Article pubs.acs.org/est
Effects of pH, Chloride, and Bicarbonate on Cu(I) Oxidation Kinetics at Circumneutral pH Xiu Yuan,† A. Ninh Pham,† Guowei Xing,† Andrew L. Rose,‡ and T. David Waite*,† †
School of Civil and Environmental Engineering, The University of New South Wales, Sydney, New South Wales 2052, Australia Southern Cross GeoScience, Southern Cross University, Lismore, New South Wales 2480, Australia
‡
S Supporting Information *
ABSTRACT: The oxidation kinetics of nanomolar concentrations of Cu(I) in NaCl solutions have been investigated over the pH range 6.5−8.0. The overall apparent oxidation rate constant was strongly affected by chloride, moderately by bicarbonate, and to a lesser extent by pH. In the absence of bicarbonate, an equilibriumbased speciation model indicated that Cu+ and CuClOH− were the most kinetically reactive species, while the contribution of other Cu(I) species to the overall oxidation rate was minor. A kinetic model based on recognized key redox reactions for these two species further indicated that oxidation of Cu(I) by oxygen and superoxide were important reactions at all pH values and chloride concentrations considered, but back reduction of Cu(II) by superoxide only became important at relatively low chloride concentrations. Bicarbonate concentrations from 2 to 5 mM substantially accelerated Cu(I) oxidation. Kinetic analysis over a range of bicarbonate concentrations revealed that this was due to formation of CuCO3−, which reacts relatively rapidly with oxygen, and not due to inhibition of the back reduction of Cu(II) by formation of Cu(II)−carbonate complexes. We conclude that the simultaneous oxygenation of Cu+, CuClOH−, and CuCO3− is the rate-limiting step in the overall oxidation of Cu(I) under these conditions.
1. INTRODUCTION Copper is an essential trace element that is vital to the viability of almost all organisms.1 In natural waters, copper typically occurs in either the cuprous (Cu(I)) or the cupric (Cu(II)) oxidation state. Copper redox chemistry in the upper water column plays a significant role in its speciation, transport, and bioavailability.2 Most previous studies3−5 have focused primarily on Cu(II), principally because Cu(I) is easily oxidized to Cu(II) by oxygen or other oxidants. However, a growing body of evidence6−11 indicates that a number of potentially important reactions, many of which are photochemically induced, may lead to Cu(I) formation and result in a significant steadystate concentration of Cu(I) in natural waters.2,12 Copper is also considered to be an important scavenger of superoxide (O2−)13,14 as it can be subjected to reduction and oxidation by O2− in natural waters.15 Zafiriou et al.16 indicated that up to 25% of inorganically complexed copper might be present as Cu(I) in sunlit seawaters provided the steady-state concentration of O2− exceeds 10−12 M. On the other hand, redox reactions of Cu(I) could result in production of reactive oxygen species (ROS), such as O2− and OH•, that may subsequently induce a cascade of radical-promoted reactions with other constituents in natural waters.17−19 As such, a better understanding of copper-catalyzed production and consumption of O2− is important in furthering insight into factors contributing to global biogeochemical cycles. © 2011 American Chemical Society
Oxygenation of micromolar and nanomolar concentrations of Cu(I) in natural seawater and synthetic NaCl solutions has been previously investigated.6,8,10,11 In general, the overall oxidation rate of Cu(I) can be described by
d[Cu(I)] = kapp[O2 ][Cu(I)] (1) dt where [Cu(I)] is the total or analytical Cu(I) concentration and kapp (in M−1·s−1) is the apparent oxidation rate constant. In oxygen-saturated systems (where [O2] ≫ [Cu(I)]), the initial decay of Cu(I) was observed to follow pseudo-first-order kinetics with a pseudo-first-order rate constant k′ = kapp[O2]0. Despite differing experimental conditions, chloride concentration has been consistently and significantly observed to influence kapp. However, large differences in calculated secondorder rate constants for individual Cu(I) species between studies suggests considerable uncertainty with regard to the actual species contributing to the overall oxidation process. For example, González-Dávila et al.6 found a significant discrepancy between the reported rate constants at micromolar and nanomolar concentrations of Cu(I) which could not be attributed entirely to the differences in the experimental conditions. More −
Received: Revised: Accepted: Published: 1527
September 27, 2011 December 16, 2011 December 20, 2011 December 20, 2011 dx.doi.org/10.1021/es203394k | Environ. Sci. Technol. 2012, 46, 1527−1535
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for 30 min with zero-grade air that had been bubbled through 1 M NaOH to remove trace CO2 present. The pH was then quickly adjusted to specific values by adding appropriate amounts of concentrated NaOH, and the solution was subsequently continuously sparged throughout the course of each experiment. Calibration curves were developed at each pH examined by measuring the absorbance at 484 nm of 50, 100, 200, 400, and 600 nM Cu(I) added to buffer solutions containing 0.05 mM BC and 0.25 mM EDTA. The method detection limit was ∼0.5 nM Cu(I). 2.3. Speciation Modeling. Cu(I) speciation was calculated with the program MINEQL+23 using the equilibrium reactions and stability constants shown in Table 1. Extrapolation of
importantly, to the best of our knowledge, none of the previous studies6,10,11,20 has convincingly explained the variation of rate constants with varying pH and bicarbonate concentrations. The weak pH dependence of the oxidation rate constant of Cu(I) has been attributed to formation of mixed copper hydroxo− chloro complexes and back reduction of Cu(II) to Cu(I).8 The effect of bicarbonate, on the other hand, has been attributed to retardation of the back reduction of Cu(II) to Cu(I)8 and/or formation of Cu(I) carbonate complexes.11,20 As such, a clear need exists to resolve the apparent knowledge gaps in our understanding of factors controlling the kinetics of inorganic Cu(I) oxidation. In this study, oxidation of nanomolar Cu(I) was investigated over the pH range 6.5−8.0 in solutions containing different concentrations of NaCl and bicarbonate to answer two particular questions: (a) which copper species are responsible for the pH, chloride, and bicarbonate dependence of the Cu(I) oxidation rate constant and (b) how important is the back reduction of Cu(II) to the kinetics of Cu(I) oxidation? To rationalize the chemical mechanisms behind the observed results and develop a more comprehensive understanding of Cu(I) oxidation in natural waters, two different modeling approaches were adopted: (a) an equilibrium speciation modeling approach, involving assessment of the contribution of various Cu(I) species to the overall oxidation and identification of the species responsible for the pH and chloride and bicarbonate concentration dependence of Cu(I) oxidation and (b) a kinetic modeling approach where the importance of various reaction pathways was examined.
Table 1. Stability Constants for Cu(I) Speciesa no.
species
logK at 25 °C, I = 0
ref
1 2 3 4 5 6 7 8 9 10 11
H+ + OH− = H2O H+ + CO32− = HCO3− 2H+ + CO32− = H2CO3 Cu+ + Cl− = CuCl(aq) Cu+ + 2Cl− = CuCl2− Cu+ + 3Cl− = CuCl32− Cu+ + Cl− + H2O = CuClOH− + H+ Cu+ + 2Cl− + H2O = CuCl2OH2− + H+ Cu+ + Cl− + 2H2O = CuCl(OH)22− + 2H+ Cu+ + H+ + CO32− = CuHCO3 Cu+ + CO32− = CuCO3−
14.0 10.30 16.70 3.10 5.68 5.02 −5.64 −7.91 −19.84 7.0. However, in the
Figure 4. Evaluation of the deduced rate constants for oxygenation of Cu(I) species in the context of Marcus Theory. The solid line corresponds to eq 12 with λ = 84 kJ·mol−1.
presence of environmentally relevant bicarbonate concentrations, formation of CuCO3− substantially increased Cu(I) oxidation rates. Although CuCl2− and CuCl32− were the dominant Cu(I) species in this pH range in both the absence and the presence of bicarbonate, their contribution to the overall rate constant was found to be negligible. Therefore, under 1533
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environmentally relevant conditions at circumneutral pH, oxidation of Cu(I) is expected to be controlled primarily by the relative proportions of Cu+, CuClOH−, and CuCO3−, which will in turn be governed largely by chloride and bicarbonate concentrations in natural waters of interest. The kinetic model developed adequately described the kinetics of Cu(I) oxidation over the range of conditions investigated in this study. In this model, oxidation of cuprous species by O2 was the most important pathway in the overall oxidation of Cu(I). While the oxidation of Cu(I) by O2− was also important under all conditions considered, back reduction of Cu(II) by O2− only became important at lower [Cl−]. Determination of values for apparent and intrinsic rate constants for oxidation of the critical Cu(I) species Cu+, CuClOH−, and CuCO3− by O2 over the pH range 6.5−8.0, whose accuracy is supported by Marcus theory, should greatly assist in understanding and predicting inorganic Cu(I) and Cu(II) transformations in natural waters. Application of the relationship determined by Marcus theory should enable further extension of these findings to systems containing other inorganic copper complexes whose stability constants and/or redox potentials are known.
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ASSOCIATED CONTENT
S Supporting Information *
Additional method details, additional justification for the kinetic model, and one supplementary table and five supplementary figures with additional supporting data. This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Corresponding Author
*Phone: +61 2 9385 5059; fax: +61 2 9313 8341; e-mail: d.waite@ unsw.edu.au.
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ACKNOWLEDGMENTS We gratefully acknowledge Christopher Miller’s assistance with the model calculations. This work was funded by the Australian Research Council Discovery Grant Scheme (DP0987188).
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