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Electrochemical investigation of arsenic redox processes on pyrite Devon Renock, and James Voorhis Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b06018 • Publication Date (Web): 02 Mar 2017 Downloaded from http://pubs.acs.org on March 3, 2017

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Environmental Science & Technology

Electrochemical investigation of arsenic redox processes on pyrite

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Authors: Devon Renock* and James Voorhis

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Department of Earth Sciences

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Dartmouth College, Hanover, New Hampshire 03755, USA

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*corresponding author: Devon Renock Phone: +.1. 603.646.3101 Fax: +.1. 603-646-3922 email: [email protected]

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Abstract

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The specific Eh-pH conditions and mechanism(s) for reduction of arsenite, As(III), by pyrite

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is incompletely understood. A fundamental question is what role the pyrite surface plays in the

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reduction process. We used electrochemical methods to evaluate the reduction of As(III) under

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controlled redox conditions. As(III) reduction to elemental As(0) occurs on the pyrite surface

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under suboxic-reducing conditions and is promoted at low pH. Remarkably, As(III) reduction on

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pyrite occurs at similar potentials to those for reduction on platinum metal suggesting a similar

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mechanism/kinetics for these surfaces. The onset for As(III) reduction at pH ≤ 3.5 coincides

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with the potential for hydrogen electroadsorption on pyrite, E ~+0.1 V (vs. RHE).

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reactions show that As(III) is reduced on pyrite at the Eh-pH predicted by the electrochemical

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study. X-ray photoelectron spectroscopy reveals that, at pH ≤ 3.5, a significant fraction of the

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surface arsenic (30-60%) has an oxidation state consistent with As(0). Here, we propose a

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mechanism whereby atomic hydrogen that forms on ferric (hydr)oxide surface layers promotes

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As(III) reduction at low Eh and pH. Insights provided by this study will have implications for

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understanding the controls on dissolved As(III) concentrations in suboxic-anoxic environments.

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Batch

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TOC art

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Introduction

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Arsenic is the cause of human health problems worldwide, most notably in Southeast Asia

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where arsenic contamination in drinking water is considered the largest mass poisoning in human

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history.1 In many cases, the source of arsenic contamination is not anthropogenic, but rather it is

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naturally occurring in aquifer environments and is mobilized by redox reactions induced by

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human-influenced changes in the hydrological and/or redox conditions.2

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comprehensive understanding of arsenic speciation and redox processes is required to improve

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upon low temperature biogeochemical models that explain and predict dissolved arsenic

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concentrations in groundwater.

Thus, a more

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Arsenite, As(III), and arsenate, As(V), adsorption to iron (hydr)oxides under oxidizing

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conditions has been extensively investigated3-11, as well as the reductive dissolution of these



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phases accompanying changes from oxidizing to suboxic and reducing conditions.12-17 The latter

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processes involve the Fe(III)-Fe(II) and As(V)-As(III) redox couples and can result in the

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mobilization of As back into groundwater. Relatively less is known about As speciation and

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redox processes under low temperature reducing conditions. It has been shown for sulfate

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reducing conditions that dissolved As concentrations can be controlled by the solubility of

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arsenic sulfides such as realgar (As4S4) and orpiment (As2S3) as well as metastable and

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amorphous phases.18-20 However, there are reducing environments in arsenic contaminated

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aquifers where arsenic sulfides are undersaturated due to low sulfate levels.16 Under these

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conditions, the pH-dependent adsorption of arsenic to iron sulfides such as pyrite (FeS2) is a

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possible removal mechanism.3, 4, 21, 22 Additionally, O’Day et al.18 suggest that dissolved As(III)

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concentrations can increase under sulfate reducing conditions when iron sulfide precipitation is

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fast. In this case, As(III) builds up in the groundwater because of differences in the solubilities

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of arsenic and iron sulfides and the limited atomic substitution of As for iron or sulfur (i.e.,

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coprecipitation of sulfides) at low temperatures. Arsenic uptake by pyrite may be an important

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process controlling As concentrations under these conditions.

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Arsenic uptake by pyrite is not well understood due in part to the technical challenges of

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carrying out these investigations under extremely low O2 conditions in order to minimize

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oxidation of the surface.

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coverage by iron (hydr)oxide phases. A study by Sun, et al.

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mechanism for As(III) on pyrite is extremely sensitive to the extent of oxidation of the surface.

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The uptake of As(III) by pyrite is complex and has been shown to occur by the formation of

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outer-sphere surface complexes3 as well as the formation of amorphous and semicrystalline

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solids on the pyrite surface with compositions similar to arsenopyrite (FeAsS), and realgar and

The result is that pyrite surfaces often have varying degrees of


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demonstrated that the uptake

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orpiment.21, 22 The presence of a reduced As(-1) phase with composition similar to arsenopyrite

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suggests that reduction of As(III) by pyrite is thermodynamically feasible over a wide pH range.

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However, the specific Eh-pH conditions and mechanism(s) of arsenic reduction by pyrite under

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anoxic conditions is not well understood. Additionally, the extent to which As(III) reduction to

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elemental As(0) occurs on pyrite has not been previously investigated despite spectroscopic

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evidence showing that As(0) is stabilized on the surface of arsenopyrite.24

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The reduction of As(III) on a pyrite surface involves two half-reactions. The cathodic

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reaction is the reduction of As(III) which is coupled to an anodic reaction which can, for

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example, be the oxidation of Fe(II) (bulk or surface bound) to Fe(III) or the oxidation of S2-

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(dissolved or surface bound) to elemental S and polysulfides. Reduction of As(III) is thought to

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occur through consecutive one electron reduction steps, but the extent to which reduction forms

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As(0), As(-1), or As(II)-As(III) (arsenic sulfides) is not known. A fundamental question is what

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role the pyrite surface plays in the reduction process. Does the surface participate directly in the

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electron transfer reaction by providing electrons for As(III) reduction or indirectly by acting as

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either a heterogeneous catalyst or by providing a conductive medium for the transport of charge

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from a sorbed reductant such as H2S to an As(III) surface complex?25-28

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The objective of this study is to use electrochemical methods to evaluate the conditions under

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which As(III) reduction is occurring on pyrite and to assess the extent to which reduction forms

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stable As(0), As(-1), and As(II)-As(III) on the surface. We evaluate the behavior of the As(III)-

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As(0) redox couple on well-characterized electrode materials to elucidate similar processes that

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may be occurring on pyrite. Our studies approach is fundamentally different from the previous

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studies in that, by using a pyrite powder microelectrode, we can control the surface potential

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(thus oxidation states) of the reactive surface of pyrite and evaluate arsenic redox processes 5 ACS Paragon Plus Environment


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occurring in situ under various solution conditions. A combination of electrochemical methods,

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batch reactions, and surface spectroscopy is used to determine the energetics of As(III) reduction

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and provide insight into the reduction pathway(s).

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Experimental

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Materials

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Research grade pyrite (FeS2) was obtained from Wards Natural Science.

Pyrite was

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determined to be phase pure by X-ray diffraction analysis. The compositional purity of pyrite

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was determined by dissolving a known mass of solids (via microwave-assisted digestion) and

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determining Fe and S concentrations using an Inductively Coupled Plasma Optical Emission

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Spectroscopy (ICPOES, Thermo Iris Intrepid II). The analyzed pyrite was determined to be

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composed almost entirely of Fe and S and have nearly ideal Fe:S stoichiometry. All steps

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involved in handling pyrite and other oxygen sensitive materials were performed in an anoxic

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chamber (Coy) which provides a 0-5 ppm O2 atmosphere maintained by flowing 5% H2 in N2 gas

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over a Pd catalyst. Pyrite cubes were crushed and ground using an agate mortar and pestle.

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Ground pyrite powder was sieved to less than 20 µm using an ultrasonic sieve and stored in the

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dark in the anoxic chamber. Arsenic was added to the experimental solutions from a 0.1 N stock

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solution of sodium arsenite (NaAs(III)O2). Electrolyte solutions were prepared from sodium

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perchlorate monohydrate (NaClO4∙H2O). The pH was adjusted using dilute perchloric acid

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(HClO4), hydrochloric acid (HCl), or sodium hydroxide (NaOH). Sulfide was added as sodium

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sulfide (Na2S·9H2O).


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Powder microelectrodes

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Electrochemical evaluation of As(III) redox reactions on pyrite was done using powder

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microelectrodes (PME) which consist of a microcavity (ca. 20 µm deep, and 100 µm in diameter)

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that is embedded in the tip of a glass capillary and packed with pyrite powder. A schematic of a

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PME is shown in Fig. S1 (SI section 1). Details concerning PME fabrication and use can be

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found in the literature29-32 and described in detail in SI section 1.

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Electrochemical evaluation

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Cyclic voltammetry and linear sweep voltammetry were used to evaluate the energetics and

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kinetics of redox processes using a PME, platinum (Pt) disk (BASi; ∅ = 1.6 mm), and glassy

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carbon disk (BASi; ∅ = 1.6 mm) working electrode. A standard 3-electrode electrochemical cell

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was utilized employing a Pt-mesh counter electrode and a Ag/AgCl reference electrode (+0.197

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V vs. reversible hydrogen electrode, RHE). All potentials are reported as V versus RHE. The

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Ag/AgCl reference was measured against a freshly made Ag/AgCl double junction reference

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electrode prior to each experiment and periodically during cycling. Any measured potential drift

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in the reference was accounted for in the data.

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VersaSTAT 3 potentiostat was used for all electrochemical experiments. Electrolyte solutions

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(0.1 M NaClO4) were purged with high purity Ar gas for 60 minutes prior to scans. Argon was

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flowed above the solution during scans. Concentrations of 1×10-2 M to 1×10-4 M As(III) were

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used in voltammetric experiments unless noted otherwise.

A Princeton Applied Research (PAR)

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PME’s were immersed in electrolyte solution (T = 25 ˚C) and allowed to equilibrate for

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approximately 30 minutes until a stable open circuit potential (EOCP) was achieved (rate of

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change ∼0.005 mV/sec). EOCP is the potential measured between the PME and the reference 7 ACS Paragon Plus Environment


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electrode when no potential is applied.

The starting potential for cyclic voltammetry

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experiments was the EOCP. EOCP was monitored throughout the course of the experiments to

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evaluate changes in the redox state of the electrode. Cyclic voltammetry involves cycling the

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electrode potential of the PME between a positive and negative potential limit. Potential scans

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were conducted at a scan rate (v) of 50 mV/s unless noted otherwise. Potential cycling was done

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until a stable cyclic voltammogram, a measure of I(A) vs. E(V), was achieved for the electrode.

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Anodic stripping experiments were used to evaluate the redox activity of pyrite after

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reactions with As(III) in batch adsorption experiments (described below). Reacted pyrite powder

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was packed into a PME in the glovebox and evaluated similarly as previously described.

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Batch adsorption experiments

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Batch experiments were performed in the anoxic chamber. Argon-purged DI water (18 MΩ-

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cm) with an Eh range of -0.1 to -0.2 V was used to make all solutions. Eh was measured at the

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start of the batch reactions and at the end using an ORP probe (Symphony, VWR) calibrated

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with an ORP standard (Thermo Scientific Orion 967901). Ionic strength was fixed initially by

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adding 0.01M NaCl.

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suspension density of 5 g FeS2/L) was reacted with As(III) ([As(III)]initial = 1×10-2 M or 1×10-4 M

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As(III)) under a wide-range of pH. Before pyrite was added to the reacting vessel, pH was

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adjusted using HCl or NaOH. After adjusting the pH and adding the pyrite the vessels were

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sealed. The suspensions were constantly agitated using a rotating mixer over a 60-hour period in

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the dark. After 60 hours, the suspensions were centrifuged at 3,750 rpm (~30 minutes). The

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centrifuged solutions were unsealed, decanted and filtered (0.2 µm syringe filter) in the anoxic

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chamber. The pH and Eh of the decanted solutions were measured in the chamber immediately

As(III) was added from a 100 mM stock solution.


Pyrite (with a

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after removal from the solid. Vessels containing residual solid were resealed and frozen and

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reacted pyrite particles were freeze dried (Labconco Freezone 2.5) for 24 hours. Freeze-dried

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powders were stored in the chamber in the dark. Some of the dried powder was used for anodic

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stripping experiments and some was reserved for spectroscopic analysis.

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X-ray photoelectron spectroscopy (XPS)

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X-ray photoelectron spectroscopy was done on reacted pyrite powder from batch

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experiments.

X-ray photoelectron spectra were obtained with a Kratos Axis Ultra XPS

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(University of Michigan - EMAL) using a monochromatized Al Kα (1486 eV) X-ray source.

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The base pressure in the analysis chamber was ~10−9 Torr. Pyrite samples were prepared by

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mounting the powder on double-sided Cu tape in the glovebox. Air exposure was minimized

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during transfer to the XPS chamber. A charge-neutralizer filament was used for all samples to

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control charging of particles that were in poor contact with the stage. Peak shifts due to surface

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charging were taken into account by normalizing energies based on the adventitious carbon peak

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at 284.5 eV. Survey and narrow-scan XPS spectra were obtained using pass energies of 160 and

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20 eV, respectively. Survey scans were used to determine the average composition of the

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surface. The semiquantitative composition of the near-surface layers was calculated from the

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peak areas of the Fe(2p), S(2p), O(1s), and As(3d) peaks and normalized by their respective

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sensitivity factor.33 Narrow-scan spectra were obtained in order to determine oxidation states of

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As, Fe and S surface species.

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Raw spectra were fit using a least-squares procedure (Casa-XPS) with peaks of convoluted

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Gaussian (80%) and Lorentzian (20%) peak shape after subtraction of a Shirley-type baseline.

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All spectra were fitted with the least number of components that would give reasonable

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agreement with the measured spectra. Binding energies for the component peaks, e.g., for the 9 ACS Paragon Plus Environment


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different oxidation states and bonding environments of As, Fe and S were identified by

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comparison with literature values (see note in SI section 3). The full-width at half-maximum

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(FWHM) of the fitted peaks were also constrained within ranges reported in the literature. The

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As(3d) spectra was modeled by four doublet peaks (3d3/2 and 3d5/2) separated by a spin-orbit

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splitting of 0.70 eV. The As(3d3/2) peak was constrained to be 2/3 the area of the As(3d5/2) peak.

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The As(3d5/2) peak positions were as follows: As(V) at 45.2±0.1 eV, As(III)-O at 44.5±0.1 eV,

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As(III) sulfide at 43.4±0.1 eV, As(II) sulfide at 43.1±0.1, and elemental As(0) at 41.9 ±0.1 eV.22,

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24, 34

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residual standard deviation between the modeled and experimental data. Fitting parameters for

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the Fe(2p3/2) and S(2p) peaks are described in Table S1 (SI section 3). A Monte-Carlo method

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was used to estimate the standard deviation of the component peak areas used in the fitting

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procedure. The program applies artificial noise to each spectrum and calculates an error matrix

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to give the variance of each fitting parameter based on the fitting constraints used.

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Results and Discussion

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Redox characteristics from cyclic voltammetry

We did not include a peak for As(-1) (41.2 eV) because its inclusion did not minimize the

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The first step in this study was to compare the redox activity of As(III) on pyrite with

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electrode materials whose As(III)-As(0) redox behavior has been previously determined. Under

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the Eh-pH range of the cyclic voltammetry (CV) experiments, the stable dissolved arsenic

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species are arsenate (H3AsO4, H2AsO4- and HAsO42-) and arsenite (HAsO2) and the possible

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redox reactions include2, 35, 36:

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H3AsO4 + 2H+ + 2e- ↔ HAsO2 + 2H2O, E˚ = +0.56 V (vs SHE)


[1]

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HAsO2 + 3H+ + 3e- ↔ As0 + 2H2O, E˚ = +0.25 V (vs SHE)

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CVs of glassy carbon, Pt and pyrite PME were acquired under identical solution conditions.

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CVs for the glassy carbon electrode in solutions containing As(III) are featureless, exhibiting no

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peaks in current that are attributable to redox reactions [1, 2] in the applied potential range of -

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0.3 to +0.70 V and at pH 3 and 7 (not shown). The lack of any faradaic current indicates that

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electron transfer between the glassy carbon electrode and As(III) is kinetically inhibited under

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these conditions.

[2]

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Representative CVs for both pyrite and Pt are shown in Fig. 1(A-D) for pH 2.5 and 7.0. The

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baseline voltammetric characteristics for pyrite are described in previous studies and peaks

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corresponding to specific redox reactions are indicated in the figures.31, 37, 38 The CV of pyrite

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acquired at pH 7.0 with 0.01 M As(III) in solution (Fig. 1A) exhibits slightly increased positive

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currents on the positive-going, or “anodic”, scan as well as on the negative-going, or “cathodic”

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scan relative to pyrite without As(III). In contrast, at pH 2.5 (Fig. 1B) the CV for pyrite in a

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solution containing As(III) shows two distinct peaks (I and II in the figure). Specifically, a broad

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current peak (peak I in Fig. 1B) appears on the first cathodic scan suggesting reduction of As(III)

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on pyrite at these potentials. The maximum current of peak II is at E ~+0.4 V and is only

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observed on anodic scans that immediately follow cathodic scans that show peak I (note that

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scan direction reverses at E = -0.3 V). Moreover, the area (and peak height) of peak II increases

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with the amount of time the applied electrode potential is held at the peak potential of peak I (not

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shown). These results indicate that a reduced arsenic phase(s) forms on the surface during the

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initial cathodic scan and oxidizes at more positive potentials on the reverse anodic scan.



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The cyclic voltammetry of a Pt electrode (Fig. 1C-D) is compared with pyrite in order to

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identify specific redox reactions of As(III) that may be occurring under identical solution

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conditions to those in Fig. 1A-B. Our approach is based on the fact that the redox reactions and

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catalytic activity for As(III) on Pt are known.39-41 The baseline CV for Pt is included in both

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figures (1C-D) for comparison with the characteristic redox features identified in the figures.40, 42,

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43

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1D) show two peaks at similar potentials to peaks I and II in Fig. 1B for pyrite. These peaks are

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also identified as I and II in Figs. 1C-D. Additionally, there is a large peak (III) that appears at E

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> +0.6 V in the anodic scans of both Pt CVs. Previous studies of Pt in the presence of As(III)

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report three peaks at similar potentials to peaks I, II, and III in Figs 1C-D, albeit at pH < 1.39, 41, 44

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Peak III is ascribed to the oxidation of As(III) to As(V) on the anodic scan per reaction [1], and

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peaks I and II are ascribed to redox reactions between As(III) in solution and As(0) per reaction

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[2]. Specifically, the appearance of peak I on the initial cathodic scan is the reduction of

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dissolved As(III) to As(0) which deposits as a solid on the Pt surface. The appearance of peak II

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on the anodic scan is due to the oxidation of As(0) back to dissolved As(III). Cabelka et al.39

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reports that the oxidation of As(0) does not appear unless preceded by a scan to sufficiently

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negative potentials to generate As(0) on the electrode. These results are consistent with our

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results for peak I, II, and III for Pt (Figs. 1C-D), and thus we posit that peaks I and II are due to

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the forward and reverse reaction [2], respectively, and peak III is due to the oxidation of As(III)

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to As(V) (the reverse reaction of [1]).

Interestingly, CVs for Pt in solutions containing As(III) at pH 7.0 (Fig. 1C) and pH 2.5 (Fig.

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Considering the forward and reverse reaction [2], the separation of anodic and cathodic peaks

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for an electrochemically reversible couple (i.e., a redox couple in which both species rapidly

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exchange electrons with the working electrode) is approximated by the relationship 0.06 V/n at 12 ACS Paragon Plus Environment

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25 °C where n is the number of electrons transferred.45 If the peak positions for I and II are

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taken at the dashed lines in Fig. 1D (pH 2.5), the peak separation, defined as ∆Ep = EpII-EpI, is

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~0.4 V which is significantly larger than the 0.06 V predicted for a 1 e- reversible reaction

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(assuming that [2] proceeds via three independent electron transfer steps). Additionally, ∆Ep

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increases with increasing scan rate (not shown). Similar ∆Ep values and the dependence of ∆Ep

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on scan rate is observed at pH 7.0. These results indicate that reaction [2] occurs irreversibly on

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Pt at these pH values (i.e., the reaction is limited by relatively slow rate of electron transfer

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between Pt and sorbed As species).

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Remarkably, peaks I and II in the CV of a pyrite electrode at pH 2.5 occurs at similar peak

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potentials to I and II on Pt suggesting that reaction [2] may also be occurring on the pyrite

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surface. If true, the similarity of the ∆Ep value (~0.4 V) between pyrite and Pt indicates that

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pyrite exhibits similar catalytic activity to Pt at pH 2.5. The fact that peaks I and II are absent

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within the applied potential range for pyrite in the presence of As(III) at pH 7.0 may be due to

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the reduction mechanism (discussed later). Moreover, it is not possible to determine the nature

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of the reduced arsenic phase(s) on pyrite based on the CV. Therefore, when discussing the

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reduced phase forming on the pyrite electrode we will use the term “Asred”.

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There are two general pathways to explain the reduction of As(III) to Asred on the pyrite

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electrode: 1) electrons are transferred directly from pyrite to As(III) that is sorbed to the surface,

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and/or 2) hydrogen sulfide (H2S) generated from the reduction of oxidized S on pyrite (via a

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reaction such as [3]) reduces As(III) on the surface.

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S0 + 2H+ + 2e- ↔ H2S, E˚ = +0.14 V (vs SHE)


[3]


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Both pathways occurring on the electrode will result in the number of moles of electrons

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transferred during reduction on the cathodic scan being equal to the moles of electrons

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transferred during oxidation on the anodic scan. The total number of electrons transferred during

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reduction and oxidation [i.e., charge(coulombs) / F(coulombs/mole e-), where F is Faraday’s

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constant] was calculated by integrating the charge under peak I and II and subtracting the

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capacitive charge for each peak (SI section 2, Fig. S2). The ratio of moles e-(peak I) : moles e-

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(peak II) is determined to be approximately 1:1 by taking the average from multiple CVs. Thus,

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both pathways above are possible provided that all of the H2S generated (in pathway 2) is

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consumed by the reduction of As(III) on the surface.

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In order to test whether H2S generated at the electrode surface is reducing As(III), CVs of

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pyrite with and without As(III) were acquired by incrementally opening the negative potential

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limit with each successive cycle and determining whether As(III) reduction occurs at potentials

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where there is no significant H2S generation (Fig. 2A-B for pyrite at pH 3.5). A slightly higher

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pH (3.5) was used because the onset of dissolved H2S generation is shifted to more negative

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potentials based on the pH dependence of reaction [3]. For pyrite alone (Fig. 2A), a region of

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increased negative and positive current is observed as broad peaks between +0.2 V and -0.3 V

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similar to pH 2.5 (Fig. 1B), however the CV lacks the sharp increase in negative current

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associated with H2S production (Fig. 1B). Next, a CV was acquired under identical scan

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conditions with As(III) in solution (Fig. 2B). The shaded area (between 0 and +0.1 V) indicates

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electrode potentials that must be reached on the cathodic scan in order to reduce As(III) to Asred

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which is then oxidized on the reverse positive-going scan at E ~+0.34 V. All scans with cathodic

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limits more negative than this range show an Asred oxidation peak at E ~+0.34 V on the

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subsequent anodic scan. Additionally, the peak current for Asred oxidation increases with each 14 ACS Paragon Plus Environment

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subsequent cycle following a more negative cathodic limit. The formation of Asred on the

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surface occurs when the cathodic scan reaches potentials that coincide with the onset of the

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broad peaks at E < +0.2 V on pyrite (Fig. 2A).

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In contrast, an identical experiment at pH 7.0 shows that Asred does not appear until the

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cathodic scan reaches E < -0.4 V, or ~0.6 V more negative than at pH 3.5 (SI section 2, Fig. S3).

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The significantly different onset potentials for As(III) reduction at pH 3.5 and 7.0 appear to be

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related to the different redox processes occurring on the pyrite surface itself at these pH values.

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At pH 7, As(III) reduction occurs at potentials coinciding with the generation of dissolved

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sulfide (H2S or HS-). Whereas, we suggest that As(III) reduction at pH ≤ 3.5 may be associated

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with the redox reaction(s) responsible for the broad peaks that precede dissolved sulfide

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formation on the cathodic scan (Figs. 1B and 2A).

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The broad peaks at E < +0.2 V at pH 3.5 (Fig. 2A) occur in the same potential range as the

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known voltammetric features of pyrite attributed to the reversible electroadsorption/desorption of

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atomic hydrogen (H) by the reaction38, 46-48:

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FeS2 + H+ + e- ↔ FeS2H

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where negative current peaks on the cathodic scan correspond to H+ reduction and formation of

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atomic H on the surface (forward reaction [4]) and positive currents on the anodic scan

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correspond to oxidation of atomic H back to H+ (reverse reaction [4]). We confirmed that the

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broad peaks in our CVs at both pH 3.5 and 2.5 are due reaction [4] based on the behavior of the

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reversible peaks at various scan rates (SI section 2, Fig. S4). Reaction [4] is analogous to the

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well-known H electroadsorption/desorption reaction that occurs on noble metals such as Pt, Ru,

317

and Ir, as well as some transition metal oxides.

[4]

For example, H electroadsorption occurs



Page 16 of 30

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immediately preceding the generation of H2 gas on cathodic scans of Pt electrodes (see baseline

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Pt CV in Fig. 1D ).49, 50 The possible mechanistic implications of this are discussed below.

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However, in order to relate As(III) reduction on a pyrite electrode to As(III) reduction under

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natural conditions we conducted a series of batch reactions.

322 323

As(III) reduction on pyrite in batch laboratory experiments

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In a natural system, the activity of the redox active species that are present control the redox

325

reactions occurring on the pyrite surface. This situation is fundamentally different from a

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potentiostat applying a potential to a pyrite electrode and driving redox reactions in an

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electrochemical cell.

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pyrite’s voltammetric behavior to its redox behavior under more natural conditions. Thus, we

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conducted batch experiments by reacting pyrite particles with As(III) at solution Eh values

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between +0.05 V and +0.22 V and pH 2.5, 3.5 and 7.0 for a total of 60 hours. Following the

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batch reactions, reacted pyrite was rinsed thoroughly under identical Eh-pH conditions, freeze-

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dried, and then packed into a PME for analysis. Potentiodynamic scans were initiated from EOCP

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(~+0.2 V) shortly after the PME was immersed in solution.

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electrochemical results, we hypothesized that if any As(III) is reduced to Asred on pyrite during

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the reactions we should be able to strip it off the surface during the anodic scan by oxidizing it

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back to dissolved As(III) (see peak II in Fig. 1B). The anodic stripping experiment is based on

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similar techniques developed for electrochemical detection of As(III) using high surface area Au

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and Pt electrodes.51 Pyrite reacted at pH 2.5 without As(III) in solution shows a CV (SI section

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2, Fig. S6) that is similar to the baseline voltammetry in Fig. 1B indicating that no significant

However, a combination of these two methods is useful for relating


Based on our previous

Page 17 of 30


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redox change of the pyrite surface occurs during the 60-hour reaction. Remarkably, pyrite

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reacted with As(III) ([As(III)] = 10-2 M) at pH 2.5 shows an oxidation peak on the first cycle of

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the CV shown in Fig. 3. Oxidation peaks are also observed for [As(III)] = 10-4 M at both pH 2.5

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and 3.5 (not shown). The peak potential at E ~+0.45 V is ~0.1 V more positive than the

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potentials from the electrochemical study, but its presence suggests that Asred forms on the pyrite

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during the batch reactions. In all cases, the oxidation peak is only observed on the first cycle

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indicating that Asred is almost entirely removed from the surface on the first scan. These results

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are in contrast to pyrite from a pH 7.0 reaction ([As(III)] = 10-4 M) that shows no distinct peak

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on the first cycle of the CV (SI section 2, Fig. S7). Instead, the CV shows a steadily rising

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current on the first scan followed by steadily diminishing currents with each successive cycle.

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The current ratio (I1st cycle / I2nd cycle) at E = +0.3 V is ~2× greater for pyrite with As(III) compared

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to a CV without As(III) suggesting that the rising current may also be due to oxidation of Asred.

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Pyrite recovered from batch reactions was characterized by XPS to determine the oxidation

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state of adsorbed Fe, S and As on the surface. The distribution of Fe and S oxidation states in the

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unreacted pyrite is consistent with vacuum fractured pyrite52, whereas after anoxic reactions the

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distribution is relatively more reduced (SI section 3, Table S1). The binding energy of the

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As(3d5/2) and As(3d3/2) peaks span a range from 39.0 to 48.0 eV (a representative fitted XPS

357

spectrum is shown in SI section 3 Fig. S8). The total concentration of As at the surface ranged

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from 2-4 % (in atomic %) for all samples reacted below pH 3.5. The balance of the surface

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composition is predominantly Fe and S (Fe:S atomic ratio ~1:3) with

Electrochemical Investigation of Arsenic Redox ... - ACS Publications

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