Reaction of FeS with Fe(III)-bearing acidic solutions

Chemical Geology 334 (2012) 131–138

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Reaction of FeS with Fe(III)-bearing acidic solutions Paul Chiriţă a,⁎, Michel L. Schlegel b a b

University of Craiova, Department of Chemistry, Calea Bucureşti, 107I, 200512 Craiova, Romania CEA, DEN/DANS/DPC/SEARS/Laboratory for the Engineering of Surfaces and Lasers, F-91191 Gif-sur-Yvette, France

a r t i c l e

i n f o

Article history: Received 3 May 2012 Received in revised form 2 October 2012 Accepted 4 October 2012 Available online 12 October 2012 Editor: J. Fein Keywords: Iron monosulfide Ferric iron Sulfur rich layer Dissolution

a b s t r a c t The reaction of FeS with Fe(III)-bearing acidic solutions (Fe(III)BAS) was probed at 25 °C and pH between 2 and 3. Initial dissolved Fe 3+ ([Fe3+]init) was varied from 0.1 mM to 1 mM, and the length of the experiments was 240 min. Except for the experiment at initial pH 2, total dissolved iron ([Fe]total) decreased immediately (within 1 min) after contact of FeS and Fe(III). Afterwards, [Fe]total increased smoothly. A progressive increase in pH values and an Eh decrease within 240 min of reaction time were also observed. The reaction order of FeS dissolution in Fe(III)BAS with respect to [H+] is estimated to 0.65 at initial pH 3.0, and increases up to 1.0 with decreasing initial pH, indicating that [H+] is an important parameter of FeS dissolution in Fe(III)BAS. In contrast, changes in [Fe3+]init have only a limited effect on the rate of FeS dissolution in Fe(III)BAS. Raman spectra of initial and reacted FeS samples reveal the accumulation of α-S8-like material on FeS surface. These results support a mechanism of FeS dissolution in Fe(III)BAS starting with the protonation of mineral surface and Fe 3+(aq) adsorption. Adsorbed protons subsequently accelerate Fe 2+ release from FeS matrix into solution. The adsorbed Fe 3+ may oxidize sulfur moieties and generate insoluble species, presumably polysulfide and elemental sulfur. The subsequent migration of Fe 2+ into solution is controlled by the formed sulfur rich layer. © 2012 Elsevier B.V. All rights reserved.

1. Introduction After iron disulfides (FeS2), iron monosulfides (FeS) are the most common sulfide minerals. FeS2 and FeS regulate and control the global geochemical cycles of iron and sulfur (Benning et al., 2000). FeS minerals occur in base metal sulfide and gold ores, and mineral wastes in many mining environments are rich in FeS (Belzile et al., 2004; Chandra et al., 2011). They can form in periodically anoxic environments (Rickard and Luther, 2007) such as salt marshes and swamps (Luther et al., 1992; Smith and Melville, 2004). In addition, they can be the by-product of sulfur reduction and metallic iron oxidation, e.g. in permeable reactive barriers, or in the near-field of nuclear waste containers (Schlegel et al., 2008). Once removed from their equilibrium conditions, e.g. by ingress of oxidative solutions, or by disposal of mine waste in atmosphere-exposed sites, they can dissolve by presumably forming Fe(III) (oxyhydr)oxide phases, and sulfur species of oxidation states intermediate between sulfide (S 2−) and sulfate (SO42−). Thus these oxidation processes play an important role in controlling the kinetics of Fe cycling. In addition, the oxidative dissolution of FeS is very often an important source of pollution because of the release of toxic impurities, such as toxic metals and arsenic, to natural solutions (Thomas et al., 1998). Finally, sulfur species in intermediate oxidation states can significantly change the oxydo-reduction (redox) conditions well away from the ⁎ Corresponding author. E-mail address: [email protected] (P. Chiriţă). 0009-2541/$ – see front matter © 2012 Elsevier B.V. All rights reserved. http://dx.doi.org/10.1016/j.chemgeo.2012.10.015

mineral surface. This change of redox conditions may alter the migration of redox-sensitive elements and radionuclides (like soluble U(VI)) out of polluted sites by promoting reducing precipitation (Descostes et al., 2010). Because of these environmental-related issues, an accurate knowledge of the mechanisms of FeS dissolution is warranted. The formation of sulfur species with intermediate oxidation states may be understood from the oxidation mechanism of sulfide (S 2−) to sulfate (SO42−). This oxidation requires a transfer of eight electrons per atom of sulfur. Because only one or two electrons are transferred per elementary step (Basolo and Pearson, 1958), complete sulfur oxidation requires several elementary steps resulting in the formation of intermediate sulfur products such as polysulfide (Sn2−), elemental sulfur (S 0), thiosulfate or sulfite (Pratt and Nesbitt, 1997; Thomas et al., 1998, 2001, 2003; Janzen et al., 2000; Belzile et al., 2004; Chirita and Descostes, 2006a,b; Chirita et al., 2008). A direct consequence of this incomplete oxidation is the formation of a sulfur rich layer (SRL) on FeS particles (Thomas et al., 1998, 2001, 2003; Janzen et al., 2000; Belzile et al., 2004). X-ray photoelectron spectroscopy (XPS) data showed that this SRL incorporates polysulfides and elemental sulfur (Thomas et al., 1998, 2001, 2003; Belzile et al., 2004). Most of laboratory studies have focused on the oxidative dissolution kinetics of FeS in oxic (Thomas et al., 1998, 2001; Mikhlin et al., 2002; Belzile et al., 2004; Chirita et al., 2008) and anoxic media (Chirita and Descostes, 2006a). Only few studies have addressed FeS oxidative dissolution in the presence of Fe 3+(aq), and the proposed reaction kinetics, mechanism and reaction products are somewhat

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inconsistent (Belzile et al., 2004). In the most extensive investigation of FeS dissolution in the presence of Fe3+(aq) (initial [Fe3+] of 2 ×10−4 and 1 ×10−3 M) at initial pH 2.50 and 2.75, Janzen et al. (2000) suggested that Fe3+ adsorption on the solid controlled FeS oxidation. Janzen et al. (2000) found that the apparent activation energy values (Ea) of FeS oxidation by Fe3+, calculated on the basis of Fe release, ranged from 23 to 63 kJ mol−1. Because of this spread, whether FeS dissolution is controlled by diffusion (for which Ea ≤40 kJ mol−1) or surface reaction (for which Ea typically equals 40 to 85 kJ mol−1; Lasaga, 1984) could not be determined. In this work we examine the kinetics and mechanism of FeS dissolution in Fe(III)-bearing acidic solutions (Fe(III)BAS) for conditions close to Fe 3+ solubility (pH between 2 and 3), by (i) monitoring the chemistry of the oxidative supernatants in contact with FeS at 25 °C and (ii) characterizing the new solid phases formed on FeS surface and resulting from the contact with oxidative solutions. 2. Methodology Synthetic FeS from Riedel-De-Haen was used as reactant. The results of previous studies (Thomas et al., 1998, 2001, 2003; Chirita and Descostes, 2006a,b; Chirita et al., 2008) have shown that the dissolution rate of synthetic FeS is important enough to yield large amounts of reaction products up to pH 3. The material was crushed under ethanol, and the supernatant suspension was removed. The solid was rinsed with fresh ethanol, and the operation repeated several times until the liquid phase was relatively clear. After drying the solid, the fraction ≤63 μm was recovered by sieving and stored as powder in an evacuated desiccator. X-ray diffraction (XRD) analysis showed the material to be troilite-2 H with residual traces of metal iron and lepidocrocite (γ-FeOOH). Lepidocrocite is known to form by air oxidation of reduced iron compounds, hence its presence may simply result from the latter stage of FeS preparation. The removal of minor impurities present in material by acid cleaning procedure used for pyrite preparation (Descostes et al., 2004) was avoided because such a treatment may cause a serious alteration of the troilite surface (Chirita and Descostes, 2006a). The specific surface area of unreacted FeS grains was determined by BET method to be 0.57 ± 0.03 m 2 g −1. Ferric chloride and hydrochloric acid used in the experiments were of reagent grade purity. pH values were measured using an Ingold InLab 400 combination electrode with a reported accuracy of ± 0.01 pH units. Before each measurement, the pH electrode was standardized against commercial pH buffers (Ingold). Deaerated acidic Fe 3+ solutions were used as oxidative media and prepared just before each experiment. The experiments were carried out in capped Erlenmeyer flasks vigorously shaken by hand every 10 min. The flasks were filled at the beginning of each experimental run with 250 mL of desired solutions made by diluting acidic (pH b1) 0.1 M FeCl3·6H2O solution to obtain an initial concentration of dissolved Fe 3+ ([Fe 3+]init) between 10 −4 and 10 −3 M. The initial pH was checked prior to each experiment and adjusted to the desired value by adding drops of concentrated NaOH or HCl solutions. Thereafter, the solutions were pre-equilibrated to the reaction temperature of 25 ± 1 °C. Dissolution experiments were initialed by adding an amount of 0.5 g FeS to the Fe 3+ solutions, followed by vigorous shaking to disperse the powder in the suspension. At specific moments, 5 mL aliquots of suspension were removed with a syringe connected to a 0.22 μm filter. The filtrates were analyzed for total dissolved iron ([Fe]total) by spectrophotometry (SPEKOL DDR spectrophotometer) using the 2,2′ dipyridyl method (λ = 522 nm) (Nacu et al., 1988). The uncertainty in measured [Fe]total is smaller than ± 5% and the determination limit equals 7 μM. [Fe(II)]total was quickly analyzed by a modified 2,2′ dipyridyl method, where the reduction of ferric iron by 10% hydroxylamine was not performed. The concentration of dissolved ferric iron ([Fe(III)]total) was estimated by subtracting

the concentration of ferrous iron ([Fe(II)]total) from the total concentration of iron ([Fe]total). An additional set of three experiments with solutions at initial (measured) pH of 3.0, 2.5 or 2.4 and [Fe 3+]init of 0.2 mM were conducted to measure the amount of dissolved sulfur ([S](aq)). For this measurement, 20 mL aliquots of suspension were withdrawn with a syringe connected to a 0.22 μm filter, fully oxidized with bromine water (Caldeira et al., 2010) and analyzed turbidimerically for dissolved sulfate using the BaSO4 method (Manescu et al., 1994). The quantification lower limit of turbidimetric method was 11 μM. Note that [S](aq) is a sum of soluble sulfur bearing species that can be oxidized by bromine water to sulfate (like S2−, SO32−, S2O32− or S4O62−) and SO42−. In four separated experiments, the supernatant Eh was monitored using a Pt electrode (ALASC Pt) coupled with a calomel reference electrode (ALASC ER 01) connected to a pH/Eh-meter (Jenway 3305). The response of Pt electrode was checked before each experiment in Zobell's solutions which were prepared from reagent grade salts (K4Fe(CN)6·3H2O, K3Fe(CN)6, and KCl) and double distilled water (Eh = 186 mV/SHE at 25 °C). Between each experiment, the Pt experiment was cleaned in concentrated HNO3. In order to verify the absence of Fe 3+ precipitation at initial pH 3.00 and 25 °C during the 240 min of experimental runs, control experiments in absence of FeS were performed at high [Fe 3+]init (10 −3 and 5 × 10 −4 M). In these control tests, samples of Fe 3+ solutions were filtered using a 0.22 μm membrane filter and [Fe]total was measured. [Fe 3+](aq) values were found to be constant during the 240 min of reaction. A dark-brown precipitate became visible only after several days (3–4 days). These observations indicate that Fe 3+ homogeneous precipitation in the conditions of our study, although thermodynamically possible (Stefansson, 2007) is kinetically hindered (Jolivet et al., 1994). Heterogeneous precipitation is still possible, however it can be ruled out by the absence of any increase in the amount of Fe hydroxide compounds detected by XRD or Raman spectroscopy. In order to keep the inert atmosphere, the reactor was flushed with a flow of nitrogen gas during pH adjustment and sampling. At the end of some experimental runs, the resulting solid residues were quickly separated from liquid phase by decantation, rinsed several times with distilled water, dried in an evacuated desiccator free of oxygen. The powders were analyzed by XRD on an X-Pert diffractometer (PANalytical) using Co Kα radiation and an X'Celerator for fast detection, a step of 0.02°2θ (Co), and a counting time of 30 s per point. Additional solid characterization was performed by point analysis of the #03 and #05 samples by micro-Raman spectrometry using a LabRam HR spectrometer (Horiba-Jobin Yvon). Analytical points were selected based on their visual aspects, and microRaman spectra were collected over the [50, 2000] cm −1 spectral range using an excitation wavelength of 785 nm and a counting time of ~ 600 s. 3. Results and interpretation 3.1. Impact of [Fe 3+]init on FeS dissolution in Fe(III)BAS The impact of dissolved Fe 3 + on FeS dissolution was examined by varying [Fe 3 +]init in the 0.1–1 mM range at 25 °C and initial pH 3. The [Fe]total variations with time are illustrated in Fig. 1. Immediately (~ 1 min) after FeS immersion in Fe 3 + solutions, [Fe]total decreased below the [Fe 3 +]init. Afterwards, [Fe]total roughly increased in all experiments, ruling out any precipitates of amorphous ferric hydroxide (Fe(OH)3(s)) as [Fe]total would have kept decreasing with time. The [Fe]total increase has been relatively smooth at [Fe 3 +]init b 4 × 10 − 4 M, whereas more fluctuating [Fe]total was observed for [Fe 3 +]init ≥ 4 × 10 − 4 M. It should be noted that, at pH 3 and [Fe 3 +]init ≤ 0.5 mM, [Fe]total values were still lower than [Fe 3 +]init even after 240 min of reaction (Table 1). However, for all

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Fig. 2. [Fe]total variation during FeS dissolution in absence of Fe3+.

Fig. 1. Effect of Fe3+(aq) concentration on the kinetics of Fe release into the supernatant at pH 3.00 and 25 °C.

[Fe 3 +]init values, similar differences (~ 0.16 mM) between the final [Fe]total and [Fe]total measured after 1 min were observed, an indication of a weak effect of [Fe 3 +]init on the rate of FeS dissolution in Fe(III)BAS. This finding is consistent with a previous study on pyrrhotite oxidation by Fe 3 + in anoxic solutions, for which a zero order with respect to [Fe 3 +](aq) was found (Belzile et al., 2004). Fig. 2 plots measured [Fe]total for control experiments at initial pH 2 and 3 and at [Fe 3+]init = 0 mM. At initial pH 2, [Fe]total is close to that observed in experiment carried out with [Fe(III)] = 0.2 mM. In contrast, at initial pH 3 [Fe]total was higher in absence of initial Fe(III). These observations suggest that although anoxic and oxidative dissolution of FeS can occur simultaneously, the overall process of FeS dissolution in Fe(III)BAS is not a simple sum of the two processes. This complex behavior probably results from control iron diffusion (i.e., FeS dissolution rate) by the SRL. In order to quantify the effect of FeS immersion in oxidant solutions the difference (Δ[Fe]) between [Fe 3+]init and [Fe]total measured after 1 min of contact time was determined (Table 2). Δ[Fe] values were observed to increase with [Fe 3+]init, and to level off around 200 μM for [Fe 3+]init ≥ 0.5 mM. This behavior is not consistent with the asymptotic behavior observed for the precipitation of hydrous ferric oxide (Stefansson, 2007; Caldeira et al., 2010). Rather, this Fe 3+ uptake would be explained by Fe 3 +(aq) adsorption on FeS surface (Descostes et al., 2010). The corresponding sorption isotherm can be easily derived from Δ[Fe] data. A plot of the amount of sorbed Fe 3+ (Γ) as a function of [Fe 3+] shows that a sorption plateau is reached for [Fe 3+]init ≥ 500 μM (Fig. 3). 3.2. pH effect on FeS dissolution in Fe(III)BAS

2 and 3. The decrease of initial pH from 3 to 2 results in a significant increase of the FeS dissolution rate (Fig. 4). Also, [Fe]total decreased immediately after FeS immersion in oxidant solution, except at pH 2, for which [Fe]total after 1 min of contact was found to be greater than [Fe 3+]init (Δ[Fe] = − 40 mM). In addition, at pH 2 the supernatant became milky after contact with FeS. The observed turbidity can be explained by the formation of colloidal sulfur (Brown and LeMay, 1983; Yao and Millero, 1996). 3.3. Analysis of dissolved sulfur Several distinct sulfur species such as SO42−, S2O32− or S4O62− can form upon oxidative dissolution of iron sulfides (Descostes et al., 2004; Lefticariu et al., 2006, 2007). Therefore, three additional experiments were performed at [Fe 3+]init = 0.2 mM by varying the pH (initial pH 2.4, 2.7, and 3) and measuring [S](aq) variations with time. Note that [S](aq) is a sum of soluble sulfur bearing species that can be oxidized by bromine water to sulfate (like S 2−, SO32−, S2O32− or S4O62−) and SO42−. However, no measurable amount of soluble sulfur could be detected using the turbidimetric method. Thus, either the sulfur products of FeS oxidation are mostly non-soluble (polysulfide or elemental sulfur) or the formed soluble species are sorbed at or near the solid surface. 3.4. Variations in pH and Eh Fig. 5 shows that pH values for all suspensions gradually increased within 240 min for all experiments. These results confirm that FeS dissolution in the presence of Fe 3+(aq) is a proton-consuming process, at least during the first 240 min. The variation of Eh in four acidic FeS suspensions is illustrated in Fig. 6. Our results show a continuous decrease of Eh with time. For example at initial pH 3 and [Fe 3+]init = 0.5 mM, Eh decreases from 899 mV (t = 0 min) to 662 mV (t = 240 min). Interestingly, no

The effect of varying initial pH on the FeS dissolution in Fe(III)BAS was investigated at [Fe 3+]init = 0.2 mM and initial pH values between Table 1 Difference Δ[Fe]aq between the final (t = 240 min) and initial (t = 0 min) [Fe]total. Run (DFeSFe(III))

Initial [Fe3+] / mM

Δ[Fe]aq / mM

#10 #12* #01 #13 #04 #09* #03 #05

0.1 0.2 0.2 0.3 0.4 0.5 0.5 1

0.060 0.016 −0.024 −0.064 −0.048 −0.096 0.100 0.072

* or ** = multiple experiments.

Table 2 Difference between the initial [Fe3+] and [Fe]total ~1 min of FeS contact with the oxidative solution. Run (DFeSFe(III))

Initial [Fe3+] / mM

Δ[Fe]a / mM

#10 #01 #12 #13 #04 #03 #09 #05

0.1 0.2 0.2 0.3 0.4 0.5 0.5 1

0.025 0.099 0.080 0.174 0.133 0.186 0.201 0.219

a Calculated as the difference between initial [Fe3+] and [Fe]total measured after 1 min of contact time.

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Fig. 3. Sorption isotherms of Fe3+ on FeS (initial pH 3.0, t = 1 min).

Fig. 5. Variation of pH in experiments of FeS dissolution in Fe(III)BAS solutions at 25 °C.

reliable redox potential could be measured after 30 min of reaction at initial pH 2 and [Fe 3+]init = 0.2 mM, as the potentiostat readouts froze with time. This was probably the result of a poisoning of the Pt electrode (Thomas et al., 1998). Because of this possible poisoning the electrode was cleaned with concentrated HNO3 before each experiment. In experiments carried out at initial pH 3, the variation of Eh could be related to the Fe 3+/Fe 2+ redox couple (Biernat and Robins, 1972; Wiersma and Rimstidt, 1984; Janzen et al., 2000; Huminicki and Rimstidt, 2009) by: 3þ

2þ

Eh ¼ 0:77 þ 0:0591lgð½Fe =½Fe Þ

ð1Þ

The total concentrations of aqueous ferric iron (i.e. [Fe(III)]total = [Fe+3](aq) +[FeCl+2](aq) +[FeOH+2](aq) +[Fe(OH)2+](aq) +[Fe2(OH)2+4] (aq)) have been computed from [Fe]total and Eh with the Visual MINTEQ 3.0 code, assuming equilibrium. Fig. 7 shows the variation of [Fe(III)]total with time at initial pH 3 and three [Fe3+]init values. It is worth noting that at pH 3, FeOH+2 is the major dissolved species for Fe3+. The calculated percentages (%) of [FeOH+2](aq) from [Fe(III)]total vary from 77.3% for [Fe3+]init =0.1 mM to 77.1% for [Fe3+]init =0.2 mM, and 76.5% for [Fe3+]init =0.5 mM. These values are in good agreement with similar concentration data published in the literature (Lam et al., 2005).

Fig. 4. Effect of pH on the kinetics of Fe release into the supernatant at 25 °C.

These computed [Fe(III)]total can be compared to experimental values obtained from dedicated experiments (#R01, #R02 and #R03) (Fig. 8). This comparison shows that the measured [Fe(III)]total is sometimes slightly lower than that calculated from the Eh values (especially for [Fe3+]init = 0.5 mM), but the general trend (decrease) of measured [Fe(III)]total is very similar to that computed. The discrepancies between experimental and calculated [Fe(III)]total suggest that thermodynamic equilibrium in the system was not reached, especially at the beginning of experiments. Also, the [Fe(III)]total decreases after 1 min of contact between FeS and Fe(III)BAS supports Fe(III) sorption on FeS surface. Finally, at initial pH 2 in the absence of Fe3+ (#R03), no aqueous Fe(III) was detected. This indicates that either lepidrocrocite present in the initial powder does not dissolve or the released Fe3+ is under limit of detection. As we can see from Figs. 7 and 8 [Fe(III)]total decreases during the experiments, and some solutions are even undersaturated with respect to amorphous Fe(OH)3(s) (dotted line) after 120 min of reaction. Thus the initial drop in [Fe]total cannot be explained by precipitation of ferric iron (hydr)oxide, again supporting surface adsorption. In addition, the subsequent increase in [Fe]total (Fig. 1) can be assigned to Fe 2+ release into solution (that decreases the solution Eh) rather than Fe 3+ desorption (that would increase solution Eh).

Fig. 6. Variation of Eh in experiments of FeS dissolution in oxidative solutions at 25 °C. Initial [Fe3+] and pH conditions are given in the figure.

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Fig. 7. Variation of computed [Fe(III)]total as a function of time at initial pH 3.00 and different initial [Fe3+].

3.5. Solid characterization X-ray diffractograms collected on the initial and Fe(III)-reacted powder did not differ significantly from each other (Fig. 9a). They revealed the predominance of FeS in the form of troilite, together with some metal iron and traces of lepidocrocite (γ-FeOOH). MicroRaman spectra obtained for the initial FeS powder emphasized the surface heterogeneity of the starting material (Fig. 9b). Several spectra displayed bands near 140, 207, and 463 cm −1 which can be attributed to α-S8 (Bourdoiseau et al., 2008), meaning that some elemental sulfur is already present. Other bands near 240 and 370 cm −1 were sometimes observed, and could be attributed to lepidocrocite (Dunn et al., 2000). In few other areas, Raman bands at 420 and 1004 cm −1 characteristics of sulfate groups, or near 1090 cm −1 typical of carbonate minerals were observed. Because such solids were not detected by XRD, they probably were limited to the outermost surface and would be easily dissolved as soon as FeS was contacted with acidic solutions. In several areas, the spectrum was almost feature-less, made only by two large bands near 250 and 750 cm −1, close to the positions expected for Fe(II,III) solids (Fig. 9). This would indicate simply that some poorly organized mixed Fe(II,III) solid is present on the surface. In contrast, the spectra collected on the #03 and #05 samples displayed essentially the Raman bands near 140, 207 and 465 cm−1 typical of α-S8, together with a few peaks near 240, and 370 cm−1 corresponding to lepidocrocite, probably as leftovers from the initial material. In conclusion, the initial material appears from the microRaman spectroscopy to be quite heterogeneous. With the onset of dissolution conditions, however, the surface appears to be gradually covered

Fig. 9. (a) X-ray diffractograms collected for the initial FeS powder and for powders reacted with Fe(III) solutions (#03 and #05). All peaks can be attributed to FeS (troilite), except ♦: lepidocrocite, and ▼: α-Fe. (b) Raman spectra collected for the initial FeS samples and powders reacted with Fe(III)-bearing solutions (samples #03 and #05).

with a material having an α-S8-like Raman signature. XRD, Raman spectroscopy and the results of the experiment #R03 suggest that the impurities present in starting material do not seem to interfere with troilite during the 240 min of the reaction. 4. Discussion

Fig. 8. Variation of measured [Fe(III)]total as a function of time at initial pH 2.00 and 3.00 and different initial [Fe3+].

The gradual increase in [Fe]total for reaction time >1 min at initial pH between 2 and 3 and [Fe 3+]init from 0.1 to 1 mM is a definite proof that FeS is dissolving. The rate of this dissolution seems to be essentially controlled by pH. This dissolution is likely the result of several distinct processes occurring upon contact between FeS and acidic oxidative solutions: (i) proton sorption and proton-promoted separation of Fe 2+ from FeS lattice, (ii) Fe 3+ sorption on FeS, (iii) Fe 3+ reduction to Fe 2+, (iv) sulfide oxidation to polysulfide, elemental sulfur and/or soluble sulfur-bearing compounds, (v) diffusion of Fe 2+ and soluble sulfur-bearing species diffusion into solution, and (vi) aqueous oxidation of released sulfur-bearing compound by Fe 3+ (Yao and Millero, 1996; Chirita and Descostes, 2006a). These

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processes can overlap in time and interfere with each other, and so assessing the real impact of each one is difficult. Hence, the challenge is identifying the prevailing process as a function of contact time between FeS and oxidant solutions. Sorption of Fe 3+(aq) on FeS surface is presumably the most rapid process, as revealed by the initial drop of [Fe]total observed at pH > 2. Moreover, the Fe 3+ sorption on FeS involves only cation binding to the surface by a few (possibly a single one) chemical bonds, or even by weak (electrostatic or Van der Waals) interactions. In contrast, processes such as Fe 2+ release into solution or Fe 3+ reduction require breaking some existing bonds, or transferring electrons, which is notoriously slow. Release of Fe 2+ from FeS lattice will prevail in the subsequent moments of mineral dissolution. This process includes the (rapid) proton sorption and then detachment of Fe 2+ from FeS lattice (the rate determining step of Fe 2+ release). The slowest processes involve electron transfer (redox reactions, Levi et al., 2003) and/or important structural transformations (ferric iron oxyhydroxides formation, Jolivet et al., 1994). The kinetic parameters of FeS reaction upon contact with Fe 3+(aq) can be derived from the model developed by Chirita and Descostes (2006b) for troilite oxidation by H2O2. According to this model, oxidative dissolution is initiated by the proton-promoted detachment of Fe 2+ from FeS lattice, followed by the oxidation of sulfide groups thus forming insoluble sulfur-bearing species, which form the SRL, and soluble sulfur-bearing species. The SRL in turn hinders diffusion of Fe 2+ from the FeS lattice into the supernatant (process akin to (v)). This model is based on two main assumptions. First, the oxidation of > SH2 groups (found at FeS/SRL interface) by an oxidant (here Fe 3+) is controlled by the oxidant diffusion across a sulfurrich layer (SRL), developing on FeS surface (Chirita and Descostes, 2006b). Second, protons can diffuse rapidly from the solution to FeS/SRL interface. 4.1. Kinetic parameters of FeS dissolution in Fe(III)BAS The experimental data were analyzed in order to determine the kinetic parameters of FeS dissolution in Fe(III)BAS. To this end, the increase in [Fe]total related to dissolution was calculated using the [Fe]total values at t = 1 min as the starting value. Fig. 10 shows that [Fe]total roughly increases as a linear function of time. This increase can be modeled using the model developed by Chirita and Descostes (2006b). The rate law of Fe2+ release from FeS lattice has the following expression: þ 1=ð1þyÞ ð1−yÞ=ð1þyÞ

½Fetotal ðtÞ−½Fetotal ðt ¼ 1minÞ ¼ ½FeΔ ¼ k′½H 

t

ð2Þ

where k′ and y are constants, and t is time (Chirita and Descostes, 2006b). Except for the experiments at initial pH 3 and [Fe 3+]init between 0.4 and 1 mM, increase in [Fe]total was correctly fitted by the model proposed by Chirita and Descostes (2006b), with coefficients of correlation in the range 0.85–1.00. The obtained rate constants, standard error of the slopes, time exponents (n = (1 − y) / (1 + y)) and reaction orders with respect to [H +] (1 / (1 + y)) are presented in Table 3. Note that this model could not account reasonably well the time evolution of [Fe]Δ values obtained from experiments at pH 3.0 and [Fe 3+]init in range 0.4– 1 mM. These findings suggest that the starting assumptions of the model proposed by Chirita and Descostes (2006b) are not valid at pH 3 and [Fe 3+]init ≥ 0.4 mM. A possible explanation is the formation of a thick SRL on FeS surface that stops the migration of oxidant from solution to mineral surface and implicitly the oxidative dissolution. Table 3 reveals that the time exponents (n values) calculated using Eq. (2) differ at pH 3 from those at lower pH. This dissimilarity may be related to the development of a SRL with properties changing with reaction time and experimental conditions (here pH) (Chirita and Descostes, 2006b). Chirita et al. (2008) suggested that, during the long-term (duration of 100 h) experiments of FeS oxidative

Fig. 10. Plots of [Fe]Δ versus tn. (a) n=0.3 and (b) n=1. [Fe]Δ values were obtained by the subtracting of [Fe]total at t=1 min from the [Fe]total measured at longer reaction times.

dissolution, the SRL undergoes a permanent process of polysulfide chains rearrangement. The rearrangement of polysulfide chains may produce new polysulfide species and elemental sulfur (Chirita et al., 2008). It is believed that the permeability of the SRL is closely related to stoichiometry and spatial arrangement of polysulfide chains (Chirita et al., 2008). For example, it is expected that short polysulfide chains may result in a more compact SRL, which would slow down FeS oxidative dissolution (Chirita et al., 2008). In contrast, in the case of the long polysulfide chains the possibilities of rearrangement are limited by steric obstructions producing a SRL with many cracks, permeable for Fe2+ or soluble sulfur-bearing reaction products (Chirita, 2009). 4.2. Fe 3+ sorption on FeS The sorption of Fe 3+(aq) to FeS surface displays a saturation plateau, a suggestion that it is limited by the availability of sorption sites. Assuming a single type of site, the variation of Γ (the concentration of sorbed Fe 3+ per unit mass of FeS) with [Fe 3+] can then be described by a simple Langmuir isotherm (Descostes et al., 2010), hence : h i 3þ Kads Fe   Γ ¼ Γ max 1 þ Kads Fe3þ

ð3Þ

where Γmax is the saturation concentration per unit mass of FeS, [Fe 3+] is the concentration of Fe3+ in solution, and Kads is the adsorption constant. A good fit could be obtained only by neglecting the experimental Γ corresponding to the lowest concentrations ([Fe3+]init = 0.1 and

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Table 3 Rate constants, time exponents and reaction orders with respect to [H+] obtained upon modeling of FeS dissolution in Fe(III)BAS. For details on the model see Chirita and Descostes (2006b). Run (DFeSFe(III))

Initial pH

Initial [Fe3+]/ M

Rate constant (k***) / μmol min−n

Time exponent (n)

Reaction order with respect to [H+]

#10* #12* #01 #08** #13 #04 #09* #03 #05 #06 #07 #02 #11 #R03 #R04

3.0 3.0 3.0 3.0 3.0 3.0 3.0 3.0 3.0 2.7 2.5 2.4 2.0 2.0 3.0

1 × 10−4 2 × 10−4 2 × 10−4 2 × 10−4 3 × 10−4 4 × 10−4 5 × 10−4 5 × 10−4 1 × 10−3 2 × 10−4 2 × 10−4 2 × 10−4 2 × 10−4 0 × 10−4 0 × 10−4

4.3 ± 0.2 4.6 ± 0.8 3.8 ± 0.2 4.0 ± 0.2 5.8 ± 0.2 n.d. n.d. n.d. n.d. 0.5 ± 0.1 0.5 ± 0.1 0.5 ± 0.1 2.8 ± 0.1 34.1 ± 0.9 5.5 ± 0.3

0.30 0.30 0.30 0.30 0.30 n.d. n.d. n.d. n.d. 1.0 1.0 1.0 1.0 0.5 0.5

0.65 0.65 0.65 0.65 0.65 n.d. n.d. n.d. n.d. 1.0 1.0 1.0 1.0 0.75 0.75

n.d. = non determinate;* or ** = multiple experiments; ***k (μmol min−n) differs from k′ (μM min−n) by units of measure.

0.2 mM) (Fig. 3). The obtained sorption parameters were Γmax = 1.48× 10−4 mol g−1 and Kads =4.64 × 10 3 L mol−1. It is important to note that these parameters must be carefully considered especially because Fe3+ sorption cannot be fully discriminated from Fe2+ dissolution. The overlap of the two processes is particularly important at low [Fe3+] when the amount of released Fe 2+ is high enough to alter the computed Γ values, and the sorption of Fe3+ is slow (rate is proportional to the concentration), so that the sorption equilibrium was not yet reached when the liquid samples were collected.

4.3. Mechanism of FeS dissolution in Fe(III)BAS The results obtained here corroborate a mechanism of FeS dissolution in the presence of Fe3+ derived from the one developed by Chirita and coworkers (Chirita and Descostes, 2006b; Chirita et al., 2008). It is reasonable to consider that the FeS dissolution in Fe(III)BAS will start with proton attack on mineral surface: þ

2þ

> FeS þ 2H ¼> H2 S−Fe

ð4Þ

3+

3þ

¼> FeS−Fe

3þ

ð5Þ

where the actual charge of the surface complex may be substantially lower by charge transfers in the solid. Note that this sorption process complies with the hypothesis of Janzen et al. (2000), i.e. that pyrrhotite oxidation by Fe3+(aq) follows an adsorption-type mechanism. The initial decrease of [Fe]total together with an insignificant decrease of Eh, indicates that Fe3+(aq) is still the major Fe species in solution during the first minutes of FeS contact with acidic Fe3+ solutions. The Eh rapidly decreases only at initial pH 2, a possible consequence of fast release of Fe2+ and dissolved sulfur species with low oxidation states. These species may speed up Fe3+(aq) reduction either in solution or at the surface. H2 S þ 2Fe

3þ ðaqÞ

¼ 2Fe

2þ ðaqÞ

þ 2H

þ ðaqÞ

0 ðsÞ

þS

ð6Þ

During the 240 min of the experiments, no dissolved sulfur-bearing compounds were detected. This observation points to either slow rates of sulfide oxidation, or to accumulation of the partially oxidized sulfur into insoluble products. In the second case, the reaction products of FeS dissolution in the presence of Fe 3+(aq) are made only by Fe 2+(aq) and a SRL of elemental sulfur (S0) and polysulfide chains (Sn2−) according to 3þ ðaqÞ

FeS þ ð2−2=nÞFe

3þ ðaqÞ

FeS þ 2Fe

2þ ðaqÞ

¼ 3Fe

0 ðsÞ

þS

ð8Þ

Reactions (7) and (8) show that the protons are not involved in overall mass balance equations of FeS oxidation into polysulfide and elemental sulfur. Therefore their concentration should not directly affect chemical equilibria corresponding to Eqs. (7) and (8). Protons, however, can speed up FeS oxidation by polarizing the sulfide groups and the surface (at the surface or in solution), thus facilitating the electron transfer from S(− II) to Fe 3 +. They can also accelerate the migration of Fe 2+ from FeS lattice into the supernatant by positively charging the mineral surface. During the reaction of FeS with oxidant solutions [Fe(III)]total decreases by more than one order of magnitude at initial pH >2 (and even by four orders of magnitude after 30 min of reaction at initial pH 2) and converges to the solubility limit of Fe(OH)3(s) (Figs. 7 and 8). The decreasing of oxidant concentration explains the quasisimilarity of measured dissolution rates and the absence of sulfur oxyanions (the products of advanced oxidation of S(− II)) after 240 min of contact between FeS and solutions.

+

As a cation, Fe directly competes with H for sorption on >FeS surface groups. This sorption is supported by our result and can be written > FeS þ Fe

and

¼ ð3−2=nÞFe

2þ ðaqÞ

2− ðsÞ

þ ð1=nÞSn

ð7Þ

5. Conclusions FeS dissolution and interaction with aqueous Fe 3+ were investigated at 25 °C and initial pH values from 2 to 3. The Fe 3+(aq) was initially removed from the solution, presumably by adsorption. A progressive increase in [Fe]total and pH values and an Eh decrease within 240 min of reaction time then occurred. [H +] is an important parameter of FeS dissolution in Fe(III)BAS. Thus, the reaction order of FeS dissolution in Fe(III)BAS with respect to [H +] is estimated to 0.65 at initial pH 3.0, and increases up to 1.0, when initial pH decreases. In contrast, Fe 3+ concentration has only a small effect on FeS dissolution rate in studied [Fe 3+]init range. Our results support a mechanism of FeS dissolution in Fe(III)BAS starting with the protonation of mineral surface and Fe 3+(aq) adsorption. Thereafter, the adsorbed protons accelerate Fe 2+ release (Chirita and Descostes, 2006a) from FeS matrix into solution. The adsorbed Fe 3+ may oxidize the sulfur moieties to insoluble species, presumably polysulfide (Sn 2−(s)) and elemental sulfur (S 0(s)). Subsequently, the formed SRL will control Fe 2 + migration into solution. The Fe 3+ sorption and gradual buildup of SRL on FeS surface make that the release of sulfur and heavy metals in oxidized swamps or acid mine drainage cannot be modeled by a simple dissolution mechanism. Clearly the complex interactions at mineral–solution interface between Fe 3+ and FeS or SRL have to be taken into account. The sulfur accumulated in the SRL limits the possible far-field impact of

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monosulfide dissolution, and possibly acts as future reaction sites for oxidoreduction processes with species such as Fe 3+. Acknowledgements The authors greatly appreciate support from IFA-CEA Program (Project C1-04). The manuscript benefited greatly from the editorial review and handling of Jeremy Fein and the helpful comments of two anonymous reviewers. References Basolo, F., Pearson, R.G., 1958. Mechanism of Inorganic Reactions: A Study of Metal Complexes in Solution. Wiley, New York. Belzile, N., Chen, Y.W., Cai, M.F., Li, Y., 2004. A review on pyrrhotite oxidation. Journal of Geochemical Exploration 84, 65–76. Benning, L.G., Wilkin, R.T., Barnes, H.L., 2000. Reaction pathways in the Fe-S system below 100°C. Chemical Geology 167, 25–51. Biernat, R.J., Robins, R.G., 1972. High-temperature potential/pH diagrams for the ironwater and iron-water-sulphur systems. Electrochimica Acta 17, 1261–1283. Bourdoiseau, J.A., Jeannin, M., Sabot, R., Remazeilles, C., Refait, P., 2008. Characterization of mackinawite by Raman spectroscopy: Effects of crystallization, drying and oxidation. Corrosion Science 50, 3247–3255. Brown, T.L., LeMay Jr., H.E., 1983. Qualitative Inorganic Analysis. Prentive-Hall Inc., New Jersey. Caldeira, C.L., Ciminelli, V.S.T., Osseo-Asare, K., 2010. The role of carbonate ions in pyrite oxidation in aqueous systems. Geochimica et Cosmochimica Acta 74, 1777–1789. Chandra, U., Sharma, P., Parthasarathy, G., 2011. High-pressure electrical resistivity, Mossbauer, thermal analysis, and micro-Raman spectroscopic investigations on microwave synthesized orthorhombic cubanite (CuFe2S3). Chemical Geology 284, 211–216. Chirita, P., 2009. Iron monosulfide (FeS) oxidation by dissolved oxygen: Characteristics of the product layer. Surface and Interface Analysis 41, 405–411. Chirita, P., Descostes, M., 2006a. Anoxic dissolution of troilite in acidic media. Journal of Colloid and Interface Science 294, 376–384. Chirita, P., Descostes, M., 2006b. Troilite oxidation by hydrogen peroxide. Journal of Colloid and Interface Science 299, 260–269. Chirita, P., Descostes, M., Schlegel, M.L., 2008. Oxidation of FeS by oxygen-bearing acidic solutions. Journal of Colloid and Interface Science 321, 84–95. Descostes, M., Vitorge, P., Beaucaire, C., 2004. Pyrite dissolution in acidic media. Geochimica et Cosmochimica Acta 68, 4559–4569. Descostes, M., Schlegel, M.L., Eglizaud, N., Descamps, F., Miserque, F., Simoni, E., 2010. Uptake of uranium and trace elements in pyrite (FeS2) suspensions. Geochimica et Cosmochimica Acta 74, 1551–1562. Dunn, D.S., Bogart, M.B., Brossia, C.S., Cragnolino, G.A., 2000. Corrosion under alternating wet and dry conditions. Corrosion 56, 470–481. Huminicki, D.M.C., Rimstidt, J.D., 2009. Iron oxyhydroxide coating of pyrite for acid mine drainage control. Applied Geochemistry 24, 1626–1634. Janzen, M.P., Nicholson, R.V., Scharer, J.M., 2000. Pyrrhotite reaction kinetics: Reaction rates for oxidation by oxygen, ferric iron, and for anoxic dissolution. Geochimica et Cosmochimica Acta 64, 1511–1522.

Jolivet, J.P., Henry, M., Livage, J., 1994. De la Solution a l'Oxyde, Condensation des Cations en Solutions Aqueuse, Chimie de Surfaces des Oxides, InterEditions et CNRS Editions. Lam, S.W., Chianga, K., Lim, T.M., Amal, R., Low, G.K.C., 2005. The role of ferric ion in the photochemical and photocatalytic oxidation of resorcinol. Journal of Catalysis 234, 292–299. Lasaga, A.C., 1984. Chemical kinetics of water-rock interactions. Journal of Geophysical Research 89, 4009–4025. Lefticariu, L., Pratt, L.M., Ripley, E.M., 2006. Mineralogic and sulfur isotopic effects accompanying oxidation of pyrite in millimolar solutions of hydrogen peroxide at temperatures from 4 to 150°C. Geochimica et Cosmochimica Acta 70, 4889–4905. Lefticariu, L., Schiminelinann, A., Pratt, L.M., Ripley, E.M., 2007. Oxygen isotope partitioning during oxidation of pyrite by H2O2 and its dependence on temperature. Geochimica et Cosmochimica Acta 71, 5072–5088. Levi, M.D., Gofer, Y., Cherkinsky, M., Birsa, M.L., Aurbach, D., Berlin, A., 2003. Electroanalytical features of non-uniformly doped conducting poly-3-(3,4,5-trifluorophenyl) thiophene films. Physical Chemistry Chemical Physics 5, 2886–2893. Luther III, G.W., Kostka, J.E., Church, T.M., Sulzberger, B., Stumm, W., 1992. Seasonal iron cycling in the salt-march sedimentary environment: The importance of ligand complexes with Fe(II) and Fe(III) in the dissolution of Fe(III) minerals and pyrite, respectively. Marine Chemistry 40, 81–103. Manescu, S., Cucu, M., Diaconescu, M.L., 1994. Chimia sanitara a mediului. Editura Medicala, Bucuresti. Mikhlin, Yu.L., Kuklinskiy, A.V., Pavlenko, N.I., Varnek, V.A., Asanov, I.P., Okotrub, A.V., Selyutin, G.E., Solovyev, L.A., 2002. Spectroscopic and XRD studies of the air degradation of acid-reacted pyrrhotite. Geochimica et Cosmochimica Acta 66, 4057–4067. Nacu, A., Mocanu, R., Onofrei, T., Gaburici, M., Simion, C., Matei, F., Balba, D., Dulman, V., Doniga, E., 1988. Chimie analitica si analiza instrumentala. Editura Institutului Politehnic Iasi, Iasi. Pratt, A.R., Nesbitt, H.W., 1997. Pyrrhotite leaching in acid mixtures of HCl and H2SO4. American Journal of Science 297, 807–820. Rickard, D., Luther III, G.W., 2007. Chemistry of iron sulfides. Chemical Reviews 107, 514–562. Schlegel, M.L., Bataillon, C., Benhamida, K., Blanc, C., Menut, D., Lacour, J.L., 2008. Metal corrosion and argillite transformation at the water-saturated, high-temperature iron–clay interface: A microscopic-scale study. Applied Geochemistry 23, 2619–2633. Smith, J., Melville, M.D., 2004. Iron monosulfide formation and oxidation in drainbottom sediments of an acid sulfate soil environment. Applied Geochemistry 19, 1837–1853. Stefansson, A., 2007. Iron(III) hydrolysis and solubility at 25°C. Environmental Science and Technology 41, 6117–6123. Thomas, J.E., Jones, C.F., Skinner, W.M., Smart, R., St, C., 1998. The role of surface sulfur species in the inhibition of pyrrhotite dissolution in acid conditions. Geochimica et Cosmochimica Acta 62, 1555–1565. Thomas, J.E., Skinner, W.M., Smart, R., St, C., 2001. A mechanism to explain sudden changes in rates and products for pyrrhotite dissolution in acid solution. Geochimica et Cosmochimica Acta 65, 1–12. Thomas, J.E., Skinner, W.M., Smart, R., St, C., 2003. A comparison of the behavior of troilite with other iron(II) sulfides; implications of structure. Geochimica et Cosmochimica Acta 67, 831–843. Wiersma, C.L., Rimstidt, J.D., 1984. Rates of reaction of pyrite and marcasite with ferric iron at pH 2. Geochimica et Cosmochimica Acta 48, 85–92. Yao, W., Millero, F.J., 1996. Oxidation of hydrogen sulfide by hydrous Fe(III) oxides in seawater. Marine Chemistry 52, 1–16.