Fenton-like copper redox chemistry revisited: Hydrogen peroxide and superoxide mediation of copper-catalyzed oxidant production

Fenton-like copper redox chemistry revisited: Hydrogen peroxide and superoxide mediation of copper-catalyzed oxidant production

Journal of Catalysis 301 (2013) 54–64 Contents lists available at SciVerse ScienceDirect Journal of Catalysis journal homepage: www.elsevier.com/loc...

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Journal of Catalysis 301 (2013) 54–64

Contents lists available at SciVerse ScienceDirect

Journal of Catalysis journal homepage: www.elsevier.com/locate/jcat

Fenton-like copper redox chemistry revisited: Hydrogen peroxide and superoxide mediation of copper-catalyzed oxidant production A. Ninh Pham, Guowei Xing, Christopher J. Miller, T. David Waite ⇑ School of Civil and Environmental Engineering, The University of New South Wales, Sydney, NSW 2052, Australia

a r t i c l e

i n f o

Article history: Received 17 October 2012 Revised 23 January 2013 Accepted 26 January 2013

Keywords: Copper Fenton reaction Reactive oxygen species Catalysis Hydroxyl radical Redox Contaminant oxidation Neurodegenerative disease

a b s t r a c t Copper toxicity has been attributed to its potential as a catalyst for oxidative damage to tissues through redox cycling between Cu(I) and Cu(II), particularly in the presence of H2O2, a by-product of oxygen metabolism. In this study, the reactions of nanomolar concentrations of Cu(I) and Cu(II) with H2O2 have been investigated in 2.0 mM NaHCO3 and 0.7 M NaCl at pH 8.0. Measurements of both the formation of the hydroxylated phthalhydrazide chemiluminescent product and the degradation of formate in the absence and presence of compounds with well-known reactivity with HO indicated that the reaction between Cu(I) and H2O2 did not result in the production of HO but involved the formation of a higher oxidation state of copper, Cu(III). The Cu(III) so-formed reacts with the substrates that were present at much slower rates compared to those of HO. The rate of formation of HO from the dissociation of Cu(III) was extremely slow at pH 8.0 with the result that HO is not an important oxidant in this system. The rapid rate of reaction of Cu(III) with Cu(I) contributes significantly to the redox cycle of copper and the associated oxidizing capacity of the Cu(I)/Cu(II)/H2O2/O2 system with exogenous input of H2O2 and O 2 exhibiting the ability to mediate ongoing copper-catalyzed production of the powerful oxidant, Cu(III). Ó 2013 Elsevier Inc. All rights reserved.

1. Introduction Copper (Cu) is an essential trace element in living systems due to its presence in a number of important enzymes that are involved in a variety of biological processes including photosynthesis, electron transport, cell wall metabolism, and oxidative stress protection [1,2]. In natural waters, Cu typically occurs in either cuprous (Cu(I)) or cupric (Cu(II)) oxidation states with transformations between these oxidation states playing a critical role in the speciation, transport, and bioavailability of this element [3]. At circumneutral pH, Cu(I) is quickly oxidized to Cu(II) on the timescale of minutes [4,5]. While organic complexation of Cu(I) is generally considered unimportant [4], almost all Cu(II) is largely complexed by dissolved organic matter and thus not immediately bioavailable [6]. Cu(II) can be reduced to Cu(I) in sunlit surface waters by reaction with photo-generated reactive oxygen species (ROS) and through direct photolysis of Cu(II) complexes. In the human body, copper is bound strongly to serum albumin or incorporated into ceruloplasmin [7] with approximately 80 mg Cu present in an average adult [8]. Copper in excess however, especially in its free hydrated form (i.e., Cu2+), can be potentially toxic to both plants and organisms by altering membrane permeability and by affecting chromatin structure, protein synthesis, and various enzyme activities [2]. Par⇑ Corresponding author. Fax: +61 2 9313 8341. E-mail address: [email protected] (T.D. Waite). 0021-9517/$ - see front matter Ó 2013 Elsevier Inc. All rights reserved. http://dx.doi.org/10.1016/j.jcat.2013.01.025

ticular attention has been given to the generation of hydroxyl radicals (HO) and/or other ROS as a result of the reaction of Cu with hydrogen peroxide (H2O2) which is ubiquitous in the cells of organisms as a by-product of oxygen metabolism [8]. In humans, several neurodegenerative diseases including Alzheimer’s and Parkinson’s disease are also characterized by modified copper homeostasis. Changes of Cu metabolism in the brain seem to contribute either directly or indirectly to increased oxidative stress, an important factor in neuronal toxicity [9]. Of the various free radicals generated in living organisms, HO is the most potent and can damage cells via non-selective oxidation of proteins, lipids, fatty acids, and nucleic acids [10]. Hydrogen peroxide (H2O2) is an intermediate in the reduction of oxygen to water and is present in natural aquatic systems at concentrations exceeding 107 M [11,12]. This relatively reactive species may influence the biogeochemistry of various transition metals and their complexes as it can act effectively as a reductant or as an oxidant. In human cells, H2O2, a by-product of oxygen metabolism, may be present at micromolar concentrations [13]. Generally, H2O2 is relatively harmless as it reacts with biomolecules at relatively low rates and specific enzymes to facilitate its removal are usually present [7]. However, in the presence of certain metals, the presence of H2O2 may lead to the formation of the highly reactive and damaging hydroxyl radical (HO). Although a large number of kinetic studies on the reaction of both Cu(I) and Cu(II) with H2O2 have been reported [7,13–37], none of the previous work has convincingly elucidated either the

A.N. Pham et al. / Journal of Catalysis 301 (2013) 54–64

nature of the oxidant formed or the transformations occurring between the various species present. Most of these studies have been undertaken in the presence of Cu complexing agents which, although typical of natural environments, complicates interpretation of the results obtained. Several studies of the copper/H2O2 system in the absence of complexing agents have been investigated either in acidic or alkaline solutions [23,24,26,27], which are non-representative of the circumneutral conditions typical of most natural aquatic environments. Additionally, possible precipitation of Cu(II) as a result of the high concentrations of Cu(II) used in many of these ligand-free studies [26,33] also creates interpretation difficulties. In general, there are two major alternate mechanisms that have been proposed for the reaction between Cu(II) and H2O2: the ‘‘free radical’’ versus ‘‘complex’’ reaction mechanisms. In the ‘‘free radical’’ mechanism, Cu(II) oxidizes H2O2 to O 2 with the Cu(I) formed in this process able to react with excess H2O2 to form HO [19,28,29]. The ‘‘complex’’ mechanism assumes that, rather than changing its oxidation state, Cu(II) forms a complex with peroxide with no formation of radical species [16,34,36]. Recent studies, however, have suggested that alternate reaction pathways may be involved [13,33]. The reaction of Cu(I) with H2O2 is even more problematic, especially with regard to identification of the active oxidizing intermediates. While a ‘‘Fenton-like’’ mechanism has commonly been proposed to describe the reaction of Cu(I) with H2O2 with evidence presented for the formation of HO [15,20,32], the results of other studies [22–24,27] suggest that, at least under some conditions, a higher oxidation state of copper (i.e., Cu(III)) rather than free HO may be the true active intermediate. In this study, the kinetics of reactions of nanomolar concentrations of Cu(I) and Cu(II) with H2O2 are investigated with the purpose of answering two particular questions: a) what mechanism is likely prevailing in the reactions of Cu(II) and Cu(I) with H2O2 at circumneutral pH and b) is HO the active oxidizing intermediate in such systems? A kinetic model is developed to describe the complex interactions occurring in the Cu(I)/Cu(II)/H2O2/O2 system over a range of Cu(I), Cu(II), and H2O2 concentrations and in the presence and absence of known HO radical scavengers. The relative importance of individual reactions in the Cu(I)/Cu(II)/H2O2/O2 system is also discussed.

2. Materials and methods

55

were made from a 30.7% w/v H2O2 solution, which had been standardized by UV spectrophotometry employing molar absorptivities of e240 = 38.3 ± 1.1 M1 cm1 and e260 = 13.0 ± 0.4 M1 cm1 [38]. Radiolabelled formic acid (H14COONa) stock solutions (0.2 mM) were prepared from a concentrated H14COONa solution. A range of potassium bromide (KBr) and 2-methylpropan-2-ol ((CH3)3COH, TBA) stock solutions were prepared in Milli-Q water. Details of additional chemicals required for Cu(I), H2O2, HO, and H14COONa measurements are given in the Supporting information (SI).

2.2. Experimental measurements Cu(I) concentrations were determined spectrophotometrically using the bathocuproine (BC) method [5,39] with details given in the SI. Upon addition of BC, Cu(I) rapidly forms a Cu(I)–(BC)2 complex which was measured colorimetrically at 484 nm. Ethylenediaminetetraacetate (EDTA, 0.25 mM) was added to the 50 lM BC solution in order to bind Cu(II), thereby preventing any BC-induced reduction of Cu(II). Disproportionation of Cu(I) (to form Cu(0) and Cu(II)) was negligible under the conditions investigated here [40]. H2O2 concentrations were measured using the Amplex Red method [41]. Briefly, in the presence of horseradish peroxidase, H2O2 oxidizes Amplex Red (AR) to form highly fluorescent resorufin (kex = 563 nm, kem = 587 nm), which was quantified using a fluorometer (Ocean Optics). EDTA was also used to bind Cu(II) thereby preventing Cu(II) from potentially interfering with the method, although no such interference has been previously reported. The magnitude of HO production was determined using the chemiluminescence-based phthalhydrazide method [42,43] details of which are given in the SI. Briefly, HO hydroxlates non-chemiluminescent phthalhydrazide (Phth) to 5-hydroxy-2,3-dihydro-1,4phthalazinedione (5-HO-Phth), which emits chemiluminescence (CL) when oxidized under alkaline conditions. CL measurements were performed using a Waterville Analytical FeLume system. Radiolabelled formic acid (H14COONa) concentrations were measured by adding 1 mL of the sample solution to 5 mL of the liquid scintillation fluid (Perkin Elmer) for analysis in a Packard TriCarb 2100 TR liquid scintillation analyzer. Various concentrations of KBr and TBA (as compounds with well-known HO reactivity) were also added to the sample solutions in order to examine their impact on the degradation rate of radiolabelled formate in the experimental system. Details of the experimental procedure are given in the SI.

2.1. Chemicals All solutions were prepared using 18 MX cm ultrapure Milli-Q water (Millipore). Analytical grade chemicals were purchased from Sigma–Aldrich (or as otherwise stated) and used without further refinement. All glassware was soaked in 5% w/v HCl for at least 1 week before use. Stock solutions were kept in dark bottles, covered in foil and refrigerated at 4 °C when not in use. All studies were performed at a controlled room temperature of 22 ± 0.6 °C. Experiments were conducted in darkness with the reactor covered in foil for the duration of the reaction. All solutions were prepared in 0.7 M NaCl buffered by 2.0 mM NaHCO3 (or as otherwise stated), with the pH adjusted to 8.0 with 1 M HCl or NaOH. All pH measurements were made using a Hanna HI9025 pH meter combined with a glass electrode and Ag/AgCl reference. The pH electrode was calibrated on the NBS scale using NIST-traceable buffer solutions (pH 4.01, 7.01 and 10.01). Cu(I) stock solutions (0.5 mM) were prepared daily according to the method described previously [5]. Cu(II) stock solutions (1 mM) were prepared weekly by dissolving an appropriate amount of CuCl2 in a 10 mM HCl solution. A range of H2O2 stock solutions

2.3. Kinetic modeling Kinetic models were fitted to the experimental data over a range of experimental conditions using the program Kintek Explorer [44]. In addition to the data collected in the current work, experimental data (measured under conditions similar to those used here) from Yuan et al. [5] (on the oxidation of Cu(I)) and Pham et al. [40] (on the reduction of Cu(II) by H2O2) were also utilized. Principal component analysis was employed following the procedure of Vajda et al. [45] using Kintecus [46] to generate Normalized Sensitivity Coefficients (NSCs) in order to eliminate a number of unimportant reactions in the system. The sensitivity of the model to changes in individual rate constant values was also assessed by examining the change in the relative difference between the experimental data and the kinetic model (defined as the relative residual, r) when one rate constant was varied while holding the others fixed at their optimal values (Kintecus was used to evaluate the kinetic model controlled using a VBA program to calculate the residual; see SI for details).

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3. Results and discussion 3.1. Cu(I) oxidation by H2O2 The decay of Cu(I), particularly in the presence of high H2O2 concentrations, was initially rapid, as shown for the oxidation of 0.4 lM Cu(I) by H2O2 in the absence and presence of O2 in Fig. 1A and B, respectively. As the concentration of Cu(II) increased with time however, the rate of Cu(I) decay was retarded due to the back reduction of Cu(II) by H2O2, which regenerates Cu(I) and lowers the apparent oxidation rate. A steady-state concentration of Cu(I) of 0.23 lM was observed in the absence of O2 (Fig. 1A). The rate of Cu(I) removal was initially faster in the presence of oxygen at low H2O2 concentrations but was similar to that in deoxygenated solutions in the presence of 45 and 90 lM H2O2. Thus, at these high concentrations of H2O2, oxygen is not a significant oxidant of Cu(I), with H2O2-mediated oxidation dominant. Assuming pseudo-first-order reaction kinetics at high H2O2 concentrations (e.g., 5–90 lM) for the first 3 min, a second-order rate constant for this initial rapid process of 79 ± 6 M1 s1 is estimated. This value is similar to that reported by Moffett and Zika [29] in seawater of 100 M1 s1 but smaller than that reported by Sharma and Millero [35] of 2.2  102 M1 s1 in 0.7 M NaCl. It should be noted that while EDTA (in this study) and ethylenediamine (in Moffett and Zika [29]) were mixed with Cu(I) solutions immediately before entering the spectrophotometric cell, in the Sharma and Millero’s [35] study, EDTA was added directly into the reaction mixture to prevent Cu(II) back reaction. Addition of EDTA to the reaction vessel will lead to rapid removal of Cu(I) by not only preventing the Cu(II) back reaction with H2O2 but also by formation of a Cu(I)-EDTA complex, which will react with oxygen and H2O2 at much faster rates than inorganic Cu(I) species [5]. Masarwa et al. [27] extrapolated the data obtained in the presence

(A) [Cu(I)], µM

0.4 0.3 0.2 0.1

1 µM H2 O2

9 µM H2 O2

5 µM H2 O2

45 µM H2 O2

0.0 5

10

15

3.2. Phthalhydrazide hydroxylation and formate oxidation in the presence and absence of compounds with well-known HO reactivity Results of 5-HO-Phth generation by 0.2 lM Cu(II) in the presence of varied concentrations of H2O2 and Phth are presented in Fig. 2. The rate of 5-HO-Phth formation appeared to be linear during the course of experiments and was estimated to be 7.2 pM s1 in the presence of 0.1 mM H2O2 and 1.0 mM Phth. Increasing the concentrations of Phth and/or H2O2 resulted in an increase in the rate of 5-HO-Phth formation (Fig. 2). It should be noted that if HO was predominantly responsible for the hydroxylation of phthalhydrazide (and produced the CL product 5-HO-Phth), then the production rate of HO should be 5  [5-HOPhth]  36 pM s1 in this instance (as 5-HO-Phth is formed in 20% yield when HO reacts with Phth). Results for the degradation of H14COO in the presence of vari ous concentrations of HCO 3 , TBA or Br , all of which were able to compete to varying degrees with H14COO for the reactive intermediate produced in this system, are given in Figs. 3 and 4. In the presence of 20 mM NaHCO3 (Fig. 3), 10 mM TBA or 1 mM Br– (Fig. 4), negligible degradation of H14COO– was observed. From   these studies, the relative reactivity of HCO 3 , TBA, Br , and HCOO with the oxidant degrading HCOO could be assessed via the kinetic model, with these relative reactivities, as discussed in Section 3.4, inconsistent with the well-known reactivity of these compounds with HO, providing evidence against significant involvement of HO in this system. 3.3. Reduction of Cu(II) by H2O2

90 µM H2 O2

0

of phenanthroline (phen) to [phen] = 0 and obtained a rate constant <500 M1 s1 at pH 5.8. Small values of standard error of the calculated rate constant over the range of 5–90 lM H2O2 also suggested that pseudo-first-order reaction kinetics is likely during peroxidation of Cu(I), at least initially, when H2O2 is in considerable excess. An additional estimate of the second-order rate constant of Cu(I) peroxidation can also be deduced from the similarity between the rate of decrease in Cu(I) concentration due to oxygenation and that due to peroxidation of 9 lM H2O2 in the deoxygenated solution (Fig. S2); that is, kCuðIÞþO2 ½O2 ½CuðIÞ  kCuðIÞþH2 O2 ½H2 O2 ½CuðIÞ. Given that kCuðIÞþO2 ½O2  ¼ 6:57  104 s1 at pH 8.0, 2 mM NaHCO3 and 0.7 M NaCl [5], at [H2O2] = 9 lM, kCuðIÞþH2 O2 ¼ 73 M1 s1 which is consistent with the estimated value above.

20

min

As can be seen in Fig. 5, Cu(I) is produced on reaction of Cu(II) with H2O2 with the rate and extent of Cu(I) formation increasing

0.1 mM H 2 O2 and [5-HO-Phth], µM

[Cu(I)], µM

(B) 0.4 No H2O2

0.3

0.2

0

5 µM H2O2

45 µM H2O2

9 µM H2O2

90 µM H2O2

5

10 min

15

20

Fig. 1. Kinetics of 0.4 lM Cu(I) removal in 0.7 M NaCl and 2.0 mM NaHCO3 at pH 8.0: (A) deoxygenated conditions and (B) oxygenated conditions ([O2]0 = 0.243 mM). Error bars are standard errors from triplicate measurements and lines represent model fits.

5 mM Phth 1 mM Phth 0.55 mM Phth 0.1 mM Phth

0.10

0.5 mM H 2 O2 0.55 mM Phth

0.05

0.00 0

20

40

60

time, min Fig. 2. Formation of 5-HO-Phth in solutions of 0.7 M NaCl, 2 mM NaHCO3 at pH 8.0 in the presence of 0.2 lM Cu(II) with different concentrations of H2O2 and Phth. Error bars are standard errors from triplicate measurements and lines represent model fits.

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0.2 µM Cu(II) 20 mM HCO3 -

0.15

0.2 µM Cu(II) 2 mM HCO3 -

0.10 0.05 0.00

(A)

0.4 µM Cu(II) 2 mM HCO3 0

8 µM H2 O2

0.1

0.0

30

60

90

4 µM H2 O2

0.2

[Cu(I)], µM

[Formate], µM

0.20

120

0

5

10

[Formate], µM

0.10

150

60 100

40

0

83 µM TBA 0.33 mM TBA

0.05

50

0

30

60

90

120

0.20 0.15 0.10 40 µM KBr

0.05

0.2 mM KBr 1 mM KBr

0

30

0

20

40

60

0

Time, min

2.5 mM TBA

time, min

[Formate], µM

80

20

0.15

0.00

200

[H2O2], nM

0.20

10 mM TBA

(B)

20

(B) 100 [Cu(I)], nM

Fig. 3. Degradation of 0.2 lM formate in solutions of pH 8.0, 0.7 M NaCl with 4 mM H2O2 and in the presence of different concentrations of Cu(II) and bicarbonate. Error bars are standard errors from triplicate measurements and lines represent model fits.

0.00

15

Time, min

Time, min

(A)

1 µM H2 O2 2 µM H2 O2

0.4 µM H2 O2

60

90

120

time, min Fig. 4. Degradation of 0.2 lM formate in solutions of pH 8.0, 0.7 M NaCl, 2.0 mM NaHCO3 with 0.4 lM Cu(II), 4 mM H2O2, and in the presence of various concentrations of (A) TBA and (B) KBr. Error bars are standard errors from triplicate measurements and lines represent model fits.

with an increase in initial H2O2 concentration. In almost all cases, particularly at high H2O2 concentrations, the Cu(I) concentration is observed to plateau after a few minutes of reaction presumably as a result of the reaction of Cu(I) with H2O2. As expected, loss of H2O2 is observed as the reaction proceeds (Fig. 5B), although the extent of loss appears to be somewhat less than the generation of Cu(I) suggesting some reformation of H2O2 through the course of the reaction. 3.4. Modeling the kinetics of Cu(I)/Cu(II) reactions with H2O2 and O2

3.4.1. General modeling approach In a relatively simple system where oxygen is the dominant oxidant, the oxidation of Cu(I) can be described by considering the

Fig. 5. (A) Formation of Cu(I) in oxygenated solutions of 0.4 lM Cu(II) and various concentrations of H2O2. (B) Formation of Cu(I) (left Y-axis) and removal of H2O2 (right Y-axis) in the presence of 0.1 lM H2O2 (circle symbols) and 0.2 lM H2O2 (square symbols) in oxygenated solutions of 0.4 lM Cu(II). Error bars are standard errors from triplicate measurements and lines represent model fits. Solution matrix: 0.7 M NaCl, 2.0 mM NaHCO3 at pH 8.0.

speciation of Cu(I) and recognizing that each Cu(I) species will react with O2 at a different rate, with this approach successfully employed by Yuan et al. [5]. While such an approach was able to be applied to a simple inorganic system where oxygenation of particular Cu(I) species was known to be the dominant process occurring [5], speciation-based kinetic modeling cannot be applied to the more complex system described here where Cu(I) oxygenation, Cu(I) peroxidation, Cu(II) reduction by H2O2, and reactions involv ing O 2 and HO all occur, as this would require estimation of species-specific reaction rates for each Cu(I) and/or Cu(II) species with  O2, H2O2, O 2 and HO . Such a system is difficult to solve uniquely due to the large number of rate constants that would need to be determined. Instead, an approach analogous to the ‘‘FeL approach’’ used by Rose and Waite [47] for description of iron redox transformations is used here where all cuprous copper species are denoted by ‘‘Cu(I)’’ and all cupric copper species are denoted by ‘‘Cu(II)’’. This model (the ‘‘CuL’’ model) is subsequently used to examine: (a) the kinetics of Cu(I) oxygenation such that the rate constant of reaction of Cu(I) and O2 could be deduced, (b) the kinetics of the reaction of Cu(II) with H2O2, (c) the kinetics of Cu(I) peroxidation, and (d) the potential involvement of HO by quantification of the rates of degradation of formate and hydroxylation of Phth in the presence of various HO scavengers. 3.4.2. Oxygenation of Cu(I) In the absence of H2O2, the initial removal of Cu(I) was principally governed by its reactions with O2 and O 2 (reactions 1 and 2, Table 1). That this is the case is clear from two lines of argument. Firstly, as the oxidation of Cu(I) by O2 is slow and rate limiting in 0.7 M NaCl solutions, only low (sub-micromolar) concentrations of H2O2 are expected to form during the course of reaction. Since the kinetics of H2O2 reactions with both Cu(II) and Cu(I) are rela-

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Table 1 Major reactions and corresponding rate constants for reaction of Cu(I) and Cu(II) with H2O2 and O2 in the absence and presence of HO scavengers.a No.

Reaction

Rate constant (M1 s1)

Ref.

Primary copper reactions 1 2

CuðIÞ þ O2 ! CuðIIÞ þ O 2

1.50 ± 0.03 (1.98 ± 0.05)  109

[40]b [48]

61 ± 1

[40]b

3 4 5 6

2Hþ CuðIÞ þ O 2 ! CuðIIÞ þ H2 O2 2OH CuðIÞ þ H2 O2 ! CuðIIIÞ  CuðIIÞ þ O2 ! CuðIÞ þ O2

8

2Hþ

CuðIIÞ þ H2 O2 ! CuðIÞ þ O 2 Cu(III) + Cu(I) ? 2Cu(II)

Reactions of Cu(III) with major scavengers 2OHþ 7 CuðIIIÞ þ H2 O2 ! CuðIIÞ þ O 2 Hþ 8  CuðIIIÞ þ HCO 3 ! CuðIIÞ þ CO3 þO2 ;OH 9 CuðIIIÞ þ Phth ! CuðIIÞ þ 0:019 5-OH-Phth þ O 2 þO2 ;Hþ 10 CuðIIIÞ þ HCOO ! CuðIIÞ þ CO2 þ O 2 11 Cu(III) + TBA ? Cu(II) + TBA 12 Cu(III) + Br ? Cu(II) + Br Ancillary reactions 13 14 15 16

þ HCOO

!

HCO 3

þ CO2 þ

O 2

2Hþ

Br þ H2 O2 ! Br þ O 2 þ  þO2 ;H

Br þ HCOO

!

3.5  109

[55]

(4 ± 2)  10

Hþ

 þO2

[48] [40]b

<1.3  105

  CO 3 þ H2 O2 ! HCO3 þ O2

CO 3

(6.6 ± 0.7)  10 460 ± 9

b



5

b



(2.2 ± 0.9)  106

bc

(7 ± 2)  106

b

(1.0 ± 0.4)  107 (4 ± 2)  107

b

4.3  105

[60]

1.1  105

[69]

4.0  109

[70]

8

[71]

,– – – –

b

4.6  10

Br þ CO2 þ O 2

a Cu(I), Cu(II), and Cu(III) represent all inorganic copper in their oxidation state +1, +2, and +3 respectively. Disproportionation of Cu(III) [55] with k = 1.0  108 M1 s1 was found to be unimportant under the experimental conditions investigated here. b Fitting parameters. c 0.019 is the specific yield coefficient of the CL 5-HO-Phth for Cu(III) based on sensitivity analysis (Fig. S8).

tively slow (as shown above for Cu(I) and shown previously by Pham et al. [40] for Cu(II)), reaction of Cu(I) with O2 will undoubtedly be the dominant Cu(I) removal pathway (i.e., k1[Cu(I)][O2]  k3[Cu(I)][H2O2], k5[Cu(II)][H2O2]). Secondly, as the rate 8 1 1 constant for reaction of O s ) is 2 with Cu(II) (k4 = 6.6  10 M only one third of that for reaction of O with Cu(I) (k = 2.0  109 2 2 M1 s1) [48], the back reduction of Cu(II) by O 2 will only become significant when more than 50% of the Cu(I) present has been oxidized to Cu(II). Thus, regardless of the mechanism by which H2O2 may react with Cu(I) and Cu(II), peroxidation of Cu(I) and back reduction of Cu(II) by both O 2 and H2O2 will not be important in the removal of Cu(I) in oxygenated solutions in the absence of added H2O2, particularly in the early stage of the oxidation process. Given that the values of k2 and k4 have been previously determined [48], only the rate constant for reaction 1 (k1) is required to model the oxygenation of Cu(I) (note that the disproportionation of O 2 is slow at pH = 8.0 [49] and therefore does not play any significant role in the kinetic model). As can be seen from Fig. 1B and Fig. S3, the model provides a satisfactory fit to the Cu(I) oxidation results for both 0.2 and 0.4 lM Cu(I) in 2.0 mM NaHCO3 and 0.7 M NaCl at pH 8.0. The value of k1 = 1.50 ± 0.03 M1 s1 deduced from model fitting is well constrained under the experimental conditions investigated here (see SI for sensitivity analysis). This ‘‘intrinsic’’ rate constant is lower than the overall Cu(I) oxidation rate constant deduced by Yuan et al. [5] (kapp = 2.6 ± 0.4 M1 s1) because of the significant contribution of O 2 to the oxidation of Cu(I) which is also accounted for in the apparent rate constant (reaction 2 in Table 1).

developed to fit the data of both Cu(I) oxygenation and Cu(II) reduction by H2O2. According to Pham et al. [40], while the values of k1 and k5 (Table 1) were well constrained from the model fits, only an upper bound for the value of k(Cu(I) + H2O2 ? Cu(II) + HO) was determined (k < 100 M1 s1). This is consistent with the apparent rate constant for reaction of Cu(I) with H2O2 of kCu(I) + 1 1 s determined in this work. At low H2O2 conH2O2  79 ± 6 M centrations (e.g., 0.4–8 lM in Pham et al. [40]), oxygenation of Cu(I) by O2 ([O2] = 0.243 mM [50], k1 = 1.50 ± 0.03 M1 s1) outcompetes peroxidation of Cu(I) (i.e., k1[Cu(I)][O2]  kCu(I) + H2O2[Cu(I)][H2O2]). Thus, the current study is consistent with the earlier conclusion that peroxidation of Cu(I) was insignificant in the study of Pham et al. [40] While the overall rate constant of Cu(II) reduction by H2O2 (to form Cu(I)) was well constrained [40], the detailed mechanism of Cu(II) reaction with H2O2 requires some clarification with two alternate pathways proposed for this reaction: the ‘‘complex’’ pathway and the ‘‘free radical’’ pathway. It has previously been shown that Cu(II) complexes are able to catalytically decompose H2O2 only if the metal coordination sphere is not saturated by ligands [34,36]. This result suggests that ternary Cu(II)–ligand-peroxo complexes are involved in the reaction. Sigel and co-workers [36] and later Lekchiri et al. [51] proposed the non-radical ‘‘complex’’ mechanism (represented by Eqs. (1)–(4) below) in which a Cu(II)-complex containing two molecules of H2O2 (e.g., CuIIL(OOH)(H2O2)) was postulated as the active species, with the reaction taking place completely within the coordination sphere of the metal ion and with the reaction shown in Eq. (4) being rate determining.

3.4.3. Reaction of Cu(II) with H2O2 Pham and co-workers [40] have recently used a ‘‘CuL’’ modeling approach to describe the reduction of Cu(II) by the naturally occurring organic matter, Suwannee River fulvic acid. Considering that H2O2 may have been a reactive species in their system, the kinetics of Cu(II) reaction with H2O2 was examined and a kinetic model was

H2 O2 Hþ þ HOO II



ð1Þ

II

Cu L þ HOO Cu LðOOHÞ II

ð2Þ

II

Cu LðOOHÞ þ H2 O2 Cu LðOOHÞðH2 O2 Þ II

II

ð3Þ 

Cu LðOOHÞðH2 O2 Þ Cu L þ O2 þ H2 O þ OH

ð4Þ

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Similarly, based on the experimental observation that the substrate oxidation rate (of quinaldine blue) is zero order with respect to its concentration in studies of the reaction between various Cu(II) complexes and H2O2 (at pH 9.1), Robbins and Drago [34] proposed that peroxide (or hydroperoxide anion, HOO) coordinates with the metal without changing its oxidation state resulting in production of the peroxo or hydrogen peroxo complex. Once formed, this copper–hydroperoxide species (CuII–OOH+), the major active oxidant in the system, reacts extremely rapidly with the substrates. These authors also concluded that other possibilities including the formation of HO by a Fenton-like reaction and a mechanism that involves a high oxidation state of copper as the active oxidant are less reasonable. A slight modification of this mechanism involving formation of a highly reactive Cu(I) intermediate was proposed by Otto et al. [31] and Schounas et al. [37] (Eqs. (5), (6)):

CuII LðOOHÞ ! CuI Lð OOHÞ

ð5Þ

CuI Lð OOHÞ þ H2 O2 ! CuII L þ O2 þ H2 O þ OH

ð6Þ

In comparison with the non-radical ‘‘complex’’ pathway, Gray [19] suggested that the reaction between Cu(II) and H2O2 produced Cu(I) and O 2 (Eqs. (7)–(10)), This ‘‘free radical’’ pathway has been promoted by a number of investigators both in the absence [13,26,28,29,33] and presence of organic ligands [14,17,18,30,32,52] with the production of Cu(I) subsequently resulting in formation of HO via a ‘‘Fenton-like’’ reaction between Cu(I) and H2O2. This mechanism and the presence of HO have been further supported by ESR [32], deoxyribose degradation [21], and dye decolorization [30,52].

H2 O2 Hþ þ HOO

ð7Þ

CuðIIÞ þ HOO ! CuðIÞ þ HO2

ð8Þ

HO2

þ

H þ

O 2

CuðIÞ þ H2 O2 ! CuðIIÞ þ HO þ OH

ð9Þ ð10Þ

Perez-Benito [33] recently proposed the rapid formation of CuIIOOH+ prior to the slow dissociation of this complex (unimolecular II + decomposition) to form Cu(I) and O 2 . However, as Cu OOH did not participate subsequently in any other important reactions, this mechanism is essentially similar to the ‘‘free radical’’ mechanism presented above. Of these options, the ‘‘free radical’’ pathway appears most consistent with the obtained experimental data in this study since this pathway (but not the ‘‘complex’’ pathway) reliably predicts the rate and extent of formation of Cu(I). As such, we adopted the ‘‘free radical’’ mechanism to describe the formation of Cu(I) in our Cu(II)/ H2O2 system. It was further assumed that the formation of a Cu(II)– H2O2 complex was either minor or would decompose rapidly to Cu(I) and O 2 under the experimental conditions investigated here. In addition, since pK a ðHO2 Þ ¼ 4:8  0:1 [53] and pKa(H2O2) = 11.6 [19] at pH 8.0, most superoxide was present in its deprotonated form (i.e., O 2 ) and hydrogen peroxide remained in its protonated form (i.e., H2O2). Reactions presented in Eqs. (7)–(9) therefore can be simplified to reaction 5 (Table 1) with apparent rate constant k5. 3.4.4. Reaction of Cu(I) with H2O2 It has often been assumed that the reaction between Cu(I) and H2O2 generates HO in a Fenton-like reaction (Eq. (10)). Since HO is a highly reactive species, this reaction is of major importance owing to its role in many biochemical processes and copper toxicity [20,21]. There is ample evidence in support of this reaction in the literature from ESR spin trapping measurements [20], reactions with known HO scavengers [21] and by identification of the prod-

59

ucts which are formed under the Cu(I)/H2O2 system compared to that in a well-defined HO generating system [15]. Other studies of the reactions of H2O2 with both Cu(I) [23,27] and its complex CuðphenÞþ 2 [22] however indicated that, under some conditions, HO may not be formed and that higher oxidation states of copper (i.e., Cu(III)) may be the true active oxidant (Eq. (11)). þ2Hþ

CuðIÞ þ H2 O2 ! CuðIIIÞ þ 2H2 O

ð11Þ

Although there is no direct evidence for the formation of Cu(III) from the reaction of Cu(I) with H2O2, production of Cu(III) by addition of HO to Cu(II) has been demonstrated by pulse radiolysis [54]. The same species could possibly be formed if an additional electron transfer step between the newly formed Cu(II) and the undissipated HO is assumed to follow an initial electron transfer from the metal–peroxo complex CuI–H2O2 (to CuII–HO) within the metal coordination sphere. The activity of Cu(III), however, is strongly pH dependent and, under some circumstances, is over an order of magnitude less reactive than HO [54]. In addition, while HO is formed during the decomposition reaction of Cu(III) in acidic solutions [54] (Eq. (12)), bimolecular decomposition of Cu(III) to form Cu(II) and H2O2 also occurs with the rate of this reaction reported to decrease with increasing pH [55] (Eq. (13)).

CuðIIIÞ CuðIIÞ þ HO þ Hþ

ð12Þ

2CuðOHÞ3 ! 2CuðIIÞ þ 4OH þ H2 O2

ð13Þ

Masarwa et al. [27] proposed that a transient complex between Cu(I) and H2O2 of the type Cu(I)–OOH may be formed and that this complex might decompose to Cu(II) + HO or Cu(III) + H2O or react directly with a substrate. Shah et al. [52] suggested the possibility of formation of different types of HO radicals in their studies of Cu(II)/succinic acid/ H2O2 systems. From the ESR spectra of spin adducts of 5,5-dimethyl-1-pyrroline-N-oxide (DMPO), they noted a slight difference in the magnetic field in the ESR spectra of the DMPO/HO adduct and suggested that HO may not diffuse into solution but, rather, may bind with the CuIIL complex (bound-HO). In addition, since HO can only travel a short distance from its site of production to its site of reaction [56], the relative efficacies of free and boundHO has also been used to explain the ineffectiveness of CuII(bpy)2 to the degradation of DNA compared to that of CuII(phen)2 under the same experimental conditions [18,57]. Ali et al. [58] showed that metal-catalyzed oxidation with Cu(II) under different oxidation conditions had a larger impact than Cu(I). This suggests that the conversion of Cu(II) to Cu(I) could result in extra oxidative steps in addition to the ‘‘Fenton-like’’ reaction of Cu(I) with H2O2. If HO is assumed to be exclusively produced from the reaction presented in Eq. (10), then the rate of HO generation is given by:

d½HO  ¼ k½CuðIÞ½H2 O2  dt

ð14Þ

For the case where reactant concentrations of 0.2 lM Cu(II) and 0.1 mM H2O2 were used, a steady-state Cu(I) concentration of 90 nM was achieved after several minutes (Fig. S4). Provided that [H2O2]0  [Cu(I)] and assuming k  79 M1 s1 (see above), Eq. (14) becomes:

d½HO  ¼ k½CuðIÞ½H2 O2   79  ð90  109 Þ  ð0:1  103 Þ dt ¼ 7:1  1010 M s1

ð15Þ

On the other hand, since HO reacts rapidly with Phth [59] with k = 5.3  109 M1 s1, in the presence of 1 mM Phth, almost all HO would be consumed by Phth. As such the actual rate of production of HO in the system should be equal to or smaller than that measured from the CL of 5-HO-Phth (assuming that a small yield of 5-

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HO-Phth could also be produced from the reaction of Cu(III), if present, with Phth [15]). The fact that only 7.2 pM s1 (Fig. 2) of 5-HO-Phth has been quantified which would require a HO production rate of 36 pM s1 (if HO was the reactive intermediate responsible for its formation) compared to the theoretical rate of 710 pM s1 deduced above suggests that HO was not an immediate product from the reaction of Cu(I) with H2O2. In addition, as the concentration of 5-HO-Phth also increased with increasing [Phth] (Fig. 2), HO was unlikely the sole Phth hydroxylating agent since Phth should be the dominant sink for HO regardless of whether it is present at 1 mM or 5 mM (i.e., the 5-HO-Phth production should be similar at both concentrations). Further to these arguments, the observation that the kinetic model derived assuming HO formation, although well able to describe the oxidation and reduction rates of copper, is unable to correctly model the magnitude of formate loss in the presence of  various concentrations of HCO 3 , TBA or Br , provides clear evidence that HO is not formed in this system. Indeed, the oxidant that is formed exhibits quite different relative reactivity between H2O2 and HCOO than HO with the kinetic model in the absence of scavenger predicting a much greater extent of HCOO oxidation than that observed experimentally (see Figs. S5 and S6 in the SI). It should be noted that, any reaction of Cu(I) and Cu(II) with the free radicals formed in the reactions of the four reducing substrates with hydroxyl radical and with CO 2 is considered to be insignificant in this study. For example, since Br (a product of Br with HO) reacts rapidly with H2O2 (k = 4  109 M1 s1 [60]) and that H2O2 is present at a high concentration of 4  103 M, even if Cu(I) or Cu(II) reacted at diffusion-controlled rates, their low concentration (2  107 M) would preclude them from being a significant sink for Br. Similarly, the reaction of O2 with CO 2 (the initial product from reaction of HCOO with HO) has a rate constant of 2  109 M1 s1 [60] and is considered to be the dominant sink for CO 2 under our experimental conditions as [O2]  [Cu]. Thus, as HO formation is not consistent with the results of this study, we propose the involvement of an alternate oxidizing intermediate, possibly Cu(III), that is capable of hydroxylating Phth to yield 5-HO-Phth, albeit with a different yield of the CL 5-hydroxy derivative than for HO. It is proposed that this alternative oxidant also reacts with various substrates present in the system (reactions 7–12, Table 1) with rate constants that are much lower than those for corresponding reactions involving HO. 3.5. Model fitting and sensitivity analysis 3.5.1. Role of HO in the system where Cu(III) is proposed to be an oxidizing intermediate While there is no doubt that HO cannot be the sole oxidizing intermediate resulting from the reaction of Cu(I) with H2O2 under the experimental conditions investigated here, slow dissociation of Cu(III) to HO might suggest that both Cu(III) and HO could act as intermediate oxidants. (It has previously been shown that the reaction of HO with Cu2+ to yield CuIIIOH2+ is an equilibrium process (i.e., Cu(OH)2+ also rapidly dissociates to form HO and Cu2+ (Eq. (16))), and that Cu(III) is also subject to hydrolysis, with the Cu(OH)3 form dominant at pH 8.0 [54,55]). This possibility is examined below by deduction of the upper limit of the rate constant for the dissociation of Cu(III) to HO at pH 8.0, determined by consideration of the previously determined hydrolysis constants for Cu(III) [55].

Cu2þ þ HO

kf ¼3108

kb ¼8103

CuðOHÞ2þ

log K ¼ 4:57

ð16Þ

The rate law equation for the production of HO by Cu(III) is given by:

d½HO  ¼ kb ½CuðOHÞ2þ   kf ½Cu2þ ½HO  < kb ½CuðOHÞ2þ  dt

ð17Þ

Additionally, Ulanski and von Sonntag [55] report the following acid–base equilibria for other Cu(III) species:

CuðOHÞ2þ þ H2 O CuðOHÞþ2 þ Hþ CuðOHÞþ2

þ H2 O CuðOHÞ3 þ H

þ

log K ¼ 3:7 log K ¼ 4:6

ð18Þ ð19Þ

Combining Eqs. (18) and (19) gives

CuðOHÞ2þ þ 2H2 O CuðOHÞ3 þ 2Hþ

log K ¼ 8:3

½CuðOHÞ2þ  ¼ 108:3  ½CuðOHÞ3 ½Hþ 2

ð20Þ ð21Þ

At pH = 8.0, the dominant Cu(III) species is Cu(OH)3 (i.e., [Cu(OH)3]  [Cu(III)], Fig. S7). Substituting Eq. (21) into Eq. (17) gives:

d½HO  < kb ½CuðOHÞ2þ  ¼ 108:3  kb  ½CuðOHÞ3 ½Hþ 2 dt d½HO  < 108:3  kb  ½CuðOHÞ3 ½Hþ 2  1:6  104 ½CuðIIIÞ dt

ð22Þ ð23Þ

Thus, the rate of HO formation as a result of Cu(III) dissociation is relatively slow at pH 8.0, with an upper limit of 1.6  104 s1. Since the major pathway for Cu(III) removal is via its reaction with Cu(I) (k6 = 3.5  109 M1 s1, Table 1), the slow rate of HO formation as a result of Cu(III) dissociation would suggest that HO would be present at an extremely low concentration and will not be an important oxidant in the present system, despite the fact that HO reacts rapidly and indiscriminately with various substrates. Meyerstein [54] also concluded that at neutral pH, Cu(III) would decompose into Cu(II) and H2O2 while in acid solutions, the mechanism involves the formation of HO. 3.5.2. Model fitting Principal component analysis undertaken on the full reaction scheme (with a total of 191 reactions, see SI) in a manner similar to that of Miller et al. [42] resulted in the subset of 16 important reactions listed in Table 1. Results of model fitting using the reaction scheme presented in Table 1 in which Cu(III) is assumed to be a sole oxidizing intermediate are shown for (i) the oxygenation and peroxidation of 0.4 lM Cu(I) (in the absence and presence of O2) in Fig. 1A and B; (ii) the formation of 5-HO-Phth in the presence of 0.2 lM Cu(II) with different concentrations of Phth and H2O2 in Fig. 2; (iii) the degradation of H14COO in the presence of various concentration of HCO 3, Br and TBA in Figs. 3 and 4; and iv) the generation of Cu(I) and removal of H2O2 at different concentrations of Cu(II) and H2O2 in Fig. 5A and B. Second-order rate constants for the various reactions that have been used in the modeling are given in Table 1. In general, except for the reactions of Cu(III) with various substrates, most of the rate constants presented in Table 1 are either well documented, or fairly constrained and in agreement with previous studies [5,40]. While the kinetics of reactions of Cu(III) with  H2O2, HCO 3 , Phth, Br , and TBA have not been previously determined at neutral pH, the rate constants of these reactions were reported to be much lower than the rate constants of corresponding reactions involving HO at lower pH. Meyerstein [54] reported values of k7 = (6 ± 3)  105 M1 s1 at pH 5.2 (cf. <1.3  105 M1 s1 here), k12 < 1.0  106 M1 s1 at pH 5.4 (cf. (4 ± 2)  107 M1 s1 here) and indicated that these rate constants were pH dependent with higher rate constants at lower pH. In a separate study, Johnson et al. [24] estimated the upper limit for the rate constant of reaction of Cu(III) with CH3OH of 4  105 M1 s1 at pH 5.2. The deduced rate constants for reactions of Cu(III) with different substrates in the kinetic model presented in Table 1 at pH 8.0 are somewhat higher than expected. While it is not possible to fully explain this discrepancy, differences in pH, electrolytes, and initial

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concentration of Cu(II) in the various studies reported may all contribute to this incongruity. For example, high concentrations of Cu(II) used in previous studies may not only result in the precipitation of Cu(II) solids but also facilitate the disproportionation of Cu(I) (to form Cu(0) and Cu(II) [4]). Similarly, different electrolytes used in experimental media may significantly affect the reactivity of various copper species. According to the reaction scheme, the rapid kinetics of the reaction between Cu(I) and Cu(III) plays a significant role in the redox cycle of copper and provides the major removal pathway of Cu(III) in the absence of substrates. Both   CO 3 and Br also contribute to the rate of degradation of HCOO when carbonate and bromide anions are present. In general, the model provides satisfactory fits to the experimental data over a range of Cu(I), Cu(II), H2O2, and substrate concentrations. The goodness of the model fits to the experimental data supports both the reaction scheme proposed in Table 1 and the modeling approach used. 3.5.3. Sensitivity analysis The deduced rate constants for primary reactions between Cu(I) with O2 and H2O2 and Cu(II) with H2O2 (reactions 1, 3, and 5 respectively, Table 1) are all well constrained, as evidenced by the relatively deep and narrow minima observed in the plot of the relative residual values (Eq. S3) against a range of rate constant value (Fig. S1A). Similarly, but to a lesser extent, the rate constants –  for the scavenging of Cu(III) by HCO 3 , Phth, HCOO , TBA, and Br (Fig. S1B and C) are also found to be reasonably constrained. The deduced rate constant for reaction of Cu(III) with H2O2, however, only represents the upper bound value (Fig. S1B). Variation of this rate constant only slightly affects the goodness of fit which suggests that the major pathway for H2O2 removal is through its reactions with Cu(I) and Cu(II). 3.5.4. Limitations of the kinetic model Despite the capability of the kinetic model in providing a relatively good description of the various reactions between Cu(I), Cu(II), O2 and H2O2 over a range of experimental conditions with a relatively well-constrained set of rate constants, the success of the model is still largely dependent on the assumption that Cu(III) is the principal oxidizing species. While the production of Cu(III) by addition of HO to Cu(II) has been previously demonstrated by pulse radiolysis [54,55], none of the previous studies have convincingly validated the hypothesis that Cu(III) is present and formed from the reaction of Cu(I) with H2O2. Although HO is unlikely to be a sole oxidant in the experimental conditions under investigation, the presence of oxidizing species other than Cu(III) cannot be ruled out. One possibility is that a transient complex in the form of Cu(I)–OOH– is formed by association of Cu(I) with H2O2 with subsequent decomposition of this into Cu(II) and HO or Cu(III) or direct reaction with a substrate[27]. The fate of this transient species will thus depend on the relative rates of these reaction pathways but, unfortunately, none of these pathways is well understood. Although the deduced apparent rates of reaction of Cu(I) and Cu(II) with H2O2 and O2 will not vary, modification of the reaction between Cu(I) and H2O2 would significantly affect the proposed mechanism and resulting rate constants. Another limit to the model is the assumption that the composite terms Cu(I), Cu(II) and Cu(III) adequately represent all inorganic copper in their +1,+2, and +3 oxidation states, respectively. While the model, in general, provides a reasonably good description of the experimental data over a range of concentrations and the presence of various substrates, care must be taken when interpreting the model results as the model, by default, is condition-specific. For example, the presence of chloride would influence the reactivity of various copper species and as such the composite ‘‘intrinsic’’

61

rate constants of various reactions with copper would need to be re-evaluated. 3.6. Catalytic nature of the copper–H2O2 reaction In the presence of excess H2O2, the two redox states of copper reach steady-state concentrations as copper is both reduced and oxidized by H2O2; this steady state is achieved within several minutes when 0.4 lM of Cu(II) reacts with micromolar concentrations of H2O2 (Fig. 1). In essence, copper is acting catalytically to degrade H2O2, with formation of the reactive intermediate Cu(III) during this process. During this catalytic process, copper is being dynamically cycled between its +1 and +2 oxidation states, with the cycling rate, RC, able to be estimated using the kinetic model. The transformation between the +1 and +2 oxidation states is complicated by the intermediacy of the +3 state, however, reduction occurs directly and through only two pathways, with H2O2 and O 2 acting as the reductants. As such, if reactions 4 and 5 in the kinetic model shown in Table 1 are altered to reactions 24 and 25 as below, the amount of Cu(II) reduced to Cu(I) can be quantified by observation of the dummy species ‘‘CuRed’’

CuðIIÞ þ O 2 ! CuðIÞ þ O2 þ CuRed 2Hþ

CuðIIÞ þ H2 O2 ! CuðIÞ þ O 2 þ CuRed

ð24Þ ð25Þ

Since the concentrations of Cu(I) and Cu(II) reach steady state, the overall rate of oxidation must be equivalent to the overall rate of reduction, which can be quantified by determining the rate of production of CuRed, denoted P(CuRed). The cycling rate can then be determined by dividing this production rate by the steady-state concentration of Cu(II), that is, RC = P(CuRed)/[Cu(II)]ss. The cycling rate was determined to be 2.2 s1 at the early stages of reaction in the formate oxidation experiments discussed in Section 3.2, with RC gradually declining during the experiment as H2O2 is consumed; although copper attains a pseudo-steady state, it is still a very dynamic system. Also of interest is the production rate of Cu(III), which is the highly reactive intermediate formed in this system. It can also be quantified by using a dummy species in the kinetic model, in this case ‘‘Cu(III)TOT’’, by modifying reaction 3 from Table 1 to be

CuðIÞ þ H2 O2 ! CuðIIIÞ þ CuðIIIÞTOT

ð26Þ

with the production rate of Cu(III) at any given time then given by the time-derivative of Cu(III)TOT. In most natural systems, including the intercellular compartments within organisms and in seawaters, lakes, rivers, and even in oxic subsurface porewaters, H2O2 is constantly produced and consumed, typically via the production of superoxide, reaching constant concentrations over short timescales. Because redox active elements such as iron and copper have been implicated in the oxidation of organic matter resulting, for example, in the transformation of riverine organic matter to both more and less refractory components [61] and biomolecules such as dopamine to highly refractory neuromelanin (the presence of which is associated strongly with the onset of Parkinson’s Disease) [62], it is of interest to examine the potential for cycling in these systems, as the formation of the oxidizing intermediate during H2O2-mediated oxidation, proposed here to be some form of Cu(III) species, is a likely initiator of organic molecule transformation. The potential importance of oxidant generation via H2O2-mediated cycling of Cu(I) and Cu(II) can be assessed with the kinetic model by examining how the cycling rate and Cu(III) production rate change as a function of copper concentration while fixing the H2O2 concentration. Total copper concentrations in surface waters are variable though average values of around 1 nM have been reported in seawaters and 40–70 nM in rivers, lakes, and groundwaters while the

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healthy brain of an adult human contains 8.8 mg of copper [63]. While all of this copper is unlikely to be available for reaction, most is expected to be present in the form of organic complexes with a significant proportion of this complexed Cu expected to be able to participate in electron transfer processes. Concentrations of H2O2 are less well known with concentrations of 10 pM to 10 lM reported in natural waters [64] and concentrations on the order of 1–10 lM measured in the brains of Gerbils [65]. The kinetic model has been run and allowed to reach a pseudo-equilibrium in the presence of a fixed concentration of H2O2 over a range of total copper concentrations (initially present as Cu(II)). The model was amended to include superoxide disproportionation which, although insignificant in the conditions used to fit the kinetic model, may compete with the Cu=O 2 reactions at low copper concentrations. Although the modeling work is only illustrative, as in any system there will likely exist additional sinks for both O 2 and/or H2O2, interesting insights can still be made. When a fixed H2O2 concentration was used, it can be seen that the Cu cycling rate is predominantly a function of H2O2 concentra-

tion; however, the production rate of the Cu(III) oxidant (P(Cu(III))) shows a more complex dependence, increasing as both the H2O2 and total copper concentration increase (Fig. 6). Of perhaps even more relevance to real-world situations are simulations we have undertaken assuming a constant production  rate of O 2 (denoted P(O2 )). Abiotic superoxide production rates 1 of up to 1.4 pM s were measured by Rose et al. [66] in upwelling waters in the Costa Rica Dome region of the equatorial eastern Pacific Ocean while Micinski et al. [67] measured light-induced superoxide production rates as high as 0.1 nM s1 in Eastern Caribbean surface waters. In comparison, Garg et al. [68] observed production rates on the order of 3 nM s1 in laboratory cultures (10,000 cells/mL) of the superoxide-producing marine raphidophyte Chattonella marina. At low concentrations of copper and low O 2 production rates, the system is very slow to reach steady state, with simulation lengths up to 104 days needed to reach pseu13 do-equilibrium at PðO M s1 and [CuTOT] = 1010 M, 2 Þ ¼ 10 whereas pseudo-equilibrium was reached within 15 min at 5 PðO M s1 and [CuTOT] = 104 M. As a result, a much high2 Þ ¼ 10

Fig. 6. Variability of (A) the copper cycling rate, RCu and (B) the production rate of Cu(III), P(Cu(III)), as a function of both the total copper concentration and a fixed concentration of H2O2. Results were obtained by running the kinetic model until steady-state conditions were achieved.

Fig. 7. Variability of (A) the steady-state H2O2 concentration, (B) copper cycling rate, RCu, and (C) the production rate of Cu(III), P(Cu(III)), as a function of both the total copper  concentration and fixed, constant production rate of O 2 , P(O2 ). Results were obtained by running the kinetic model until steady-state conditions were achieved.

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63

Fig. 8. Degradation of 4 mM H2O2 by (A) 0.2 lM, (B) 2 lM, and (C) 20 lM Cu(II) in 0.7 M NaCl, 2 mM NaHCO3 at pH 8.0. Solid symbols are experimental data, with the error estimated by the standard deviation of the H2O2 concentration determined at 240, 250, and 260 nm from a single experiment (in panel (A), a replicate is shown with open symbols). In all panels, the predicted decay from the kinetic model in Table 1 is shown as a solid line.

er minimum concentration of copper was considered in this modeling work (1010 M cf. 1018 M used in the fixed H2O2 work). Although such durations are clearly not relevant to any natural process, it does serve to highlight important system behavior. It is seen that the steady-state H2O2 concentration and the Cu cycling rate exhibit similar trends with both obtaining their maxima at low Cu concentration and high PðO 2 ) (Fig. 7A and B). This is as expected, as the copper cycling rate is limited by the production rate of O 2 relative to the concentration of copper. Likewise, higher production of O 2 leads to more H2O2; however, higher concentrations of copper increase the rate of H2O2 destruction resulting in lower steady-state concentrations. The more interesting trend is that observed for P(Cu(III)), which is seen to be almost independent of the concentration of copper, with P(Cu(III))  0.25  P(O 2 ), regardless of the copper concentration (Fig. 7C). In this case, copper is shown to exhibit high catalytic activity in converting O 2 , a mildly reactive species, into the more reactive Cu(III) species. The conditions described in the modeling studies are difficult to examine experimentally. The catalytic nature of the Cu–H2O2 system has however been verified by examining the ability of Cu to degrade an excess of H2O2 (4 mM H2O2 with 0.2, 2, and 20 lM Cu(II)). In this work, H2O2 concentrations were monitored using its absorbance at 240, 250 and 260 nm (e = 38.3, 22.7 and 13.0 M1 cm1, respectively) [38], with the interference due to absorbance of copper species in this wavelength range corrected by subtraction of the final spectrum obtained after H2O2 degraded to such an extent that its absorbance was negligible, with this correction only being significant at 20 lM Cu. Regardless of the small uncertainty introduced due to copper absorption, the timescale of H2O2 removal still clearly demonstrates that Cu is able to catalytically degrade H2O2, even when present at a concentration 2  104 higher than Cu. As can be seen in Fig. 8, although H2O2 is always completely degraded, the rate at which H2O2 degrades is dependent upon the total copper concentration and is well predicted by the kinetic model in Table 1 (note that these data were not considered in the kinetic model fitting process).

the detailed mechanism nor the nature of the oxidizing species generated has been convincingly elucidated, particularly at circumneutral pH. In this study, we have revisited the kinetics of reactions of nanomolar concentrations of Cu(I) and Cu(II) with H2O2 in the absence and presence of O2 in 2 mM NaHCO3, 0.7 M NaCl and at pH 8.0. The reaction between Cu(II) and H2O2, under the experimental conditions investigated here, was found to occur via a ‘‘free radical’’ mechanism in which Cu(II) was utilized through a oneelectron redox process that converted H2O2 into the intermediate 1 1 reactive radical species O s . 2 with a rate constant of 460 ± 9 M Measurements of both hydroxylated phthalhydrazide CL product and the degradation of radiolabelled formate in the absence and presence of known HO scavengers indicated that the reaction between Cu(I) and H2O2 did not result in the production of HO but involved the formation of a higher oxidation state of copper, Cu(III), with a rate constant of 61 ± 1 M1 s1. This intermediate oxidizing species reacts with the various substrates that were present at rate constants several orders of magnitude smaller than those of reactions of HO with the same substrates. Formation of HO from the dissociation of Cu(III) was determined to be extremely slow at pH 8.0 with an upper bound to the rate constant of 1.6  104 s1; as such, HO is not an important oxidant in this system. A kinetic model using composite terms Cu(I), Cu(II), and Cu(III) (which represent all inorganic copper species in their oxidation states +1, +2, and +3, respectively) has been developed and shown to satisfactorily describe the kinetics of the Cu(I)/Cu(II)/H2O2/O2 system over a range of conditions and in the presence of various substrates. The success of the approach supports the possibility of using a single entity for each oxidation state of copper and helps to minimize the number of unknown kinetic constants. This provides a framework for modeling the dynamics of more complicated Cu systems (such as that in the presence of organic ligands) and subsequently assists in predicting the likely extent of oxidative stress and copper toxicity in natural waters and environments of significance to human health.

4. Conclusions and implications The redox transformations of Cu(I) and Cu(II) species play a critical role both in the speciation, transport, and bioavailability of Cu in natural waters and in the toxicity of copper, particularly, as a result of the harmful oxidants generated within cells as a result of these transformations. In the presence of partially reduced oxygen species such as H2O2 and O 2 , which are by-products of oxygen metabolism in cells, redox cycling of Cu may result in the generation of highly reactive and damaging intermediate(s). While numerous previous studies have focused on such systems, neither

Acknowledgment This work was funded by the Australia Research Council Discovery Grant Scheme (DP0987188).

Appendix A. Supplementary material Supplementary data associated with this article can be found, in the online version, at http://dx.doi.org/10.1016/j.jcat.2013.01.025.

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