CORRESPONDENCE/REBUTTAL pubs.acs.org/est
Comment on “Oxidation of Sulfoxides and Arsenic(III) in Corrosion of Nanoscale Zero Valent Iron by Oxygen: Evidence against Ferryl Ions (Fe(IV)) as Active Intermediates in Fenton Reaction” recently published paper1 questioned our earlier report2 that hydroxyl radical (OH•) is not the only oxidant produced when Fe[II] reacts with hydrogen peroxide (H2O2) at circumneutral pH values. While we are pleased by the interest in our research, we have serious reservations about the paper for the following reasons: 1. The Authors Failed to Consider Different Forms of Fe[IV]. On the basis of experiments in which either ozone or Fe[VI] were used to generate oxidants, the authors concluded that Fe[IV] always reacts with DMSO exclusively through a two-electron reaction. This conclusion is questionable because the authors did not consider species such as Fe(OH)22þ and Fe(OH)3 þ which likely account for a significant fraction of the Fe[IV] formed under circumneutral pH conditions. 3,4 A two-electron transfer was observed in experiments in which ozone was used to produce FedO2þ at pH values from 2 to 5.1 Data at higher pH values, which showed products of one-electron reactions, were relegated to the Supporting Information section. The failure to consider the one-electron reactions of DMSO in the circumneutral pH range, where the clearest evidence for a species other than OH • has been obtained, undermines the paper’s claims. Fe[IV] species react with aliphatic alcohols, aldehydes, and ethers by both one- and two-electron transfer reactions.4,5 The lack of PMSO2 and TMSO 2 production at pH values between 2 and 5 simply indicates that a twoelectron transfer process did not occur. Contrary to the authors’ assertion, it does not prove that OH• was the only oxidant present in the system. Results from experiments in which Fe[IV] was generated from Fe[VI] at pH 8-9 were used as further evidence against the presence of an oxidant other than OH •. Fe[VI] decomposition is complex and the role of Fe[IV] in the process is still being resolved. Therefore, it is premature to conclude that the same Fe[IV] species are produced when Fe[VI] is reduced or Fe[II] is oxidized. Nonetheless, the data presented by Pang et al. 1 were inconsistent with their presumption that FedO2þ acted as the oxidant because Fe[II] should not have been able to compete with DMSO for FedO2þ based on known rate constants. 4,5 2. The Reactivity of DMSO Was Inconsistent with the Presence of Only OH•. Pang et al. 1 observed no sulfone production when sulfoxides were exposed to oxidants produced by iron nanoparticles at circumneutral pH values. They used this finding as evidence against the formation of oxidants other that OH •. We demonstrated that DMSO can only outcompete methanol for the oxidant produced by iron nanoparticles at pH 7 when [DMSO]/[methanol] ratios are high.6 The low reactivity of the oxidant with DMSO is inconsistent with known reaction rates of these compounds with OH• 7 or FedO 2þ. 4,8 Recently collected data from competition experiments with DMSO and methanol provide
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Figure 1. Production of HCHO from oxidation of DMSO and methanol (MeOH) by Fenton’s reagent ([Fe(II)]0 = 100 μM; [H2O2]0 = 1 mM; [DMSO]0 = [MeOH]0 = 200 mM; [PIPES]0 = 1 mM at pH 7; reaction time = 60 min).
further evidence that an oxidant other than OH• is produced by the Fenton reaction under circumneutral pH conditions (Figure 1). The oxidant produced at pH 3 (i.e., OH•) was scavenged efficiently by DMSO, whereas methanol outcompeted DMSO at pH 7, contrary to expectations if FedO2þ or OH• were the main oxidants at circumneutral pH values. 3. The Explanation for the pH-Dependence of Product Yields Was Inconsistent with Previous Findings. Pang et al. 1 suggested that the observed pH-dependence of product formation was attributable to changes in the reaction between the compounds and OH •. However, this explanation fails to consider an extensive body of previously published research. For example, changes in product yields were not observed over a wide pH range when benzene or phenol were oxidized by OH• formed by H 2 O2 photolysis. 9,10 The authors also suggested that the pH-dependence observed in the Fenton system was due to a reaction between Fe[II] and a carbon-centered radical. This explanation is unlikely because the relative rates of reactions between carboncentered radicals and O 2 are much faster than with Fe2þ (e.g., ∼10 9 M -1 s-1 and ∼10 5 M -1s -1 with the radical produced by phenol oxidation, respectively11 ) under the conditions studied. In conclusion, the exact structure of the oxidant produced when Fe[II] reacts with H2O2 at circumneutral pH values is uncertain. However, a preponderance of evidence suggests that
Published: March 04, 2011 3177
dx.doi.org/10.1021/es104399p | Environ. Sci. Technol. 2011, 45, 3177–3178
Environmental Science & Technology
CORRESPONDENCE/REBUTTAL
an oxidant other than OH• is formed under conditions typically encountered in engineered treatment systems. Christina K. Remucal,† Changha Lee,‡ and David L. Sedlak§,* †
Institute of Biogeochemistry and Pollutant Dynamics, Swiss Federal Institute of Technology (ETH), Z€urich, Switzerland
‡
School of Urban and Environmental Engineering, Ulsan National Institute of Science and Technology (UNIST), Ulsan, South Korea
§
Department of Civil and Environmental Engineering, University of California at Berkeley, Berkeley, California 94720, United States
’ AUTHOR INFORMATION Corresponding Author
*Phone: (510) 643-0256; fax: (510) 643-5264; e-mail: sedlak@ berkeley.edu
’ REFERENCES (1) Pang, S.; Jiang, J.; Ma, J., Oxidation of sulfoxides and arsenic(III) in corrosion of nanoscal zero valent iron by oxygen: Evidence against ferryl ions (Fe(IV)) as active intermediates in fenton reaction. Envrion. Sci. Technol. 2010(ASAP). (2) Keenan, C. R.; Sedlak, D. L. Factors affecting the yield of oxidants from the reaction of nanoparticulate zero-valent iron and oxygen. Environ. Sci. Technol. 2008, 42, 1262–1267. (3) Jacobsen, F.; Holcman, J.; Sehested, K. Reactions of the ferryl ion with some compounds found in cloud water. Int. J. Chem. Kinet. 1998, 30, 215–221. (4) Pestovsky, O.; Bakac, A. Reactivity of aqueous Fe(IV) in hydride and hydrogen atom transfer reactions. J. Am. Chem. Soc. 2004, 126, 13757–13764. (5) Pestovsky, O.; Stoian, S.; Bominaar, E. L.; Shan, X.; Munck, E.; Que, L.; Bakac, A. Aqueous FeIV=O: Spectroscopic identification and oxo-group exchange. Angew. Chem., Int. Ed. 2005, 44, 6871–6874. (6) Keenan, C. R.; Goth-Goldstein, R.; Lucas, D.; Sedlak, D. Oxidative stress induced by zero-valent iron nanoparticles and Fe(II) in human bronchial epithelial cells. Environ. Sci. Technol. 2009, 43, 4555–4560. (7) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17, 513–886. (8) Pestovsky, O.; Bakac, A. Aqueous ferryl(IV) ion: Kinetics of oxygen atom transfer to substrates and oxo exchange with solvent water. Inorg. Chem. 2006, 45, 814–820. (9) Weir, B. A.; Sundstrom, D. W.; Klei, H. E. Destruction of benzene by ultraviolet light-catalyzed oxidation with hydrogen peroxide. Hazard. Waste Hazard. Mater. 1987, 4 (2), 165–176. (10) Castrantas, H. M.; Gibilisco, R. D. UV destruction of phenolic compounds under alkaline conditions. ACS Symp. Ser. 1990, 422, 77–99. (11) Chen, R.; Pignatello, J. Role of quinone intermediates as electron shuttles in Fenton and photoassisted Fenton oxidations of aromatic compounds. Environ. Sci. Technol. 1997, 31, 2399–2406.
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