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APPLIED CHEMISTRY Mechanistic Study of the ZEA Organic Pollutant Degradation System: Evidence for H2O2, HO•, and the Homogeneous Activation of O2 by FeIIEDTA Derek F. Laine, Alexander Blumenfeld, and I. Francis Cheng* Department of Chemistry, UniVersity of Idaho, Moscow, Idaho 83844-2343
The ZEA (zero valent iron, ethylenediaminetetraacetic acid (EDTA), and air) organic pollutant degradation system has been previously shown to degrade a variety of organic pollutants and chemical warfare agent surrogates; however, mechanistic details and reactive intermediates formed in this system have not been identified. It is hypothesized that the ZEA system produces reactive oxygen species (H2O2, HO•) by the reduction of oxygen by FeIIEDTA(aq). This hypothesis is examined through an electrochemical model of the ZEA system. A carbon basket electrode is used as the reducing agent in place of Fe(0). The FeIIIEDTA complex (0.5 mM) is electrochemically reduced to FeIIEDTA at an applied potential of -120 mV (vs Ag/ AgCl) under aerobic conditions. Hydrogen peroxide was observed to form in the presence of the metal complex with a maximum concentration reaching 0.139 mM H2O2 after 3 h of electrolysis. In the absence of FeEDTA, 0.04 mM H2O2 is obtained by the direct reduction of O2 at the electrode surface. Electron resonance spectroscopy (ESR), along with 5,5-dimethyl-1-pyrroline-N-oxide (DMPO) as a spin trap and methanol as a radical scavenger confirms the formation of HO• produced via the Fenton reaction in the electrochemical system. Hydroxyl radical attack on EDTA caused the degradation of FeII/IIIEDTA to a steady-state concentration of 0.14 mM from 0.5 mM as observed by HPLC. The pH of the electrolysis solution increased from 2.64 to 9.25 during 6 h of reductive electrolysis which is indicative of the consumption of H+ during the reduction of O2 to form H2O2. These experiments provide evidence that the ZEA system uses atmospheric O2 to produce reactive oxygen species including those that deeply oxidize organics under room temperature and pressure conditions. Introduction The use of O2 as an abundant and versatile oxidant under ambient conditions offers considerable green chemistry potential.1–3 However, at room temperature and pressure, the slow reactivity of triplet oxygen imposes kinetic limitations.4,5 Excluding high temperature processes, O2 can be activated by its transformation to a singlet state by excitation with light, or by univalent reduction.4–6 Most low temperature O2 activation processes are found in biological systems, especially those based on iron containing metalloenzymes such as cytochrome P-450.7–9 These enzymatic systems have mastered the use of oxygen for selective oxidations in biological processes under ambient conditions.10–17 Reduction of O2 can produce kinetically facile oxygen species, including H2O2, O2•-, and HO• which are known collectively as reactive oxygen species (ROS).18 Through various reaction routes these species are able to degrade organic compounds at low temperatures and pressures to CO2/HCO3-/CO32-, C2O42- and other low molecular weight acids and this provides a route for “green” pollutant destruction.1–3,19–21 There are relatively few examples of low temperature, nonbiological O2 activation systems, and even fewer involve application toward pollutant degradation and environmental remediation.1–3,7,10 One such process is known as the ZEA organic pollutant degradation system.1–3 This system is both unique and simple in that it consists of only zero valent iron (20-40 mesh), ethylenediaminetetraacetic acid (EDTA), and air (ZEA) in an aqueous slurry containing the target organic * To whom correspondence should be addressed. E-mail: ifcheng@ uidaho.edu. Tel.: (208) 885-6387. Fax: (208) 885-6173.
pollutant. With the system exposed to the atmosphere, a variety of compounds, including 4-chlorophenol, pentachlorophenol, phenol, and malathion have been nonselectively and deeply oxidized to low molecular weight acids and CO2 under ambient conditions.1–3 It has been hypothesized in a previous report that the ZEA system may operate by three possible mechanisms.1 These include (1) the heterogeneous activation of O2 by Fe(0), (2) homogeneous activation of O2 by FeIIEDTA, and (3) heterogeneous activation producing ferryl as the oxidization agent.1 However, which mechanism the ZEA system operates has not been resolved and the oxidizing agents that the system produces to degrade organic pollutants has not been determined. In this work, the homogeneous O2 activation mechanism is proposed (reactions 1–6) and examined. In this case, the formation of HO• and H2O2 in the ZEA system is hypothesized to be driven primarily by the FeIIEDTA complex. This assumption is based off of literature sources that indicate many Fe(II) complexes are able to reduce O2 to ROS.22–31 In the ZEA system, FeIIEDTA forms as a result of the presence of EDTA which complexes Fe2+ produced from the dissolution of Fe(0) (reactions 1 and 2). In turn, FeIIEDTA reduces O2 as in reactions 3 and 4.29 The next logical steps are the formation of H2O2 from reduction of O2 or O2•- (reactions 4 and 5) followed by the Fenton reaction releasing HO• (reaction 6). It is this radical with E0 ) 2.02 V vs SHE (•OHaq/OHaq-)32 that is thought to be responsible for the nonselective and deep oxidation of organic pollutants in the ZEA system.33–38 The formation of H2O2 and HO• in the ZEA system has yet to be confirmed.
10.1021/ie701676q CCC: $40.75 2008 American Chemical Society Published on Web 07/31/2008
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Reactions 3–6 provide a reasonable pathway for the homogeneous formation of HO•. In the presence of an excess of reducing agent, in this case Fe(0), the sequence of reactions 3–6 can redox cycle several times through the reduction of FeIIIEDTA to FeIIEDTA. Fe(0) f Fe2+ + 2e2+
Fe
(1)
+ EDTA f Fe EDTA II
(2)
Fe EDTA + O2 f Fe EDTA - O2 II
II
•-
(3) •-
Fe EDTA - O2 f Fe EDTA + O2 II
III
(4)
FeIIEDTA + O2•- + 2H+ f FeIIIEDTA + H2O2
(5)
FeIIEDTA + H2O2 f FeIIIEDTA + HO- + HO•
(6)
In order to examine the possibility for the reduction of O2 by FeIIEDTA to ROS via the mechanism of reactions 3–6, a carbon basket electrode is employed in place of Fe(0). The potential of this electrode is held at a point where it is capable of reducing FeIIIEDTA. This allows redox cycling between the FeII/ IIIEDTA oxidation states as it undergoes electro-reduction then oxidation by dissolved O2 via reactions 3–6. Since the ZEA system has been shown to degrade multiple aqueous organic compounds, the degradation of EDTA in the electrochemical system was monitored as evidence of the proposed homogeneous activation mechanism (reactions 3–6).1–3 Beneficial to this method is the ability to control the oxidation state of the iron in the EDTA complex demonstrating the importance of FeIIEDTA in the ZEA system and in the production of ROS formed from O2 reduction. The reactions involved in the electrochemical cell are shown in Figure 1. Experimental Details Chemicals and Solutions. All chemicals were used as obtained without further purification. Solutions were prepared using deionized water that was further purified by passage of house distilled water through a purification cartridge (Barnstead, model D8922, Dubuque, IA). A stock solution of 12 mM EDTA was prepared by dissolution of Na2H2EDTA · 2H2O (J.T. Baker, NJ, 99.9%). A stock solution of 14 mM Fe3+ was prepared by dissolution of Fe(NO3)3 · 9H2O (Fisher, NJ, 99.3%) in deionized water and pH adjusted with HNO3 to ca. pH 1.0. The FeIIIEDTA complex was formed from equal molar amounts of Fe3+ and EDTA prior to analysis. Powdered starch, KCl (1.5 M stock solution), and KI were obtained from Fisher (certified ACS grade, NJ). The 1 M KI stock solutions were made fresh daily and concentrated starch stock solutions were made fresh weekly. Ammonium heptamolybdate, (NH4)6Mo7O24 · 4H2O, was obtained from Sigma-Aldrich (99.98%, St. Louis, MO) and was dissolved in 0.5 M H2SO4 (EMD, ACS grade) for a 1 mM stock solution. Concentrated HCl was obtained from EMD (ACS grade, Darmstadt Germany) and was used to make 1 M stock solution. Peroxidase test strips where obtained from EM science, Gibbstown, New Jersey. Methanol was obtained from Fisher (ACS reagent grade, Fair Lawn, NJ). The 5,5-dimethyl-1pyrroline-N-oxide (DMPO) solution was obtained from Fisher (high purity, Fair Lawn, NJ). A 90 mM DMPO stock solution was made up in deionized water and kept under a N2 atmosphere, refrigeration, and in the dark. The DMPO stock solution was purified periodically by adding activated charcoal (Fisher, Fair Lawn, NJ) with stirring for several minutes followed by filtration.39 This step was repeated until no contaminate electron resonance spectroscopy (ESR) signals were observed.
Figure 1. Reaction scheme in the electrochemical oxygen activation system mediated by FeII/IIIEDTA. This system acts as a chemical analog to the ZEA organic pollutant degradation system. Table 1. Various Reaction Conditions during Controlled Potential Electrolysisa I II III IV a
-120 +120 -120 -120
mV, mV, mV, mV,
O 2, O2 , N2 , O2 ,
0.5 0.5 0.5 0.0
mM mM mM mM
FeEDTA FeEDTA FeEDTA FeEDTA
All solutions contain 100 mM KCl.
Procedure. All electrode potentials are reported versus the Ag/AgCl/3 M NaCl reference electrode (BAS, West Lafayette, IN). A bulk electrolysis cell (description below) was filled with 0.5 mM FeIIIEDTA and 100 mM KCl solution. Oxygen gas was bubbled in the electrochemical cell for 30 min prior to and throughout electrolysis to ensure saturation. A potential of -120 mV was applied to the carbon basket working electrode for bulk electrolytic controlled potential coulometry using a Bioanalytical Systems CV-50W potentiostat (West Lafayette, IN). Under these conditions, FeIIIEDTA is reduced to FeIIEDTA. After 1 h, the electrolysis was terminated and samples were taken for analysis (described below). A new solution was made and the electrolysis was allowed to run for 2 h. This procedure was repeated in triplicate for each of 6 h. The electrochemical cell was kept under room temperature and pressure conditions and exposed to the laboratory atmosphere. The pH of the solution was measured before and after each electrolysis by an Accumet model 20 pH/conductivity meter. Control experiments include bubbling of nitrogen gas (99.99%, Oxarc, Spokane Washington) to deoxygenate the solution and the application of a positive potential to the working electrode so that FeIIIEDTA reduction does not occur. To investigate reactions 3–6, several control experiments were conducted under various reaction conditions. For clarity, these reaction conditions are presented in Table 1 as conditions I, II, III, and IV. Hereafter, discussion will refer to these numbers to represent the conditions of each experiment. Analyses. FeII/IIIEDTA was quantified by an HPLC method developed previously.40 H2O2 was quantified by the starch iodide method,41 employing a UV-vis spectrophotometer (HewlettPackard, model 8453). In this method, hydrogen peroxide oxidizes iodide ions (77 mM KI) to iodine/triiodide according to reactions 7 and 8. Triiodide, in the presence of starch solution (38.5 mM) and ammonium molybdate catalyst (0.077 mM), will form a complex with an intense blue color (570 nm). After each hour of electrolysis, an aliquot was taken from the reaction mixture and added to the starch reagents. H2O2(aq) + 3I- - (aq) + 2H+(aq) f I3-(aq) + 2H2O(aq) (7) I3- + starch f blue complex
(8)
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The color formation process is time dependent; therefore, a 20 min delay between mixing and absorbance measurements for both calibration standards and samples was established. The HO• was qualified using electron resonance spectroscopy (ESR) with DMPO as a radical trap. Spectra were taken with a Bruker 6/1 spectrometer with a resonant frequency of 9.861 GHz, microwave power of 20.2 mW, modulation frequency at 100 KHz, modulation amplitude 1.00 G, sweep width 100.000 G, time constant 163.840 ms, conversion time 163.840 ms, sweep time 167.772 s, receiver gain 2.00 × 105, number of data points (resolution) 1024, and number of scans 8. At appropriate time intervals, 7 mL of electrolysis mixture was taken from the cell during electrolysis and added to a vial containing 1 mL of the spin trap DMPO (final concentration of 11.04 mM). This sample was immediately placed into an aqueous sample cell (Bruker, AquaX high sensitive aqueous sample cell) and tuned for ESR measurements. To show the affect of radical scavenging and therefore verify the formation of HO• radicals in the electrolysis cell, 2.625 mL of electrolysis mixture was taken at 3 h of electrolysis and added to a vial containing methanol (final concentration of 20% v/v) and DMPO (final concentration of 8.28 mM). This procedure was also applied to the initial electrolysis mixture before electrolysis to show the possibility of nucleophilic addition of methanol to DMPO. Electrochemical Cell. The electrochemical cell consisted of a high surface area reticulated vitreous carbon basket working electrode (BAS, West Lafayette, IN), height 50 mm, diameter 40 mm, and thickness 5 mm, placed into a 100 mL beaker. A graphite rod counter electrode (99% Alfa Aesar, Ward Hill, MA) was isolated from the bulk solution by a 12 mm glass tube with an agar/saturated KCl-porous frit salt bridge (Ace glass, Inc., Vineland, NJ). The counter electrode solution was a 100 mM KCl solution and the reference electrode consisted of Ag/AgCl/3 M NaCl (BAS, West Lafayette, IN). Results and Discussion Degradation of EDTA. The production of HO• from the sequence of reactions 3–6 drives the oxidative damage to EDTA.3 The loss of EDTA is demonstrated in a liquid chromatographic analysis of the controlled-potential electrolysis solutions.3 Under condition I of Table 1, the 0.5 mM FeII/ IIIEDTA complex is redox cycled between the electrode and O2 (see Figure 1). This action produces HO• which deeply oxidizes EDTA. An absorption peak (258 nm) at a retention time of 2.2 min was observed and assigned to FeIIIEDTA.40 Under condition I, this peak decreased with electrolysis time from a concentration of 0.5 mM to a steady state value of 0.14 mM after 4 h of electrolysis (Figure 2C). Also observed is the formation of a peak that is most likely an intermediate (retention time ca. 1.9 min, data not shown) after the first hour of electrolysis, which decays thereafter. Reactions 3 and 4 in the proposed mechanism suggest that EDTA degradation may only occur when FeEDTA is in the 2+ oxidation state and in the presence of O2. This was qualified by performing control experiments where the experiment was repeated under conditions II and III in Table 1. Figure 2A shows that when the working electrode potential is adjusted positive of the FeIIIEDTA reduction potential (condition II), no EDTA degradation occurs, despite the presence of O2. Figure 2B shows that when the reaction mixture is deoxygenated by nitrogen purge (condition III), no degradation is observed, although the electrode drives the reduction of FeIIIEDTA to FeIIEDTA. These control experiments indicate that only when FeIIEDTA is formed
Figure 2. Degradation of FeIIIEDTA by electrochemical oxygen activation under various initial conditions (as described in Table 1). All experiments were conducted in 0.1 M KCl with a carbon basket electrode. Each data point represents the average of three runs with the error bars indicating one standard deviation unit. (A) Condition II, +120 mV, O2, 0.5 mM FeEDTA. (B) Condition III, -120 mV, N2, 0.5 mM FeEDTA. (C) Condition I, -120 mV, O2, 0.5 mM FeEDTA.
in the presence of O2 does EDTA degradation occur. Each curve in Figure 2, when taken together, gives support for the proposed mechanism for the ZEA system (reactions 1–6). In particular, oxygen activation by FeIIEDTA in reactions 3 and 4 are supported by these results. The incomplete degradation of FeIIIEDTA (from 0.5 to 0.14 mM) in Figure 2C may be associated with an increase in solution pH from 2.64 to 9.25 after 6 h of electrolysis. Increases in pH are expected because of the consumption of H+ in reaction 5, and production of HO- in reaction 6. The ability of FeIIEDTA to reduce O2(aq) has been shown to be pH dependent with the optimum value at pH 3 as indicated by previous studies.2,3,23,29,33,42 At pH 8 and above the rate of O2, reduction by FeIIEDTA drops off rapidly.29 The optimal complex for O2 reduction is presumed to be a protonated form, FeIIEDTA(H) whereas above pH 8 the hydroxylated forms of FeEDTA (FeIIEDTA(OH) and FeIIEDTA(OH)2) tend to predominate (see Figure 3). In terms of the former, protonation of one of the carboxylate arms of EDTA will create an open metal-centered site for the coordination of O2.29,43 Hydrolyzed forms of FeIIEDTA may prevent coordination of O2 (Figure 3).23,29 Previous work on the ZEA system has shown that increased EDTA concentration reduces the rate of its degradation.3 On the basis of another previous study, it was noted that this is most likely due to a kinetic barrier to O2 and/or H2O2 reductions at high [EDTA]:[Fe] ratios.44 It was hypothesized in that study that a high [chelate]:[Fe] ratio prevents binding of H2O2 to the metal coordination sphere. A later study confirmed this hypothesis as Fenton reaction activity is recovered when Ca2+ introduced to solutions of high [EDTA]:[Fe] ratios.45 The calcium ion binds excess EDTA (β ) 10.69) with displacing Fe2+ (β ) 14.32) or Fe3+(β ) 25.1) from that chelate. Formation of Hydrogen Peroxide. In the proposed mechanism of reactions 1–6, H2O2 is suggested to form following the reduction of O2. The result of electrolysis at -120 mV (condition I of Table 1) is shown in Figure 4A where a maximum concentration of 0.139 mM H2O2 is reached after 3 h of electrolysis. Figure 4B shows that when FeIIIEDTA is absent from the reaction mixture, i.e. condition IV, some H2O2 is observed to form due to the direct reduction of O2 at the electrode surface (reaction 10). O2 + 2e- + 2H+ f H2O2 E0 ) 0.695 V vs SHE
(9)
This reaction reaches a near steady-state concentration of 0.04 mM H2O2 after 3 h of electrolysis. The control experiments in
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Figure 3. FeIIEDTA speciation diagram showing the dominant species at pH ) 3. Speciation was constructed using Hyperquad Speciation and Simulation with 0.5 mM FeIIEDTA as the initial condition.
outlined in the sequence of reactions 3–6 can be clarified by using the electrochemical analog of the ZEA system shown in Figure 1. Under reaction condition I, HO• is detected by employing DMPO as a spin trap along with ESR spectroscopy. The addition of HO• to DMPO is shown below in reaction 10. The ESR signal assigned to the product of this reaction (DMPO•-HO, compound 11) has a well represented 1:2:2:1 quartet with aN ) aH ) 14.9 G from the hyperfine splitting from both the β-hydrogen as well as the nitroxide nitrogen.49,50 Figure 4. Formation of hydrogen peroxide by electrochemical oxygen activation under various initial conditions (as described by Table 1). All experiments were conducted in 0.1 M KCl with a carbon basket electrode. Each data point represents the average of three runs with the error bars indicating one standard deviation unit. (A) Condition I, -120 mV, O2, 0.5 mM FeEDTA. (B) Condition IV, -120 mV, O2, 0.0 mM FeEDTA. (C) Condition III, -120 mV, N2, 0.5 mM FeEDTA.
Figure 4C and D (conditions III and V, respectively) show negligible H2O2 production as expected from the proposed mechanism. It is apparent from the results of Figure 4 that mediated reduction of oxygen by the FeII/IIIEDTA redox couple produces H2O2, i.e. reactions 3–5. The production of H2O2 during electrolysis was further confirmed by the color formation of peroxidase based analytical test strips that were dipped in the reaction mixture periodically during electrolysis. The transient nature of FeIIEDTA induced H2O2 production in Figure 4A may be explained by the consumption of H2O2 via reaction 6 and loss of FeEDTA via oxidation by HO• (Figure 2). The steady-state concentration of 0.04 mM H2O2 after 2 h of direct electrolysis of O2 may be explained by the pH dependence of reaction 9. The electrolysis solution was observed to increase and plateau to pH 10 after 2 h thereby decreasing the electrochemical driving force for O2 reduction to H2O2. In a previous study it was observed that the rate of EDTA degradation by HO• via reactions 1–6 was strongly dependent on the [EDTA]:[FeII/III] ratio.3 In this work, the affect was not significant on the production of H2O2. The observation in the previous study may be attributed to the affect of EDTA on the rate of Fe(0) dissolution.46–48 Formation of HO• The ZEA organic pollutant degradation system is hypothesized to generate HO• that, in turn, can deeply oxidize organic compounds (reaction 6).1–3 The plausibility of generating HO• from the reduction of O2 via the mechanism
The 1:2:2:1 quartet is indeed observed in the system of Figure 1 when the reaction mixture was added to DMPO after a 3 h electrolysis period (Figure 5A). However, Figure 5B shows that the same 4-line spectra is obtained without electrolysis and that the signal arises when a solution of FeIIIEDTA, KCl, and O2 (aq) is added to DMPO. This adventitious signal has been given some attention in the literature and is generally accepted to be from the nucleophilic addition of water to DMPO catalyzed by Fe3+ (reaction 11) producing the same DMPO•-HO adduct (compound 11) as in reaction 11 (compound 11).49,51
Alternative evidence for HO• was found by adding the electrolysis mixture to DMPO along with methanol as a radical scavenger.49 The HO• reacts with CH3OH producing the hydroxymethyl radical (reaction 12). In turn, this radical combines with DMPO forming DMPO•-CH2OH which has a characteristic ESR signal (compound 14 in Reaction 13). Figure
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Figure 7. Kinetic trace of the formation of DMPO-HO•. The magnetic field was fixed on the second peak of the 1:2:2:1 quartet. (A) Formation of DMPO-HO• (compound 11) after 3 h of electrolysis of 0.5 mM FeIIIEDTA due to addition of hydroxyl radicals. (B) Formation of DMPO-HO• (compound 12) from a solution of 0.5 mM FeIIIEDTA before electrolysis due to water addition.
Figure 5. Comparison of the ESR signal obtained by (A) adding the electrolysis mixture under condition I to DMPO after 3 h of electrolysis to produce compound 11 and (B) adding 0.5 mM FeIIIEDTA and 100 mM KCl to DMPO to produce compound 12 (no electrolysis). Both solutions were saturated with O2.
Figure 6. (A) 0.5 mM FeIIIEDTA reaction mixture added to DMPO and methanol before electrolysis: (b) compound 12; (O) compound 15. (B) Reaction mixture added to DMPO and methanol after 3 h of electrolysis under condition I; (b) compound 11; (O) compound 14.
6A shows the results of adding the reaction mixture (FeIIIEDTA, KCl, O2) to a solution of DMPO and CH3OH without electrolysis. This produced a mixture of signals that include the 4-line spectra from reaction 11 (compound 12, represented as b in Figure 6A) as well as the product from the nucleophilic addition of CH3OH to DMPO shown below in reaction 14 (compound 15, represented as O in Figure 6A) However, when
the reaction mixture is added to DMPO/CH3OH after electrolysis
when HO• is expected to form, an entirely different spectra is obtained as shown in Figure 6B. Figure 6B shows a mixture of signals due to HO• competitively reacting with both DMPO and CH3OH (reaction 10 vs reactions 11 and 13) where the signal due to compound 11 is mixed with a 6-line signal from compound 14 (labeled as O in Figure 6B). Figure 6A shows a 6-line spectrum attributable to DMPO-OCH3 (reaction 14, compound 15). In Figure 6B, another 6-line spectrum is produced but this one is from DMPO-CH2OH (compound 14) from reactions 12 and 13. Compound 14 can be distinguished from compound 15 through their unique ESR 6-line signals which have different relative coupling constants as indicated in Figure 6. From the results of Figure 6, compound 15 gave coupling constants aN ) 15.0 G and aH ) 10.2 G while compound 14 gave coupling constants of aN ) 15.8 G and aH ) 22.8 G. These values are consistent with the literature.49,52 To further distinguish reactions 10 and 11, the rate of increase in the intensity of the 1:2:2:1 quartet of Figure 5 was monitored by measuring the increase in ESR intensity of the second peak in the ESR spectrum. Reaction 10 is faster than reaction 11 with a reported rate constant of about 109 M-1 s-1.49 Figure 7A shows the ESR signal rapidly increases from the formation of compound 11 in reaction 10 and corresponds to Figure 5A. Figure 7B demonstrates that reaction 11 is much slower to produce the same product (compound 12) corresponding to the signal of Figure 5B. The results of Figure 7 are reproducible and confirm that there are different mechanisms by which the DMPO•-OH radical can be formed, as indicated by reactions 10 and 11. With the formation of •CH2OH by hydroxyl radical attack on methanol, it is confirmed that the system of Figure 1 does in fact produce HO•. Therefore, reaction 6 of the proposed ZEA pollutant degradation system is supported. The provided results give evidence for a plausible mechanism (reactions 1–6) for the ZEA organic pollutant degradation system. With the mechanism of the ZEA reaction established, a route to further advance the ZEA system as it applies to a green method of environmental remediation can be anticipated. Future investigations will be aimed at improving the rate at which the ZEA reaction produces ROS from dissolved oxygen. Acknowledgment D.F.L. acknowledges the summer support from the Malcolm Renfrew Scholarship Fund. This investigation was supported by NSF award number BES-0328827. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the author(s) and do not necessarily reflect the views of the National Science Foundation.
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ReceiVed for reView December 10, 2007 ReVised manuscript receiVed May 20, 2008 Accepted May 27, 2008 IE701676Q
