Quantifying Fenton reaction pathways driven by self-generated H₂O₂ on pyrite surfaces

Abstract

Oxidation of pyrite (FeS₂) plays a significant role in the redox cycling of iron and sulfur on Earth and is the primary cause of acid mine drainage (AMD). It has been established that this process involves multi-step electron-transfer reactions between surface defects and adsorbed O₂ and H₂O, releasing sulfoxy species (e.g., S₂O₃²⁻, SO₄²⁻) and ferrous iron (Fe²⁺) to the solution and also producing intermediate by-products, such as hydrogen peroxide (H₂O₂) and other reactive oxygen species (ROS), however, our understanding of the kinetics of these transient species is still limited. We investigated the kinetics of H₂O₂ formation in aqueous suspensions of FeS₂ microparticles by monitoring, in real time, the H₂O₂ and dissolved O₂ concentration under oxic and anoxic conditions using amperometric microsensors. Additional spectroscopic and structural analyses were done to track the dependencies between the process of FeS₂ dissolution and the degradation of H₂O₂ through the Fenton reaction. Based on our experimental results, we built a kinetic model which explains the observed trend of H₂O₂, showing that FeS₂ dissolution can act as a natural Fenton reagent, influencing the oxidation of third-party species during the long term evolution of geochemical systems, even in oxygen-limited environments.

Introduction

The chemical and electrical properties of pyrite (FeS₂) – the most common iron disulfide in the Earth’s crust – make it a promising material for the construction of photovoltaic panels¹, as well as for wastewater remediation^2,3,4,5,6,7. In the last decades, a number of experimental and field studies have addressed the oxidative dissolution of pyrite, providing an accurate characterization of the process^{8,9,10,11,12,13,14,15,16,17}. Recent investigations have shown that, aside from the iron and sulfoxy species released during pyrite dissolution, ROS are always present as transient by-products, both under oxic and anoxic conditions^{18,19,20,21,22,23,24,25,26}, and providing an important pathway to oxidize third-party species through oxygen evolution reaction (OER)^{18,27,28,29,30}.

The rupture of S-Fe and S-S bonds over pyrite surface by mechanical fracture or during the dissolution process induces the formation of the non-stoichiometric defect sites (i.e., the dangling bonds) that trigger adsorption reactions^31,32,33. The iron surface species formed cause a reduction in the band gap of the pyrite from 0.86 electron volts (eV) in the bulk to 0.55 eV on the surface³⁴, catalyzing the dissociation of adsorbed oxygen (O₂) and water (H₂O) molecules and leading to the formation of H₂O₂ (equations 1,2,3,4)^19,25,26,35:

At the same time, pyrite releases Fe²⁺into solution than can catalyze the Fenton and Haber-Weiss reactions, leading to the generation of other ROS (equations 5)^4,22,25,26:

Although this sequence of reactions describes the basic features of H₂O₂ formation during pyrite dissolution (equations 1,2,3,4) and its further degradation to secondary ROS in solution (equations 5,6,7,8), various important aspects remain unclear, for instance: (1) can free radicals be formed by mechanism other than the photoactivation (i.e. by mechanical bond fracture or non-stoichiometry dissolution of pyrite)? (2) can the formation of H₂O₂ occur in strictly anoxic conditions overcoming the energy required to split the water molecule and the further release of O₂? and (3) can the process of ROS generation by pyrite be sustained over long periods of time?.

Real-time measurements of this process are made difficult by ROS reactivity and the subsequent redox transformations of iron and sulfur species. Spectroscopic and fluorescence methods are commonly used to measure the concentration of H₂O₂; however, these methods usually need the presence of dyes or chelating agents that are not well suited for the kinetic analysis of transient phases. In this study, we measured the real-time generation and decomposition of H₂O₂ and dissolved O₂ induced by pyrite surfaces under different boundary conditions (i.e., dark/light, oxic/anoxic) to investigate the kinetic role of ROS during pyrite dissolution. In addition, we analyzed both the chemical evolution of dissolved species and surface pyrite oxidation with spectroscopy (UV-Vis, XPS), Cyclic Voltammetry (CV) and high-resolution transmission electron microscopy (HTREM) and performed specific experiments to evaluate the persistence of H₂O₂ formation. Based on our results, we developed a kinetic model for the coupling between pyrite dissolution and H₂O₂ generation/degradation through the Fenton reaction. When combined with the observed trends this model leads to the definition of constraints on the overall process of ROS oxidation mechanism induced by pyrite surfaces.

Results

H₂O₂ and O₂ evolution: general pathway

Figure 1 shows the concentration of H₂O₂ as a function of time in aqueous suspensions of pyrite starting at circumneutral pH, under oxic and anoxic conditions. H₂O₂ concentration increased at the beginning of the experiment until a maximum value was reached, and decreased thereafter asymptotically towards a nearly stationary, residual value ([H₂O₂]>200 nM) still measurable at the end of the experiment (~22 h). This coupled generation-decay response was generally observed in every experiment, although there were variations in the particular shape of the curves and, some experiments showed characteristic shoulders or secondary maxima at intermediate stages of the process. The overall process followed a sigmoidal trend, suggesting a strong interaction between the H₂O₂ generation and degradation rates characteristic of an autocatalytic process. H₂O₂ formation and decomposition evolved more slowly under anoxic than under oxic conditions. We did not observe any correlation between pyrite loading and the H₂O₂ yield between experimental runs under our experimental conditions (i.e., 0.5–0.3 g/L, unbuffered neutral pH) because of several factors (e.g., kinks, steps, lattice anisotropy) can determine the variability of the reactive surface between samples^36,37,38,39.

Figure 1: H₂O₂ evolution induced by pyrite slurries in unbuffered water.

Inset: particle loading (g/L) (a) oxic conditions, (b) anoxic conditions.

The evolution of O₂ under oxic conditions was characterized by an asymptotic decrease followed by a slight increase and a steady stable period at the end of the experiment (Supplementary Fig. S1). The opposite trend was observed under anoxic conditions (i.e., an initial increase in O₂ followed by an asymptotic decrease). To clarify the role of O₂ over the formation of H₂O₂, we monitored simultaneously the evolution of O₂ and H₂O₂ under oxic and anoxic conditions (Fig. 2). In oxic conditions (Fig. 2a), the amount of H₂O₂ increased whereas O₂ was rapidly consumed at the beginning of the experiment. Under anoxic conditions O₂ and H₂O₂ formed concomitantly (Fig. 2b).

Figure 2: Simultaneous trends of O₂ and H₂O₂ obtained from pyrite slurries in unbuffered water.

The bottom figure shows the first two hours of the reaction, under (a) oxic (particle loading = 1.75 g/l) and (b) anoxic conditions (particle loading = 1.09 g/L).

During the experiments, pH became acidic under both oxic and anoxic conditions. Overall, pH dropped rapidly towards a nearly constant value approximately 2 to 3 pH units lower then at the initial pH value. The initial drop was more pronounced under anoxic than under oxic conditions (2 hours vs 10 hours, respectively), and in both cases the initial decline in pH values was accelerated with increasing pyrite loading (Supplementary Fig. S2).

The Fenton reaction in anoxic conditions: spectroscopic experiments

To determine if the Fenton reaction was actually the mechanism to degrade the H₂O₂ forming under anoxic conditions, we monitored the production of Fe³⁺and OH^· species in pyrite suspensions under anoxic conditions and in the dark (Fig. 3). First, we evaluated the intensity of the absorbance bands in the UV-Vis range of 260–700 nm to identify the formation of dissolved Fe³⁺-complexes that coexist in solution in the range 300–450 nm^40,41,42,43 (Fig. 3a). The first absorbance peak (λmax ~375 nm) evolved after three hours of reaction and was shifted to higher wavelengths (around 390–400 nm), as the reaction proceeded. Second, we monitored the formation of short-lived radicals, mainly OH^· by measuring the decrease in the light absorption spectrum (λmax = 590 nm) of Crystal Violet used as a dye probe⁴⁴ (Fig. 3b). The absorption spectrum of CV showed a rapid decline indicating that OH^· formation occurs concomitantly from the start-up of the experiment.

Figure 3: Spectroscopic monitoring of Fe³⁺and OH^· from pyrite slurries under anoxic conditions.

(a) Absorption bands detected from aqueous pyrite suspension in anoxic conditions and in the dark showing Fe³⁺-complexes signatures (particle loading = 0.28 g/L). The numbers inserted over the absorption bands show the reaction time. Spectra were registered in real time using a liquid waveguide capillary flow cell (LWCC; path length: 250 cm; WPI), connected to the batch reactor by a peristaltic pump; (b) Degradation of CV solution upon pyrite aqueous reaction under anoxic conditions and in the dark (particle loading = 0.12 g/L, [CV]₀ = 225.5 μM).

Tracking the reversibility of H₂O₂ generation: cycling experiments

We performed these experiments to better understand the long-term evolution of H₂O₂ during pyrite dissolution in an open system. In order to use the same physical pyrite particles and in the same geometry, pyrite microparticles were adhered over silicone strips to form a thin film which was placed on the internal wall of the reactor as in the spectroscopic experiments. We monitored the evolution of H₂O₂ until the observable amount of H₂O₂ attained a constant value or a concentration of zero (Fig. 4). At that point, we replaced the solution inside the batch reactor with distilled H₂O, and monitored the evolution of H₂O₂ again (we repeated this procedure twice). After a full cycle of H₂O₂ generation and decay, the formation of H₂O₂ was found to resume after the addition of fresh distilled H₂O, although the maximum H₂O₂ yield was consistently lower than in the previous cycle (Fig. 4a). Assuming a zero-order kinetic model, the initial observed rate of the H₂O₂ formation (k_{0 obs}) showed a slight decrease between the first and the second cycle, which was more prevalent in the third cycle (Fig. 4b). This decrease could be attributed to the formation of oxide patches on pyrite surface that partially block some of the iron reactive centers.

Figure 4: Recurrence of H₂O₂ formation by pyrite slurry after H₂O renewal.

(a) H₂O₂ evolution from the same pyrite slurry after renewing H₂O twice, in oxic-open conditions (pyrite load particle = 0.33 g/L, ΔpH₁cycle = 6.8–3.8, ΔpH₂cycle = 6.8–6.1, ΔpH₃cycle = 6.8–6.7); (b) Initial observed rate of H₂O₂ formation assuming a zero kinetic order rate.

Surface characterization (CV, XPS and HRTEM)

Results of cyclic voltammetry under anoxic conditions revealed the anodic peaks associated with H₂O oxidation by one electron and by two electron transfer (Supplementary Fig. S3). Interestingly, in the cathodic counterpart, the peak assigned to iron reduction is split into two peaks (0.1 V NHE and 0.2VNHE, respectively), indicating that a fraction of ≡Fe³⁺is reduced in a nearly spontaneous manner. The analyzed XPS spectra of the (001) face of pyrite after aqueous reaction in the presence and absence of dissolved O₂ (g) resulted in significant differences, showing major surface oxidation under oxic conditions with the subsequent formation of patches (Supplementary Fig. S4). In fact, XPS analysis of the (100) face of the pyrite after anoxic reaction only showed appreciable changes in the O1s orbital showing a shift to lower binding energies associated with an increase of the hydroxyl contribution (−OH) (Supplementary Fig. S5). In order to facilitate the identification of ≡S²⁻ and ≡Fe³⁺dangling bonds, one sample was ion-sputtered, which promoted the breakage of the S-S dimers (as in the grinding procedure) (Supplementary Fig. S6). The formation of iron oxidation patches was also observed with HRTEM. Figure 5 shows an image of the pyrite surface after 22 hours of reaction in a micromolar solution of H₂O₂ in absence of O₂(g). The presence of discrete oxidation patches was observed in the uppermost area of the micrograph (Fig. 5a). The FFT (Fast Fourier transform) proved that the lattice fringe spacing of low contrast clusters (~0.25 nm) were consistent with ferrihydrite nanocrystals (Fig. 5b), viewed down [001]. Additional HRTEM images (Supplementary Fig. S7) also showed the interplanar spacing characteristic of two-line ferrihydrite^45,46. The FFT of the crystalline part showed the interplanar spacing of pyrite but also of goethite, suggesting that ferrihydrite can be a precursor of goethite formation (Fig. 5c)⁴⁷.

Figure 5: Identification of iron oxide patches in pyrite surface by HTREM.

(a) HTREM image showing the formation of secondary products over a lamella of pyrite after 22 hours immersed in a micromolar solution of H₂O₂ under anoxic conditions; (b) FFT of ferrihydrite patches; (c) FFT of the crystalline part showing spacing characteristic of pyrite and goethite.

Fitting the experimental data: kinetic modeling of H₂O₂ evolution

Our experimental data suggests that the observed trend of H₂O₂ is the result of the coupling between the H₂O₂ generation by iron defect sites on the pyrite surface and the H₂O₂ decomposition by the Fenton reaction. Based on these results, we built a kinetic model that allowed us to analyze the process in terms of elementary reactions and to determine the specific rate constant of the H₂O₂ formation in oxic and anoxic conditions by fitting the experimental data. In summary, the model describes the net amount of H₂O₂ according with the following expression (Supplementary, Modeling approaches):

Where k₁, k₂ and k₃, represent the rate constants of each Fenton reaction step (Supplementary, Table S1). As shown in the equation, the H₂O₂ formation was calculated assuming a first order dependence on the reactive surface and, both iron defect sites act simultaneously, but in anoxic conditions the H₂O₂ produced by Fe²⁺sites is limited by the O₂ derived by the Fenton reaction. Minimization between experimental curves and the values calculated from equation 9 were made with a non-linear least squares approach using the Marquardt algorithm, with k_oxic and k_anoxic, the specific rate constants of H₂O₂ surface generation, as adjustable parameters (Table 1).

Table 1 Rate constants obtained by fitting experimental data.

Figure 6 presents the fitted values corresponding to the evolution of H₂O₂ in oxic and anoxic conditions together with the model derived results. Under oxic conditions (Fig. 6a), when pH was higher, the ferrous surface iron was oxidized by dissolved O₂ forming H₂O₂ according the first order rate. Deviations of the linearity started to occur due to a change in the reaction stoichiometry by the simultaneous decrease in pH and increase in dissolved Fe²⁺and H₂O₂. From this point, the Fenton reaction became important and the oxidation rate of Fe²⁺by O₂ slowed down leading to a minor O₂ consumption. As a result, reactions forming ROS -catalyzed by Fe²⁺and Fe³⁺− became effective for H₂O₂ degradation, which rapidly decreased, while iron species followed an opposite trend. Under anoxic conditions (Fig. 6b), the formation of H₂O₂ proceeded more slowly. The decomposition of H₂O₂ was also retarded because the concentration of dissolved [Fe²⁺] supplied to solution by pyrite dissolution was lower. The Fenton reaction was initiated when pH values dropped below 4.5 and the ratio H₂O₂/Fe²⁺increased, catalyzing the H₂O₂ decomposition similar to the oxic experiments. The analysis of ROS derived from the model in both oxic and anoxic conditions (Fig. 6c and d) shows that, the first reactive species formed was OH^· acting as a chain initiator, forming additional free radicals. The majority of the OH^· reacted with H₂O₂, generating HO₂^·, the conjugate acid of O₂^·−, which is the limiting reagent to assist the redox cycling of Fe³⁺/Fe²⁺forming O₂ and helping to buffer the pH drop

Figure 6: Experimental and fitting curves of H₂O₂ evolution, together with model-derived trends of pyrite dissolution products and secondary ROS.

(a) experimental and model curves of H₂O₂ together with Fe²⁺/ Fe³⁺, SO₄²⁻, pH and O₂ model trends, under oxic conditions (pyrite particle loading = 0.71 g/L, A/V₀ = 1 m²/L, pH₀ = 7, [O₂]₀ = 232 μM); and, (b) anoxic conditions (pyrite particle loading = 2.16 g/L, A/V₀ = 3.16 m²/L, pH₀ = 7, [O₂]₀ = 0 μM). Time evolution of ROS obtained with the previous models under: (c) oxic conditions; and (d) anoxic conditions. ^*The evolution of O₂^·− was omitted from the plot because model calculations gave negligible concentrations.

Contrary to the expectation, the maximum amount of H₂O₂ measured was significantly different between experimental runs, and independent of particle loading. This was likely because the effect of increasing the reactive surface was twofold: (i) an increase in the formation of H₂O₂, and (ii) an increase in the release of dissolved Fe²⁺, accelerating the decomposition of H₂O₂ by the Fenton reaction (Fig. 7). We plotted H₂O₂ and Fe²⁺evolution as predicted by the model for a set of different values of pyrite reactive surface area (in Fig. 7a). As a result of this coupling effect, higher values of reactive surface area tend to increase the amount of secondary ROS (i.e., OH^·/O⁻, HO₂^·/O₂^·−, O₂) rather than stabilize the presence of H₂O₂ in solution (in Fig. 7b)^2,48.

Figure 7: Influence of pyrite reactive surface area on H₂O₂ evolution.

(a) 3D plot showing the H₂O₂ and Fe²⁺simulated trends at different values of the surface area of pyrite per volume of H₂O (A₀/V from 1 to 2.5, m²/L). (b) Summation of ROS generated by the Fenton reaction for the lowest and the highest values of the reactive surface (Initial conditions: k_oxic = 2.33 × 10⁻⁴, k_anoxic = 6.46 × 10⁻⁹ mol/m²s –average values of the rate constants estimated by the model-, pyrite particle loading = 1 g/L, pH₀ = 7, [O₂]₀ = 0 μM).

Discussion

Aqueous suspensions of pyrite form H₂O₂, in the presence and absence of dissolved O₂, following a generation-decay trend (Fig. 1). The asymmetric shape observed in the experimental curves indicated that the apparent rate of H₂O₂ generation was significantly faster than the apparent H₂O₂ degradation. Previous studies^22,25 have suggested that in the presence of dissolved O₂, pyrite slurries form H₂O₂ by an electron transfer between the ferrous iron defect sites (Fe²⁺−S) and the adsorbed O₂ through a Habber-Weiss reaction, involving the O₂^·− radical formation (equations 1 and 2). This hypothesis is consistent with the observation that there is an inverse relationship between H₂O₂ and O₂ at the beginning of the oxic experiment (Fig. 2a). The mechanism of H₂O₂ formation in anoxic conditions remains more controversial. Some studies have suggested that in absence of O₂ the formation of H₂O₂ is driven by the oxidation of adsorbed H₂O catalyzed by the pyrite surface^4,19,24,26, whereas other studies have considered that this reaction is unlikely due to energetic considerations^15,25. A possible reconciliation comes from considering the presence of ≡Fe³⁺dangling bonds generated from the cleavage of S-S bonds. Briefly, the rupture of S-S bonds generate S⁻ species which are highly instable and, to compensate its charge disequilibrium, donate one electron to the nearest iron atom by the autoredox reaction^{26,31,32,49,50,51,52,53,54}:

These dangling bonds could decrease the energy requirements for the chemisorption of H₂O molecules, some authors even talk about “ferryl” iron dangling bonds ≡Fe^4+ 55, leading to H₂O splitting into H⁺and OH^·, with the subsequent formation of H₂O₂, as described by¹⁹ equations 3 and 4. Our experiments showed that H₂O₂ was generated in the absence of both dissolved O₂ and light (Fig. 1b). Besides, the estimated ratio of S/Fe (<2) in unreacted pyrite particles (Supplementary Fig. S8) indicated the presence of S vacancies.

Experiments with pyrite slurries under both oxic and anoxic conditions showed a sudden drop of pH during the first hours (Supplementary Fig. S2). This increase of H⁺concentration together with the progressive accumulation of H₂O₂ and Fe²⁺in solution is expected to trigger the Fenton reaction. UV-Vis spectroscopy confirmed the presence of Fe³⁺-complexes during anoxic pyrite dissolution (Fig. 3a) whose maximum wavelengths were compatible with the following reaction⁴¹:

The detection of both Fe³⁺-complexes and OH^· in solution (Fig. 3) indicated that the Fenton reaction occurred even in the absence of dissolved O₂, supporting the idea that the suite of Fenton reactions conditioned the decay period of H₂O₂ curves. Additionally, the formation of OH^· concomitantly with the formation of H₂O₂ from the start-up of the experiment under anoxic conditions (Fig. 3b) together with the splitting of the cathodic peak associated with the nearly spontaneous iron reduction (Supplementary Fig. S3), suggest that ≡Fe³⁺dangling bonds, actually catalyze the H₂O oxidation by one single electron, forming OH^· as described by equation 3.

In principle, the overall process of aqueous pyrite oxidation -under oxic conditions- involves only O_{2 (g)} consumption, according to:

However, considering a free radical mechanism, a simultaneous uptake and release of dissolved O₂ it to be expected, because it can be both a product and a reactant. In oxic experiments, O₂ concentration rapidly dropped and then reached a more constant value (Fig. 2a). Since the amount of Fe²⁺in solution is presumably low at these early stages, the consumption of O₂ during this period could be attributed to Fe²⁺oxidation at surface sites to form the superoxide anion (O₂^·−) and to the subsequent production of H₂O₂ (equations 1 and 2). An interesting result was the formation of O_{2 (aq)} as a by-product in anoxic experiments (Fig. 2b and Supplementary Fig. S1b). Although O₂ is not a direct product in the equation 3, since H₂O₂ is formed, the formation of O₂ can occur via several pathways such as the Fenton-like reaction, the “catalase-like reaction”²⁵ or via non-radical disproportionation of H₂O_2.

XPS and HRTEM results reported herein showed the formation of Fe³⁺−O patches over pyrite surfaces. The implication of these Fe³⁺−patches during aqueous pyrite oxidation is not clear. Some studies argued that the electron cycling of Fe²⁺and Fe³⁺between unoxidized and oxidized areas favors the electron transfer from the surface of the pyrite to molecular O₂, increasing the oxidation rate of pyrite⁵⁶. However, this mechanism is based on atmospheric oxidation of pyrite and does not explain the formation of H₂O₂. In contrast, other studies have suggested that the formation of these Fe³⁺−patches interrupts the redox cycling of Fe²⁺/Fe³⁺, thereby inhibiting the formation of H₂O₂ and decreasing the rate of pyrite dissolution²⁵. Our cycling experiments showed that most of the defect sites of the pyrite microparticles remained active at the end of the experiment, even when H₂O₂ was no longer observable in solution (Fig. 4). Moreover, it is expected that oxidized patches will desorb at low pH during the pyrite dissolution process. HRTEM images showed an appreciable increment of the pyrite alteration layer (Supplementary Fig. S9), suggesting that these Fe³⁺−patches failed to completely block the surface renewal during pyrite dissolution.

Based on our experimental results, we propose a model that explains the generation-decay trend of H₂O₂ in terms of a kinetic competition between (1) the formation of H₂O₂ by the self-oxidation of iron-sulfur cluster defect sites; and, (2) the degradation of H₂O₂ by the Fenton reaction triggered by pyrite dissolution. Accordingly, the evolution of H₂O₂ in solution can be summarized as follows:

At the beginning of the model run (t ≈ 0), the amount of H₂O₂ in solution was controlled by the surface reaction and the second term tended toward zero. However, when the Fenton reaction started, H₂O₂ was progressively degraded to secondary ROS, in solution (i.e., OH^·, HO₂^·/O₂^·−, O₂). At the end of the model run the amount of H₂O₂ remained constant within a value close to zero and equation 13, became:

The model allowed us to estimate the rate constants of H₂O₂ formation under both oxic and anoxic conditions (Table 1). Peak production of H₂O₂ was shifted towards longer times in anoxic conditions (Supplementary Fig. S10), pointing to slower oxidation kinetics of pyrite in the absence of O₂, a result also supported by our XPS analysis (Supplementary Figs S4 and 5) and by the sulfate and iron released by pyrite dissolution (Supplementary Fig. S11). When compared with experimental data, the model reproduces qualitatively well the observed trends for pH, O₂, Fe_total and SO₄²⁻, as these parameters are calculated using the overall equations of pyrite dissolution^17,57,58 (Fig. 6a–b). An additional feature of the model is that provides a way to reconcile the classical dissolution approach with the free radical assumption, opening a new pathway to analyze the flux of secondary ROS resulting from the degradation of H₂O₂ through the Fenton reaction (Fig. 6c–d). OH^· radicals were the first species produced, rapidly decaying to HO₂^·/O₂^·−. These ROS species counterbalances the decrease in pH and promoted the so-called “Fenton like” reaction, which resulted in the formation of O₂(g). Fe²⁺regenerated through this sequence, and also through the dissolution of pyrite, makes the Fenton reaction more efficient⁵⁹. Therefore, the disappearance of H₂O₂ in solution was likely due to a fast transformation into ROS, catalyzed by dissolved Fe²⁺, rather than by the cessation of the generation mechanism itself. This result is consistent with the continuous production of H₂O₂ and OH^·23,26,60 during pyrite dissolution. We hypothesize that as a result of the progressive acidification of the solution together with iron recycling by the Fenton reaction, the abiotic dissolution of pyrite microparticles can be considered as a natural and auto-catalytic Fenton reagent, useful to understanding long-term oxidation processes even in oxygen limited environments.

Methods

Sample preparation and characterization

Natural pyrite samples (Logroño, Spain) were ground to obtain particles with average diameter of 1.4 μm, Laser Diffraction Particle Size Analyzer (LS13320) and BET (Brunauer – Emmett – Teller) surface area of 1.46 m²/g (multi-point N₂ adsorption). Prior to use, freshly ground pyrite microparticles were washed by sonication in ethanol (96%) and hydrochloric acid (HCl 0.25 M), to remove organics and oxide surface coatings. Samples were then rinsed with deoxygenated deionized water (MilliQ) and dried in a low vacuum chamber purged with nitrogen (N₂) until used. Minor and trace elements in the acid washed samples were evaluated by scanning electron microscopy (SEM) using X-ray mapping (XRM) (Supplementary Fig. S8). X-ray diffraction (XRD) was used to assess the presence of secondary phases and to characterize the degree of structural disorder in the samples using the Rietveld method (Supplementary Fig. S12).

Batch kinetic experiments

The kinetics of H₂O₂ formation and degradation on pyrite slurries was investigated in batch reactors utilizing amperometric sensors for H₂O₂ (ISO-HPO-100, World Precision Instruments, Inc.) and dissolved O₂ (Unisense DK), and a glass electrode for pH determination (Vernier FPH-BTA). Batch reactors were designed to fit with microsensors, spectroscopic probes, ports for pyrite inlet under N₂ atmosphere, and valves. The valves allowed fluid circulation and solution sampling in a closed system configuration.

The production of ferric iron (Fe³⁺) and hydroxyl radical (OH^·) species by pyrite slurries in absence of dissolved O₂ were measured using UV-vis spectroscopy by monitoring absorption bands at specific wavelengths. To prevent data masking due to particle absorption, the pyrite microparticles were deposited onto silicone strips as a thin film adhered to the inner reactor walls. Spectroscopic data were collected with a fiber optic UV-Vis spectrometer (Black-comet, Stellarnet or USB4000, Ocean Optics) and acquired with the SpectraWiz^® or loggerpro3 codes. In addition, total dissolved iron and sulfate released during pyrite dissolution were measured at different time intervals using inductively coupled plasma mass spectrometry (ICP-MS) and ion chromatography (IC).

All experiments were made under continuous magnetic stirring (500 rpm), at room temperature (T° ∼ 22 °C) and in the dark, unless other conditions are specified. Due to the autocatalytic nature of Fenton chain reactions, kinetic experiments were carried out in unbuffered distilled water. Further details of the batch reactors, electrochemical sensors, experimental procedures and test analysis are given as supplementary material (Supplementary Figs S13–18).

Surface analysis

Cyclic Voltammetry (CV) was employed as an additional way to assess the sequence of redox reactions involving free radicals during H₂O adsorption at pyrite interface. CV was performed using Pt/Pyrite-Np’s/ Nafion©/ electrodes in N₂ purged solutions vs Ag/Cl 3 M KCl. A detailed description of the experimental set-up is given as Supplementary Material. X-ray photoelectron spectroscopy (XPS) was used to analyze the surface oxidation states of (001) faces of pyrite (single-crystal, ~1 cm² × 2 mm) after aqueous reaction in oxic and anoxic conditions. Platelets parallel to (001) faces were cut (1 cm² × 2 mm) and cleaned following the same procedure described above and allowed to react with water in oxic and anoxic conditions. Samples were dried and stored in N₂-filled tubes until introduction into the XPS vacuum chamber. The oxidation states were analyzed at three different stages of the dissolution process: (t₁) unreacted sample; the pyrite crystal was acid-washed to generate oxide-free surfaces by removing the normal contaminants, carbon, nitrogen, and oxygen (C, N, O) due to atmospheric oxidation; (t₂): sample after 22 hours immersed in oxic water. (t₃): sample after 22 hours immersed in anoxic water. Prior to this stepped analysis, the XPS spectra of a pyrite surface was analyzed after argon ion (Ar⁺) sputtering to verify the formation of Fe³⁺and S²⁻ surface species as occurs during mechanical fracture by preferential sulfur removal. XPS Spectra were collected from a take-off angle of 90° relative to the sample surface in a Thermo Scientific K-Alpha ESCA analyzer using monochromatic Al Kα (1486.6 eV) radiation and pass energies of 100 eV and 20 eV for survey spectra and narrow region spectra, respectively. Spectra were aligned by setting the C1s peak to a binding energy of 285 eV⁶¹. Deconvolution and fitting of experimental data were done with the XPSpeak4.2 software (http://www.phy.cuhk.edu.hk/~surface/XPSPEAK/). The Shirley method was used for background subtraction and the binding energies of the species identified were assigned using values taken from literature (Supplementary Tables S2 and 3).

High resolution Electron Transmission Microscopy (HTREM) analysis were performed to identify nano-domains of secondary oxidation products at the pyrite interface. HTREM images were acquired on a JEOL JEM- 3011 microscope with accelerating voltage of 200 kV using a Gatan Ultrascan 1000 CCD camera and Digital Microgrograph software. Data processing was performed with the GADDS and image-J codes. Pyrite lamellas were prepared using a focused ion beam (FIB) with a high resolution JEOL JSM-6700 f.

Kinetic model

The model was run using the computer code Copasi 4.8 (COmplex PAthway SImulator)⁶². We assume that the experimental trend of H₂O₂, in the presence and absence of O₂ (g), is shaped by three main processes: (i) the rate of H₂O₂ generation by the iron defect sites on the surface of pyrite particles; (ii) the production rates of Fe²⁺and SO₄²⁻ by pyrite dissolution; and, (iii) the kinetics of H₂O₂ degradation by the Fenton reaction. Fits to the experimental curves that describe the evolution of H₂O₂ were performed by using the Marquardt algorithm employing as adjustable parameters the rate constants of the H₂O₂ formation. A detailed explanation of the modeling set-up is included in the supplementary material.