Kinetics of Homogeneous and Surface-Catalyzed Mercury(II

Article pubs.acs.org/est

Kinetics of Homogeneous and Surface-Catalyzed Mercury(II) Reduction by Iron(II) Aria Amirbahman,*,† Douglas B. Kent,‡ Gary P. Curtis,‡ and Mark C. Marvin-DiPasquale‡ †

Department of Civil and Environmental Engineering, University of Maine, 5711 Boardman Hall, Orono, Maine 04469, United States U.S. Geological Survey, Water Resources Division, 345 Middlefield Road, Menlo Park, California 94025, United States

‡

S Supporting Information *

ABSTRACT: Production of elemental mercury, Hg(0), via Hg(II) reduction is an important pathway that should be considered when studying Hg fate in environment. We conducted a kinetic study of abiotic homogeneous and surface-catalyzed Hg(0) production by Fe(II) under dark anoxic conditions. Hg(0) production rate, from initial 50 pM Hg(II) concentration, increased with increasing pH (5.5−8.1) and aqueous Fe(II) concentration (0.1−1 mM). The homogeneous rate was best described by the expression, rhom = khom [FeOH+] [Hg(OH)2]; khom = 7.19 × 10+3 L (mol min)−1. Compared to the homogeneous case, goethite (α-FeOOH) and hematite (α-Fe2O3) increased and γ-alumina (γ-Al2O3) decreased the Hg(0) production rate. Heterogeneous Hg(0) production rates were well described by a model incorporating equilibrium Fe(II) adsorption, rate-limited Hg(II) reduction by dissolved and adsorbed Fe(II), and rate-limited Hg(II) adsorption. Equilibrium Fe(II) adsorption was described using a surface complexation model calibrated with previously published experimental data. The Hg(0) production rate was well described by the expression rhet = khet [>SOFe(II)] [Hg(OH)2], where >SOFe(II) is the total adsorbed Fe(II) concentration; khet values were 5.36 × 10+3, 4.69 × 10+3, and 1.08 × 10+2 L (mol min)−1 for hematite, goethite, and γalumina, respectively. Hg(0) production coupled to reduction by Fe(II) may be an important process to consider in ecosystem Hg studies.

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INTRODUCTION

Direct Hg(II) reduction by Fe(II) may be important in Ferich freshwater environments with sufficiently low organic matter concentrations in which aqueous Hg(II) species dominate. This includes environments such as groundwater and ferrallitic soils, where Fe(III)- and Al-oxyhydroxides, clays, and other aluminosilicates control the Hg content in the soil profile.21 In at least one study, the presence of Hg(0) in a contaminated groundwater and aquifer material has been attributed to Fe(II).22 Mercury occurrence in groundwater has received relatively little attention, even though groundwater can act as a potentially important source of Hg to lakes and other water bodies.23−25 In a series of studies of an unconfined quartz sand aquifer in coastal New Jersey, USA, Barringer and co-workers observed groundwater total Hg concentrations in excess of the U.S. EPA’s maximum contaminant limit of 2 μg L−1 in ∼600 domestic wells.26−28 It was proposed that reductive dissolution of Fe(III) oxyhydroxides associated with the aquifer material brought about by septic-system effluent led to Hg mobilization in this aquifer system. In a granitic aquifer in coastal Maine, USA, 19% of shallow wells contained total Hg concentrations

Mercury (Hg) is a global contaminant that poses great health risks to humans and wildlife. The biogeochemical Hg cycle involves pathways into water bodies from the atmosphere and soil, with nonpoint sources constituting the majority of Hg loading in most environments.1 Hg speciation, which determines its reactivity and bioavailability, is dominated by organically or inorganically complexed Hg(II) species, Hg(0), and mono- and dimethyl Hg. Even though Hg enters most aquatic systems predominantly as inorganic Hg(II), organic species are the most bioavailable.2 The methylated Hg production rate depends in large part on the concentration and speciation of Hg(II).3−5 Mercury(II) can also be reduced to the volatile Hg(0), resulting in the decrease of the Hg(II) concentration and the consequent decrease of the bioavailable organic Hg concentration. The presence of Hg(0) has been documented in different aquatic systems.6−9 The dark reduction of Hg(II) to Hg(0) can be microbially catalyzed,10,11 or abiotic, involving Fe(II)12−14 or dissolved organic matter (DOM).15−17 The DOM-catalyzed Hg(II) reduction likely takes place by the DOM’s redox-active hydroquinone−quinone couple.18 At higher concentrations, it has been argued that DOM can stabilize Hg(II) by possibly stabilizing the HgS nanoparticles.19,20 © 2013 American Chemical Society

Received: Revised: Accepted: Published: 7204

April 4, 2013 June 2, 2013 June 3, 2013 June 3, 2013 dx.doi.org/10.1021/es401459p | Environ. Sci. Technol. 2013, 47, 7204−7213

Environmental Science & Technology

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>2 μg L−1.29 The occurrence of high Hg concentrations in this system results from the reducing environment in local groundwater rather than a point source. In a recent study of a 17-m-long sediment core from Lake Tulane, Florida, USA, a drastic increase in Hg flux following the last deglaciation was attributed to the rise in groundwater brought about by global sea-level rise.30 It was argued that flooding the oxidized soils surrounding the lake resulted in the reductive dissolution of Fe(III) oxyhydroxides and mobilization of the adsorbed Hg. The objective of this paper is to develop a kinetic model for Hg(II) interactions with Fe(II) in an environment that represents that of mineral soil water and groundwater. In particular, we have developed a coupled equilibrium Fe(II) surface complexation model and kinetic rate law to quantify the abiotic homogeneous and heterogeneous Hg(II) reduction by Fe(II). The model considers aqueous speciation, adsorption and desorption, and oxidation−reduction reactions.

of the desired value during all experiments. After a 30-min equilibration, aliquots of 0.05 M FeSO4 were added and the suspensions were mixed for ∼18 h with a Teflon-coated magnetic stirbar, while being purged with N2 gas at a low flow rate to achieve adsorptive equilibrium between Fe(II) and the mineral oxyhydroxides. The N2 gas was scrubbed of residual O2 by bubbling into a suspension of Fe(III)/Fe(II) inside the glovebag33 and delivered in the suspension through a FEP tube. To avoid chemical reduction in the glovebag, Fe(III) oxyhydroxide stock solutions were stored outside. Speciation calculations showed that in all experiments initial Fe(II) concentrations were below the solubility limit of Fe(OH)2(S). This reactor configuration yielded reproducible experimental results for goethite suspensions as long as the FEP and other Teflon components of the reactor were allowed to outgas in the glovebag for at least 24 h. Reduction experiments were initiated by adding an aliquot of a 50 nM HgCl2 stock solution stored at 4 °C in 2% trace metal grade HNO 3 to the mixed suspension for an initial concentration of 50 pM Hg(II). The total Cl− concentration in the experiments was ∼10−5 M. Oxygen-free N2 was continuously bubbled into the suspension at a flow rate of 250 mL min−1 and the evolved gaseous Hg(0) was collected on an in-line gold-coated sand trap. At specific times, the gas flow was shut down and the trap was quickly removed and replaced by a new trap; gas flow resumed after placement of the new trap. The Hg(0) on the trap was thermally desorbed in an Ar stream and detected on a Tekran 2500 fluorescence detector. The detection limit was 0.04 pmol (∼1 pM in our system) based on 3× one standard deviation for the peak areas of blank gold traps. A five-point calibration was performed routinely with Hg(II) concentrations spanning from 1 to 50 pM. Modeling. Sets of equilibrium and kinetic reactions were used to describe the interaction of Hg(II) with Fe(II) in solution and on mineral surfaces used here under different geochemical conditions. The reactions were modeled using the computer program PHREEQC that performs chemical speciation, surface-complexation, and kinetic calculations.34 The nature of the reactions (equilibrium or kinetic), the corresponding equilibrium or rate constants, total concentrations of the components, and the time intervals at which concentrations are calculated were specified, as required by the program. For aqueous species, PHREEQC uses the equilibrium constants from the database files of MINTEQA2 (Supporting Information Table S1). We obtained or derived the equilibrium constants for H+ and Fe(II) surface coordination and total concentrations of surface sites from the existing data in the literature. The rate equations included Hg(II) adsorption, desorption, and homogeneous and surface-catalyzed reduction reactions. A MATLAB routine was developed to fit the rate equations to the experimental data by minimizing a weighted least-squares objective function with respect to the parameter values. The latter consisted of rate constants for Hg(II) adsorption and desorption, as well as the homogeneous and surface-catalyzed reduction. The MATLAB routine estimated the rate constants for multiple experimental conditions by running the PHREEQC files sequentially for each set of conditions. Each of the PHREEQC files contained the set of equations and model parameters for Hg(0) production at different Fe(II) and mineral concentrations, and pH values. The simulated Hg(0) concentrations were compared to the observed concentrations and the rate constants were adjusted to achieve minimum error.

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MATERIALS AND METHODS Chemicals. Goethite was synthesized based on the method of Atkinson et al.31 A Fe(NO3)3 solution was adjusted to pH 12 with CO2-free NaOH and aged at ∼50 °C for approximately 3 days in a high-density polyethylene container. Hematite was synthesized as described in Bargar et al.32 Briefly, a Fe(ClO4)3 + HClO4 solution was filtered (0.22 μm) and added to a 0.016 N HClO4 solution at 100 °C and kept in a preheated ∼100 °C oven for approximately 5 days in a borosilicate container. Both Fe(III) oxyhydroxides were freeze-dried after a thorough dialysis against deionized water (Milli-Q). Pyrogenic γ-Al2O3 (γ-alumina) was used as received from the manufacturer (Degussa). The identity and crystallinity of the Fe(III) oxyhydroxide minerals were confirmed with X-ray powder diffraction. None of the most intense peaks of any other Fe(III) oxyhydroxide minerals was detected in either the goethite or hematite samples. Specific surface area measurements were conducted by the N2−BET technique (Micromeritics Tristar 2000) after outgassing at 110 °C. The samples had no detectable microporosity. The surface area measurement results are reported in Table 1. Table 1. Properties of Minerals mineral α-FeOOH (goethite) α-Fe2O3 (hematite) γ-Al2O3 (γ-alumina)

specific surface area (m2 g−1)a

site density (nm−2)b

45

1.68

43 97

2.07 2.3 (weak), 0.079 (strong)

This work. bFrom refs 47, 48, and 52, for goethite, hematite, and γalumina, respectively.

a

Experimental Setup. Mercury(II) reduction experiments were conducted at 20−23 °C in 125-mL fluorinated ethylene propylene (FEP) bottles fitted with inflow and outflow ports that had been acid-washed and treated with 5% BrCl overnight. The 50-mL suspensions were prepared in an anaerobic glovebag (Coy Laboratory Products, Grass Lake, MI) with H2−N2 atmosphere by adding aliquots of mineral oxyhydroxides to a solution of 10−3 M MOPS buffer and 10−1 M NaNO3. Either 1 M HNO3 or NaOH was added to adjust the pH to the desired value. The pH remained within 0.1 unit 7205

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The model also performed a local sensitivity analysis for each parameter, which allows the user to observe the significance of each parameter with respect to the simulated values and to identify the most sensitive parameters. To obtain equilibrium H+ and Fe(II) surface coordination constants for the Fe(III) minerals, we fit the existing data using computer program FITEQL.35 As input, FITEQL uses known equilibrium constants and total or free concentrations of chemical components and species, including those of adsorption sites.3

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Hg(0) production rate increased with increasing pH and Fe(II) concentration. The reaction exhibited an overall second-order behavior and the pH trend exhibited a first-order dependence on the concentration of the first Fe(II) hydrolysis species, FeOH+ (Table S1), rhom = khom[FeOH+][Hg(OH)2 ]

(1)

where rhom is the homogeneous Hg(0) production rate and khom is the observed second-order homogeneous reduction rate constant (M−1 min−1). The best fit to the data was provided by khom = 7.19 × 10+3 M−1 min−1 (Table 1). In this study, the Fe2+ species is assumed to be nonreactive,37 and the Fe(OH)20 concentration is negligible over the pH range of our experiments and, therefore, was not included. Even though other hydrolysis species of Hg(II), such as Hg(OH)+ and Hg(OH)3−, may possess different reactivity with respect to Fe(II), a rate expression involving any one aqueous species other than Hg(OH)2 could not describe the observed pH trend, as the concentrations of these species are very low in the pH range studied here.38 A similar pH dependence with respect to dissolved Fe(II) speciation has been previously observed for reduction of O2,39 Cr(VI),40 and U(VI),41 and in each case was attributed to enhanced reactivity of FeOH+ compared to Fe2+ species. The enhanced reactivity of FeOH+, despite its presence at significantly lower concentrations than Fe2+ in these experiments, is due to its higher metal basicity and lower reduction potential; E0 for Fe3+/Fe2+ = 0.77 V and for FeOH2+/FeOH+ = 0.34 V.42 This lower reduction potential is brought about by the OH− ligand that donates electron density to Fe(II) through σand π-systems.43 The homogeneous reaction mechanism is described in detail in the Supporting Information (SI). Reduction of Hg(II) to Hg(0) is a two-step electron transfer process, with the first electron transfer step from FeOH+ resulting in inorganic Hg(I) production as the rate-determining step. The formation of Hg(I) intermediate as Hg22+ is thermodynamically favored over direct reduction to Hg(0) because E0 for the Hg2+/Hg22+ couple (0.91 V) is greater than that for the Hg2+/Hg(0)(l) couple (0.85 V).44 Reduction of Hg(I) to Hg(0) (E0 = 0.80 V)44 may then take place via a one-electron transfer reaction from another FeOH+ species, resulting in oxidation of two moles of Fe(II) for every mole of Hg(0) produced. In this work, the presence of Fe(II) in an extreme excess compared to Hg(II) precludes the determination of the correct reaction stoichiometry with respect to Fe(II). Alternatively, in water, disproportionation of inorganic Hg(I) may result in Hg(0) production.6,44 In either case, the firstorder rate expression with respect to FeOH+ and Hg(OH)2 concentrations (eq 1) implies that (a) back reactions for electron transfer and disproportionation are relatively slow and (b) Hg(I) reaches a steady-state concentration (see SI). Steadystate concentration of an intermediate complex or species, such as Hg(I), may be assumed when it accumulates at a lower level compared to the reactants and the products.45 Anions, especially OH−, tend to complex more strongly with Hg(II) than inorganic Hg(I), and, as such, Hg22+ is only marginally stable against disproportionation and may be assumed to exist at low concentrations.44 Iron(III) is a product of the reduction reaction, and its precipitation can potentially lead to enhanced Hg(0) production via autocatalysis, which is brought about by Fe(II) coordination with the surface of the newly formed Fe(III)

RESULTS AND DISCUSSION

Homogeneous Reduction. Figure 1a−c shows homogeneous Hg(0) production at pH values between 6.3 and 8.1 and total Fe(II) concentrations between 0.1 and 1.0 mM. The

Figure 1. Homogenous Hg(0) production kinetics for initial [Hg(II)] = 50 pM and total [Fe(II)] = (a) 0.1 mM, (b) 0.5 mM, and (c) 1 mM. The number next to each curve is the pH. The lines are model fit to the data based on the rate expressions and constants given in Table 2. Aqueous speciation equilibrium reactions are given in Table S1. RMSE = 0.276. 7206



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Table 2. Rate and Surface Equilibrium Constantsa k or log K reaction stoichiometry and rate expression homogeneous Hg(II) reduction (kinetic) rhom = khom [FeOH+] [Hg(OH)2]b surface H+ adsorption (equilibrium)c >SOH + H+ ⇔ >SOH2+ >SOH ⇔ >SO− + H+ surface Fe(II) adsorption (equilibrium)c >SOH + Fe2+ ⇔ >SOFe+ + H+ >SOH + Fe2+ + H2O ⇔ >SOFe(II)OH + 2H+ Hg(II) adsorption (kinetic) >SOH + Hg(OH)2 → >SOHg+ + H2O+ OH− rads = kads [>SOH] [Hg(OH)2]b Hg(II) desorption (kinetic) >SOHg+ + 2H2O → >SOH + Hg(OH)2 + H+ rdes = kdes [Hgads]b surface-catalyzed Hg(II) reduction (kinetic) rhet = khet [>SOFe(II)T] [Hg(OH)2]b

homogeneous

γ-alumina

goethite

hematite

6.19 ± 0.06d −10.41 ± 0.12d

8.41 ± 0.06d −10.57 ± 0.16d

7.7e −10.2e

−1.29 ± 0.05d −9.85 ± 0.06d

−0.48 ± 0.10d −9.18 ± 0.14d

−3.70 ± 0.06e and −0.66 ± 0.06e

8.98 × 10+3

1.13 × 10+4

1.14 × 10+1

1.89 × 10−2

1.50 × 10−2

3.66 × 10−3

3.12 × 10+3

1.08 × 10+2

7.19 × 10+3

4.69 × 10+3

All constants are reported for I = 0. khom, kads, and khet are second-order rate constants (M min ); kdes is first-order rate constant (min−1). Reported as intrinsic equilibrium surface binding constants conditional on the site densities in Table 1 and the DDL Coulombic term. dRefit using data in refs 47 and 48 for goethite and hematite, respectively (see SI). Errors are two standard deviations from the mean as determined by FITEQL. e Equilibrium binding constants for weak and strong sites obtained from ref 52. a

−1

b

−1

c

to obtain these fits are shown in Table 2. In all cases, the model represents the data well. For goethite and hematite, the best fits for the pH-dependent Fe(II) adsorption data were obtained using two different Fe(II) surface species: >SOFe(II)+ and >SOFe(II)OH, where >SO represents an adsorption site. The reaction stoichiometry and intrinsic equilibrium binding constants for these species are shown in Table 2. The contribution of each of these surface species to total Fe(II) adsorption is shown in Figure S1, where Fe(II) adsorption is dominated by >SOFe(II)+ species at the lower pH and >SOFe(II)OH species at the higher pH range. Other authors obtained a reasonable fit to the experimental data using only the >SOFe(II)OH species.46 However, since we have described the surface-catalyzed Hg(II) reduction kinetics based on the total adsorbed Fe(II) and not individual surface species, surface speciation of Fe(II) does not affect the magnitude of the Hg(II) reduction rate constants. For goethite (Figure S1a), Fe(II) adsorption data from Dixit and Hering47 and H+ binding from Liger et al.48 were refit using the DDL model. In our Hg(0) production experiments with goethite, the total adsorption site concentrations ranged from 0.014 to 0.084 mM and total Fe(II) was 1 mM, resulting in a near-complete coverage of goethite surface sites in the circumneutral pH range. Dixit and Hering47 used a higher adsorption site concentration of 0.448 mM and a lower Fe(II) concentration of 0.215 mM, resulting in a partial coverage of the goethite surface (up to ∼0.45). Although the Fe(II) coverage in the data set used to calibrate the Fe(II) adsorption model were somewhat lower than those used in our study, the overall range of chemical conditions is sufficiently close to justify our using this Fe(II) adsorption model. Characterization of sorbed Fe(II) species via direct spectroscopic techniques is complicated by the fact that it requires an excess Fe(II) loading, such that precipitation of Fe(OH)2(S) will result. Proposed surface Fe(II) speciation, such as the ones used here, are based on the best fits to the Fe(II) adsorption data.49 Further, mixed-valent Fe with overlapping dorbitals, such as Fe(II) sorbed onto the surface of Fe(III) oxyhydroxide, can lead to interfacial electron transfer within the

hydroxide. However, provided the stoichiometrically insignificant concentration of Fe(III) produced here (maximum 10−11 M), Fe(III) hydroxide precipitation would be insignificant, and as such, autocatalytic Hg(II) reduction should be ruled out. Peretyazhko et al.13 conducted similar experiments to reduce 20 pM of Hg(II), albeit at a significantly lower total dissolved Fe(II) concentration of 1.5 μM, and concluded that dissolved Fe(II) does not contribute to homogeneous Hg(II) reduction. We simulated their experiments using the rate expression in eq 1 and observed that, under such experimental conditions, the total reduced Hg(II) would be close to 0.2 pM at pH 7.8, which is the detection limit of their fluorescence detector. Thus, their finding of minimal Hg(0) production by aqueous Fe(II) is consistent with the Hg(0) production rates observed in our experiments. Heterogeneous Reactions. Fe(II) Adsorption to Mineral Surfaces. To quantify the role of mineral surfaces in catalyzing Hg(II) reduction by Fe(II), the extent and pH dependence of Fe(II) adsorption onto these surfaces must be known. A quantitative knowledge of the density and reactivity of mineral surface functional groups is also necessary to characterize the extent of Hg(II) adsorption onto mineral surfaces. To this end, we used the existing data on surface H+ and Fe(II) coordination from the literature. The programming platform used in this work, PHREEQC, only supports the diffuse double layer (DDL) model for surface complexation models with a Coulombic term. Therefore, for cases where the authors had generated surface complexation models with alternative Coulombic terms to describe their surface H+ and Fe(II) coordination data, we refit their experimental data to obtain equilibrium constants for surface interactions compatible with the DDL Coulombic term. Figure S1 shows the experimentally measured and modelcalculated Fe(II) adsorption edges for the three minerals used here. For all minerals, adsorption increases with increasing pH, with no adsorption at pH < 5 and maximum adsorption at pH > 7.5. The specific surface areas and site densities in Table 1 were used in fitting the data, and the equilibrium constants used 7207



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mineral and oxidation of the sorbed Fe(II).46,50 Mössbauer spectroscopy has shown that at below Fe(II) saturation surface coverage, interfacial electron transfer into the hematite structure results in Fe(II) oxidation.51 It is only at higher than saturation surface coverage levels that a stable sorbed Fe(II) phase forms on hematite, but the nature of this species is not clear.50 In light of this, we have conducted experiments at higher than Fe(II) saturation coverage, and constructed our Hg(0) production rate model around the total sorbed Fe(II) concentration and not any specific surface Fe(II) species. For hematite, the H+ and Fe(II) adsorption data from Liger et al.48 were refit using the DDL model (Figure S1b). The best fit was obtained using the same two reactions used for goethite (Table 2). Hematite adsorption site concentrations in our Hg(0) production experiments were 0.017 and 0.035 mM and the total Fe(II) concentration was 1 mM, which, similar to goethite, resulted in a near-complete coverage of goethite surface sites in the circumneutral pH range. Liger et al.,48 however, used an adsorption site concentration of 0.2 mM and a total Fe(II) concentration of 0.160 mM, resulting in a partial surface coverage (up to ∼0.75). Despite the somewhat different Fe(II) surface coverage, similarity in the overall range of chemical conditions between the calibration data set and those used in our study justifies applying this model to describe Fe(II) adsorption in our experiments. For Fe(II) adsorption onto γ-alumina (Figure S1c), we used the original constants by Villaseñor Nano and Strathmann52 as they were developed using the DDL model (Table 2). These authors obtained the best fit to their experimental data with a single reaction stoichiometry but with adsorption sites with two different binding affinities (called “strong” and “weak” adsorption sites) (Table 2). The Fe(II) surface complexation model was calibrated at an adsorption site concentration of 1.83 mM and total Fe(II) concentrations of 0.02−0.5 mM. Our Hg(II) reduction experiments were conducted over a similar range in chemical conditions justifying use of this model to described Fe(II) adsorption in our experiments. Mercury(0) Production in the Presence of Oxide Surfaces. Time−concentration plots for surface-catalyzed Hg(0) production by Fe(II) are shown in Figures 2−4 for goethite, hematite, and γ-alumina, respectively. Similar to the homogeneous reduction experiments, Hg(0) production rate increases with pH (Figures 2, 3, 4b,c) and total Fe(II) concentration (Figure 4a). Hg(0) production rates, however, vary depending on solid type and solid/liquid ratio for the same initial Fe(II) concentration. Under the experimental conditions here, Hg(0) production rate increases with increasing solid/liquid ratio for hematite and goethite (Figures 2 and 3), but decreases with increasing γ-alumina solid:liquid ratio (Figure 4b and c) compared to the homogeneous reduction (Figure 1). Unlike the case for homogeneous experiments, the time− concentration plots for surface-catalyzed Hg(0) production for a given mineral cannot be modeled with an overall secondorder rate expression with respect to Hg(II) and adsorbed Fe(II) concentrations, and a single rate constant (not shown here). This suggests that Hg(0) production involves several steps and processes that include surface interactions in addition to the homogeneous reaction. Conceptually, these steps are (a) transport of Hg(II) to the surface, (b) two-step reduction of Hg(II) to Hg(0) following its interaction with adsorbed Fe(II), and (c) partitioning of Hg(0) into the gas phase. In addition to reduction by adsorbed Fe(II), Hg(II) adsorbs directly to the mineral surface, decreasing the Hg(II) aqueous concentration.

Figure 2. Goethite-catalyzed Hg(0) production kinetics for initial [Hg(II)] = 50 pM, total [Fe(II)] = 1 mM, and goethite total surface area of (a) 5 m2 L−1, (b) 10 m2 L−1, and (c) 30 m2 L−1. The number next to each curve is the pH. The lines are model fit to the data based on the rate expressions and constants given in Table 2. Aqueous speciation equilibrium reactions are given in Table S1. RMSE = 0.324.

In our model, steps a and c are assumed to be fast considering the suspension’s high agitation and N2 purging rates, as well as lack of detectable microporosity and, therefore, intraparticle diffusion of Hg(II) or Fe(II). These are shown to be reasonable assumptions for modeling our kinetic data. The overall model considers simultaneous homogeneous and surface-catalyzed Hg(II) reduction by Fe(II). Model presentation starts with Hg(II) adsorption at the surface that is shown to decrease the Hg(II) reduction rate, followed by surface-catalyzed Hg(II) reduction. Mercury(II) Adsorption to Mineral Surfaces. The model accounting only for Hg(II) reduction by aqueous and sorbed Fe(II) yielded poor fits to experimental data collected across the range of chemical conditions in this study. Significant adsorption of Hg(II) has been reported on Fe(III) and Al oxyhydroxides.53−57 Adsorption of Hg(II) to the metal 7208



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Figure 3. Hematite-catalyzed Hg(0) production kinetics for initial [Hg(II)] = 50 pM, total [Fe(II)] = 1 mM, and hematite total surface area of (a) 5 m2 L−1 and (b) 10 m2 L−1. The number next to each curve is the pH. The lines are model fit to the data based on the rate expressions and constants given in Table 2. Aqueous speciation equilibrium reactions are given in Table S1. RMSE = 0.509.

oxyhydroxides diminishes the aqueous Hg(II) concentration, thereby decreasing the Hg(0) production rate. Surface complexation models have been used to describe equilibrium Hg(II) interaction with goethite55 and hydrous ferric oxide.56 Both studies used total Hg(II) concentrations in the 10−5 to 10−3 M range, which is considerably higher than what is found in most natural environments and used in this study. To explain the pH adsorption edge, both studies invoked surface species >SOHg+ and >SOHgOH.55,56 Modeling Hg(II) adsorption onto hydrous ferric oxide with only the >SOHg+ surface species also provided reasonable fits under most experimental conditions.36 More recently, Kim et al.57 used surface spectroscopy to propose that Hg(II) forms inner-sphere surface compelxes that are predominantly bidentate (binuclear) with goethite, and monodentate and bidentate with Al (hydr)oxides. They also observed evidence for surface coordination of Hg(I), as Hg22+, with γ-alumina, but not with goethite. For several reasons, the existing Hg(II) equilibrium adsorption models55,56 were not used. First, Hg(II) was added at t = 0 and adsorption equilibrium onto Fe(III) and Al oxyhydroxides is not instantaneous. For example, equilibrium with respect to adsorption of Hg(II), present at μM concentrations, on goethite was reached within 18 h followed by continued slow uptake.58 Second, equilibrium models calibrated at high adsorbate loadings may underestimate the extent of adsorption at extremely low adsorbate loadings, where adsorption is dominated by low-abundance, high-affinity binding sites.59−61 Third, existing equilibrium Hg(II) models calibrated in single-adsorbate systems55,56 may not provide

Figure 4. γ-Alumina-catalyzed Hg(0) production kinetics for initial [Hg(II)] = 50 pM and (a) total [Fe(II)] = 0.1, 0.5, and 1 mM at pH = 7.2 ± 0.1 and γ-alumina SSA = 951 m2 L−1, (b) [Fe(II)] = 1 mM, γalumina total surface area =317 m2 L−1, and (c) [Fe(II)] = 1 mM and γ-alumina total surface area = 3]170 m2 L−1. The number next to each curve in b and c is the pH. The lines are model fit to the data based on the rate expressions and constants given in Table 2. Aqueous speciation equilibrium reactions are given in Table S1. RMSE = 0.813.

reliable predictions of adsorption in the presence of high adsorbed concentrations of Fe(II). The adsorption step was instead described with an overall second-order expression with respect to available surface sites, >SOH, and Hg(OH)2 concentrations, and the desorption step was described with a first-order expression with respect to adsorbed Hg rads = kads[>SOH][Hg(OH)2 ]

(2)

rdes = kdes[Hgads]

(3)

where rads and rdes are the adsorption and desorption rates, respectively, kads (M−1 min−1) and kdes (min−1) are their corresponding rate constants, and Hgads represents the total concentration of adsorbed Hg(II). Including the concentration of available adsorption sites improved the ability of the model 7209



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to fit experimental data by allowing more extensive Hg(II) adsorption to occur at higher solid/liquid ratios. Rate constants that produced the best fits to the data are shown in Table 2. Yin et al.62 also used second-order adsorption and first-order desorption rate expressions to model the adsorption kinetics of Hg(II), in the ppm range, to soils. The rate constants in Table 2 project that equilibrium with respect to Hg(II) adsorption on goethite, hematite, and γalumina would be achieved within minutes to hours, as expected from the many previous studies of metal ion adsorption on these solids (Figure S2). Rate constants were similar between goethite and hematite and considerably lower for γ-alumina (Table 2). Fits were insensitive to assumptions about Hg(II) adsorption reaction stoichiometry, so the desorption rate was expressed in terms of total adsorbed Hg(II). We have also assumed that as a metastable species, Hg(I) would be present at a negligible, but steady-state concentration (see SI). A review of literature on Hg(II) interaction with mineral surfaces shows that Fe(III) minerals are more effective sorbents than Al (hydr)oxides with respect to Hg(II)63 and other cations.64 Kim et al.57 observed a significantly higher Hg(II) adsorption density onto goethite compared to γ-alumina. This is consistent with our modeling results, where the extent of Hg(II) adsorption at equilibrium calculated by the model is considerably greater for goethite and hematite than for γalumina (Figure S2). Surface-Catalyzed Reduction. The general reaction stoichiometry for the surface-catalyzed Hg(II) reduction by Fe(II) can be written as

production is via electron transfer to Hg(I) from a surfacebound Fe(II) species or Hg(I) disproportination, khet = ((ka1kr)/ (ka−1 + kr)). Assuming that the precursor surface complex formation is the rate-determining step and ka−1 ≪ kr,12 khet = ka1. In eq 4, n = 0 or 1, representing >SOFe+ and >SOFeOH surface species, respectively. Based on the pH dependence of the reduction kinetics of Hg(II) and other trace metal and organic contaminants, and the strong correlation between reduction rates and the concentration of the >FeOFeIIOH species, Charlet and co-workers have proposed that surfacecatalyzed reduction by Fe(II) involves electron transfer from the hydroxylated surface species that can be attributed to its higher electron density.13,48,49,65 Such correspondence between Hg(0) production rates and >SOFeOH surface species concentrations (Figure S1) was also observed in our data, as both increased with increasing pH. However, we also obtained good fits to the data using the total adsorbed Fe(II) concentration. The model fits shown in Figures 2−4 are based on the total adsorbed Fe(II) concentration, rhet = khet[>SOFe(II) T ][Hg(OH)2 ]

(II)+ where [>SOFe(II) ] + [>SOFe(II)(OH)+] for T ] = [>SOFe (II)+ goethite and hematite and [>SOFe(II)+ W ] + [>SOFeS ] for γalumina. khet values for the three minerals are listed in Table 2. Using total adsorbed Fe(II) obviates the need to rely on surface speciation derived from fits of adsorption models to macroscopic experimental data and allows for the estimation of the Hg(0) production rate in natural soils by measuring the total adsorbed Fe(II) concentration using chemical extraction techniques (e.g., Heron et al.66). Finally, that surface-catalyzed Hg(0) production is observed at pH < 6, where the modeled concentration of >FeOFe(II)OH surface species is very low (Figure S1a and b), suggests that >FeOFe(II)+ species also contributes to Hg(II) reduction, albeit at a lower rate than >FeOFe(II)OH. The rate constants for the homogeneous, hematite-, and goethite-catalyzed reductions are within the same order of magnitude (Table 2). However, the observed rates in the presence of the Fe(III) oxyhydroxides (Figures 2 and 3) are higher than those for the homogeneous system (Figure 1). This difference is due to the smaller concentration of aqueous FeOH+ compared to surface-bound Fe(II) species at a given pH. The similarity in rate constants between the homogeneous and Fe(III) oxyhydroxide-catalyzed reduction suggests that the reduction potentials for their respective couples are similar. Using a linear free energy relationship (LFER), the reduction potentials for adsorbed Fe(II) and goethite and lepidocrocite (γ-FeOOH) systems were estimated at 0.48 and 0.44 V, respectively.67 These values are close to the reported reduction potential for the FeOH2+/FeOH+ couple of 0.34 V. Reduction potentials for structural Fe(II) in magnetite, ilmenite and several Fe-silicate minerals are in the range of 0.33−0.65 V.68,69 This similarity in reduction potentials between Fe(II) adsorbed on goethite and minerals with structural Fe(II) suggests that Fe(II)-bearing minerals may exhibit second-order Hg(II) reduction rate constants similar to those found in this study. The khet value for γ-alumina-catalyzed reduction, on the other hand, is more than an order of magnitude smaller than those for Fe(III) oxyhydroxides-catalyzed and homogeneous reduction (Table 2), suggesting a higher reduction potential for

>SOFe(II)(OH)1n− n + Hg(OH)2 k1a

Xooa Y >SOFe(II)(OH)1n− n ···Hg(OH)2 k −1 kr

→ >SOFe(III)(OH)11 −+ nn +

1 Hg 2 + + OH− 2 2

(4)

The first reaction represents the kinetically controlled reversible Hg(II) interaction with the surface-bound Fe(II) that results in the formation of a precursor surface complex with ka1 and ka−1 as the formation and dissociation rate constants, respectively. The second reaction represents a one-electron reduction of Hg(II) to Hg(I) as Hg22+ with kr as rate constant for the irreversible reaction. In the reaction scheme shown in eq 4, the dissociation of the reaction products from the surface is assumed to be fast and irreversible. Inorganic Hg(I) is marginally stable in the presence of anions, such as OH−, and its further reduction to Hg(0) is thermodynamically favorable.44 Hg(I) can undergo a one-electron reduction by dissolved Fe(II), surface-bound Fe(II), or by disproportionation6,12 as shown in the SI. To derive the rate expression for surface-catalyzed Hg(0) production, we have assumed a steady-state concentration for the precursor complex in eq 4. This assumption implies that the precursor surface complex is present at a lower concentration than the reactants and products.45 In a treatment similar to that performed for the homogeneous reduction rate expression shown in the SI, the surface-catalyzed Hg(0) production rate, rhet, may be written as − n) + rhet = khet[>SOFe(II)(OH)(1 ][Hg(OH)2 ] n

(6)

(5)

where khet is the observed second-order surface-catalyzed reduction rate constant (M−1 min−1). Assuming that Hg(0) 7210



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Fe(II) adsorbed on γ-alumina. In contrast to the Fe(III) oxyhydroxide-catalyzed reduction, where solid concentration increased the reduction rate due to increased Fe(II) adsorption, the reduction rate decreased as the γ-alumina concentration increased (Figure 4b and c). Based on our model calculations, increased γ-alumina concentration resulted in increased Hg(II) adsorption, decreasing the rate of both the homogeneous and heterogeneous Hg(II) reduction pathways. Similar behavior has been observed for Cr(VI) reduction by Fe(II), where increased γ-alumina concentration decreased the reduction rate, in contrast to what was observed with Fe(III) (hydr)oxide, SiO2, and clays, where the reduction rate increased with solid concentration.67 Buerge and Hug67 also observed a marked decrease in goethite-catalyzed Cr(VI) reduction rate as the suspension was titrated with Al. The inhibiting effect of Al may have important implications for Hg(0) production in natural systems because of the importance of Al-substituted goethite in soils and aquifer sediments,70,71 and the relative abundance of Al (hydr)oxide species, compared to Fe(III) oxyhydroxides, in soils of temperate environments. Peretyazhko et al.13 studied the kinetics of Hg(II) reduction in the presence of 1.5 μM Fe(II) and 5 mg L−1 hematite. Applying the rate expressions and constants in Table 2 to their experimental conditions, the simulated Hg(0) concentrations are within 30−50% of the experimental data (Figure S3). Given the significant differences between the Fe(II) and hematite concentrations, this level of agreement is reasonable. According to our model calculations, nearly all of Fe(II) was adsorbed to the hematite surface and, therefore, nearly all of the Hg(0) was generated via the surface-catalyzed pathway. Environmental Implications. Results of this study suggest that in some freshwater sand and gravel aquifers, Hg(II) can be effectively reduced to Hg(0) by Fe(II) within several hours. Fe(II) adsorption on hematite and goethite accelerated the Hg(0) production rate over that observed with aqueous Fe(II). Thus, the presence of nanoparticulate Fe(III) oxyhydroxide coatings on the aquifer material,71,72 especially close to the oxic−anoxic boundary where redox potential is often controlled by Fe(II) production and oxidation, would accelerate Hg(0) production. pH-Dependent Fe(II) speciation in the aqueous phase and Fe(II) coordination at the surface was shown to control Hg(0) production rate. Use of Fe(II) at a circumneutral pH, as a ferrous salt or elemental Fe(0), in remediation of Hgcontaminated waters may also prove feasible and cost-effective. The latter provides a steady source of Fe(II) that is associated with an oxidized Fe(III) layer that is present as a corrosion product, and as such, may greatly accelerate the Hg(II) reduction rate. In contrast to Fe(III) oxyhydroxides, the Hg(0) production rate was slower in the presence of γ-alumina than what was observed in aqueous Fe(II) solutions. Thus, Hg(0) production in sediments with low Fe(III) content may be less important than in sediments with high Fe(III) content. The fact that the Hg reported in anoxic groundwater and freshwater sediments is rarely completely in Hg(0) form9,22 may be because of degassing of Hg(0) or, more likely, the decrease in Hg(II) reduction rate due to presence of high concentrations of complexing ligands, such as Cl− that has been shown to decrease the abiotic Hg(II) reduction rate by increasing the electron density of Hg(II). 14,16 At Cl − concentrations in the lower ppm range (typical of many freshwaters), however, where Hg(II) speciation is dominated by Hg(OH)2, Hg(II) reduction by Fe(II) would be plausible. Dissolved sulfide (S(-II)) binds Hg(II) highly favorably, and as

such, Hg(II) reduction by Fe(II) in sulfidic environments is not expected to be significant. Also, there is a growing body of literature on the formation of HgS nanoparticles that are stabilized by DOM in sulfidic waters.19,20 Alternatively, any Hg(0) produced that is transported into oxic sediments may be oxidized by sediment-bound MnO2 that is energetically capable of oxidizing Hg(0) to Hg(II) (E0 for MnO2/Mn2+ couple = 1.23 V). The kinetics of this reaction have not been adequately studied. However, in the subsurface environment, where Hg(0) may persist in dissolved or adsorbed form, Hg(0) oxidation by MnO2 may be plausible. The modeling framework developed in this study can accommodate such additional factors as the presence of competing ligands, as well as Hg(0) oxidation. Further model development and application to natural systems require obtaining the appropriate calibration data.

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ASSOCIATED CONTENT

S Supporting Information *

Reaction mechanism for homogeneous Hg(II) reduction, a table of equilibrium binding constants used in modeling the kinetic data, a figure of pH edges for Fe(II) adsorption onto mineral surfaces, a table summarizing the experimental conditions used in equilibrium modeling for Fe(II) adsorption onto these surfaces, a figure simulating Hg(II) adsorption to the mineral surfaces, and a figure simulating data from Peretyazhko et al.13 on Hg(0) production kinetics. This information is available free of charge via the Internet at http://pubs.acs.org/

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AUTHOR INFORMATION

Corresponding Author

*Phone: +1-207-581-1277; fax: +1-207-581-3888; e-mail: [email protected];. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS Funding was provided by the USGS through the National Research Program and Toxic Substances Hydrology Program, and the Trustee Professorship at the University of Maine. A.A. is grateful to the University of Maine for granting him sabbatical leave in 2011. Michelle Arias, Chris Fuller, Evangelos Kakouros, Nazila Kaviani, Le Kieu, Matthias Kohler, Carl Lamborg, and Fred Murphy are acknowledged for their assistance. Matthias Kohler also kindly provided the hematite sample. Andrea Foster and four anonymous reviewers provided valuable critiques of the manuscript. Any use of trade, product, or firm names is for descriptive use only and does not imply endorsement by the U.S. Government.

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REFERENCES

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