the total Fe(III) in the system will exist as an Fe(III)-HCOOH complex. Considering the dramatic effect that simple organics (e.g., oxalate, citrate) can have on the redox cycling of iron (34, 38), it is reasonable to assume that the oxidation of carboxyl radicals by free Fe(III) and the Fe(III)-HCOOH complex occurs at significantly different rates. Combined with the uncertainty in the kinetics of formation of the Fe(III)-HCOOH complex, we can only conclude that the rate constants used in reactions 20 and 21 represent a composite figure that includes both the speciation of Fe(III) and Fe(III)-HCOOH as well as their respective rate constants for reaction with CO2• -. The similarity between the simulated and observed profiles for the other species, H2O2 and HCOOH, provides support for this assumption. Extending this approach to lower HCOOH concentrations, the values of k20 and k21 should decrease in proportion to the fraction of total Fe(III) that is complexed by HCOOH. Simulations of experiments at lower initial HCOOH con-

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FIGURE 6. Experimental HCOOH degradation observed for airsaturated (O) and deaerated (b) solutions. Model simulations of HCOOH concentrations for air-saturated (---) and deaerated (s) solutions are also shown. (A) [H2O2]0 ) 1.1 mM, [Fe(II)]0 ) 200 µM, [HCOOH]0 ) 2 mM; (B) [H2O2]0 ) 2.2 mM, [Fe(II)]0 ) 200 µM, [HCOOH]0 ) 2 mM; (C) [H2O2]0 ) 2.2 mM, [Fe(II)]0 ) 100 µM, [HCOOH]0 ) 2 mM. centrations (200 µM) do in fact generate more accurate fits to experimental data when k20 and k21 are adjusted according to the fraction of Fe(III)-HCOOH present. The rate constants used in simulations of experiments using different initial HCOOH concentrations are identified in Table 2, and simulated data for the respective values are shown in Figures 2-5. In experiments at higher H2O2 concentrations (2.2 mM), one would expect the difference in rates for the reactions of Fe(III) and Fe(III)-HCOOH with CO2• - to have less of an effect on the Fe(II) profile, as any ferrous iron generated would be immediately reoxidized by H2O2. Fe(II) profiles for experiments with initial [H2O2]0:[Fe(II)]0 mole ratios of 11:1 and 22:1 show such behavior, as the model closely matches the observed data, especially in the later stages of the profile. The difference in the early stages of these simulations can again be explained by Fe(III) speciation as the kinetics of complex formation would have a much more dramatic effect in the earlier stages of the reaction. Aerated vs Deaerated Fenton Systems. The experimental data and model results shown in Figures 2-5 show the significant effect that oxidants such as oxygen can have on the degradation efficiency of the Fenton system. Comparison

between air-saturated and argon-saturated experiments with similar starting conditions shows considerably more rapid decomposition of HCOOH in the absence of air (Figure 6). The higher Fe(II) concentrations observed in the deaerated experiments, which allow for higher steady-state concentrations of the hydroxyl radical, are the reason for such improvement and can be traced back to the organic radicals produced by the initial attack of •OH on HCOOH. The carboxyl radicals produced by this initial reaction are known to be strong reducing agents (21), and the type of oxidant available will have an important effect on the chain cycle of the Fenton reactions. In air-saturated systems the carboxyl radicals reduce oxygen to superoxide, a species that both terminates and propagates the Fenton cycle through reactions with Fe(II) and Fe(III) (reactions 4 and 6, Table 1). Under anoxic conditions, on the other hand, the dominant oxidant available is Fe(III), and its reduction by the carboxyl radicals extends the chain cycle by regenerating Fe(II). Thus the difference in organic decomposition for similar initial concentrations under air-saturated and argon-saturated conditions is explained by the fact that, in the presence of oxygen, organic radicals are converted into species that can both oxidize and reduce Fe, while in the absence of oxygen, only reducing species are available. While the issues raised above may have significant implications for the performance of the Fenton process in treatment systems, it is important to recognize that the effects of oxygen on the Fenton system largely depend on the properties of the organic radicals present. Walling (7) first recognized the ability of various organic radicals to propagate [through reduction of Fe(III)] or terminate the Fenton chain cycle [through dimerization or oxidation of Fe(II)], noting that the affinity for either pathway was subject to the radical’s structure and ease of oxidation or reduction. The carboxyl radicals produced from the reaction of •OH with formic acid are known to be excellent reducing agents, and as expected, the results show that degradation efficiency improves in the absence of oxygen. However, radical species produced by reaction of •OH with other organic compounds exhibit a wide range of reactivities and redox potentials (7, 31, 35), including some species that, like superoxide, may act as both oxidizing and reducing agents. The hydroxycyclohexadienyl radical, widely accepted as the initial product of •OH addition to aromatic compounds, is known to reduce O2 and Fe(III) in producing hydroxylated products. However, in acidic or basic solutions the species readily undergoes water elimination to the phenoxyl radical, a strong oxidizing agent (35-37, 39). Other radical species, such as those produced from reaction of hydroxyl with carbonyl compounds, are inert to oxidation by O2 or Fe(III) and instead are readily reduced by Fe(II) (7). Knowledge of the degradation mechanism of the target substrate, especially radical intermediates and byproducts and their reactivity toward other species in the Fenton system, is therefore vital in the application of the kinetic model to determine the effect of operating conditions such as the presence of oxygen.

Literature Cited (1) Fenton, H. J. H oxidation of tartaric acid in the presence of iron. J. Chem. Soc. 1894, 65, 899-910. (2) King, D. W.; Farlow, R. Role of carbonate speciation on the oxidation of Fe(II) by H2O2. Mar. Chem. 2000, 70, 201-209. (3) Tang, W. Z.; Huang, C. P. Stoichiometry of Fenton’s reagent in the oxidation of chlorinated aliphatic organic pollutants. Environ. Technol. 1997, 18, 13-23. (4) Wadley, S.; Waite, T. D. Fenton processes applied to water and wastewater treatment. In Advanced Oxidation Processes in Water and Wastewater Treatment; Parsons, S., Ed.; IWA Press: London, 2004; pp 1-31. (5) Teel, A. L.; Warberg, C. R.; Atkinson, D. A.; Watts, R. J. Comparison of mineral and soluble iron Fenton’s catalysts for the treatment of trichloroethylene. Water Res. 2001, 35, 977-984. VOL. 39, NO. 13, 2005 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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(6) Gallard, H.; De Laat, J. Kinetic modeling of Fe(III)/H2O2 oxidation reactions in dilute aqueous solution using atrazine as a model organic compound. Water Res. 2000, 34, 3107-3116. (7) Walling, C. Fenton’s reagent revisited. Acc. Chem. Res. 1975, 8, 125-131. (8) Yoon, J.; Lee, Y.; Kim, S. Investigation of the reaction pathway of OH radicals produced by Fenton oxidation in the conditions of wastewater treatment. Water Sci. Technol. 2001, 44, 15-21. (9) Haber, F.; Weiss, J. J. The catalytic decomposition of hydrogen peroxide by iron salts. Proc. R. Soc. A 1934, 134, 332-351. (10) Barb, W. G.; Baxendale, J. H.; George, P.; Hargrave, K. R. Reactions of ferrous and ferric ions with hydrogen peroxide. Part I. The ferrous ion reaction. Trans. Faraday Soc. 1951, 97, 462-500. (11) Bossman, S. H.; Oliveros, E.; Go¨b, S.; Siegwart, S.; Dahlen, E. P.; Payawan, L., Jr.; Straub, M.; Wo¨rner, M.; Braun, A. M. New evidence against hydroxyl radicals as reactive intermediates in the thermal and photochemically enhanced Fenton reactions. J. Phys. Chem. A 1998, 102, 5542-5550. (12) Wink, D. A.; Wink, C. B.; Nims, R. W.; Ford, P. C. Oxidizing intermediates generated in the Fenton reagent: Kinetic arguments against the intermediacy of the hydroxyl radical. Environ. Health Perspect. 1994, 102 (Suppl. 3), 11-15. (13) Kremer, M. L. Mechanism of the Fenton reaction. Evidence for a new intermediate. Phys. Chem. Chem. Phys. 1999, 1, 35953605. (14) Buda, F.; Ensing, B.; Gribnau, M. C. M.; Baerends, E. J. DFT study of the active intermediate in the Fenton reaction. Chem. Eur. J. 2001, 7, 2775-2783. (15) Ensing, B.; Buda, F.; Blo¨chl, P. E.; Baerends, E. J. A Car-Parrinello study of the formation of oxidizing intermediates from Fenton’s reagent in aqueous solution. Phys. Chem. Chem. Phys. 2002, 4, 3619-3627. (16) De Laat, J.; Gallard, H. Catalytic decomposition of hydrogen peroxide by Fe(III) in homogeneous aqueous solution: Mechanism and kinetic modeling. Environ. Sci. Technol. 1999, 33, 2726-2732. (17) Kwan, W. P.; Voelker, B. M. Decomposition of hydrogen peroxide and organic compounds in the presence of dissolved iron and ferrihydrite. Environ. Sci. Technol. 2002, 36, 1467-1476. (18) Rivas, F. J.; Beltra´n, F. J.; Frades, J.; Buxeda, P. Oxidation of p-hydroxybenzoic acid by Fenton’s reagent. Water Res. 2001, 35, 387-396. (19) Kang, N.; Lee, D. S.; Yoon, J. Kinetic modeling of Fenton oxidation of phenol and monochlorophenols. Chemosphere 2002, 47, 915924. (20) De Heredia, J. B.; Torregrosa, J.; Dominguez, J. R.; Peres, J. A. Kinetic model for phenolic compound oxidation by Fenton’s reagent. Chemosphere 2001, 45, 85-90. (21) Flyunt, R.; Schuchmann, M. N.; Von Sonntag, C. A common carbanion intermediate in the recombination and protoncatalysed disproportionation of the carboxyl radical anion, CO2*-, in aqueous solution. Chem. Eur. J. 2001, 7, 796-799. (22) Allen, A. O. The Radiation Chemistry of Water and Aqueous Solutions; Van Nostrand: Princeton, NJ, 1961; pp 117-140. (23) Buxton, G. V.; Sellers, R. M. Acid dissociation constant of the carboxyl radical. J. Chem. Soc., Faraday Trans. 1 1973, 69, 555559.

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9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 39, NO. 13, 2005

(24) Greenberg, A. E., Clesceri, L. S., Eaton, A. D., Eds. Standard Methods for the Examination of Water and Wastewater, 18th ed.; 1992. (25) Fortune, W. B.; Mellon, M. G. Determination of iron with o-phenanthroline. Ind. Eng. Chem., Anal. Ed. 1938, 10, 60-64. (26) Bader, H.; Sturzenegger, V.; Hoigne´, J. Photometric method for the determination of low concentrations of hydrogen peroxide by the peroxidase catalyzed oxidation of N,N-diethyl-p-phenylenediamine (DPD). Water Res. 1988, 22, 1109-1115. (27) Voelker, B. M. Ph.D. Thesis, Swiss Federal Institute of TechnologysZurich, Zurich, Switzerland, 1994. (28) Braun, W.; Herron, J. T. Acuchem/Acuplot; National Bureau of Standards: Gaithersburg, MD, 1986. (29) Ianni, J. C. Kintecus v3.0, Windows Version, http://www.kintecus.com, 2002. (30) Baxendale, J. H.; Khan, A. A. The pulse radiolysis of pnitrosodimethylaniline in aqueous solution. Int. J. Radiat. Phys. Chem. 1969, 1, 11-24. (31) Staehelin, J.; Hoigne´, J. Decomposition of ozone in water in the presence of organic solutes acting as promoters and inhibitors of radical chain reactions. Environ. Sci. Technol. 1985, 19, 12061213. (32) Ilan, Y.; Rabani, J. On some fundamental reactions in radiation chemistry: Nanosecond pulse radiolysis. Int. J. Radiat. Phys. Chem. 1976, 8, 609-611. (33) Atkins, P. W. Physical Chemistry; Oxford University Press: Oxford, U.K., 1995. (34) Sedlak, D. L.; Hoigne´, J. The role of copper and oxalate in the redox cycling of iron in atmospheric waters. Atmos. Environ. 1993, 27A, 2173-2185. (35) Raghaven, N. V.; Steenken, S. Electrophilic reaction of the OH radical with phenol. Determination of the distribution of isomeric dihydroxycyclohexadienyl radicals. J. Am. Chem. Soc. 1980, 102, 3495-3499. (36) Anderson, R. F.; Patel, K. B.; Stratford, M. R. L. Radical spectra and product distribution following electrophilic attack by the OH* radical on 4-hydroxybenzoic acid and subsequent oxidation. J. Chem. Soc., Faraday Trans. 1 1987, 83, 3177-3187. (37) Klein, G. W.; Bhatia, K.; Madhavan, V.; Schuler, R. H. Reaction of hydroxyl radicals with benzoic acid. Isomer distribution in the radical intermediates. J. Phys. Chem. 1975, 79, 1767-1774. (38) Zepp, R. G.; Faust, B. C.; Hoigne, J. Hydroxyl radical formation in aqueous reactions (pH 3-8) of iron(II) with hydrogen peroxide: The photo-Fenton reaction. Environ. Sci. Technol. 1992, 26, 313-319. (39) Maskos, Z.; Rush, J. D.; Koppenol, W. H. The hydroxylation of the salicylate anion by a Fenton reaction and Γ-radiolysis: A consideration of the respective mechanisms. Free Radical Biol. Med. 1990, 8, 153-162.

Received for review October 18, 2004. Revised manuscript received April 7, 2005. Accepted April 8, 2005. ES048378A