ARTICLE pubs.acs.org/est
Strong Enhancement on Fenton Oxidation by Addition of Hydroxylamine to Accelerate the Ferric and Ferrous Iron Cycles Liwei. Chen,† Jun. Ma,†,‡,* Xuchun. Li,† Jing. Zhang,† Jingyun. Fang,† Yinghong. Guan,† and Pengchao. Xie† † ‡
State Key Laboratory of Urban Water Resource and Environment, Harbin Institute of Technology, Harbin 150090, PR China; National Engineering Research Center of Urban Water Resource and Environment, Harbin 150090, PR China.
bS Supporting Information ABSTRACT: The Fenton system generates reactive species with high oxidation potential such as hydroxyl radicals (HO•) or ferryl via the reaction between Fe (II) and H2O2. However, a number of drawbacks limit its widespread application including the accumulation of Fe (III) and the narrow pH range limits, etc. The aim of this study is to propose a much more efficient Fenton-HA system which is characterized by combining Fenton system with hydroxylamine (NH2OH), a common reducing agent, to relieve the aforementioned drawbacks, with benzoic acid (BA) as the probe reagent. The presence of NH2OH in Fenton’s reagent accelerated the Fe (III)/Fe (II) redox cycles, leading to relatively steady Fe (II) recovery, thus, increased the pseudo first-order reaction rates and expanded the effective pH range up to 5.7. The HO• mechanism was confirmed to be dominating in the Fenton-HA system, and the generation of HO• was much faster and the amount of HO• formed was higher than that in the classical Fenton system. Furthermore, the major end products of NH2OH in Fenton-HA system were supposed to be NO3 and N2O.
’ INTRODUCTION The Fenton’s reagent (H2O2 and Fe (II)), which was first reported over 100 years ago,1 generated hydroxyl radicals (HO•)2 to initiate free radical chain reactions.3 Although the mechanism related to the formation of reactive species is still controversial because ferryl might also be involved in Fenton system,4 it is generally accepted that HO• is the major oxidant.5 The Fenton system is attractive because of its fast reaction rates, low toxicity, and simplicity to control. Nevertheless, the Fenton system has some intrinsic drawbacks that limit its widespread application, such as strict pH range limits, high H2O2 dosage, and the accumulation of ferric oxide sludge, which causes the decline of oxidation rates and requires a separation step.5 Some of these shortcomings may be alleviated by technological developments. Consequently, researches have been focusing on the optimization of Fenton system in recent years. Several transition metals were selected as the substitutes for ferrous iron in the oxidative treatment.6 Meanwhile, the heterogeneous oxidation technologies were widely investigated to avoid Fe (III) sludge formation and to expand the effective pH range.7 Another effort was to add some chelating agents into the Fenton system to alleviate the ferric iron precipitation.8 A third effort was to combine Fenton’s reagent with other technologies to enhance the generation of HO•. Some researchers introduced UV radiation9 into Fenton system, which accelerated the circulation of Fe (III)/Fe (II) and the oxidation reactivity. Others adopted electrochemistry10 in Fenton system, but the optimum pH ranges were even narrower than that in classical Fenton system. Some scientists added quinone11 and r 2011 American Chemical Society
humic acid12 to Fenton system, which acted as catalyst to facilitate the transformation of Fe (III) to Fe (II). The major drawback of the Fenton system was the accumulation and precipitation of Fe (III), which could further result in the decline of reaction rates and narrow optimal pH range. For the purpose of reducing Fe (III) to Fe (II) whose solubility is much higher than ferric iron, we could take into account some reducing agents with low reaction rates to reactive species. As a common reductive chemical to reduce Fe (III), hydroxylamine (NH2OH) was adopted in many applications such as total iron determination.13 Besides, it was frequently used to modify the PAN fiber14 and to activate catalyst15in advanced oxidation. However, NH2OH has never been reported to react directly with Fenton’s reagent in literature for it is generally referred as an antioxidant.16 NH2OH was employed in this study for its selective reactivity with HO•.17 In addition, the reductive product of NH2OH and Fe (III) are generally inorganic substances like N2, N2O, NO2, and NO3.1719 The objective of this study is to explore the role of NH2OH in the redox cycles of Fe (II)/Fe (III) in Fenton’s reagent and the mechanism of the Fentonhydroxylamine (Fenton-HA) system. As a HO• probe,20 benzoic acid (BA) was selected as the model compound. The addition of NH2OH to Fenton-HA system could accelerate the redox cycles of Fe (III) to Fe (II). Hence, compared Received: September 16, 2010 Accepted: March 29, 2011 Revised: March 28, 2011 Published: April 06, 2011 3925
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Table 1. Pseudo First-Order Rate Constants in Fenton System at pH 3.0a Kobs1
t1/2
(μM)
(mM)
(min1)
(min)
10.0 20.0
0.4 0.4
0.151 0.262
40.0
0.4
10.0
0.04
10.0 10.0 10.0
[Fe(II)]0 [H2O2]0
a
Kobs2 R12 (min1)
Table 2. Pseudo First-Order Rate Constants in Fenton-HA System at pH 3.0
t1/2
[Fe(II)]0
[H2O2]0
[NH2OH]0
Kobs
t1/2
(μM)
(mM)
(mM)
(min1)
(min)
R2
43.322 0.957 17.773 0.984
10.0 20.0
0.4 0.4
0.4 0.4
0.136 0.254
5.097 2.729
0.993 0.996
(min)
R22
4.590 0.982 2.646 0.915
0.016 0.039
0.520
1.333 0.927
0.140
4.951 0.998
40.0
0.4
0.4
0.541
1.281
0.997
0.023
30.137 0.899
0.012
57.762 0.968
10.0
0.04
0.4
0.020
34.657
0.982
0.2
0.098
7.073 0.959
0.016
43.322 0.940
10.0
0.2
0.4
0.082
8.453
0.993
1.0
0.186
3.727 0.749
0.019
36.481 0.940
10.0
1.0
0.4
0.208
3.332
0.994
2.0
0.248
2.795 0.660
0.015
46.210 0.992
10.0
2.0
0.4
0.277
2.502
0.994
Note: the first stage was within 1.5 min
with Fenton system, the new system notably alleviated the Fe (III) accumulation, enhanced the reaction rates and expanded the effective pH range. The HO• was confirmed as the dominating reactive species and the generation of HO• was much faster in the Fenton-HA system due to the participation of NH2OH. In addition, the NO3 and N2O were proposed as the dominant end products of NH2OH.
’ MATERIALS AND METHODS Materials. The following chemicals are reagent grade. Benzoic acid, hydroxylamine hydrochloride (99.999%), acetic acid, sodium acetate, monosodium phosphate, sodium hydrogen phosphate, 5, 5-dimethyl-1-pyrolin-N-oxide (DMPO), potassium hydroxide, hydrogen-peroxide oxidoreductase (POD), and N, N-diethyl-p-phenylenediamine (DPD) were supplied by SigmaAldrich. Hydrogen peroxide (H2O2, 35% v/v, stab.) was from Alfa Aesar. Sodium nitrite, ferrous sulfate, perchloric acid, tertbutyl alcohol, ascorbic acid, and sodium chloride were purchased from Sinopharm Chemical Reagent Co., Ltd. Methanol (Tedia), acetone (Tedia), methyl tert-butyl ether (MTBE from Fisher) and phosphoric acid (Dikma) were of HPLC grade. Procedures. All experiments were carried out at 20 ( 0.5 °C with 250 mL triangular flasks under constant stirring with a PTFE-coated magnetic stirrer in Milli-Q water (18.2 MΩ 3 cm). BA, NH2OH 3 HCl and Fe (II) with desired concentrations were spiked in 100 mL perchloric acid or acetic acid-sodium acetate buffer. Each run was switched on by adding the desired dosage of H2O2. The pH changed less than (0.2 units during the process. Samples were withdrawn at set intervals and quenched by sodium nitrite (for BA) or ascorbic acid (for p-hydroxybenzoic acid (PHBA)) before analysis. The quenching experiments employed tert-butyl alcohol as the quencher, which were introduced in excess immediately after the addition of H2O2. EPR experiments were conducted and run in duplicate; the trend was similar to those shown in Figure 3. The detailed parameters are shown inSupporting Information (SI) Text S4. The chloride ion introduced into the experiments was from sodium chloride. No obvious effect of chloride ion on BA oxidation was observed when its concentration was below 2 mM as demonstrated in SI Figure S4. Sample Analysis. The concentrations of BA and PHBA were analyzed on HPLC (Waters), equipped with 2487 dual λ absorbance detector, a reverse-phase C18 column, a 1525 binary HPLC pump, and a 717 plus autosampler. The pH was measured by a pH Meter (Ultrabasic 7 from Denver Instrument). The Fe (III)
concentrations were measured on UVvis spectrometer (Varian Cary 300 UVvis spectrometer) at 300 nm21 using 10 cm quartz sampling cells. The H2O2 concentrations were measured by colorimetric method using DPD22 on Varian Cary 300 UVvis spectrometer. The NH2OH samples were derivatized to acetone oxime by reaction with acetone and extracted by MTBE. The acetone oxime concentrations were measured by GC23 (gas chromatograph) 7890A (Agilent) using FID detector with heater of 300 °C and N2 as the carrier gas. The total nitrogen (TN) was measured by combining the TNM-1 with a TOC-V CPN analyzer (Shimadzu). The ammonia nitrogen was measured by Nessler’s reagent and organic nitrogen was equal to the difference between Kjeldahl nitrogen and ammonia nitrogen. The NO2 and NO3 concentration measurements were carried out on a Dionex ICS-3000 ion chromatograph system. The detailed procedures of sample analysis are shown in SI Text S4.
’ RESULTS AND DISCUSSION The Comparison of Fenton System with Fenton-HA System. The pseudo first-order rate constants (Kobs) in Fenton
system under different Fe (II) and H2O2 dosages during the process are shown in Table 1 which exhibits a two-stage oxidation. The oxidation of organic compounds in classical Fenton system generally proceeds through two stages: a fast stage and a much slower stage.24 The reaction rate of the fast stage depends on the initial Fe (II) concentration, and its rate-limiting step is shown in eq 1 or 2. The slower stage is due to the accumulation of Fe (III) and the bad recovery of Fe (II) for eqs 3 and 6 are sandare much slower than eq 1 or 2 (the oxidation product of ferryl is also Fe (III)). Meanwhile, the reaction of Fe (III) and H2O2 via eq 3 generates HO2•, whose oxidation ability is far inferior to HO• and ferryl. This two-stage phenomenon, however, did not fit the FentonHA system shown in Table 2. The whole process in Fenton-HA system fits well with pseudo first-order kinetics model. The Kobs and the half time (t1/2) in Fenton-HA system were similar to that in the fast stage in classical Fenton system under the same condition. It seems that the addition of NH2OH to Fenton system effectively avoided the Fe (III) accumulation and kept the oxidation reactivity of the fast stage in Fenton system. Besides this, NH2OH as a reducing reagent seems not to compete obviously with BA for reactive species (HO• or ferryl) under this pH.
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FeðIIÞ þ H2 O2 f FeðIIIÞ þ OH þ HO•
ð1Þ
FeðIIÞ þ H2 O2 f FeðIVÞðFeO2þ Þ þ H2 O
ð2Þ
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Figure 1. Effect of pH with buffer on the Ct/C0 of BA at the reaction time of 10 min in Fenton system and Fenton-HA system at [BA]0 = 40.0 μM; [H2O2]0 = 0.4 mM; [Fe (II)]0 = 10.0 μM; [NH2OH]0 = 0.4 mM.
FeðIIIÞ þ H2 O2 f FeðIIÞ þ HO•2 þ Hþ
ð3Þ
HO• þ H2 O2 f HO•2 þ H2 O
ð4Þ
HO• þ FeðIIÞ f FeðIIIÞ þ OH
ð5Þ
FeðIIIÞ þ HO•2 f FeðIIÞ þ O2 Hþ
ð6Þ
FeðIIÞ þ HO•2 þ Hþ f FeðIIIÞ þ H2 O2
ð7Þ
HO•2 þ HO•2 f H2 O2 þ O2
ð8Þ
Figure 1 presents the profile of the variation BA concentration at different pH buffered from 2.0 to 6.0 in Fenton and Fenton-HA systems. A maximum value for the Fenton system reactivity was observed at pH 3.0, being identical to those reported in literature.3,5 When the solution pH was higher than 5.0, the oxidation of BA in the Fenton system could be neglected. The Fenton-HA system showed the same optimal pH range from 3.0 to 4.0. The oxidation reactivity in Fenton-HA system was also gradually decreased along with the increase of pH in solution. No obvious BA oxidation could be seen under pH 6.0 in the FentonHA system. But the BA oxidation in the Fenton-HA system at pH 5.7 was still better than that achieved in the Fenton system at the optimal pH 3.0. Hence, compared to Fenton system, the FentonHA system has a wider effective pH range. The species of Fe (II) ion is dominant in strong acid solution. As pH increases, the Fe (OH)2 begins to form which is more reactive than Fe (II) ion.25 That is the reason why both the oxidation of BA in the two systems was lower at pH 2.0. The narrow optimal pH range observed in the Fenton system is primary due to the Fe (III) precipitation. The Fe (III) begins to precipitate above pH 3.0 in the form of oxyhydroxides,26,27 which are considerably less Fenton reactive and the precipitated species do not redissolve readily. Although with similar optimal pH range, it should be noted that the Fenton-HA system could widen the effective pH range with certain degree. Some
Figure 2. Effect of reaction time on the Fe (III) concentration and the Ct/C0 of BA and H2O2 at pH 3.0 in Fenton system and Fenton-HA system at [BA]0 = 40.0 μM, [H2O2]0 = 0.4 mM, [Fe (II)]0 = 10.0 μM, [NH2OH]0 = 0.4 mM.
materials could also expand the pH range by enhancing the solubility of ferric at near neutral pH, but the increased efficiencies were not better than that achieved at optimal pH in the Fenton system, which is different from Figure 1. Hence, in addition to the possibility of complexing with Fe (III) species, NH2OH might also relieve the precipitation of Fe (III) through other ways like accelerating the recovery of Fe (II) and avoiding the accumulation of Fe (III). The Role of Hydroxylamine in the Fenton-HA System. Based on the aforementioned data and analysis, we could infer a possibility that the addition of NH2OH to the Fenton system might effectively alleviate the accumulation of Fe (III), consequently increased the oxidation reactivity. In order to better understand the role of NH2OH in the Fenton-HA system, the valent state variation of iron in Fenton and Fenton-HA systems at the optimal pH 3.0 were measured on UVvis spectrometer and the results are shown in Figure 2. The presence of BA has some interference on the measurement, so we use tert-butanol (the tert-butanol dosage calibration is shown in SI Text S1) instead of BA to consume HO• for tert-butanol and NH2OH have no absorption under such wavelength range of UV light In the Fenton system, more than 70% of the total Fe was transformed into Fe (III) within 1.0 min, and almost all of the Fe (II) was transformed into Fe (III) after 2.0 min. The trend of Fe (III) concentration increase was consistent with two pseudo firstorder kinetic stages. The slow stage in the Fenton system was due to the accumulation of Fe (III).5,24 Nearly 30% of Fe (II) was transformed into Fe (III) in the Fenton-HA system in 30 s, and then the concentration of Fe (III) kept relatively steady throughout the runs. It should be noted that the concentration of Fe (III) reduced from 10 μM to less than 1 μM within 10 s when mixed with 0.4 mM NH2OH without H2O2 (data not shown). Hence, the Fe (III) concentration in Figure 1 was just a steady state concentration that resulted from a dynamic circulation and could be changed in Fenton-HA system with the variation of NH2OH and H2O2 concentration. It suggests that the ferric reduction was accelerated by NH2OH and the redox cycles of Fe (II)/Fe (III) were in dynamic equilibrium in the Fenton-HA system. Owing to the effective recovery of Fe (II) and the excess of H2O2, the reaction between H2O2 and Fe 3927
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Environmental Science & Technology (II) could occur continuously, resulting in sustained generation of HO• or ferryl, higher oxidation rates and no visible slow rate stage in the Fenton-HA system. As the Fe (II) concentrations in Fenton-HA were relatively steady, the production of reactive species and the BA oxidation were merely relevant to the concentration of H2O2, which means that the BA oxidation in the Fenton-HA system could be fitted as pseudo first-order model. The assumption agreed well with the fitting results with good correlation coefficients (R2) as shown in Table 2. Since the concentration of Fe (II) was relatively steady during the treatment and the H2O2 dosage was much higher than that of Fe (II) in the Fenton-HA system, the eq 1 or 2 could continuously occur to increase the H2O2 consumption. This deduction could be verified by the measurement of H2O2 concentration at set intervals in both systems as shown in Figure 2. The H2O2 concentrations in the Fenton system were almost constant after 30 s. The H2O2 consumption was very low in classical Fenton system for the bad recovery rate of Fe (II). That is why the enhancement of H2O2 dosage did not increase the Kobs and the t1/2 obviously in Fenton system (Table 1). On the contrary, the concentration of H2O2 in the Fenton-HA system decreased continually. The result was consistent with the aforementioned inference, which further implies that the addition of NH2OH in the Fenton-HA system accelerated the recovery of Fe (II), and thus promoted the consumption of H2O2, the generation of reactive species and the oxidation of BA. Mechanism Discussion. In addition to the classical freeradical chain reaction mechanism in Fenton system, there is still another possibility of Fenton system pathways: high-valent oxorion (ferryl) species mechanism.28 The two pathways might be compatible; we should test and verify which mechanism is the major one. The tert-butanol is a strong HO• scavenger, but a weak ferryl scavenger, which is frequently used to distinguish the two mechanisms.29,30 SI Figure S1 shows the effect of tertbutanol on the oxidation of BA. The presence of excessive radical scavenger inhibited a majority of the BA concentration decrease in both Fenton and Fenton-HA systems, which were in accordance with the HO• mechanism in these oxidation processes. In addition, the NH2OH could rapidly reduce Fe (VI) and Fe (IV) (ferryl) to Fe (II).19 Even if the ferryl occurred in the Fenton-HA system, it would be reduced by NH2OH before its reaction with BA. Hence, the ferryl species mechanism could be almost excluded in the Fenton-HA system. These results indicated that the HO• oxidation could be the dominating process responsible for the BA oxidation in the Fenton-HA system. As the addition of NH2OH significantly enhanced the circulation of Fe (III) to Fe (II) in the Fenton-HA system, the generation of HO• should be much higher in the Fenton-HA system than that in the Fenton system. If this result could be verified, the mechanism of Fe (II)/Fe (III) circulation accelerated by NH2OH and the fact that HO• is dominated in the Fenton-HA system would be further consolidated. An attempt to characterize the first intermediates formed during the oxidation of BA is a validated method to investigate the mechanism, for the m-, o-, p-hydroxybenzoic acid being the main products after HO• attacked BA. The concentration of PHBA has been frequently used as an indicator to estimate the quantities of the generated HO•.20,31 The variations of PHBA concentration in the Fenton system and the Fenton-HA system at different set intervals are illustrated in SI Figure S2. The
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Figure 3. Comparison of the intensity of DMPOOH adducts signals at pH 3.0 in Fenton system and Fenton-HA system at [H2O2]0 = 20.0 mM, [Fe (II)]0 = 50.0 μM, [NH2OH]0 = 10.0 mM.
concentration of PHBA was gradually increased in the Fenton system, and reached about 2.15 μM at 20.0 min, whereas the PHBA formed in the Fenton-HA system reached its maximum concentration of about 2.34 μM at 6.0 min, and then began to decompose. This phenomenon in the Fenton-HA system might be due to the higher production of HO•, which reacted with PHBA rapidly to form other products. It means that the presence of NH2OH in Fenton’s reagent enormously accelerated the generation of HO•. Electron paramagnetic resonance (EPR) experiment was performed to detect HO• by measuring the intensity of the DMPOOH adducts signal32 (Figure 3). The specific spectrum (quartet lines with peak height ratio of 1:2:2:1) were both obtained in the Fenton and Fenton-HA systems. The intensity of the DMPOOH signal in the Fenton-HA system was at least 6 times stronger than that in the classical Fenton system, which meant that the addition of NH2OH could enhance the generation of DMPOOH concentration. As the DMPOOH intensity is proportional to the amount of HO•,32 the HO• generation in the Fenton-HA system was much higher than that in the Fenton system. This result further verified the mechanism that the addition of NH2OH to Fenton system enhanced the generation of HO•. Although the Fenton-HA system could enhance the formation of HO•, which is primarily generated from H2O2 by eq 1, whether the utilization efficiency of H2O2 (means how much consumed H2O2 are directly transformed to HO•) being stronger in the Fenton-HA system was still in doubt. From the point of view of BA oxidation, the utilization of H2O2 in classical system was higher. While, PHBA was accumulated during the complete runs in the Fenton system (SI Figure S2), PHBA was generated and coupled with the decomposition contemporary in the Fenton-HA system. It means that a mass of HO• produced in the Fenton-HA system participated in the oxidation of intermediate products. Hence, employing the BA oxidation to count the utilization of H2O2 was not sufficient. The utilization efficiency of H2O2 in the Fenton-HA system needs further investigation in our future study. Anyway, it is indicative that the addition of NH2OH to Fenton system accelerates the decomposition of H2O2 and the generation of HO•. 3928
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Moreover, as NH2OH is a typical reducing agent, it might react with HO• under some circumstances. The pKa of NH2OH is 13.7433 and 5.96,34 respectively. When the solution pH is below 5.96, the rate of HO• radical reacting with NH2OH (the primary form of NH2OH in acid solution is NH3OHþ) is dramatically reduced.17 This fact interprets the results related to the effect of different pH on BA oxidation in the Fenton-HA system shown in Figure 1. The Fe (III) was effectively reduced to Fe (II) by NH3OHþ, thereby the Fe (III) accumulation was alleviated, and the amounts of Fe (III) precipitation was reduced and the oxidation reactivity was obviously sustained in the Fenton-HA system. As the pH was higher than 4.0, the residual BA concentration gradually increased since the NH2OH began to react with HO•. When the initial pH was higher than 6.0, the role of NH2OH in the Fenton-HA system was the termination agent more than the accelerating agent. FeðIIIÞ þ NH2 OH f NH2 O• þ FeðIIÞ þ Hþ
ð9Þ
2NH2 O• f N2 þ H2 O
ð10Þ
FeðIIIÞ þ NH2 O• f NHO þ FeðIIÞ þ Hþ
ð11Þ
2NHO f N2 O þ H2 O
ð12Þ
5FeðIIIÞ þ 2H2 O þ NH2 O• f NO3 þ 5FeðIIÞ þ 6Hþ ð13Þ NO3 þ NH2 OH þ OH f NO þ NO2 þ H2 O ð14Þ The fate of NH2OH in the Fenton-HA system was significant since it is not only a typical reducing agent, but also a kind of toxic compound and it contains nitrogen element, whose regulations of transfer and transformation are arresting in environmental research. SI Figure S3 shows the variation of NH2OH concentration, which could continuously decrease in the Fenton-HA system. Generally, the end products of NH2OH when reacted with Fe (III) in solution were N2, N2O, NO2, and NO3, as shown in eqs 914.18,19,17,35 In order to ascertain the end product of NH2OH in the Fenton-HA system, the dosages of BA, H2O2, Fe (II), and NH2OH were enlarged 10-fold. When the reaction time was extended to 4 h, NH2OH was undetectable. The concentration of TN in solution was constant throughout the runs, which excluded the possibility of producing N2. The concentration of organic nitrogen and NO2 were all negligible. Besides, the concentration of NO3 at the reaction time of 4 h was 2.41 mM. Hence, the most likely other residual end product of NH2OH was supposed to be N2O, a kind of greenhouse gases and its solubility in water is quite high.36 In previous studies of Butler,37 the maximum N2O was attained at pH 3.0, which was also the pH selected in most of our experiments. The analysis of N2O was only semiquantitative by the gas chromatography (the method is shown in SI Text S3). It should be noted that NH2OH is a kind of toxic compound38 and the addition of NH2OH in Fenton system is far from practical application. Our work has just put forward an interesting phenomenon and gives a preliminary interpretation that NH2OH could be adopted to alleviate Fe (III) accumulation, thus accelerate the reaction rates and expand effective pH range. The ultimate objective of our study is to find an environmentally
friendly chemical with feeble reaction rate with HO• at a wider pH range to alleviate Fe (III) accumulation and precipitation, thus to reduce the cost and accelerate the rates of HO• generation in Fenton system. In addition to the attempt to find other chemicals, the investigation to relieve or avoid the toxicity of NH2OH and its end products (i.e., N2O, NO3) in the FentonHA system in a further study also makes sense.
’ ASSOCIATED CONTENT
bS
Supporting Information. Additional details. This material is available free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected].
’ ACKNOWLEDGMENT We thank Dr J M. Shen for his technical help with EPR analysis. This study was supported by the Natural Science Foundation of China (50821002) and the Science and Technology Ministry of China (2009ZX07424-005; 2009ZX07424-006; 2008ZX07421-002). The support by 863 High Tech Scheme (2009AA06Z310) and the State Key Laboratory of Urban Water Resource and Environment (2010DX10) are greatly appreciated. ’ REFERENCES (1) Fenton, H. J. H. Oxidation of tartaric acid in presence of iron. J. Chem. Soc. Trans. 1894, 65, 899–910. (2) Haber, F.; Weiss, J. The catalytic decomposition of hydrogen peroxide by iron salts. Proc. Roy. Soc. A. 1934, 134, 332–351. (3) 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, 47, 462–500. (4) Walling, C. Fenton’s reagent revisited. Acc. Chem. Res. 1975, 8, 125–131. (5) Pignatello, J. J.; Oliveros, E.; MacKay, A. Advanced oxidation processes for organic contaminant destruction based on the Fenton reaction and related chemistry. Crit. Rev. Environ. Sci. Technol. 2006, 36 (1), 1–84. (6) Nieto-Juarez, J. I.; Pierzchla, K.; Sienkiewicz, A.; Kohn, T. Inactivation of MS2 coliphage in Fenton and Fenton-like systems: role of transition metals, hydrogen peroxide and sunlight. Environ. Sci. Technol. 2010, 44 (9), 3351–3356. (7) Pham, A. L.-T.; Lee, C.; Doyle, F. M.; Sedlak, D. L. A silicasupported iron oxide catalyst capable of activating hydrogen peroxide at neutral pH values. Environ. Sci. Technol. 2009, 43 (23), 8930–8935. (8) Sun, Y.; Pignatello, J. J. Chemical treatment of pesticide wastes. Evaluation of iron (III) chelates for catalytic hydrogen peroxide oxidation of 2, 4-D at circumneutral pH. J. Agric. Food. Chem. 1992, 40 (2), 322–327. (9) Zepp, R. G.; Faust, B. C.; Hoigne, J. Hydroxyl radical formation in aqueous reactions (pH 38) of iron(II) with hydrogen peroxide: the photo-Fenton reaction. Environ. Sci. Technol. 1992, 26 (2), 313–319. (10) Oturan, M. A.; Oturan, N.; Lahitte, C.; Trevin, S. Production of hydroxyl radicals by electrochemically assisted Fenton’s reagent Application to the mineralization of an organic micropollutant, pentachlorophenol. J. Electroanal. Chem. 2001, 507 (12), 96–102. (11) Duesterberg, C. K.; Waite, T. D. Kinetic modeling of the oxidation of p-hydroxybenzoic acid by Fenton’s reagent: implications of the role of quinones in the redox cycling of iron. Environ. Sci. Technol. 2007, 41 (11), 4103–4110. 3929
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