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 Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b02067 • Publication Date (Web): 07 Jul 2016 Downloaded from http://pubs.acs.org on July 7, 2016

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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†, 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

Submitted to Environmental Science and Technology

*Corresponding author. Phone: +82-052-217-2812; Fax.: +82-052-217-2809; E-mail: [email protected] 1 ACS Paragon Plus Environment

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TOC/Abstract art

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Abstract

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This study reports that the combination of Cu(II) with hydroxylamine (HA) (referred to

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herein as Cu(II)/HA system) in situ generates H2O2 by reducing dissolved oxygen, subsequently

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producing reactive oxidants through the reaction of Cu(I) with H2O2. The external supply of

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H2O2 to the Cu(II)/HA system (i.e., the Cu(II)/H2O2/HA system) was found to further enhance

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the production of reactive oxidants. Both the Cu(II)/HA and Cu(II)/H2O2/HA systems effectively

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oxidized benzoate (BA) at pH between 4 and 8, yielding a hydroxylated product, p-

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hydroxybenzoate (pHBA). The addition of a radical scavenger, tert-butanol, inhibited the BA

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oxidation in both systems. However, electron paramagnetic resonance (EPR) spectroscopy

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analysis indicated that •OH was not produced under either acidic or neutral pH conditions,

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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

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intensively studied for the degradation of refractory organic contaminants in water and

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wastewater.1 The decomposition of hydrogen peroxide (H2O2) by the catalytic redox cycle of

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Fe(III)/Fe(II) is known to produce reactive oxidants such as hydroxyl radical (OH) and ferryl ion

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(Fe[IV]) that are capable of oxidizing organic compounds. Although there has been a long-term

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controversy on the identity of the reactive oxidant from the Fenton reaction (i.e., •OH vs.

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Fe(IV)),2,3 recent studies suggested that both •OH and Fe(IV) are produced with the dominant

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oxidant shifting from •OH to Fe(IV) as pH increases from acidic to neutral values.4-7 •OH is a

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nonselective oxidant that rapidly reacts with a broad spectrum of organic and inorganic

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compounds,8 whereas Fe(IV) oxidizes a relatively limited range of compounds.9-11 Besides low

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solubility of iron, the shift of the main oxidant to Fe(IV) is another factor that limits the

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applicability of the Fenton (-like) reactions at neutral pH.

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Similar to iron, copper can also convert H2O2 into reactive oxidants via the catalytic redox

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cycle of Cu(II)/Cu(I).12-14 Previous reports suggest that the nature of reactive oxidants from the

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copper-catalyzed Fenton-like system may also be pH-dependent; •OH and cupryl ion (Cu[III])

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are dominantly produced under acidic and neutral/alkaline conditions, respectively.7,15 However,

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the production of •OH under neutral/alkaline conditions cannot be completely excluded based on

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the observations that •OH scavengers such as tert-butanol were found to inhibit the oxidation of

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the target compounds and compounds such as benzoate from forming hydroxylated products.7,15

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In both the Fe(III)/H2O2 and the Cu(II)/H2O2 systems, the reduction of oxidized metal ion

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(i.e., Fe(III) and Cu(II)) by H2O2 is the rate-limiting step for the production of reactive

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oxidants.16,17 Approaches frequently employed to accelerate the reductive conversion of Fe(III) 4 ACS Paragon Plus Environment

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into Fe(II) in the Fe(III)/H2O2 system include UV light irradiation (photo-Fenton)18,19 and

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electricity application (electro-Fenton)20,21 . Recently, the addition of hydroxylamine (HA), a

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reducing agent, was demonstrated as a suitable method for facile Fe(III) reduction.22 The use of

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HA in the Fe(III)/H2O2 system accelerated the oxidation of benzoic acid by more than an order

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of magnitude, expanding the effective pH range up to 5.7. However, similar to most iron-based

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Fenton (-like) systems, the Fe(III)/H2O2/HA system still exhibits the optimal activity around pH

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3−4, and the system activity dramatically decreases as pH increases above 6.

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Considering the similarities between the Fe(III)/H2O2 and Cu(II)/H2O2 systems, one can

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postulate that HA also can be instrumental in enhancing the efficiency of Cu(II)/H2O2 system by

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accelerating the reduction of Cu(II). The fact that Cu(II)/H2O2 system efficiently functions at

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neutral pH is particularly appealing; note that i) Cu(II) has higher solubility at neutral pH than

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Fe(III); the solubility values of Cu(II) and Fe(III) at pH 7 are ca. 10−5 and 10−10 M,

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respectively,7,23 and ii) the Cu(II)/H2O2 system exhibits substantial activity toward oxidizing

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organic contaminants such as phenol, benzoate, diclofenac, and carbamazepine at neutral pH

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(interpreted as the contribution of •OH in those studies).7,24 In addition, one can further postulate

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that the combination of Cu(II) with HA can produce reactive oxidants without the external

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supply of H2O2 25 because Cu(I) is known to generate H2O2 by reducing dissolved oxygen (O2).

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Despite the anticipated benefits of using HA in copper-catalyzed Fenton-like reactions, no

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previous studies have evaluated Cu(II)/H2O2/HA and Cu(II)/HA systems for the degradation of

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organic contaminants.

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The objectives of this study are two-fold. First, we assess the potential of the copper-based

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Fenton-like systems with HA (i.e., the Cu(II)/H2O2/HA and the Cu(II)/HA systems) for the

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oxidation of select organic compounds; benzoate was selected as a main target compound 5 ACS Paragon Plus Environment

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because it has been used as a •OH probe compound and its oxidation mechanism by •OH is well-

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known.9,26,27 These systems are compared to the conventional Cu(II)/H2O2 system. Second, we

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evaluate the nature of the oxidants produced by the Cu(II)-catalyzed Fenton-like reaction under

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different pH conditions. The pH-dependent behaviors of copper-catalyzed Fenton-like systems

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were explained by the speciation of Cu(II) and Cu(III) complexes and their reactions. We expect

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that Cu(II)/H2O2/HA system shall be particularly useful for studying the nature of oxidants

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because the production of reactive oxidants is expected to be enhanced compared to the

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Cu(II)/H2O2 system. In the Cu(II)/H2O2 system concentrations of the short-lived reactive species

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are too low to be accurately assessed. This study suggests new oxidation technologies using

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copper and HA that are potentially applicable to degradation of refractory organic compounds at

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neutral pH. In addition, this study improves understanding of the chemistry of copper-based

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Fenton-like reactions, particularly providing insight into the pH-dependent nature of reactive

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oxidants generated by the reactions.

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Materials and Methods

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Reagents. All chemicals were of reagent grade and used without further purification. High purity

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(99.99%) gases such as nitrous oxide (N2O) and oxygen (O2) were used for some experiments.

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All solutions were prepared using 18.2 MΩ·cm Milli-Q water from a Millipore system. The

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stock solution of Cu(II) (10 mM) was prepared using cupric sulfate, and stored at 4°C until use.

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The stock solution of HA (500 mM) was prepared daily.

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Experimental Setup and Procedure. Experiments to evaluate the kinetics of organic compound

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oxidation in copper-based Fenton-like systems were conducted in a 100-mL Pyrex flask at room 6 ACS Paragon Plus Environment

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temperature (20 ± 2°C). No pH buffers were used for experiments at pH 3−5 because the pH

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variations before and after the reaction were minor. Phosphate and borate buffers (1 mM) were

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used for neutral (pH 6.5−8.0) and alkaline (9.0–10) pH ranges, respectively. The phosphate

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buffer can affect the Cu(II) speciation at pH 5−7 (Figure S1 in the supporting information). The

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initial pH was adjusted using 1 N HClO4 and 1 N NaOH solution. The reaction was initiated by

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adding an aliquot of stock solutions of H2O2 or HA, to a pH-adjusted solution containing organic

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compounds and Cu(II). Samples were withdrawn at predetermined time intervals and filtered

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using a 10-mL glass syringe and a 0.45-µm nylon syringe filter. Ethylenediaminetetraacetic acid

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(EDTA) (4 mM) was immediately added to quench the reaction. The experiments were

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conducted in duplicate, and the average values with the standard deviations (error bars) are

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presented.

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Analytical Methods. Benzoic acid (pKa = 4.2) or benzoate (BA) was analyzed by high

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performance liquid chromatography (HPLC) (UltiMate 3000, Dionex Co.) with UV absorbance

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detection at 255 nm. Separation was performed on a Dionex - Acclaim C-18 column (250 mm ×

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4.6 mm, 5 µm) using nitric acid solution (10 mM) and neat acetonitrile as eluents at a flow rate

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of 1.0 mL/min. p-Hydroxybenzoate (pHBA) and other oxidation products were analyzed on a

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LC/orbitrap MS/MS system. The analyses were performed using a rapid separation liquid

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chromatography (RSLC) (UltiMate 3000, Dionex Co.) coupled with a quadrupole-Orbitrap mass

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spectrometer (Q-Exactive, Thermo Fisher Scientific Inc.). Detailed procedures are described in

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the supporting information (Text S1). Ammonium (NH4+), nitrite (NO2−), and nitrate (NO3−) ions

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were analyzed by ion chromatography (IC) (ICS-3000, Dionex Co.) with conductivity detection.

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The separation of NH4+ was performed on an IonPac CS-17 cationic column (4 mm × 250 mm) 7 ACS Paragon Plus Environment

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using methanesulfonic acid (6 mM) as the eluent at a flow rate of 1.0 mL/min. For the analysis of

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NO2− and NO3−, an IonPac AS-9 anionic column (4 mm × 250 mm) was employed using

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carbonate solution (9.0 mM) as the eluent. The concentrations of total organic carbon (TOC) and

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total nitrogen (TN) were determined by a TOC/TN analyzer (TOC-5000A, Shimadzu Co.). N2O

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was analyzed using gas chromatography (GC 7820A, Agilent Co.) with the electron capture

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detector (ECD); a Porapak Q (80/100 mesh) column was used with high purity N2 as a carrier

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gas at a flow rate of 35 mL min-1. The concentrations of Cu(I) and H2O2 was

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spectrophotometrically determined by the neocuproine method28 and the titanium sulfate

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method29, respectively. Formaldehyde (HCHO) was analyzed by HPLC after DNPH

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derivatization.30 EPR spectroscopy was used to detect •OH, using 5,5-dimethyl-1-pyrroline N-

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oxide (DMPO) as a spin-trapping agent.31 EPR signals of the DMPO-OH spin adduct were

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obtained on a CW/Pulse EPR system (ELEXYS E580, Bruker Co.) with a 9.64 GHz microwave

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(0.94 mW) at a modulation frequency of 100 kHz and a modulation amplitude of 2.0 G.

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Results

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Degradation of BA by the Cu(II)/HA System at Neutral pH. The oxidative degradation of BA

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by the Cu(II)/HA system was examined at pH 7 under different aeration conditions (Figure 1a).

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Neither Cu(II) nor HA alone changed the concentration of BA. The combination of Cu(II) with

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HA under N2 condition (Cu(II)/HA/N2) did not degrade BA either; the Cu(II)/HA/N2 system did

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not produce H2O2 (data not shown), indicating that dissolved oxygen is the precursor of H2O2.

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The combination of Cu(II) with HA in the presence of oxygen degraded BA by more than 70%

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in 4 h. The Cu(II)/HA system with no aeration (open to the atmosphere) exhibited similar degree

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of BA degradation to that with O2 aeration (Cu(II)/HA/O2). The concentrations of H2O2 and Cu(I) 8 ACS Paragon Plus Environment

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were monitored in the Cu(II)/HA and Cu(II)/HA/O2 systems (Figures 1b and 1c). The

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Cu(II)/HA/O2 system produced higher concentration of H2O2 than the Cu(II)/HA system (Figure

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1b), but produced lower concentration of Cu(I) (Figure 1c). The products produced during the

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BA degradation by the Cu(II)/HA system were analyzed; compounds including

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hydroxybenzoates, dihydroxybenzoates, nitrobenzene, nitrobenzoates, and nitro-

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hydroxybenzoates were identified, and the pathways of BA oxidation were presented (refer to

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Figures S2−S4 in the supporting information for details).

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Decomposition of HA. To examine the decomposition of HA, variation in concentrations of TN

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and nitrogenous products22 (NH4+, NO2−, NO3−) was monitored in the Cu(II)/HA system (Figure

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2). The TN concentration decreased by 97% in 4 h, exhibiting the pseudo-first order decay (k =

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0.0216 min−1). However, the production of NH4+, NO2−, and NO3− was very minor throughout

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the entire reaction. A small amount of N2O was also detected in the headspace of the reactor

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(Figure S5 in the supporting information). In the Cu(II)/H2O2/HA system, the decrease in TN

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concentration was faster than that in the Cu(II)/HA system (Figure S6 in the supporting

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information).

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Cu(II)/H2O2, Cu(II)/HA, Cu(II)/H2O2/HA Systems. The oxidative degradation of BA was

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compared in three systems: Cu(II)/H2O2, Cu(II)/HA, and Cu(II)/H2O2/HA. The rate of

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degradation of BA was found to be Cu(II)/H2O2 < Cu(II)/HA < Cu(II)/H2O2/HA at pH 7 (Figure

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3a); the combination of H2O2 with HA (i.e., the H2O2/HA system) did not degrade BA. The

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Cu(II)/H2O2/HA system exhibited the synergistic enhancement of BA degradation; the

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degradation rate of BA by the Cu(II)/H2O2/HA system was greater than the simple sum of those 9 ACS Paragon Plus Environment

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by the Cu(II)/H2O2 and Cu(II)/HA systems (also refer to Figure 3b). The BA degradation

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experiments were performed at different pH values from 3 to 10 in each system. The resulting

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pseudo-first order rate constants for the BA degradation were depicted as a function of pH

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(Figure 3b). Overall, the circumneutral pH conditions favored the BA degradation; the BA

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degradation rate was lower under acidic pH conditions than alkaline pH conditions. At almost all

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pH values, the BA degradation rate was Cu(II)/H2O2 < Cu(II)/HA < Cu(II)/H2O2/HA.

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Meanwhile, a very slow degradation of BA was observed by the H2O2/HA system under acidic

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pH conditions; consistent with the recent report that •OH is produced by the acid-catalyzed

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reaction of H2O2 with HA.32 We also observed that the decomposition of H2O2 in the

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Cu(II)/H2O2/HA system accelerated with increasing pH (Figure S7 in the supporting

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information).

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The pHBA formation was monitored during the BA oxidation by the Cu(II)/H2O2,

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Cu(II)/HA, Cu(II)/H2O2/HA, and H2O2/HA systems at different pH values (Figure 3c, and

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Figures S8−S10 in the supporting information). pHBA formed after 30 min of reaction time

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(Figure 3c), consistent with the rate of BA degradation (Figure 3b) except for the

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Cu(II)/H2O2/HA system. The Cu(II)/H2O2/HA system exhibited much lower pHBA

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concentrations at pH 5-7 than expected, which may be attributed to the secondary oxidation of

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pHBA to dihydroxybenzoates during the oxidation of BA (Figures S2−S4). pHBA was rapidly

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formed in the initial 10 min, and then degraded as the reaction proceeded in the Cu(II)/H2O2/HA

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system (Figures S2-1c and S2-3), whereas the pHBA concentrations in the other systems

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continuously increased over the entire reaction time (Figures S2-1a, b, and d).

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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 10 ACS Paragon Plus Environment

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S11 in the supporting information). Overall, the HCHO production exhibited similar trends to the

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BA oxidation (Figures 3b and 3c), except for the H2O2/HA system which exhibited relatively

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higher yields of HCHO.

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Effect of tert-Butanol on BA Degradation by the Cu(II)/H2O2/HA System. The addition of

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tert-butanol inhibited the BA degradation by the Cu(II)/H2O2/HA system (Figure 4). Without

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tert-butanol, BA was almost completely degraded in 2 h. However, in the presence of 1, 10, and

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100 mM tert-butanol, the BA degradation efficiencies in 2 h were 92.7, 57.8, and 10.6%,

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respectively. The BA degradation rate (pseudo first-order rate constant) decreased by 1.6, 2.3,

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and 8-fold in the presence of 1, 10, and 100 mM tert-butanol, respectively (the inset in Figure 4).

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EPR Spectroscopy. The EPR technique with DMPO as a spin-trapping agent was used to

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identify the production of •OH in the systems of Cu(II)/H2O2/HA, H2O2/HA, Cu(II)/HA, and

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Cu(II)/H2O2 at pH 3 and 7 (Figure 5). At pH 3, the Cu(II)/H2O2/HA and H2O2/HA systems

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exhibited the signal of DMPO-OH spin adduct, 1:2:2:1 quartet lines with hyperfine constants of

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aN = aH = 14.9 G31 (Figure 5a). A notable observation is that the signal intensity of the H2O2/HA

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system was much higher than that of the Cu(II)/H2O2/HA system. The Cu(II)/HA, and

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Cu(II)/H2O2 systems did not generate noticeable signals. At pH 7, no signals were obtained any

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of the systems tested.

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Discussion

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Production of Reactive Oxidants by Cu(II) in Combination with HA. The Cu(II)/HA system

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produces reactive oxidants via the reaction of in situ generated Cu(I) and H2O2. Primarily, HA 11 ACS Paragon Plus Environment

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reduces Cu(II) to Cu(I). The two-electron oxidation of HA into N2O has been postulated based

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on the observed stoichiometry between Cu(II) consumed and the total gas produced (reaction

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1).33 However, the little production of N2O in this study (Figure S5) suggests that the one-

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electron oxidation of HA into N2 may be more favored (reaction 2). The kinetics for this reaction

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are unknown, but the second-order rate constant is estimated to be at least higher than 103 M−1 s−1;

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10 µM Cu(II) was completely reduced to Cu(I) in 5 s by the addition of 0.2 mM HA under

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anoxic conditions (data not shown).

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NH2OH + 2Cu(II)

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NH2OH + Cu(II)



1/2N2O + 1/2H2O + 2Cu(I) + 2H+ 1/2N2 + H2O + Cu(I) + H+



(1) (2)

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Subsequently, the Cu(I) produced reduces O2 into H2O2 by two single-electron transfer reactions

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(reactions 3 and 4).13,34 →

Cu(II) + O2•−

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Cu(I) + O2

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Cu(I) + O2•− + 2H+



(k3 = 3.1 × 104 M−1 s−1 at pH 6−8)34 (k4 = 2.0 × 109 M−1 s−1)35

Cu(II) + H2O2

(3) (4)

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Another possibility is HA directly reduces O2 into H2O2 (Reactions 5). However, this reaction

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appears to be minor due to the low reaction rate.

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2NH2OH + O2



N2 + H2O2 + 2H2O

(k5 = 9.4 × 10−2 M−1 s−1 at pH 7)36 (5)

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Finally, the reaction of Cu(I) with H2O2 (the Fenton-like reaction; reaction 6) produces reactive

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oxidants such as •OH and Cu(III), capable of oxidizing BA (reaction 7). In this series of

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reactions, most of HA is liberated as N2O and N2 gases without leaving residual nitrogenous

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products in the solution (Figure 2). In fact, HA can also serve as the scavenger of reactive

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oxidants while it is decomposed during the reaction. However, the role of HA in the oxidant

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scavenging appears to be less important than its role in the acceleration of oxidant production;

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note that the BA degradation by the Cu(II)/H2O2/HA system is much greater than that by the

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Cu(II)/H2O2 system (Figure 3). →

Cu(II) + •OH + OH− or Cu(III) + 2OH−

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Cu(I) + H2O2

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(k6 = 4 × 105 M−1 s−1 at pH 6−8)37

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BA + •OH or Cu(III)

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Based on the sequence of reactions described above, the Cu(II)/HA system in the absence of



(6)

Products

(7)

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O2 is not likely to produce reactive oxidants, which is consistent with the observation that the

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Cu(II)/HA/N2 system does not degrade BA (Figure 1a). Meanwhile, the Cu(II)/HA/O2 system

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did not make a significant difference in the BA degradation rate compared to the Cu(II)/HA

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system open to the atmosphere (Figure 1a), which explains the trade-off effect between H2O2 and

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Cu(I). The O2 aeration increases the steady-state concentration of H2O2 (Figure 1b), but

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decreases the concentration of Cu(I) (Figure 1c) (refer to reactions 3 and 4).

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Comparison of the Cu(II)/H2O2, Cu(II)/HA, and Cu(II)/H2O2/HA Systems. The external

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supply of H2O2 to the Cu(II)/HA system (the Cu(II)/H2O2/HA system) greatly accelerates the BA

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degradation (Figure 3a) by removing the rate-limiting factor for the production of reactive

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oxidants in the Cu(II)/HA system (i.e., in situ generation of H2O2 by Cu(I)). The Cu(II)/H2O2

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system showed the slowest BA degradation rate among the three systems (i.e., the Cu(II)/H2O2,

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Cu(II)/HA, and Cu(II)/H2O2/HA systems), which is related to the slow reductive conversion of

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Cu(II) into Cu(I) by H2O2 (reaction 8) compared to the reduction of Cu(II) by HA (reaction 1 or

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2).

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Cu(II) + H2O2



Cu(I) + HO2•− (↔ O2•− + H+) + H+

(k8 = 6 can be attributed to the precipitation of Cu(II); the

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formation of tenorite (CuO) is favored at pH > 6 (refer to Figure S1 in the supporting

271

information). However, the decomposition of H2O2 in the Cu(II)/H2O2/HA system did not

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decelerate at alkaline pH; in fact, decomposition of H2O2 accelerated (Figure S7). The

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Cu(II)/H2O2 system also showed increasing rates of the H2O2 decomposition with increasing

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pH.24 In addition, according to a previous study7, some target compounds such as Reactive Black

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5 and As(III) showed even greater degradation at alkaline pH by the Cu(II)/H2O2 system using

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0.1 mM Cu(II). These observations collectively imply that the speciation of Cu(II) does not

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necessarily limit the production of reactive oxidants by the Fenton-like reactions; the effect of

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Cu(II) speciation on the formation of Cu(II)-peroxo complex (reaction 9) may be minor. The

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decreasing rates of BA degradation at alkaline pH is believed to be associated with the reactivity

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change of the responsible oxidant (Cu(III) species) due to the pH-dependent speciation.

281 282

Nature of Reactive Oxidants Produced by Copper-Catalyzed Fenton-like Reactions. There

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is a debate about the identity of reactive oxidants produced by the Fenton reaction (i.e., the

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reaction of Fe(II) with H2O2 forming •OH vs. Fe(IV))2,3. Investigators have provided different

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views on the production of •OH vs. Cu(III) by the copper-catalyzed Fenton-like reaction.7,12-15,24

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Recent studies have claimed that Cu(III) rather than •OH is the dominant oxidant under neutral 15 ACS Paragon Plus Environment

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and alkaline pH conditions. This claim is based on observations of the oxidation kinetics on

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different organic compounds and the effect of •OH scavengers on these compounds are

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inconsistent with the known reactivity of •OH with these compounds.7,15

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A lack of signal in the EPR spectra at pH 7 confirmed that the production of •OH is not

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important in the copper-catalyzed Fenton-like systems at neutral pH. There are a few cases that

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oxidants other than •OH induce the DMPO-OH signal have been reported41, but the literature has

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never been reported that •OH does not induce the DMPO-OH signal. A notable observation is

294

that the DMPO-OH signals appear negligible, even at pH 3, except for in the Cu(II)/H2O2/HA

295

and H2O2/HA systems (Figure 5a), suggesting that •OH is not produced by the copper-catalyzed

296

Fenton-like reaction at acidic pH either. Although the Cu(II)/H2O2/HA system generated the

297

DMPO-OH signal, its intensity was even lower than that of the H2O2/HA system (which was

298

recently found to produce •OH at acidic pH32). Considering that the BA degradation rates in the

299

Cu(II)/H2O2/HA system are much higher than those in the H2O2/HA system (Figure 3b) at acidic

300

pH values, we conclude that the Cu(II)/H2O2/HA system does not produce •OH as a main

301

reactive oxidant, and the minor DMPO-OH signal observed in the Cu(II)/H2O2/HA system at pH

302

3 may result from the partial contribution of the H2O2/HA reaction. On the contrary, the

303

Fe(III)/H2O2/HA system produced higher intensity DMPO-OH signals than the H2O2/HA system

304

at pH 3 (Figure S12 in the supporting information), indicating that the Fenton reaction (iron-

305

catalyzed) proceeds by a different reaction mechanism possibly yielding •OH. Several studies

306

have presented evidence for •OH from the reaction of complexed Cu(II) with H2O2 based on

307

EPR analysis using several spin-trapping agents including DMPO.41-43 However, most of these

308

studies employed organic ligands, and •OH may indeed be formed depending on the type of

309

Cu(II) complexes. Similar to the iron-based Fenton-like reactions25, 44-46, the ligand coordinated 16 ACS Paragon Plus Environment

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310

to Cu(II) can affect the electron-transfer mechanism of Fenton (-like) reactions (one- vs. two-

311

electron transfer), determining the resultant reactive oxidant (•OH vs. Cu(III), possibly

312

complexed forms with the ligand). If the postulation that Cu(III) rather than •OH is dominantly produced by the copper-

313 314

catalyzed Fenton-like reaction in all pH ranges were true, Cu(III) would be responsible for the

315

hydroxylation of BA into hydroxybenzoates including pHBA (Figure 3c and Figures S2−S4) as

316

well as the substantial reactivity with tert-butanol (Figure 4). Similar observations have been

317

reported in a previous study; the copper-catalyzed Fenton-like reaction at pH 8 hydroxylates

318

phthalhydrazide and rapidly reacts with various substrates including formate, bromide, and tert-

319

butanol (k = 106−107 M−1 s−1).15 However, the results of Figures 3 and 4 indicate that Cu(III)

320

exhibits a lesser degree of hydroxylating BA and scavenging by tert-butanol than •OH. The

321

yields of pHBA in the Cu(II)-catalyzed Fenton-like systems (i.e., Cu(II)/H2O2, Cu(II)/HA, and

322

Cu(II)/H2O2/HA systems) appear relatively low compared to those in the H2O2/HA system where

323



324

are the lowest (the values differ from those in the Cu(II)-catalyzed Fenton-like systems by

325

several orders of magnitude) (Figure 3b), whereas the differences of pHBA concentrations in

326

between the H2O2/HA system and the Cu(II)-catalyzed Fenton-like systems are relatively small

327

(Figure 3c, also refer to Figures S2). The hydroxylation of BA by Cu(III) may proceed via the

328

hydrogen-abstraction from the aromatic ring followed by the oxygen-transfer. A similar

329

hydroxylation mechanism has been suggested for the reaction of Fe(IV).47 The oxidant

330

scavenging effect of tert-butanol (Figure 4) is also lower than expected by •OH. Based on the

331

known second-order rate constants for reactions of •OH with BA and tert-butanol (kBA = 5.9 ×

332

109 M−1 s−1 and kt-BuOH = 6.6 × 108 M−1 s−1)8, the •OH scavenging efficiencies (kt-BuOH[t-

OH is the dominant reactive oxidant. Note that the BA degradation rates in the H2O2/HA system

17 ACS Paragon Plus Environment

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333

BuOH]/(kBA[BA] + kt-BuOH[t-BuOH])) by 1, 10, and 100 mM tert-butanol are 53, 92, and 99%,

334

respectively. However, the BA degradation rate decreased by 0.0286 min-1, 0.0196 min-1, and

335

0.0056 min-1 in the presence of 1, 10, and 100 mM tert-butanol, respectively (the inset of Figure

336

4).

337

Cu(III) should undergo the pH-dependent speciation (i.e., Cu3+, Cu(OH)2+, Cu(OH)2+,

338

Cu(OH)3, Cu(OH)4− etc. for Cu(III)-hydroxo complexes)7,12, even if the stability constants of

339

these Cu(III) species are still uncertain; the range of pKa values for Cu(OH)2+ and Cu(OH)2+

340

have been speculated to be < 3.5 and 4−6, respectively.12 Metal-hydroxo complexes have

341

decreasing oxidation power with increasing pH as they shift to the forms with more hydroxo

342

ligands. The occurrence of less reactive Cu(III) species (e.g., Cu(OH)3, Cu(OH)4−) may be

343

responsible for the decreasing rate of BA degradation at alkaline pH (Figure 3b). Previous

344

studies have shown that the degradation of phenol and pharmaceutical compounds by the

345

Cu(II)/H2O2 system decelerates at alkaline pH,7,24 which may suggest the occurrence of less

346

reactive Cu(III)-hydroxo complexes.

347

It is interesting to compare the reactivity of Cu(III) and Fe(IV) produced by the Fenton (-

348

like) reactions. Reports in the literature have proposed that the Fenton reaction (iron-catalyzed)

349

produces Fe(IV) at neutral pH, which is capable of oxidizing As(III), some primary alcohols

350

(methanol and ethanol), and a reactive dye.2,7,9-11 However, this Fe(IV) species was effective in

351

oxidizing neither aromatic compounds (phenol and benzoate) nor secondary and tertiary alcohols

352

(isopropanol and tert-butanol).7,9-11 The results in this study and in the literature collectively

353

indicate that the Cu(III) species from the Fenton-like reaction appears to have higher reactivity

354

than the Fe(IV) species produced under analogous conditions.

355 18 ACS Paragon Plus Environment

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Environmental Implications. The combination of Cu(II) with HA (i.e., the Cu(II)/HA and

357

Cu(II)/H2O2/HA systems) provides a new potential oxidation process for the degradation of

358

recalcitrant organic contaminants in wastewater and reclaimed water. The main oxidant, the

359

Cu(III) species, appears to have nonselective reactivity with a wide spectrum of organic

360

compounds (similar to •OH, but probably less reactive), although the information about the

361

nature and reactivity of Cu(III) is still very limited. In addition, these systems using Cu(II) and

362

HA are effectively applicable for a broad pH range (approximately 4−8). Conversely, iron-

363

catalyzed Fenton (-like) systems (even with the addition of HA22) do not cover a broad pH range.

364

Concerns about the potential toxicity of HA are mitigated by the fact HA mostly decomposed

365

into inert gases without forming residual products in solution. However, to reduce the

366

environmental risk from residual Cu(II) (the maximum contaminant level for copper set by the

367

USEPA is 1.3 mg/L = 20.5 µM)48 and recycle the copper catalyst, immobilization of Cu(II) by

368

coupling with nanofiltration or using heterogeneous copper catalysts would be necessary, which

369

warrants a further study. In addition, the performance of Cu(II)/HA and Cu(II)/H2O2/HA systems

370

is expected to be influenced by the type and concentration of copper-chelating compounds that

371

exist in the water to be treated (both negative and positive effects are possible) because oxidants

372

of different reactivity can be generated depending on the ligand. More studies are needed on the

373

nature of complexed Cu(III) species.

374 375

Acknowledgments

376

This work was supported by the National Research Foundation of Korea (NRF) grants funded by

377

the Korean government (MSIP) (NRF-2015R1A5A7037825 and NRF-

378

2015R1A2A1A15055840), and in part by the KIST-UNIST partnership program (1.150091.01). 19 ACS Paragon Plus Environment

Environmental Science & Technology

379 380

Supporting Information Available

381

Calculated speciation of Cu(II) (Figure S1), analytical procedures of pHBA and other products

382

using the RSLC/orbitrap MS/MS system (Text S1), product analysis of BA degradation in the

383

Cu(II)/HA system (Text S2 and Figures S2−S4), production of N2O in the Cu(II)/HA system

384

(Figure S5), variation in TN concentration during the decomposition of HA in the

385

Cu(II)/H2O2/HA system (Figure S6), decomposition of H2O2 in the Cu(II)/H2O2/HA system at

386

different pH values (Figure S7), production of pHBA by Cu(II)/H2O2, Cu(II)/HA,

387

Cu(II)/H2O2/HA, and H2O2/HA systems (Figures S8−S210), production of HCHO from

388

methanol oxidation by Cu(II)/H2O2, Cu(II)/HA, Cu(II)/H2O2/HA, and H2O2/HA systems (Figure

389

S11), and EPR spectra obtained by spin trapping with DMPO in Fe(III)/H2O2 and

390

Fe(III)/H2O2/HA systems (Figure S12): This material is available free of charge via the Internet

391

at http://pubs.acs.org.

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392

References

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enhanced production of reactive oxidants by nanoparticulate zero-valent iron and ferrous ion in the

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presence of oxygen: Yield and nature of oxidants. Water Res. 2015, 86, 66-73.

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benzoic acid. Isomer distribution in the radical intermediates. J. Phys. Chem. 1975, 79, (17), 1767-

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hydroxyl radical production by phagocytic cells. BBA-Gen. Subjects 1980, 628, 90-97.

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E., Standard methods for the examination of water & wastewater, 21st, ed. American Public Health

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Association, 2005.

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1943, 15, (5), 327-328.

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seawater and freshwater. Mar. Chem. 1990, 30, 71-88.

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radicals with direct in vivo EPR detection: a comparison of 5,5-dimethyl-1-pyrroline-N-oxide and

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5-diethoxyphosphoryl-5-methyl-1-pyrroline-N-oxide as spin traps for •OH and SO4•−. Free Radic.

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Biol. Med. 1999, 27, (3–4), 329-333.

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radical via the activation of hydrogen peroxide by hydroxylamine. Environ. Sci. Technol. 2015, 49,

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bicarbonate on Cu(I) oxidation kinetics at circumneutral pH. Environ. Sci. Technol. 2012, 46, (3),

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1527-1535.

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Reactions with organic Cu complexes. Environ. Sci. Technol. 2000, 34, (6), 1036-1042.

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hydroxylamine. J. Electroanal. Chem. Interfacial Electrochem. 1975, 59, (3), 255-260.

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activation by halide Ions. Monatshefte für Chemie / Chemical Monthly 2001, 132, (12), 1477-1492.

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(39) Sharma, V. K.; Millero, F. J., Oxidation of copper(I) in seawater. Environ. Sci. Technol. 1988,

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(40) Sedlak, D. L.; Hoigné, J., The role of copper and oxalate in the redox cycling of iron in

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(46) Keenan, C. R.; Sedlak, D. L., Ligand-enhanced reactive oxidant generation by nanoparticulate

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zero-valent iron and oxygen. Environ. Sci. Technol. 2008, 42, (18), 6936-6941.

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(47) Sychiov, A. Y.; Isac, V. G., Iron compounds and mechanisms of homogeneous catalysis of O2

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(48) Agency, U. S. E. P. National primary drinking water regulations; United States, 2009.

514

http://water.epa.gov/drink.

27 ACS Paragon Plus Environment

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1.2

0.8

Cu(II) HA Cu(II)/HA/N2

0.6

Cu(II)/HA/O2

[H2O2] (mM)

(a)

Cu(II)/HA 0.4

0.2

Cu/HA/O2

0.9

(b)

Cu/HA (open to atmosphere)

0.6 0.3 0.0

(open to atmosphere)

Cu/HA/O2 [Cu(I)] (mΜ)

[BA]/[BA]0

1.0

Page 28 of 32

0.0

(c)

Cu/HA

0.04

(open to atmosphere)

0.02

0.00 0

60

120

180

240

0

60

Reaction time (min)

120

180

240

Reaction time (min)

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).

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5 TN NH4+

Products conc. (mM)

4

NO2NO3-

3

2

1

0 0

60

120

180

240

Reaction time (min)

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).

29 ACS Paragon Plus Environment

[BA]/[BA]0

0.8

0.6

Cu(II)/H2O2 Cu(II)/HA Cu(II)/H2O2/HA

0.4

0.2

0.0 0

60

120

180

Reaction time (min)

240

1e-1

Page 30 of 32

2.0

(b)

(c)

Cu(II)/H2O2 Cu(II)/HA Cu(II)/H2O2/HA

1.5 1e-2

[pHBA] (µM)

(a)

-1

1.0

Pseudo first order rate constant, k (min )

Environmental Science & Technology

1e-3

Cu(II)/H2O2

H2O2/HA

1.0

0.5

Cu(II)/HA Cu(II)/H2O2/HA H2O2/HA

1e-4

0.0 3

4

5

6

7

8

9

pH

10

3

4

5

6

7

8

9

10

pH

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)).

30

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1.0

0.05

0.8 Control

[BA]/[BA]0

with 10 mM T-BuOH

0.6

with 100 mM T-BuOH

k (min-1)

0.04

with 1 mM T-BuOH

0.03 0.02 0.01 0.00 0

1

10

100

T-BuOH (mM)

0.4

0.2

0.0 0

60

120

180

240

Reaction time (min) Figure 4. Effect of tert-butanol 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).

31 ACS Paragon Plus Environment

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Intensity (a. u.)

Cu(II)/H2O2/HA

H2O2/HA

Cu(II)/HA

Cu(II)/H2O2

3220

3240

3260

3280

3300

Magnetic Field (G) 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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