Activation of Oxygen and Hydrogen Peroxide by Copper(II) Coupled

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Activation of Oxygen and Hydrogen Peroxide by Copper(II) Coupled with Hydroxylamine for Oxidation of Organic Contaminants Hongshin Lee,†,§ Hye-Jin Lee,§ Jiwon Seo,§ Hyung-Eun Kim,§ Yun Kyung Shin,‡ Jae-Hong Kim,† and Changha Lee*,§ †

Department of Chemical and Environmental Engineering, Yale University, New Haven, Connecticut 06511, United States Southeast Sea Fisheries Research Center, National Fisheries Research and Development Institute (NFRDI), 397-68 Sanyangilju-ro, Tongyeong-si, Gyeongsangnam-do 53085, Republic of Korea § School of Urban and Environmental Engineering, and KIST-UNIST-Ulsan Center for Convergent Materials (KUUC), Ulsan National Institute of Science and Technology (UNIST), 50 UNIST-gil, Ulsan 44919, Republic of Korea Environ. Sci. Technol. 2016.50:8231-8238. Downloaded from pubs.acs.org by DURHAM UNIV on 08/27/18. For personal use only.

‡

S Supporting Information *

ABSTRACT: This study reports that the combination of Cu(II) with hydroxylamine (HA) (referred to herein as Cu(II)/HA system) in situ generates H2O2 by reducing dissolved oxygen, subsequently producing reactive oxidants through the reaction of Cu(I) with H2O2. The external supply of H2O2 to the Cu(II)/ HA system (i.e., the Cu(II)/H2O2/HA system) was found to further enhance the production of reactive oxidants. Both the Cu(II)/HA and Cu(II)/H2O2/HA systems effectively oxidized benzoate (BA) at pH between 4 and 8, yielding a hydroxylated product, p-hydroxybenzoate (pHBA). The addition of a radical scavenger, tert-butyl alcohol, inhibited the BA oxidation in both systems. However, electron paramagnetic resonance (EPR) spectroscopy analysis indicated that •OH was not produced under either acidic or neutral pH conditions, suggesting that the alternative oxidant, cupryl ion (Cu(III)), is likely a dominant oxidant.

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INTRODUCTION Advanced oxidation technologies based on the Fenton and Fenton-like reactions have been intensively studied for the degradation of refractory organic contaminants in water and wastewater.1 The decomposition of hydrogen peroxide (H2O2) by the catalytic redox cycle of Fe(III)/Fe(II) is known to produce reactive oxidants such as hydroxyl radical (•OH) and ferryl ion (Fe(IV)) that are capable of oxidizing organic compounds. Although there has been a long-term controversy on the identity of the reactive oxidant from the Fenton reaction (i.e., •OH vs Fe(IV)),2,3 recent studies suggested that both • OH and Fe(IV) are produced with the dominant oxidant shifting from •OH to Fe(IV) as pH increases from acidic to neutral values.4−7 •OH is a nonselective oxidant that rapidly reacts with a broad spectrum of organic and inorganic compounds,8 whereas Fe(IV) oxidizes a relatively limited range of compounds.9−11 Besides low solubility of iron, the shift of the main oxidant to Fe(IV) is another factor that limits the applicability of the Fenton (-like) reactions at neutral pH. Similar to iron, copper can also convert H2O2 into reactive oxidants via the catalytic redox cycle of Cu(II)/Cu(I).12−14 Previous reports suggest that the nature of reactive oxidants from the copper-catalyzed Fenton-like system may also be pHdependent; •OH and cupryl ion (Cu(III)) are dominantly © 2016 American Chemical Society

produced under acidic and neutral/alkaline conditions, respectively.7,15 However, the production of •OH under neutral/alkaline conditions cannot be completely excluded based on the observations that •OH scavengers such as tertbutyl alcohol were found to inhibit the oxidation of the target compounds and compounds such as benzoate from forming hydroxylated products.7,15 In both the Fe(III)/H2O2 and the Cu(II)/H2O2 systems, the reduction of oxidized metal ion (i.e., Fe(III) and Cu(II)) by H2O2 is the rate-limiting step for the production of reactive oxidants.16,17 Approaches frequently employed to accelerate the reductive conversion of Fe(III) into Fe(II) in the Fe(III)/H2O2 system include UV light irradiation (photo-Fenton)18,19 and electricity application (electro-Fenton).20,21 Recently, the addition of hydroxylamine (HA), a reducing agent, was demonstrated as a suitable method for facile Fe(III) reduction.22 The use of HA in the Fe(III)/H2O2 system accelerated the oxidation of benzoic acid by more than an order of magnitude, expanding the effective pH range up to 5.7. Received: Revised: Accepted: Published: 8231

April 26, 2016 July 2, 2016 July 7, 2016 July 7, 2016 DOI: 10.1021/acs.est.6b02067 Environ. Sci. Technol. 2016, 50, 8231−8238

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Environmental Science & Technology

Figure 1. (a) Oxidative degradation of BA and (b, c) concentrations of H2O2 and Cu(I) in the Cu(II)/HA system under different aeration conditions ([BA]0 = 0.1 mM, [Cu(II)]0 = 0.1 mM, [HA]0 = 5 mM, pH 7).

understanding of the chemistry of copper-based Fenton-like reactions, particularly providing insight into the pH-dependent nature of reactive oxidants generated by the reactions.

However, similar to most iron-based Fenton (-like) systems, the Fe(III)/H2O2/HA system still exhibits the optimal activity around pH 3−4, and the system activity dramatically decreases as pH increases above 6. Considering the similarities between the Fe(III)/H2O2 and Cu(II)/H2O2 systems, one can postulate that HA also can be instrumental in enhancing the efficiency of Cu(II)/H2O2 system by accelerating the reduction of Cu(II). The fact that Cu(II)/H2O2 system efficiently functions at neutral pH is particularly appealing; note that (i) Cu(II) has higher solubility at neutral pH than Fe(III); the solubility values of Cu(II) and Fe(III) at pH 7 are ca. 10−5 and 10−10 M, respectively,7,23 and (ii) the Cu(II)/H2O2 system exhibits substantial activity toward oxidizing organic contaminants such as phenol, benzoate, diclofenac, and carbamazepine at neutral pH (interpreted as the contribution of •OH in those studies).7,24 In addition, one can further postulate that the combination of Cu(II) with HA can produce reactive oxidants without the external supply of H2O225 because Cu(I) is known to generate H2O2 by reducing dissolved oxygen (O2). Despite the anticipated benefits of using HA in copper-catalyzed Fenton-like reactions, no previous studies have evaluated Cu(II)/H2O2/HA and Cu(II)/HA systems for the degradation of organic contaminants. The objectives of this study are 2-fold. First, we assess the potential of the copper-based Fenton-like systems with HA (i.e., the Cu(II)/H2O2/HA and the Cu(II)/HA systems) for the oxidation of select organic compounds; benzoate was selected as a main target compound because it has been used as a •OH probe compound and its oxidation mechanism by •OH is well-known.9,26,27 These systems are compared to the conventional Cu(II)/H2O2 system. Second, we evaluate the nature of the oxidants produced by the Cu(II)-catalyzed Fenton-like reaction under different pH conditions. The pHdependent behaviors of copper-catalyzed Fenton-like systems were explained by the speciation of Cu(II) and Cu(III) complexes and their reactions. We expect that Cu(II)/H2O2/ HA system shall be particularly useful for studying the nature of oxidants because the production of reactive oxidants is expected to be enhanced compared to the Cu(II)/H2O2 system. In the Cu(II)/H2O2 system concentrations of the short-lived reactive species are too low to be accurately assessed. This study suggests new oxidation technologies using copper and HA that are potentially applicable to degradation of refractory organic compounds at neutral pH. In addition, this study improves

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MATERIALS AND METHODS Reagents. All chemicals were of reagent grade and used without further purification. High purity (99.99%) gases such as nitrous oxide (N2O) and oxygen (O2) were used for some experiments. All solutions were prepared using 18.2 MΩ·cm Milli-Q water from a Millipore system. The stock solution of Cu(II) (10 mM) was prepared using cupric sulfate, and stored at 4 °C until use. The stock solution of HA (500 mM) was prepared daily. Experimental Setup and Procedure. Experiments to evaluate the kinetics of organic compound oxidation in copperbased Fenton-like systems were conducted in a 100 mL Pyrex flask at room temperature (20 ± 2 °C). No pH buffers were used for experiments at pH 3−5 because the pH variations before and after the reaction were minor. Phosphate and borate buffers (1 mM) were used for neutral (pH 6.5−8.0) and alkaline (9.0−10) pH ranges, respectively. The phosphate buffer can affect the Cu(II) speciation at pH 5−7 (Figure S1 in the Supporting Information (SI)). The initial pH was adjusted using 1 N HClO4 and 1 N NaOH solution. The reaction was initiated by adding an aliquot of stock solutions of H2O2 or HA, to a pH-adjusted solution containing organic compounds and Cu(II). Samples were withdrawn at predetermined time intervals and filtered using a 10 mL glass syringe and a 0.45 μm nylon syringe filter. Ethylenediaminetetraacetic acid (EDTA) (4 mM) was immediately added to quench the reaction. The experiments were conducted in duplicate, and the average values with the standard deviations (error bars) are presented. Analytical Methods. Benzoic acid (pKa = 4.2) or benzoate (BA) was analyzed by high performance liquid chromatography (HPLC) (UltiMate 3000, Dionex Co.) with UV absorbance detection at 255 nm. Separation was performed on a Dionex Acclaim C-18 column (250 mm × 4.6 mm, 5 μm) using nitric acid solution (10 mM) and neat acetonitrile as eluents at a flow rate of 1.0 mL min−1. p-Hydroxybenzoate (pHBA) and other oxidation products were analyzed on a LC/orbitrap MS/MS system. The analyses were performed using a rapid separation liquid chromatography (RSLC) (UltiMate 3000, Dionex Co.) coupled with a quadrupole-Orbitrap mass spectrometer (Q8232

DOI: 10.1021/acs.est.6b02067 Environ. Sci. Technol. 2016, 50, 8231−8238

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Environmental Science & Technology Exactive, Thermo Fisher Scientific Inc.). Detailed procedures are described in the SI (Text S1). Ammonium (NH4+), nitrite (NO2−), and nitrate (NO3−) ions were analyzed by ion chromatography (IC) (ICS-3000, Dionex Co.) with conductivity detection. The separation of NH4+ was performed on an IonPac CS-17 cationic column (4 mm × 250 mm) using methanesulfonic acid (6 mM) as the eluent at a flow rate of 1.0 mL min−1. For the analysis of NO2− and NO3−, an IonPac AS-9 anionic column (4 mm × 250 mm) was employed using carbonate solution (9.0 mM) as the eluent. The concentrations of total organic carbon (TOC) and total nitrogen (TN) were determined by a TOC/TN analyzer (TOC-5000A, Shimadzu Co.). N2O was analyzed using gas chromatography (GC 7820A, Agilent Co.) with the electron capture detector (ECD); a Porapak Q (80/100 mesh) column was used with high purity N2 as a carrier gas at a flow rate of 35 mL min−1. The concentrations of Cu(I) and H2O2 was spectrophotometrically determined by the neocuproine method28 and the titanium sulfate method,29 respectively. Formaldehyde (HCHO) was analyzed by HPLC after DNPH derivatization.30 EPR spectroscopy was used to detect •OH, using 5,5-dimethyl-1pyrroline N-oxide (DMPO) as a spin-trapping agent.31 EPR signals of the DMPO−OH spin adduct were obtained on a CW/Pulse EPR system (ELEXYS E580, Bruker Co.) with a 9.64 GHz microwave (0.94 mW) at a modulation frequency of 100 kHz and a modulation amplitude of 2.0 G.

Figure 2. Variations in TN, NH4+, NO2−, and NO3− concentrations during the decomposition of HA in the Cu(II)/HA system ([BA]0 = 0.1 mM, [Cu(II)]0 = 0.1 mM, [HA]0 = 5 mM, pH 7).

Cu(II)/H2O2, Cu(II)/HA, Cu(II)/H2O2/HA Systems. The oxidative degradation of BA was compared in three systems: Cu(II)/H2O2, Cu(II)/HA, and Cu(II)/H2O2/HA. The rate of degradation of BA was found to be Cu(II)/H2O2 < Cu(II)/HA < Cu(II)/H2O2/HA at pH 7 (Figure 3a); the combination of H2O2 with HA (i.e., the H2O2/HA system) did not degrade BA. The Cu(II)/H2O2/HA system exhibited the synergistic enhancement of BA degradation; the degradation rate of BA by the Cu(II)/H2O2/HA system was greater than the simple sum of those by the Cu(II)/H2O2 and Cu(II)/HA systems (also refer to Figure 3b). The BA degradation experiments were performed at different pH values from 3 to 10 in each system. The resulting pseudo-first order rate constants for the BA degradation were depicted as a function of pH (Figure 3b). Overall, the circumneutral pH conditions favored the BA degradation; the BA degradation rate was lower under acidic pH conditions than alkaline pH conditions. At almost all pH values, the BA degradation rate was Cu(II)/H2O2 < Cu(II)/HA < Cu(II)/H2O2/HA. Meanwhile, a very slow degradation of BA was observed by the H2O2/HA system under acidic pH conditions; consistent with the recent report that •OH is produced by the acid-catalyzed reaction of H2O2 with HA.32 We also observed that the decomposition of H2O2 in the Cu(II)/H2O2/HA system accelerated with increasing pH (Figure S7 in the SI). The pHBA formation was monitored during the BA oxidation by the Cu(II)/H2O2, Cu(II)/HA, Cu(II)/H2O2/ HA, and H2O2/HA systems at different pH values (Figure 3c, and Figures S8−S10 in the SI). pHBA formed after 30 min of reaction time (Figure 3c), consistent with the rate of BA degradation (Figure 3b) except for the Cu(II)/H2O2/HA system. The Cu(II)/H2O2/HA system exhibited much lower pHBA concentrations at pH 5−7 than expected, which may be attributed to the secondary oxidation of pHBA to dihydroxybenzoates during the oxidation of BA (SI Figures S2−S4). pHBA was rapidly formed in the initial 10 min, and then degraded as the reaction proceeded in the Cu(II)/H2O2/HA system (SI Figures S8c and S10), whereas the pHBA concentrations in the other systems continuously increased over the entire reaction time (SI Figures S8a, b, and d). The production of HCHO by the oxidation of methanol was also examined in the Cu(II)/H2O2, Cu(II)/HA, Cu(II)/H2O2/ HA, and H2O2/HA systems at different pH values (Figure S11 in the SI). Overall, the HCHO production exhibited similar

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RESULTS Degradation of BA by the Cu(II)/HA System at Neutral pH. The oxidative degradation of BA by the Cu(II)/HA system was examined at pH 7 under different aeration conditions (Figure 1a). Neither Cu(II) nor HA alone changed the concentration of BA. The combination of Cu(II) with HA under N2 condition (Cu(II)/HA/N2) did not degrade BA either; the Cu(II)/HA/N2 system did not produce H2O2 (data not shown), indicating that dissolved oxygen is the precursor of H2O2. The combination of Cu(II) with HA in the presence of oxygen degraded BA by more than 70% in 4 h. The Cu(II)/HA system with no aeration (open to the atmosphere) exhibited similar degree of BA degradation to that with O2 aeration (Cu(II)/HA/O2). The concentrations of H2O2 and Cu(I) were monitored in the Cu(II)/HA and Cu(II)/HA/O2 systems (Figure 1b and c). The Cu(II)/HA/O2 system produced higher concentration of H2O2 than the Cu(II)/HA system (Figure 1b), but produced lower concentration of Cu(I) (Figure 1c). The products produced during the BA degradation by the Cu(II)/HA system were analyzed; compounds including hydroxybenzoates, dihydroxybenzoates, nitrobenzene, nitrobenzoates, and nitro-hydroxybenzoates were identified, and the pathways of BA oxidation were presented (refer to Figures S2−S4 in the SI for details). Decomposition of HA. To examine the decomposition of HA, variation in concentrations of TN and nitrogenous products22 (NH4+, NO2−, NO3−) was monitored in the Cu(II)/HA system (Figure 2). The TN concentration decreased by 97% in 4 h, exhibiting the pseudo-first order decay (k = 0.0216 min−1). However, the production of NH4+, NO2−, and NO3− was very minor throughout the entire reaction. A small amount of N2O was also detected in the headspace of the reactor (Figure S5 in the SI). In the Cu(II)/ H2O2/HA system, the decrease in TN concentration was faster than that in the Cu(II)/HA system (Figure S6 in the SI). 8233

DOI: 10.1021/acs.est.6b02067 Environ. Sci. Technol. 2016, 50, 8231−8238

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Figure 3. (a) Oxidative degradation of BA, and (b) pseudo first-order rate constants for the degradation of BA and (c) production of pHBA by different systems as a function of pH ([BA]0 = 0.1 mM, [Cu(II)]0 = 0.1 mM, [HA]0 = 5 mM, [H2O2]0 = 10 mM, pH 7 for (a), reaction time = 30 min for (c)).

trends to the BA oxidation (Figure 3b and c), except for the H2O2/HA system which exhibited relatively higher yields of HCHO. Effect of tert-Butyl Alcohol on BA Degradation by the Cu(II)/H2O2/HA System. The addition of tert-butyl alcohol inhibited the BA degradation by the Cu(II)/H2O2/HA system (Figure 4). Without tert-butyl alcohol, BA was almost

Figure 5. EPR spectra obtained by spin trapping with DMPO in different systems at acidic pH ([DMPO]0 = 10 mM, [Cu(II)]0 = 0.1 mM, [HA]0 = 5 mM, [H2O2]0 = 10 mM, pH 3, Reaction time = 10 min).

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DISCUSSION Production of Reactive Oxidants by Cu(II) in Combination with HA. The Cu(II)/HA system produces reactive oxidants via the reaction of in situ generated Cu(I) and H2O2. Primarily, HA reduces Cu(II) to Cu(I). The two-electron oxidation of HA into N2O has been postulated based on the observed stoichiometry between Cu(II) consumed and the total gas produced (reaction 1).33 However, the little production of N2O in this study (SI Figure S5) suggests that the one-electron oxidation of HA into N2 may be more favored (reaction 2). The kinetics for this reaction are unknown, but the second-order rate constant is estimated to be at least higher than 103 M−1 s−1; 10 μM Cu(II) was completely reduced to Cu(I) in 5 s by the addition of 0.2 mM HA under anoxic conditions (data not shown).

Figure 4. Effect of tert-butyl alcohol on BA degradation by the Cu(II)/ H2O2/HA system. The inset represents pseudo first-order rate constants for degradation of BA ([BA]0 = 0.1 mM, [Cu(II)]0 = 0.1 mM, [HA]0 = 5 mM, pH 7).

completely degraded in 2 h. However, in the presence of 1, 10, and 100 mM tert-butyl alcohol, the BA degradation efficiencies in 2 h were 92.7, 57.8, and 10.6%, respectively. The BA degradation rate (pseudo first-order rate constant) decreased by 1.6, 2.3, and 8-fold in the presence of 1, 10, and 100 mM tert-butyl alcohol, respectively (the inset in Figure 4). EPR Spectroscopy. The EPR technique with DMPO as a spin-trapping agent was used to identify the production of •OH in the systems of Cu(II)/H2O2/HA, H2O2/HA, Cu(II)/HA, and Cu(II)/H2O2 at pH 3 and 7 (Figure 5). At pH 3, the Cu(II)/H2O2/HA and H2O2/HA systems exhibited the signal of DMPO−OH spin adduct, 1:2:2:1 quartet lines with hyperfine constants of aN = aH = 14.9 G31 (Figure 5a). A notable observation is that the signal intensity of the H2O2/HA system was much higher than that of the Cu(II)/H2O2/HA system. The Cu(II)/HA, and Cu(II)/H2O2 systems did not generate noticeable signals. At pH 7, no signals were obtained any of the systems tested.

NH 2OH + 2Cu(II) → 1/2N2O + 1/2H 2O + 2Cu(I) + 2H+ (1)

NH 2OH + Cu(II) → 1/2N2 + H 2O + Cu(I) + H+ (2)

Subsequently, the Cu(I) produced reduces O2 into H2O2 by two single-electron transfer reactions (reactions 334 and 435).13,34 8234

DOI: 10.1021/acs.est.6b02067 Environ. Sci. Technol. 2016, 50, 8231−8238

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Environmental Science & Technology Cu(I) + O2 → Cu(II) + O2•−(k 3 = 3.1 × 104 M−1s−1at pH 6−8)

the formation of Cu(II)-peroxo complex and the subsequent ligand-to-metal charge transfer (reactions 9 and 10).24,37

(3)

Cu(II) + H 2O2 ↔ Cu(II) − OOH + H+

Cu(I) + O2•− + 2H+ → Cu(II) + H 2O2 (k4 = 2.0 × 109M−1s−1)

(4)

Cu(II) − OOH → Cu(I) + HO2•( ↔ O2•− + H+)

Another possibility is HA directly reduces O2 into H2O2 (Reactions 536). However, this reaction appears to be minor due to the low reaction rate.

(10)

Although the stability constant for Cu(II)-peroxo complex is unknown, the equilibrium shift toward the formation of Cu(II)peroxo complex with increasing pH (reaction 8) should be responsible for the pH-dependence of reaction 8. The Cu(II)/H2O2/HA and Cu(II)/HA systems exhibited the optimal BA degradation at pH 6 (Figure 3b). The increasing rates of BA degradation at pH < 6 indicate that the RDS for the oxidant production in these two systems accelerates with increasing pH. Reaction 3 appears to be the RDS for the Cu(II)/HA system, which is believed to accelerate with pH.17,34,39 The mechanism for the pH-dependence of reaction 3 is uncertain, but may be associated with the speciation of Cu(I) (e.g., the contribution of Cu(I)-hydroxo species such as Cu(OH) and Cu(OH)2− that react faster with oxygen).34 The increasing rate of BA degradation by the Cu(II)/H2O2/HA system at pH < 6 can be explained by the pH-dependence of reaction 6, which becomes the RDS in this system because H2O2 and Cu(I) (via the instant reduction of Cu(II) by HA) are initially provided. Similar to reaction 3, reaction 6 can also accelerate with pH by the contribution of Cu(I)-hydroxo complexes. Little information about the stability constants for Cu(I)-hydroxo complexes is available, and their concentrations can be very low compared to that of Cu+. However, the contributions of these minor species cannot be neglected when the difference in rate constants is great; note that the Fenton reaction of Fe(OH)+ (i.e., its reaction with H2O2) is faster by almost 5 orders of magnitude than that of Fe2+ (kFe(OH)+ = 5.9 × 106 M−1 s−1 and kFe2+ = 63 M−1 s−1).40 All three systems exhibited decreasing rates of BA degradation at alkaline pH (Figure 3b). Decreasing rates of BA degradation at pH > 6 can be attributed to the precipitation of Cu(II); the formation of tenorite (CuO) is favored at pH > 6 (refer to Figure S1 in the SI). However, the decomposition of H2O2 in the Cu(II)/H2O2/HA system did not decelerate at alkaline pH; in fact, decomposition of H2O2 accelerated (SI Figure S7). The Cu(II)/H2O2 system also showed increasing rates of the H2O2 decomposition with increasing pH.24 In addition, according to a previous study,7 some target compounds such as Reactive Black 5 and As(III) showed even greater degradation at alkaline pH by the Cu(II)/H2O2 system using 0.1 mM Cu(II). These observations collectively imply that the speciation of Cu(II) does not necessarily limit the production of reactive oxidants by the Fenton-like reactions; the effect of Cu(II) speciation on the formation of Cu(II)-peroxo complex (reaction 9) may be minor. The decreasing rates of BA degradation at alkaline pH is believed to be associated with the reactivity change of the responsible oxidant (Cu(III) species) due to the pH-dependent speciation. Nature of Reactive Oxidants Produced by CopperCatalyzed Fenton-like Reactions. There is a debate about the identity of reactive oxidants produced by the Fenton reaction (i.e., the reaction of Fe(II) with H2O2 forming •OH vs Fe(IV)).2,3 Investigators have provided different views on the production of •OH vs Cu(III) by the copper-catalyzed Fentonlike reaction.7,12−15,24 Recent studies have claimed that Cu(III) rather than •OH is the dominant oxidant under neutral and

2NH 2OH + O2 → N2 + H 2O2 + 2H 2O (k5 = 9.4 × 10−2M−1s−1at pH 7)

(5)

Finally, the reaction of Cu(I) with H2O2 (the Fenton-like reaction; reaction 637) produces reactive oxidants such as •OH and Cu(III), capable of oxidizing BA (reaction 7). In this series of reactions, most of HA is liberated as N2O and N2 gases without leaving residual nitrogenous products in the solution (Figure 2). In fact, HA can also serve as the scavenger of reactive oxidants while it is decomposed during the reaction. However, the role of HA in the oxidant scavenging appears to be less important than its role in the acceleration of oxidant production; note that the BA degradation by the Cu(II)/H2O2/ HA system is much greater than that by the Cu(II)/H2O2 system (Figure 3). Cu(I) + H 2O2 → Cu(II) + •OH + OH−or Cu(III) + 2OH− (k6 = 4 × 105M−1s−1at pH 6−8)

BA + •OH or Cu(III) → products

(6) (7)

Based on the sequence of reactions described above, the Cu(II)/HA system in the absence of O2 is not likely to produce reactive oxidants, which is consistent with the observation that the Cu(II)/HA/N2 system does not degrade BA (Figure 1a). Meanwhile, the Cu(II)/HA/O2 system did not make a significant difference in the BA degradation rate compared to the Cu(II)/HA system open to the atmosphere (Figure 1a), which explains the trade-off effect between H2O2 and Cu(I). The O2 aeration increases the steady-state concentration of H2O2 (Figure 1b), but decreases the concentration of Cu(I) (Figure 1c) (refer to reactions 3 and 4). Comparison of the Cu(II)/H2O2, Cu(II)/HA, and Cu(II)/ H2O2/HA Systems. The external supply of H2O2 to the Cu(II)/HA system (the Cu(II)/H2O2/HA system) greatly accelerates the BA degradation (Figure 3a) by removing the rate-limiting factor for the production of reactive oxidants in the Cu(II)/HA system (i.e., in situ generation of H2O2 by Cu(I)). The Cu(II)/H2O2 system showed the slowest BA degradation rate among the three systems (i.e., the Cu(II)/ H2O2, Cu(II)/HA, and Cu(II)/H2O2/HA systems), which is related to the slow reductive conversion of Cu(II) into Cu(I) by H2O2 (reaction 8) compared to the reduction of Cu(II) by HA (reaction 1 or 2). Cu(II) + H 2O2 → Cu(I) + HO2•( ↔ O2•− + H+) + H+(k 8 =