Electrochemical Investigation of Arsenic Redox Processes on Pyrite

Subscriber access provided by University of Newcastle, Australia

Article

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

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

Environmental Science & Technology is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 30

1

Environmental Science & Technology

Electrochemical investigation of arsenic redox processes on pyrite

2

3

4

5

Authors: Devon Renock* and James Voorhis

6 7 8 9 10

Department of Earth Sciences

11

Dartmouth College, Hanover, New Hampshire 03755, USA

12 13 14 15 16 17 18 19 20 21 22

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

1 ACS Paragon Plus Environment


Page 2 of 30

23

Abstract

24

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

25

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

26

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

27

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

28

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

29

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

30

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

31

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

32

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

33

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

34

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

35

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

36

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

37

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

38 39 40


Batch

Page 3 of 30


41 42

TOC art

43

Introduction

44

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

45

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

46

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

47

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

48

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

49

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

50

upon low temperature biogeochemical models that explain and predict dissolved arsenic

51

concentrations in groundwater.

Thus, a more

52

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

53

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



Page 4 of 30

54

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

55

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

56

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

57

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

58

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

59

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

60

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

61

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

62

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

63

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

64

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

65

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

66

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

67

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

68

process controlling As concentrations under these conditions.

69

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

70

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

71

oxidation of the surface.

72

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

73

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

74

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

75

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

76

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


23

demonstrated that the uptake

Page 5 of 30


77

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

78

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

79

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

80

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

81

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

82

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

83

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

84

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

85

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

86

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

87

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

88

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

89

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

90

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

91

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

92

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

93

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

94

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

95

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

96

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

97

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

98

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

99

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


Page 6 of 30

100

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

101

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

102

and provide insight into the reduction pathway(s).

103

Experimental

104

Materials

105

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

Pyrite was

106

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

107

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

108

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

109

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

110

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

111

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

112

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

113

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

114

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

115

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

116

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

117

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

118

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

119

sulfide (Na2S·9H2O).


Page 7 of 30

120


Powder microelectrodes

121

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

122

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

123

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

124

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

125

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

126

Electrochemical evaluation

127

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

128

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

129

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

130

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

131

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

132

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

133

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

134

in the reference was accounted for in the data.

135

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

136

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

137

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

138

used in voltammetric experiments unless noted otherwise.

A Princeton Applied Research (PAR)

139

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

140

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

141

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


Page 8 of 30

142

electrode when no potential is applied.

The starting potential for cyclic voltammetry

143

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

144

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

145

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

146

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

147

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

148

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

149

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

150

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

151

Batch adsorption experiments

152

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

153

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

154

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

155

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

156

adding 0.01M NaCl.

157

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

158

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

159

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

160

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

161

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

162

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

163

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

Page 9 of 30


164

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

165

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

166

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

167

stripping experiments and some was reserved for spectroscopic analysis.

168

X-ray photoelectron spectroscopy (XPS)

169

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

170

experiments.

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

171

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

172

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

173

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

174

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

175

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

176

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

177

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

178

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

179

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

180

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

181

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

182

As, Fe and S surface species.

183

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

184

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

185

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

186

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


Page 10 of 30

187

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

188

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

189

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

190

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

191

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

192

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,

193

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,

194

24, 34

195

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

196

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

197

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

198

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

199

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

200

Results and Discussion

201

Redox characteristics from cyclic voltammetry

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

202

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

203

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

204

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

205

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

206

redox reactions include2, 35, 36:

207

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


[1]

Page 11 of 30


208

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

209

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

210

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

211

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

212

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

213

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

214

these conditions.

[2]

215

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

216

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

217

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

218

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

219

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

220

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

221

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

222

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

223

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

224

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

225

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

226

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

227

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

228

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



Page 12 of 30

229

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

230

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

231

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

232

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

233

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

234

43

235

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

236

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

237

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

238

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

239

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

240

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

241

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

242

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

243

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

244

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

245

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

246

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

247

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

248

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.

249

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

250

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

251

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

Page 13 of 30


252

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

253

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

254

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

255

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

256

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

257

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

258

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

259

between Pt and sorbed As species).

260

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

261

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

262

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

263

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

264

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

265

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

266

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

267

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

268

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

269

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

270

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

271

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

272

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


[3]


Page 14 of 30

273

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

274

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

275

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

276

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

277

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

278

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

279

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

280

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

281

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

282

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

283

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

284

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

285

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

286

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

287

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

288

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

289

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

290

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

291

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

292

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

293

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

294

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

295

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

Page 15 of 30


296

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

297

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

298

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

299

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

300

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

301

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

302

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

303

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

304

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

305

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

306

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

307

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

308

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

309

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

310

FeS2 + H+ + e- ↔ FeS2H

311

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

312

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

313

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

314

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

315

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

316

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

318

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

319

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

320

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

321

natural conditions we conducted a series of batch reactions.

322 323

As(III) reduction on pyrite in batch laboratory experiments

324

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

326

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

327

electrochemical cell.

328

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

329

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

330

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

331

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

332

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

333

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

334

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

335

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

336

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

337

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

338

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

339

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


340

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

341

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

342

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

343

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

344

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

345

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

346

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

347

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

348

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

349

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

350

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

351

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

352

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

353

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

354

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

355

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

356

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

358

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

359

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

Electrochemical Investigation of Arsenic Redox Processes on Pyrite

Recommend Documents