Pyrite oxidation in air-equilibrated solutions: An electrochemical study

Chemical Geology 470 (2017) 67–74

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Pyrite oxidation in air-equilibrated solutions: An electrochemical study a,⁎

Paul Chiriță , Michel L. Schlegel a b c

MARK

b,c

University of Craiova, Department of Chemistry, Calea Bucuresti 107I, Craiova 200478, Romania CEA, DEN, DPC/SEARS/LISL, Bât. 391, F-91191 Gif-sur-Yvette, France University of Evry-Val d'Essonne, LAMBE, F-91025 Evry, France

A R T I C L E I N F O

A B S T R A C T

Keywords: Pyrite oxidation Oxygen Electrochemical techniques Reaction mechanism

Oxidative dissolution of pyrite (FeS2) in air-equilibrated solutions was investigated with electrochemical techniques (polarization measurements and electrochemical impedance spectroscopy) at different solution pH (1 to 5) and temperatures (25 to 40 °C). We have found that, at a given temperature, the oxidation current density (jox) varies only slightly from initial pH 1 to 4, but was significantly lower at initial pH 5, while oxidation potential (Eox) decreases when pH increases. At the investigated temperatures, the reaction order with respect to hydrogen ions has been found to be zero, suggesting that H+ is not involved into the rate-determining step of the oxidation rate. The jox values increase with temperature, whereas Eox potentials are hardly affected by temperature from 25 to 40 °C. Activation energies (Ea) were estimated from polarization data and vary between 57 ± 28 kJ/mol (initial pH 5) and 19 ± 3 kJ/mol (initial pH 4). The obtained values indicate that the rate-determining step of the overall oxidative dissolution of FeS2 involves a surface reaction. The importance of surface reaction is also supported by the electrochemical impedance spectroscopy measurements which show that interfacial transfer of electrons is dominated by surface charge transfer and the elementary surface reaction which controls the mineral oxidative dissolution is the electron transfer from the conduction band of pyrite to adsorbed O2.

1. Introduction

Furthermore, Fe2 + released by dissolution can be oxidized by O2 to Fe3 + which acts as an oxidant of pyrite

Pyrite (FeS2) oxidation is important in a wide range of environmental processes and technological applications, and, as such, has been extensively investigated (Lowson, 1982; Rimstidt and Vaughan, 2003; Rickard and Luther, 2007; Murphy and Strongin, 2009). Knowledge of the oxidation mechanisms and rates are essential to better model biogeochemical cycling of iron and sulfur, formation of acid mine drainage (AMD), and mineral processing. Ferric iron (Fe3 +) and oxygen (O2) have been recognized as the most important oxidants of pyrite in natural systems (Lowson, 1982; Lehner et al., 2007). Early dissolution studies have demonstrated that Fe3 + is the most effective oxidant of pyrite, especially under acidic pH conditions. However, the dissolved concentration of this species steadily decreases with pH, thereby diminishing its relative importance (Nicholson et al., 1988). In contrast, O2 is fairly abundant in most subsurface aerobic environments, hence it may be considered as a potent oxidant in AMD and hydrometallurgical processes (Williamson and Rimstidt, 1994). Pyrite oxidation by O2 usually leads to endproducts such as dissolved sulfate, protons and Fe2 +, according the mass balance equation

FeS2 + 14 Fe3 + + 8 H2 O = 15 Fe2 + + 2 SO4 2 − + 16 H+

FeS2 + 7 2 O2 + H2 O = Fe2 + + 2 SO4 2 − + 2 H+

⁎

(1)

Corresponding author. E-mail address: [email protected] (P. Chiriță).

http://dx.doi.org/10.1016/j.chemgeo.2017.08.023 Received 2 March 2017; Received in revised form 27 August 2017; Accepted 28 August 2017 Available online 01 September 2017 0009-2541/ © 2017 Elsevier B.V. All rights reserved.

2+

(2) 3+

is oxidized to Fe , and Fe3 + Thus, a cycle appears in which Fe subsequently oxidizes pyrite, releasing additional protons and Fe2 + (Singer and Stumm, 1970). At near-neutral pH, Fe3 + precipitation somewhat hinders this Fe2 +/Fe3 + cycling (Nicholson et al., 1990). The overall processes of pyrite oxidation can then be written as

FeS2 + 15 4 O2 + 5 2 H2 O = FeOOH + 2 SO4 2 − + 4 H+

(3)

FeS2 + 15 4 O2 + 7 2 H2 O = Fe(OH)3 + 2 SO4 2 − + 4 H+

(4)

The precipitated Fe(III) phases can inhibit the mineral oxidation by shielding the reaction center from the oxidant (O2) (Nicholson et al., 1990; Perez-Lopez et al., 2007). Moses et al. (1987) reported that in the presence of dissolved O2 and at pH > 3.9, thiosulfate and polythionate species are present in the reaction system. Also elemental sulfur was found on the reacted surface of pyrite at pH 2 in the presence of O2(aq) (Demoisson et al., 2008; Sun et al., 2015; Qiu et al., 2016). However, other studies (e.g. Descostes et al., 2004; Gartman and Luther, 2014) did not detect all of these species, and so their formation is still somewhat controversial.

Chemical Geology 470 (2017) 67–74

P. Chiriță, M.L. Schlegel

2000 and 3000), immersed in 1 M HNO3 for 60 s, rinsed several times with O2-free distilled water and degreased with acetone. The reference electrode was a saturated calomel electrode (SCE) and the counter electrode was platinum foil. A Luggin capillary, positioned at 1 mm from the working electrode, was connected to the reference electrode via a 3 M KCl solution. All potentials in this paper are quoted relative to the Standard Hydrogen Electrode (SHE). Potentiodynamic polarization, EIS measurements and data analysis were carried out with an electrochemical workstation (ZAHNER Elektrik IM6e) controlled by Thales software. The initial pH (pHi), the pH after the immersion of pyrite electrode (pHai) and the final pH (pHf) of solutions were measured using a combination glass electrode connected to a pH/millivoltmeter (Consort C534). Before each measurement, the glass electrode was calibrated against two commercial buffers (pH 4.01 and pH 7.00). For each experiment, the reactor was filled with 250 mL of desired solution, introduced in a thermostated water bath, and equilibrated with air at the desired temperature. No stirring was performed, except for the slow air bubbling. The working electrode was equilibrated in the oxidizing solutions for > 40 min to assure steady states. For the potentiodynamic polarization the potential was scanned from − 250 mV below the open circuit potential (OCP) to 250 mV above OCP at a scanning rate of 1 mV/s, and the results were processed using the Thales software. First, the cathodic and anodic Tafel slopes (βc and βa) were determined from tangents to the linear regions of the cathodic and anodic curves. Next, the oxidation potential (Eox) and oxidation current density (jox) (Table 1) were obtained from the intersection point of these tangents. The computational uncertainties on Eox ( ± 0.007 V), current density ( ± 16%), cathodic and anodic Tafel slopes ( ± 0.04 and ± 0.025 V/dec, respectively) were obtained from duplicate tests performed at 30 °C and in the pH range of 1 to 5. Impedance measurements were carried out at OCP with a sinusoidal signal of 10 mV amplitude and over a frequency range of 10 mHz–3 MHz.

Although the aqueous oxidation of pyrite in the presence of O2 has been studied using several techniques, including spectroscopic (Bailey and Peters, 1976; Moses et al., 1987; Moses and Herman, 1991; Nesbitt and Muir, 1994; Caldeira et al., 2010), aqueous batch (Williamson and Rimstidt, 1994; Kamei and Ohmoto, 2000; Descostes et al., 2004; Gleisner et al., 2006) and electrochemical measurements (Ahlberg and Broo, 1997; Holmes and Crundwell, 2000; Liu et al., 2008; Savage et al., 2008; Liu et al., 2009; Constantin and Chirita, 2013), some key aspects (rate law, rate determining step(s), reaction intermediates and reaction products) of the pyrite oxidative dissolution under mildly acidic conditions in aerated solutions can be further clarified. Williamson and Rimstidt (1994) compiled the most reliable data available in the literature and derived a rate law of pyrite oxidation as a function of pH (over the pH range of 2–10) and concentration of dissolved O2 ([O2(aq)]). They showed that the rate of aqueous oxidation of pyrite by O2 is inversely proportional to the concentration of protons (reaction order − 0.11) and directly proportional to the square root of [O2(aq)]. The experimental observations supported an electrochemical mechanism whereby anodic and cathodic reactions occur on the pyrite surface (Williamson and Rimstidt, 1994). Holmes and Crundwell (2000) further developed the electrochemical equations for anodic and cathodic reactions at the pyrite surface. By applying the condition that the net accumulation of electrons in pyrite is zero, they could derive a theoretical relationship for the rate of pyrite oxidative dissolution. This rate order could reproduce the reaction orders observed experimentally. In addition, several studies have demonstrated that the initial step of pyrite oxidation can be successfully investigated by electrochemical methods at the surface of pyrite electrodes (Peters and Majima, 1968; Biegler and Swift, 1979; Meyer, 1979; Ahlberg and Broo, 1997; Kelsall et al., 1999). The results of these studies can be translated directly in terms of pyrite interfacial potential and its impact on the dissolution rate and mechanism. Yet, most of these studies are restricted to acidic conditions and thus their applications to environmentally relevant conditions remain problematic. The purpose of this study is to obtain the electrochemical kinetics parameters and identify the reaction pathway of pyrite oxidative dissolution in air-equilibrated solutions over pH range 1 to 5 and temperature range of 25 to 40 °C. The pH and temperature ranges thus extend the conditions previously investigated by Holmes and Crundwell (2000). The electrochemical parameters for pyrite oxidative dissolution (oxidation current densities, oxidation potentials, cathodic and anodic Tafel slopes) were determined by Tafel method (Brett and Oliveira Brett, 1993). Information about physical and chemical processes occurring in the pyrite/solution interface was provided by electrochemical impedance spectroscopy (EIS) (Bryson and Crundwell, 2014). Based on experimental results, we analyzed the interaction of pyrite with dissolved O2 and identified the rate determining step as the electrons transfer from the conduction band of pyrite to O2.

3. Results and interpretation 3.1. pH measurements The pHai drift with respect to pHi is within uncertainties. The deviation of pHf with respect to pHai is also small at low pHi (Fig. 1). At higher pHi values of 4 and 5, the pH slightly drifts to higher values. This increase could be explained by the non-oxidative dissolution of S(-II) presents on the electrode surface (Kamei and Ohmoto, 2000) and/or the Table 1 Summary of the electrochemical kinetic parameters for pyrite in aerated HCl solutions with pH in the range of 1.0 to 5.0 and at temperatures between 25 and 40 °C.

2. Materials and methods Natural pyrite of unknown origin was used for electrochemical experiments. Pyrite chemical analysis was performed by combustion (sulfur) and ICP-OES (other elements) (S.C. Prospectiuni Geologice S.A.) and yielded a S:Fe atomic ratio close to 2. Impurities detected include Mn (52.8 mg/kg), Co (283.7 mg/kg), Ni (848.6 mg/kg); Cu (53.9 mg/kg) Zn (93.5 mg/kg) Pb (29.5 mg/kg) Cd (0.4 mg/kg) Ag (0.9 mg/kg) As (146.3 mg/kg), and Sb (1.0 mg/kg). X-ray diffraction confirmed the pyrite nature of the starting material, and showed the presence of only quartz as a detectable crystalline contaminant. All electrochemical experiments were performed in a three-electrode cell with aerated distilled water adjusted to the desired initial pH (pHi) by adding aliquots of HCl solutions. The working electrode was prepared from a pyrite monolith cut into a cubic shape and sealed with epoxy resin, exposing an area of 2 cm2. Before each experiment, the working electrode was mechanically polished with emery paper (600, 68

pH

Temperature/°C

jox/μA cm− 2

Eox vs. SHE/V

βa/V

βc/V

1.0 2.5 3.5 4.0 5.0 1.0 2.5 3.5 4.0 5.0 1.0 2.5 3.5 4.0 5.0 1.0 2.5 3.5 4.0 5.0

25 25 25 25 25 30 30 30 30 30 35 35 35 35 35 40 40 40 40 40

1.04 1.31 0.72 1.18 0.23 1.30 1.32 1.43 1.23 0.44 1.73 1.94 1.43 1.49 0.31 2.24 1.82 1.61 1.67 0.88

0.561 0.516 0.460 0.442 0.366 0.573 0.520 0.477 0.459 0.372 0.569 0.519 0.468 0.431 0.354 0.577 0.524 0.467 0.427 0.366

0.17 0.16 0.18 0.20 0.17 0.15 0.18 0.19 0.19 0.14 0.17 0.16 0.18 0.21 0.08 0.14 0.15 0.18 0.17 0.18

− 0.12 − 0.14 − 0.16 − 0.19 − 0.17 − 0.11 − 0.13 − 0.17 − 0.19 − 0.16 − 0.11 − 0.13 − 0.17 − 0.22 − 0.10 − 0.11 − 0.14 − 0.17 − 0.20 − 0.18

Chemical Geology 470 (2017) 67–74

P. Chiriță, M.L. Schlegel

Fig. 1. Variation of pHf versus pHai at different pHi. The temperature has no coherent influence on the observed deviations.

Fig. 3. The effect of pH on the OCP of pyrite at pH between 1.0 and 5.0 and temperatures in the range of 25 to 40 °C.

proton adsorption on the mineral surface (Weerasooriya and Tobschall, 2005).

Table 2 The slopes of Eox with respect to pH (pH between 1.0 and 5.0), αFeS2 − αO2 and αFeS2.

3.2. Polarization measurements 3.2.1. Effect of pH The influence of temperature on pyrite oxidation in air-equilibrated solutions was studied by potentiodynamic polarization measurements at pHi from 1 to 5 (Fig. 2). At a given temperature, the oxidation potential (Eox) decreases when pHi increases (Fig. 3). The slopes of Eox with respect to pH are summarized in Table 2. The obtained slopes are different from that predicted by Holmes and Crundwell (2000) for pyrite polarized in oxygen-bearing solutions, i.e. 0.038. The predicted value can be computed with Eq. (5) (derived from Eq. (15) of Holmes and Crundwell (2000) for the case of O2 as unique oxidant)

Eox = 2.303

k O2 RT [O2 ][H+]0.64 ⎞⎟ log ⎛⎜ k F FeS 2 ⎝ ⎠

Temperature/°C

Slope

25 30 35 40

− 0.048 − 0.048 − 0.053 − 0.053

± ± ± ±

0.005 0.006 0.006 0.005

αFeS2 − αO2

αFeS2

− 0.21 − 0.20 − 0.28 − 0.29

0.29 0.30 0.22 0.21

± ± ± ±

0.09 0.10 0.08 0.06

± ± ± ±

0.09 0.10 0.08 0.06

where [O2] and [H+] are concentrations, F is the Faraday constant, R is the gas constant, T is the temperature. kO2 and kFeS2 are rate constant for the reactions of oxygen cathodic reduction

O2 + 4 H+ + 4e− = 2 H2 O

(6)

and pyrite anodic oxidation

FeS2 + 8 H2 O = Fe2 + + 2 SO4 2 − + 16 H+ + 14e− (5)

(7)

The obtained differences can be caused by different charge transfer Fig. 2. Potentiodynamic cathodic and anodic polarization curves recorded for pyrite in aerated HCl solutions with pH 1.0, 2.5, 3.5, 4.0, 5.0 and at (a) 25 °C; (b) 30 °C; (c) 35 °C and (d) 40 °C.

69

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coefficients (αFeS2 and αO2). Starting from the relations of Holmes and Crundwell (2000) for the current density of the pyrite dissolution (jFeS2) and the current density of the reduction of oxygen (jO2),

jFeS2 = k FeS2 [H+]−0.5 exp ⎛ ⎝

jO2

α FeS2 FE ⎞ RT ⎠

Temperature/°C

nH

25 30 35 40

0.01 0.00 0.03 0.05

(8)

−(1 − αO2 )FE ⎞ = −k O2 [O2 ][H+]0.14 exp ⎛ RT ⎠ ⎝ ⎜

Table 3 The reaction order with respect to proton concentration (nH) for each studied temperature. The values after ± represent two standard deviations for nH.

⎟

(9)

± ± ± ±

0.06 0.01 0.02 0.01

(αFeS2 and αO2 are the corresponding charge transfer coefficients) and following the procedure described by Holmes and Crundwell (2000) for the evaluation of mixed potential (or oxidation potential) (less imposing condition αFeS2 = αO2 = 0.5) we obtain the following equation

Eox = −(0.038 (1 + α FeS2 − αO2 ))pH + (0.059 (1 + α FeS2 − αO2)) log

k O2 [O2] k FeS2

(10)

From the slopes of Eox with respect to pH we have estimated the differences between αFeS2 − αO2 (Table 2). It is known that the shift of charge transfer coefficient from 0.5 occurs in the case of semiconductor electrodes (Brett and Oliveira Brett, 1993; Duinea et al., 2016), such as pyrite. Considering that αO2 is 0.5 we can estimate αFeS2 (Table 2). The obtained values are shifted towards 0 and indicate that the activated complex has predominantly the structure of oxidized species (Brett and Oliveira Brett, 1993). The βc and βa values slightly fluctuate over the pH range from 1 to 5, with no significant trend (Table 1). Within uncertainties, the obtained values of βa and βc are in the ranges of those obtained by Liu et al. (2008) under similar aerated conditions. In contrast, jox varies only slightly from pHi 1 to 4 at each studied temperature, but was systematically significantly lower at pHi 5. This finding indicates that the kinetics of pyrite oxidative dissolution at pH 5, and so the mechanism, differs from that at lower pH. Therefore, jox values from pH 1 to 4 were used first to determine the reaction order with respect to proton concentration (nH) (Fig. 4). At temperatures between 25 and 40 °C the reaction order equals zero within uncertainties (Table 3), suggesting that H+ has a negligible effect on the oxidation rate. This absence of dependence indicates that both cathodic and anodic halfreactions are equally affected by pH at pH ≤ 4.0. The obtained nH is higher than that found by Williamson and Rimstidt (1994) (i.e., − 0.11); is similar to that reported by McKibben and Barnes (1986) (i.e., 0), and is in the range of reaction orders reported by Liu et al. (2008) (i.e., − 0.08 to 0.25). The observed differences in nH are probably related to the various methods used to estimate the reaction rates (or equivalents).

Fig. 5. Arrhrenius plots (-ln (jox) versus T− 1) for the temperature range 25–40 °C. Activation energies are listed in Table 3 (pH in the range of 1.0 to 5.0).

3.2.2. Effect of temperature The jox values increase with temperature, consistent with an increase in the reaction rates. In contrast, Eox potentials are hardly affected by temperature from 25 to 40 °C (Table 1). Thus, from Eq. (5), the kO2/kFeS2 ratio can be considered invariant, meaning a unique activation energy (Ea) can be calculated for each pH value. The Ea values were obtained by a linear regression of -ln(jox) values vs T− 1 (Fig. 5). The obtained values vary between 57 ± 28 kJ/mol at pHi 5 and 19 ± 3 kJ/mol at pHi 4 (Table 4 and Fig. 6). The high uncertainty at pHi 5 is caused by the scattering of jox (Fig. 5). All the values obtained in our study are smaller than 121.1 kJ/mol, the highest Ea reported by Schoonen et al. (2000), and they are closer to the value of 56.7 ± 7.5 kJ/mol reported by McKibben and Barnes (1986). In fact, the discrepancy in Ea is in line with the spread reported in the literature (Lowson, 1982; Schoonen et al., 2000). Still, within uncertainty all these values are contained in the [22, 40] kJ/mol range, slightly greater than the activation energy for diffusion of aqueous species (Dove and Rimstidt, 1994). Thus, the rate-determining step of the overall oxidative dissolution of pyrite likely involves a surface reaction. Although Fe (III) oxyhydroxides may develop on the pyrite surface, at least at pH ≥ 3, their effect on dissolution rates does not show up during our experiments. Since our measurements lasted about one hour, any formed layer is likely not thick enough for diffusion to have a significant impact, and so the Ea value can be attributed to a surface reaction. Table 4 The activation energy (Ea) at pH in the range of 1.0 to 5.0. The values after ± represent two standard deviations for Ea.

Fig. 4. Pyrite oxidative dissolution current density dependence on pH at temperatures in the range of 25 to 40 °C.

70

Initial pH

Ea/kJ/mol

1.0 2.5 3.5 4.0 5.0

40.1 21.3 37.7 19.1 56.8

± ± ± ± ±

1.6 9.2 16.2 3.2 28.2

Chemical Geology 470 (2017) 67–74

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et al., 2012)

d sl =

(12)

where A is the exposed surface of the electrode, ε0 the permittivity of vacuum (8.85 × 10− 14 F/cm), and εsl, the dielectric constant of the surface layer. We have assumed that εsl is 10 (Ghahremaninezhad et al., 2012). At 25 °C the estimated thickness of the surface layer equal 27 ± 68 (pHi 1) and 460 ± 850 nm (pHi 5). The high uncertainties on dsl illustrate the limited impact, if any, of surface layers on interfacial reactivity, and precludes any definite conclusion as to the properties of these layers. The irregularities observed in Nyquist plots at middle frequencies (for example Fig. 7d,e,i,j) may result from inhomogeneity of the electrode surface (Chirita and Schlegel, 2015). Also, possible secondary processes such as adsorption of reaction products or intermediates on oxidized surface of pyrite can be responsible for the minor inductive features observed for some Nyquist plots at high frequencies (Fig. 7d,e,i,j) (Lasia, 1999).

Fig. 6. Variation of Ea with pH (1.0 to 5.0). Error bars represents the two standard deviation.

3.3. Electrochemical impedance spectroscopy measurements

4. Discussion

The oxidative dissolution of pyrite in air-equilibrated solutions at pH values from 1 to 5 was investigated by EIS at 25 and 40 °C (Fig. 7). A very important advantage of EIS is that the measurements at open circuit potential do not induce any change of the electrode surface by artificial polarization (Gao et al., 2014). Thus, EIS results illustrate the surface structural properties of “natural” pyrite/solution interface. The general shape of the curves in the Nyquist representation is very similar for all measurements, indicating that practically no change in the structure of pyrite/solution interface occurred within the range of reaction conditions. The impedance spectra exhibit two capacitive loops in the frequency range of 10 mHz to 3 MHz. The low-frequency capacitive loop can be assigned to charge transfer resistance (Rct) at the interface between pyrite and the solution (Aghzzaf et al., 2012; Constantin and Chirita, 2013). Because it is not an ideal semicircle, the capacitance associated to double layer (Qdl) is better described by a constant phase element (CPE) (Saleh, 2006; Qafsaoui et al., 2013; Constantin and Chirita, 2013; Chirita and Schlegel, 2015). The capacitive loop (Csl) at high frequencies may correspond to a surface layer developed on pyrite surface during oxidative dissolution (Constantin and Chirita, 2013). This layer can incorporate Fe oxyhydroxide and elemental sulfur (Demoisson et al., 2008) and is characterized by surface layer resistance (Rsl). The impedance parameters calculated by fitting the experimental data to equivalent circuit shown in Fig. 8 (where Rs is the solution resistance) are given in Table 5. It is important to note the high uncertainty on fitted Csl values (between 58 and 260%). The reaction resistance (Rr) can be defined as the following sum (Tuken et al., 2006; Chirita et al., 2015)

Rr = R sl + R ct

εsl ε 0 A Csl

4.1. Evaluation of the initial reaction rate from jox The jox values obtained from the Tafel plots significantly differ from previous electrochemical estimates (Holmes and Crundwell, 2000; Liu et al., 2008). In our study, jox values correspond to OCP conditions, i.e. the real conditions under which the oxidation rates of pyrite were determined in aerated aqueous batch dissolution experiments. The j values (Fig. 2) indicate that only minor amounts of pyrite were oxidized by the electrical current during each electrochemical experiment. For example, in the case of the highest current density (28.7 μA cm− 2, recorded at E = 0.844 V; pHi 1 and 40 °C) the amount of pyrite oxidized cannot be higher than 0.15 μmol, after approximately 1 h. The small variation of pH during the overall oxidation process indicates that only minor amounts of material are dissolved by proton attack (Kamei and Ohmoto, 2000). This is consistent with the known predominance of oxidation reactions in the dissolution process of pyrite. Therefore, it is reasonable to assume that the electrode surface remained largely unchanged after potentiodynamic polarization experiments (approximately 1 h), and the reaction rates obtained by the conversion of the obtained current densities according to Eq. (13) are similar to the initial rates of pyrite oxidation (rpyrite):

rpyrite (μmol m2 s) =

i ox 0.0015F

(13) −2 −1

s . Eq. (13) where F = 96,486 C/mol and r is expressed in μmol m is valid if the oxidation products are both Fe(III) and SO42 −. This is the de facto situation especially at pH > 3 (Peiffer and Stubert, 1999; Descostes et al., 2004). The derived (initial) rates best describe the aqueous oxidation of pyrite and can further be used with confidence in the geochemical modeling. Unlike the initial rates obtained in aqueous batch experiments, those obtained by Eq. (13) are not affected by incomplete oxidation of sulfur to soluble species, non-redox dissolution of Fe2 + (Badica and Chirita, 2015) or Fe3 + precipitation.

(11)

Eq. (11) indicates that the overall reaction progress (which is inversely proportional to Rr) can be controlled by the rate of charge transfer (surface reaction) if Rsl < Rct, mass transfer (diffusion) across surface layer if Rsl > Rct or both if Rsl = Rct. With the exception of the values obtained at pHi 4, Rct is higher than Rsl by at least one order of magnitude (Table 5) which confirms that the rate-determining step is related to a surface reaction. Even at pHi 4, where Rct/Rsl ratios equal 3 at 25 °C and 5 at 40 °C, respectively, Rsl and Rct are still significantly different so we can consider that the overall rate is limited by a surface reaction. Hence, the non-monotonic variations of Ea with pHi would rather be explained by the reactivity changes (for example by hydrolysis) of the surface sites involved in the rate determining step. Although the values of Csl are affected by large uncertainties, they may be used to estimate the thickness (dsl) of the surface layer formed on pyrite surface during mineral oxidation using (Ghahremaninezhad

4.2. Rate determining step Our findings show that pyrite oxidation by O2(aq) is not dependent on [H+] (nH = 0) and the rate is likely controlled by a surface reaction. The elementary surface reaction which controls the mineral oxidative dissolution can be cathodic or anodic, but does not involve protons. We can assume that this elementary reaction is the electron transfer from the conduction band (CB) of pyrite to adsorbed O2:

> py‐CB…O2 = > py‐CB+…O2 − 71

(14)

Chemical Geology 470 (2017) 67–74

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Fig. 7. Nyquist impedance spectra of pyrite in air-equilibrated solutions for pH ranging between 1.0 and 5.0, at 25 and 40 °C. The insert figures within the Nyquist plots (a) and (f) show the magnified curves in the high-frequency region.

Considering Eq. (14) as a rate-determining step is in in good agreement with the study by Gartman and Luther (2014) showing that pyrite oxidation by O2(aq) in sea water is first order with respect to both pyrite and O2 concentrations. Also, our results are in agreement with the results reported by Chirita et al. (2014) suggesting that the ratedetermining step does not involve sulfate and that the activated complex of pyrite oxidative dissolution in the presence of O2(aq) is of ionic nature (perhaps py-CB+ δ…O2-δ). According to the estimated values of αFeS2 which indicate that the activated complex has a structure closer to that of the oxidized species, it results that δ is shifted towards 1. Association with surface sites would polarize and destabilize O2 molecules, thereby accelerating their reductive dissociation (Rozgonyi and Stirling, 2015; Dos Santos et al., 2016). Depending on the pHi and temperature, the anodic surface sites may have distinct reactivities. If, for example, the oxygen binding site at the pyrite surface involves Fe

Fig. 8. Equivalent electrical circuit used to fit the impedance spectra obtained for pyrite electrode in air-equilibrated solutions.

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Table 5 Electrochemical parameters obtained from EIS measurements on the pyrite electrode in aerated HCl solutions at 25 and 40 °C and pH between 1.0 and 5.0. The errors were estimated by Thales software. pH

Temp./°C

Rs/Ω cm2

Error (%)

Csl/nF cm− 2

Error (%)

Rsl/Ω cm2

Error (%)

Qdl/μF cm− 2

Error (%)

ndl

Error (%)

Rct/kΩ cm2

Error (%)

1 2.5 3.5 4.0 5.0 1 2.5 3.5 4.0 5.0

25 25 25 25 25 40 40 40 40 40

2.08 6.3 20.3 50.64 51.66 2.42 5.02 14.3 42.5 44.82

8.4 7.8 5.7 10.9 11.5 4.3 9.5 5.5 9.6 11.1

332.65 63.65 22.65 23.8 19.1 391.05 76.65 34.75 38.1 29.05

250.3 70.3 259.7 102.7 184.8 199.8 58.0 179.9 207.2 71.6

22.94 478.8 3450 8042 5464 21.5 340.2 1803.8 3162 2882

12.5 8.3 6.5 4.8 7.5 12.7 8.2 8.0 5.1 10.4

46.26 29.8 12.85 42.55 8.96 47.55 16.35 14.2 92.55 14.55

3.2 1.3 3.7 12.9 3.5 4.6 2.6 6.2 7.9 5.1

0.771 0.755 0.737 0.801 0.753 0.721 0.692 0.707 0.841 0.694

0.5 0.4 0.4 0.8 0.6 0.8 0.2 0.8 0.5 0.4

29.58 23.86 52.3 22.5 122.2 10.23 15.4 18.18 15.1 29

28.3 9.6 5.2 14.4 27.0 14.9 4.7 1.8 23.7 27.4

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(III), then it can hydrolyze with increasing pH, and this hydrolysis may explain the non-monotonic variations in estimated activation energies. At pH < 4, the dominant Fe(III) species is FeOH2 + while at pH > 4 the dominant species is Fe(OH)2+ (Descostes, 2001). This change can explain the decrease of the current densities observed at pH 5. 5. Conclusions The aqueous oxidation of pyrite was investigated in air-equilibrated solutions over the pH range 1–5 and temperatures from 25 to 40 °C by potentiodynamic polarization and EIS. It was found that, at a given temperature, Eox decreases when pH increases, while jox varies only slightly from pHi 1 to 4, but was significantly lower at pHi 5. At studied temperatures the reaction order equals zero, suggesting that H+ has a negligible effect on the oxidation rate. The jox values increase with temperature, whereas Eox potentials are hardly affected by temperature from 25 to 40 °C. The obtained Ea values vary between 57 ± 28 kJ/ mol (pHi 5) and 19 ± 3 kJ/mol (pHi 4). These values indicate that the rate-determining step of the overall oxidative dissolution of pyrite involves a surface reaction. This finding is also supported by the EIS measurements which show that interfacial transfer of electrons is dominated by surface charge transfer (Rct). Our findings confirm that the elementary surface reaction which controls the mineral oxidative dissolution is the electron transfer from CB of pyrite to adsorbed O2. Our study underlines the interest of the electrochemical methods for the evaluation of the initial rates of the minerals aqueous oxidation. The calculated initial rates are more reliable than those obtained by aqueous batch experiments because they are not affected by incomplete oxidation of sulfur to soluble species, non-redox dissolution of Fe2 + or Fe3 + precipitation. Acknowledgements This article is dedicated to Cristina Constantin, our deceased colleague who has been an outstanding support for electrochemical experiments. This work was supported by the IFA-CEA Collaborative Program (Project C1-04) and by a grant of the Romanian National Authority for Scientific Research, CNDI–UEFISCDI, project number 51/ 2012. We thank Professor J. Donald Rimstidt and an anonymous reviewer for their valuable comments and suggestions. References Aghzzaf, A.A., Rhouta, B., Steinmetz, J., Rocca, E., Aranda, L., Khalil, A., Yvon, J., Daoudi, L., 2012. Corrosion inhibitors based on chitosan-heptanoate modified beidellite. Appl. Clay Sci. 65–66, 173–178. Ahlberg, E., Broo, A.E., 1997. Electrochemical reaction mechanisms at pyrite in acidic perchlorate solutions. J. Electrochem. Soc. 144, 1281–1285. Badica, C.E., Chirita, P., 2015. An electrochemical study of the oxidative dissolution of iron monosulfide (FeS) in air-equilibrated solutions. Electrochim. Acta 178, 786–796. Bailey, L.K., Peters, E., 1976. Decomposition of pyrite in acids by pressure leaching and anodization: the case for an electrochemical mechanism. Can. Metall. Q. 15, 333–344.

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