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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Environmental Science & Technology

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


TOC/Abstract art

2 ACS Paragon Plus Environment

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Abstract

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

3

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

4

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

5

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

6

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

7

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

8

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

9

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

10

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

11

suggesting that the alternative oxidant, cupryl ion (Cu[III]), is likely a dominant oxidant.

12



13 14

Introduction Advanced oxidation technologies based on the Fenton and Fenton-like reactions have been

15

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

16

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

17

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

18

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

19

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

20

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

21

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

22

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

23

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

24

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

25

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

26

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

27

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

28

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

29

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

30

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

31

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

32

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

33

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

34

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

35

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

38

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

39

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

40

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

41

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

42

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

43

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

44

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

45

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

46

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

47

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

48

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

51

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

52

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

53

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

54

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

55

organic contaminants.

56

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

58

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

62

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

63

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

65

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

66

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

69

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

70

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

71

oxidants generated by the reactions.

72 73

Materials and Methods

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

75

(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

77

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

78

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

79 80

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

84

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

85

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

87

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

88

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

90

(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

92

presented.

93 94

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

95

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

96

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

99

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

100

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

101

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

110

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.

118 119

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

132

hydroxybenzoates, dihydroxybenzoates, nitrobenzene, nitrobenzoates, and nitro-

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

134

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

135 136

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

138

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

139

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

141

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

142

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

143

information).

144 145

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

146

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

147

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

148

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

150

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

152

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

153

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

154

(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

156

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

159

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

160

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

161

information).

162

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

164

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

165

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

166

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

167

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

168

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

169

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

170

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

171

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

172 173

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

176

higher yields of HCHO.

177 178

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

180

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%,

182

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

183

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

184 185

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

186

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

189

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

190

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

192

of the systems tested.

193 194

Discussion

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

196

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

198

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

199

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

200

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

201

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

202

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)

205

NH2OH + Cu(II)

→

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

→

(1) (2)

206

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

207

(reactions 3 and 4).13,34 →

Cu(II) + O2•−

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Cu(I) + O2

209

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)

210

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

211

appears to be minor due to the low reaction rate.

212

2NH2OH + O2

→

N2 + H2O2 + 2H2O

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

213

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

214

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

215

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

216

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

217

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

218

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

220

Cu(II)/H2O2 system (Figure 3). →

Cu(II) + •OH + OH− or Cu(III) + 2OH−

221

Cu(I) + H2O2

222

(k6 = 4 × 105 M−1 s−1 at pH 6−8)37

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BA + •OH or Cu(III)

224

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

226

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

227

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

228

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

229

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

230

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

231 232

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

233

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

234

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

235

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

236

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

237

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

238

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

239

2).

240

Cu(II) + H2O2

→

Cu(I) + HO2•− (↔ O2•− + H+) + H+

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

270

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

272

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

273

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

274

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

275

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

276

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

277

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

278

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

279

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

280

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

283

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

284

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

285

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

286

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


287

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

288

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

289

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

290

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

291

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

292

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

293

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

Page 16 of 32

Page 17 of 32


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



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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Page 19 of 32


356

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


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.


Page 20 of 32

Page 21 of 32


392

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(2) Goldstein, S.; Meyerstein, D.; Czapski, G., The Fenton reagents. Free Radic. Biol. Med. 1993,

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(3) Goldstein, S.; Meyerstein, D., Comments on the mechanism of the “Fenton-Like” reaction. Acc.

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Degradation of 32 emergent contaminants by UV and neutral photo-Fenton in domestic wastewater

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Fenton processes. Water Res. 2006, 40, (9), 1841-1847. 23 ACS Paragon Plus Environment


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Comninellis, C., Electrochemical destruction of chlorophenoxy herbicides by anodic oxidation and

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electro-Fenton using a boron-doped diamond electrode. Electrochim. Acta 2004, 49, (25), 4487-

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4496.

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(22) Chen, L.; Ma, J.; Li, X.; Zhang, J.; Fang, J.; Guan, Y.; Xie, P., Strong enhancement on Fenton

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oxidation by addition of hydroxylamine to accelerate the ferric and ferrous iron cycles. Environ.

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Sci. Technol. 2011, 45, (9), 3925-3930.

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(23) Stumm, W.; Morgan, J. J., Aquatic chemistry: An introduction emphasizing chemical

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equilibria in natural waters. New York: Wiley, 1981.

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(24) Lee, H.-J.; Lee, H.; Lee, C., Degradation of diclofenac and carbamazepine by the copper(II)-

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catalyzed dark and photo-assisted Fenton-like systems. Chem. Eng. J. 2014, 245, 258-264.

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(25) Kim, H.-H.; Lee, H.; Kim, H.-E.; Seo, J.; Hong, S. W.; Lee, J.-Y.; Lee, C., Polyphosphate-

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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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(26) Klein, G. W.; Bhatia, K.; Madhavan, V.; Schuler, R. H., Reaction of hydroxyl radicals with

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

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1774.


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(27) Sagone, A. L.; Decker, M. A.; Wells, R. M.; Democko, C., A new method for the detection of

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

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(28) Eaton, A. D.; Franson, M. A. H.; Association, A. P. H.; Association, A. W. W.; Federation, W.

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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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(29) Eisenberg, G., Colorimetric determination of hydrogen peroxide. Ind. Eng. Chem. Anal. Ed.

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

468

(30) Zhou, X.; Mopper, K., Determination of photochemically produced hydroxyl radicals in

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

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(31) Timmins, G. S.; Liu, K. J.; Bechara, E. J. H.; Kotake, Y.; Swartz, H. M., Trapping of free

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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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(32) Chen, L.; Li, X.; Zhang, J.; Fang, J.; Huang, Y.; Wang, P.; Ma, J., Production of hydroxyl

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

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(17), 10373-10379.

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(33) Anderson, J. H., The copper-catalysed oxidation of hydroxylamine. Analyst 1964, 89, (1058),

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357-362. 25 ACS Paragon Plus Environment


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(34) Yuan, X.; Pham, A. N.; Xing, G.; Rose, A. L.; Waite, T. D., Effects of pH, chloride, and

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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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(35) Voelker, B. M.; Sedlak, D. L.; Zafiriou, O. C., Chemistry of superoxide radical in seawater:

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

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(36) Tomat, R.; Rigo, A.; Salmaso, R., Kinetic study on the reaction between O2 and

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

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seawater. Environ. Sci. Technol. 1987, 21, (8), 804-810.

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(38) Perez-Benito, F. J., Copper(II)-catalyzed decomposition of hydrogen peroxide: Catalyst

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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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22, (7), 768-771.

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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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atmospheric waters. Atmos. Environ. Part A 1993, 27, (14), 2173-2185.

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(41) Ozawa, T.; Hanaki, A., The first ESR spin-trapping evidence for the formation of hydroxyl

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(42) Verma, P.; Shah, V.; Baldrian, P.; Gabriel, J.; Stopka, P.; Trnka, T.; Nerud, F., Decolorization

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of synthetic dyes using a copper complex with glucaric acid. Chemosphere 2004, 54, (3), 291-295.

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cuprous ion and hydrogen peroxide: A spin-trapping study. Arch. Biochem. Biophys. 1995, 316, (1),

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(44) Belanzoni, P.; Bernasconi, L.; Baerends, E. J., O2 Activation in a dinuclear Fe(II)/EDTA

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2009, 113, (43), 11926-11937.

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compounds by nanoparticulate zero-valent iron and ferrous ion in the presence of oxygen. Environ.

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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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1183-1209.

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

514

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



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).


Page 29 of 32


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).


[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 )


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

ACS Paragon Plus Environment

Page 31 of 32


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).



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).


Page 32 of 32

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

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