Mechanisms of Sb(III) Oxidation by Pyrite-Induced Hydroxyl Radicals

Subscriber access provided by NEW YORK MED COLL

Article

Mechanisms of Sb(III) oxidation by pyriteinduced hydroxyl radicals and hydrogen peroxide Linghao Kong, Xingyun Hu, and Mengchang He Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 25 Feb 2015 Downloaded from http://pubs.acs.org on February 25, 2015

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 32

Environmental Science & Technology

1

Mechanisms of Sb(III) oxidation by pyrite-induced hydroxyl radicals

2

and hydrogen peroxide

3 4

Linghao Kong, Xingyun Hu, Mengchang He*

5

State Key Laboratory of Water Environment Simulation, School of Environment, Beijing Normal

6

University, Beijing 100875, P.R. China.

7

8

* Corresponding author:

9

Mengchang He

10

State Key Laboratory of Water Environment Simulation

11

School of Environment

12

Beijing Normal University

13

Beijing 100875, P.R. China.

14

Tel : +86-10-5880 7172

15

Fax: +86-10-5880 7172

16

E-mail: [email protected] (MC He)

1

ACS Paragon Plus Environment


17

TOC ART

18

19 20

2


Page 2 of 32

Page 3 of 32


21

ABSTRACT

22

Antimony (Sb) is an element of growing interest, and its toxicity and mobility are

23

strongly influenced by redox processes. Sb(III) oxidation mechanisms in pyrite

24

suspensions were comprehensively investigated by kinetic measurements in oxic and

25

anoxic conditions and simulated sunlight. Sb(III) was oxidized to Sb(V) in both

26

solution and on pyrite surfaces in oxic conditions; the oxidation efficiency of Sb(III)

27

was gradually enhanced with the increase of pH. The pyrite-induced hydroxyl radical

28

(·OH) and hydrogen peroxide (H2O2) are the oxidants for Sb(III) oxidation. ·OH is the

29

oxidant for Sb(III) oxidation in acidic solutions, and H2O2 becomes the main oxidant

30

in neutral and alkaline solutions. ·OH and H2O2 can be generated by the reaction of

31

previously existing FeIII(pyrite) and H2O on pyrite in anoxic conditions. The oxygen

32

molecule is the crucial factor in continuously producing ·OH and H2O2 for Sb(III)

33

oxidation. The efficiency of Sb(III) oxidation was enhanced in surface-oxidized pyrite

34

(SOP) suspension, more ·OH formed through Fenton reaction in acidic solutions, but

35

Fe(IV) and H2O2 was formed in neutral and alkaline solutions. Under the illumination

36

of simulated sunlight, more ·OH and H2O2 were produced in the pyrite suspension,

37

and the oxidation efficiency of Sb(III) was remarkably enhanced. In conclusion,

38

Sb(III) can be oxidized to Sb(V) in the presence of pyrite, which will greatly

39

influence the fate of Sb(III) in the environment.

3



40

Page 4 of 32

INTRODUCTION

41

Antimony (Sb) is widely used in a variety of industrial products (approximately

42

1.4×105 tons each year), such as flame retardants, batteries, alloys, and catalysts.1 The

43

exploitation and utilization of Sb results in elevated concentrations of Sb in many

44

soils and waters, especially around mining and smelting areas2,

45

ranges.4 It is considered a priority pollutant by the Environmental Protection Agency

46

of the United States (USEPA) and the Council of the European Communities.5, 6 In the

47

natural environment, Sb(V) is the predominant species and exists as Sb(OH)-6 in oxic

48

environments, and Sb(III) primarily occurs as Sb(OH)3 and is more stable and toxic

49

under anoxic conditions between pH 2 and 10.1, 7 The mobility and toxicity of Sb

50

greatly depend on its oxidation state. Therefore, understanding Sb speciation in soil

51

and aquatic systems is important for assessing its fate and risk.

3

and at shooting

52

Sb(III) can react with relevant substances when it enters into the natural

53

environment. Several studies have previously been conducted on the oxidation of

54

Sb(III) to Sb(V), including dissolved molecular oxygen and hydrogen peroxide,8, 9

55

iodate,10 natural minerals(Fe and Mn oxyhydroxides, etc)11-15 and humic acid.16

56

Notably, the natural minerals exist widely and have large surface areas, acting as

57

strong oxidants in transforming Sb(III) to Sb(V). Mn oxyhydroxides can directly

58

oxidize Sb(III). Wang et al. (2012)13 investigated a synthetic manganite (γ-MnOOH)

59

for Sb(III) oxidation in aqueous suspensions, and the oxidation of Sb(III) by

60

manganite occurred on a time scale of minutes. Fe oxyhydroxides can catalyze Sb(III)

61

oxidation by dissolved oxygen14 or light11. Photo-induced oxidation of Sb(III) on 4


Page 5 of 32


62

goethite was investigated recently, the results indicated that significant amounts of

63

Sb(III) were oxidized to Sb(V) when the suspension was exposed to light.11 Hence,

64

oxidation on mineral surfaces is an important reaction in the environmental fate of

65

Sb(III).

66

Pyrite (FeS2) is the most abundant natural sulfur mineral in the earth's crust. A

67

series of complex reactions occur upon addition of pyrite to O2-free water; many free

68

radicals and oxidants, such as the hydroxyl radical (·OH) and hydrogen peroxide

69

(H2O2), are formed (Table 1). Numerous studies have reported the degradation of

70

various organic pollutants by pyrite-induced ·OH and H2O2 in the environment.17-24

71

Simultaneously, pyrite is also associated with stibnite; thus, the speciation of Sb may

72

also be influenced by pyrite, which is widespread in soils, waters and sediments.

73

Therefore, we deduce that pyrite-induced H2O2 and ·OH could oxidize Sb(III) to

74

Sb(V). However, to our knowledge, no studies have examined the reaction of Sb(III)

75

with pyrite or other sulfide minerals.

76

The present study aimed to investigate the reaction between Sb(III) and pyrite and

77

further study the reaction mechanisms under specific conditions (oxic, anoxic,

78

surface-oxidized pyrite (SOP) and simulated solar light). Understanding the

79

mechanism of Sb(III) oxidation by pyrite is helpful in clarifying the fate and

80

geochemical cycling of Sb in the environment.

81

MATERIALS AND METHODS

82

Reagents and Materials 5



83

Antimony trioxide, Sb2O3 (99.999%) and pyrite (naturally occurring mineral) were

84

obtained from Alfa Aesar (Johnson Matthey, USA). HCl, KOH, H2O2, FeSO4,

85

Fe2(SO4)3, tert-Butyl alcohol (TBA), thiourea, ascorbic acid and KBH4 are guaranteed

86

reagent (GR) grade, 2-Morpholinoethanesulfonic Acid (MES), K2B4O7, H2BO3, citric

87

acid monohydrate, phenanthroline, CH3COONH4 and CH3COOH are analytical

88

reagent

89

Chemical Reagent Co. Ltd. (Shanghai, China). Catalase from bovine liver was

90

purchased from Sigma-Aldrich (St. Louis, MO, USA). Sb(III) stock solution was

91

prepared by dissolving Sb2O3 in 2 mol L-1 HCl. All aqueous solutions were prepared

92

using deionized water treated with a Milli-Q water purification device (Millipore),

93

stored at 4°C and protected from light.

(AR) grade. The above chemicals were purchased from Sinopharm

94

Pyrite was ground and sieved to a 200 mesh powder, washed with 1 M HCl to

95

remove surface oxidation layers, rinsed three times with deoxygenated deionized

96

water and dehydrated with ethanol, and dried and stored in a closed vial under a N2

97

atmosphere. Then, the pyrite was exposed to air at room temperature for 2 weeks, and

98

SOP was obtained. BET surface areas of the pyrite and SOP were 0.214 ± 0.021 and

99

0.281 ± 0.025 m2 g-1, respectively. The pyrite and SOP were characterized by scanning

100

electron microscopy (SEM), energy dispersive spectrometer (EDS) (S-4800, Hitachi,

101

Japan) and X-ray diffraction (XRD) (X' Pert PRO MPD, PANalytical, Holland)

102

analysis (Figures S1). X-ray photoelectron spectroscopy (XPS) (ESCALAB 250Xi,

103

Thermofisher, Japan) was performed for evidence of the speciation of Sb (Figures S4)

104

and Fe (Figures S6) on the surface of pyrite and SOP, respectively. The XPS spectra 6


Page 6 of 32

Page 7 of 32


105

for Sb 3d5/2 and Fe 2p3/2 were fitted using XPSPEAK41 software.

106

Oxidation Reaction Experiment in the Dark

107

Sb(III) oxidation in pyrite suspension under oxic and anoxic conditions were

108

investigated at pH 3, 5, 7 and 9 under oxic and anoxic conditions with 20 µM Sb(III)

109

and 0.25 g L-1 pyrite. In anoxic experiments, solutions were purged with N2 (99.9%)

110

for at least 6 h to ensure that oxygen was excluded. The quenching experiments were

111

performed with 200 mM TBA or 0.075 g L-1 catalase, 20 mM Sb(III) and 0.25 g L-1

112

pyrite to investigate pyrite-induced ·OH and H2O2 on Sb(III) oxidation at pH 3, 5, 7

113

and 9. In addition, 20 µM Sb(III) oxidation experiments by 50 µM H2O2, 20 µM Fe(II)

114

or 20 µM Fe(III) were performed at pH 3, 5, 7 and 9 in the absence of pyrite,

115

respectively.

116

All solutions were prepared in a background of 12 mM KCl. Acidic solution (pH 3)

117

was prepared via small additions of 0.1 M HCl or KOH, circumneutral solution (pH 5

118

and 7) and alkaline solution (pH 9) were kept constant using 2 mM MES/KOH and

119

0.5 mM K2B4O7/H2BO3 buffer solution, respectively. The experiments with 200 mL

120

solution were performed in triplicate in 250 mL polyethylene bottles and shaken at

121

180 rpm at 25°C for 6 h, the suspensions were taken out at selected time intervals. For

122

the analysis of dissolved Sb(III) and total Sb (TSb) in the presence of pyrite, 5 mL

123

aliquots were removed with a plastic syringe and filtered through a 0.22 µm cellulose

124

acetate membrane immediately and then diluted into polyethylene bottles. For the

125

analysis of Sb(III) and TSb in H2O2, Fe2+ and Fe3+ experiments, 1 mL aliquots were

126

removed to 1ml 1:1 HCl to dissolve solid phase and then diluted into polyethylene 7



127

bottles. The samples was immediately analyzed.

128

Photo-oxidation Reaction Experiment

129

A diagram of the reactor is shown in Figures S2. Batch photo-induced oxidation

130

experiments were conducted in 500 mL beakers. During each experiment, 200 mL of

131

solution was placed in a beaker containing given concentrations of pyrite and Sb(III).

132

The initial Sb(III) concentration was 20 µmol L-1, and the pyrite concentration was

133

0.25 g L-1. The control experiment in the dark was conducted by aluminum foil

134

coverage. A 500 W long-arc xenon lamp (35 cm×1 cm) (Shanghai Jiguang Special

135

Lighting Appliance Factory, China) with a lamp cover used to irradiate visible light

136

(λmax = 400~700 nm) was placed above the beaker. The reaction temperature was

137

maintained at 25 ± 2°C by cold water circulation. Then, the suspension was exposed

138

to light for 2 h. Samples were taken at selected time intervals and filtered through

139

0.22 µm cellulose acetate membranes for analysis. All of the experiments were

140

conducted in triplicate.

141

Analysis

142

TSb and Sb(III) were selectively analyzed using a hydride generation-atomic

143

fluorescence spectrometer (HG-AFS) (AFS 9700 Titan Instrument Co. Ltd., Beijing,

144

China); 10% (m/V) thiourea and ascorbic acid were used to reduce Sb(V) to Sb(III).

145

The conditions for the hydride generation system were as follows: carrier solution 5%

146

(V/V) HCl, 2% (m/V) potassium borohydride and 0.5% (m/V) potassium hydroxide.

147

The concentration of Sb(III) was determined after adding 10% citric acid for masking

148

the Sb(V) signal with the same hydride generation conditions as the determination of 8


Page 8 of 32

Page 9 of 32


149

TSb. Sb(V) concentration was calculated as the difference between the concentration

150

of TSb and Sb(III). The calibration curves were obtained from 0 to 50 µg L-1 of TSb.

151

The relative standard deviation was 2%, obtained in 10 measurements performed per

152

day. The method detection limit was 0.01 µg L-1. Aliquot samples were diluted in 6

153

mol L-1 hydrochloric acid to fit within the working range.

154

For Fe(II) analysis, 2 mL of solution was taken and added to 10 mL centrifuge

155

tubes. Then, 0.2 mL of 2 g L−1 phenanthroline and 1 mL of CH3COONH4-CH3COOH

156

buffer solution were added to the centrifuge tubes. After dilution to 10 mL and 15 min

157

of color development, the absorbance of the samples at 510 nm was analyzed on a

158

spectrophotometer (Shimadzu UV-2500) using a 5 cm quartz cuvette.25, 26

159

Electron spin resonance (ESR) measurements

160

To verify the existence of hydroxyl radicals in pyrite solutions, ESR measurements

161

were performed on a Bruker EleXsys E500 EPR spectrometer (Germany) at room

162

temperature (25 ± 1°C) using 50 µL capillary tubes. Typical instrument settings were

163

as follows: sweep width 100.0 G, power attenuation 13.0 dB, modulation amplitude of

164

2 G and sweep time 40 s. The ESR signals were recorded from the samples for 320 s

165

at pH 3, 6 and 9 in the dark, and a 500 W high-pressure mercury lamp was used to

166

investigate the existence of hydroxyl radicals in the presence of light.

167

RESULTS AND DISCUSSION

168

Sb(III) oxidation by pyrite under oxic and dark conditions

169

To investigate Sb(III) oxidation in pyrite solutions, the kinetics experiments were 9



170

conducted at pH 3, 5, 7 and 9 (Figure 1). The concentration of Sb(V) increased with

171

increasing reaction time (Figure 1a), indicating that Sb(III) was oxidized to Sb(V) in

172

the presence of pyrite. Figure 1b further exhibits the Sb(III)/TSb ratio in pyrite

173

solution. The Sb(III)/TSb ratio decreased with increasing reaction time, and the

174

oxidation efficiencies of Sb(III) were gradually enhanced as the solution pH ranged

175

from acid to alkaline. However, in the absence of pyrite, no significant Sb(III)

176

oxidation was observed in 6 h reactions (Figure S3).

177

The concentrations of TSb in solution decreased with increasing reaction time

178

(Figure 1a), which illustrated that a portion of the Sb was adsorbed by pyrite; the

179

adsorption amount increased gradually as pH ranged from acidic to alkaline. To verify

180

the species of Sb adsorbed on pyrite, XPS analysis was performed at pH 5 after a 6 h

181

reaction. The Sb 3d5/2 spectrum was collected and analyzed to determine the state of

182

Sb on pyrite (Figure S4). Two peaks of 531.0 and 532.3 eV binding energies for Sb(III)

183

and Sb(V), respectively, were obtained from the original Sb 3d5/2. The majority of Sb

184

adsorbed on the pyrite surface was Sb(V); the quantitative analysis showed that 73.0%

185

of the Sb on pyrite existed as Sb(V). The experimental results confirmed that Sb(III)

186

was oxidized to Sb(V) in both solution and on the surface of pyrite.

187

Mechanisms on the formation of H2O2 and ·OH by pyrite in oxic aqueous solution

188

were investigated systematically in previous studies.27, 28 Sulfur-deficient defects are

189

common on pyrite.29, 30 In aerobic conditions, FeII(pyrite) can react with oxygen to form

190

FeIII(pyrite) and H2O2 at a sulfur-deficient defect site (Equ. 8, 9); then, FeIII(pyrite) reacts

191

with the adsorbed H2O at a sulfur-deficient defect site on the pyrite surface to form 10


Page 10 of 32

Page 11 of 32


192

·OH (Equ. 10), and two ·OH combine to produce H2O2 (Equ. 25). Therefore, we

193

believe that the pyrite-induced ·OH and H2O2 could oxidize Sb(III).

194

The ESR confirmed the existence of ·OH in pyrite suspension under aerobic

195

conditions, and the concentrations of ·OH, which were highest at pH 3 and lowest at

196

pH 9, decreased with increasing pH (Figure S5). Fe(II) can be released to the solution

197

in the process of pyrite oxidation (Equ. 5, 6). Approximately 2.5 µM Fe2+ was found

198

in pyrite solution at pH 3 after a 6 h reaction; thus, more ·OH could form via a Fenton

199

reaction in acidic solution, but ·OH formation was inhibited as pH increased (Equ. 15,

200

16). There was a discrepancy between •OH concentrations and Sb(III) oxidation at

201

different pH values. Hence, ·OH is likely not the only oxidant for Sb(III) oxidation,

202

and other oxidants must exist. Once ·OH formed, it could rapidly combine to form

203

H2O2 (Equ. 25). To investigate the effect of H2O2 on Sb(III) oxidation, Sb(III)

204

oxidation experiments by H2O2 were conducted (Figure S6). At pH 3 and pH 5, H2O2

205

could not react with Sb(III), but the oxidation of Sb(III) occurred at pH 7 and 9, and

206

the oxidation efficiency of Sb(III) was higher at pH 9 than pH 7. Previous studies

207

showed that H2O2 was an important and effective oxidant in the process of Sb(III)

208

oxidation in alkaline aqueous environment with half-lives of 117 and 11 days at pH 8

209

and pH 9 with H2O2 concentrations 10-6 M, respectively8, 9. According to our

210

experiments above, we speculated that ·OH and H2O2 could be the main oxidants for

211

Sb(III) oxidation in acid solutions and alkaline solutions, respectively, in pyrite

212

suspensions.

213

To further verify the mechanisms of Sb(III) oxidation by pyrite-induced ·OH and 11



214

H2O2, quenching experiments were conducted at pH 3, 5 and 9 by adding tert-butyl

215

alcohol (TBA) and catalase into the pyrite suspension (Figure 2). TBA is a widely

216

used scavenger of ·OH26,

217

peroxide to water and molecular oxygen but also efficiently scavenge ·OH,32 with the

218

constant rate of 2.6×1011 M-1 s-1. At pH 3, Sb(III) oxidation was completely quenched

219

by TBA. At pH 5, the contribution of ·OH in Sb(III) oxidation was approximately

220

73.1%, but at pH 9, this proportion dropped to approximately 5.9%. These results are

221

consistent with the ESR results, which implies that ·OH is the main oxidant for Sb(III)

222

at low pH. The addition of catalase conspicuously decreased the oxidation efficiency

223

of Sb(III) at pH 5 and 9, and the results implied that H2O2 is the key oxidant in Sb(III)

224

oxidation in pyrite suspensions in alkaline conditions.

225

Sb(III) oxidation by pyrite under anoxic and dark conditions

31

. Catalase can not only efficiently convert hydrogen

226

The experiments were conducted in anoxic conditions to investigate the effect of

227

oxygen on Sb(III) oxidation in pyrite suspension. The oxidation efficiencies of Sb(III)

228

under two conditions (oxic or anoxic) at pH 5 are shown in Figure 3. The oxidation

229

efficiency of Sb(III) was 37.9% in oxic conditions, whereas it dropped to 15.6% at pH

230

5 in anoxic conditions. The results indicated that the presence of oxygen enhances the

231

oxidation efficiency of Sb(III) in pyrite suspension. In anoxic conditions, the observed

232

rate constant of Sb(III) (Kobs = 0.0010 min-1) was slightly lower than that in oxic

233

conditions (Kobs = 0.0012 min-1) in the first 2 h, but an obvious inhibitory effect was

234

observed after 2 h (Kobs = 0.0001 min-1).

235

In oxic conditions, a cycle between FeII(pyrite) and FeIII(pyrite) exists (Equ. 8-10). 12


Page 12 of 32

Page 13 of 32


236

However, in anoxic conditions, the oxidation process of FeII(pyrite) to FeIII(pyrite) cuts off,

237

and the production of ·OH and H2O2 is inhibited. However, the oxidation of Sb(III)

238

was still occurring in anoxic conditions in the first two hours, implying that ·OH and

239

H2O2 could be generated by pyrite in the absence of molecular oxygen. The original

240

FeIII(pyrite) can still react with H2O to produce ·OH and H2O2, and once the original

241

FeIII(pyrite) is exhausted, the oxidation reaction stops. Hence, molecular oxygen is the

242

crucial factor in continuously producing ·OH and H2O2.

243

Comparison of SOP and pyrite on Sb(III) oxidation

244

Pyrite oxidation is an important process in the geochemical cycles of sulfur and

245

iron. In aerobic environments, pyrite is thermodynamically unstable and can be

246

oxidized by dissolved molecular oxygen, hydrogen peroxide, or dissolved ferric iron

247

(Equ. 5-7). Oxidation of pyrite surfaces may occur upon exposure to atmospheric

248

O2.33, 34 Our study also showed that the ferric and ferrous (hydr)oxides were formed

249

on the surface of SOP, which were verified by XPS analysis of Fe 2p spectra and EDS

250

analysis of SOP (More details were exhibited in Supporting information, Figure S1,

251

S7). To compare the difference and mechanism of Sb(III) oxidation by pyrite and SOP,

252

experiments were performed in solutions of 0.25 g L-1 pyrite and SOP containing 20

253

µM Sb(III) (Figure 4). The oxidation efficiencies of Sb(III) by SOP were enhanced

254

compared with pyrite at pH 3, 5, 7, and 9.

255

More Fe(II) is released to solution in SOP suspension. Approximately 19 µM Fe(II)

256

was found in the SOP solution at pH 3, whereas nearly no Fe(II) existed at pH 9. To

257

investigate the effect of Fe(II) on Sb(III) oxidation, experiments of 20 µM Sb(III) in 13



258

the presence of 20 µM Fe2+ were carried out at pH 3, 5, 7 and 9 under oxic and dark

259

conditions (Figure S8). At pH 3 and 5, the oxidation efficiencies of Sb(III) were not

260

obvious, but at pH 7 and 9, rapid oxidation of Sb(III) occurred, and the oxidation

261

efficiencies of Sb(III) increased with increasing pH. The rapid oxidation of Sb(III)

262

was due to the formation of H2O2 and Fe(IV) in the presence of Fe(II). Fe(II) was not

263

stable in alkaline condition, it was rapidly oxidized by dissolved oxygen, leading to

264

the formation of H2O2 (Equ. 11, 12). In addition, Fe(II) and H2O2 can form an

265

intermediate (INT), and the INT could lead to formation of Fe(IV) by the outer-sphere

266

electron-transfer reaction in neutral and alkaline solutions (Equ. 17-21). The results

267

further verified the effect of Fe(II) on Sb(III) oxidation. However, Sb(III) oxidation in

268

SOP suspension was also enhanced in acidic conditions, which was attributed to the

269

formation of ·OH through a Fenton reaction.

270

Notably, SOP has higher surface area, which is about 1.3 times higher than the

271

surface area of pyrite, this may be due to the large surface area of formed ferrous and

272

ferric (hydr)oxides on the surface of SOP.35 In order to meaningfully compare the

273

oxidation reactivity of these two forms of pyrite, Sb(III) oxidation by pyrite and SOP

274

was normalized by the surface area (Table 2). SOP showed higher oxidation activity

275

for Sb(III), the pyrite oxidation process leads to the transformation of inactive

276

Fe(II)-S to more active Fe(II)-O compounds on the surface of SOP.

277

The experiments of Fe(III) on Sb(III) oxidation was also investigated at pH 3, 5, 7,

278

and 9. Little Sb(V) was observed in a 6 h reaction in the presence of 20 µM Fe(III)

279

and 20 µM Sb(III) (Figure S9); thus, the influence of Fe(III) on Sb(III) oxidation was 14


Page 14 of 32

Page 15 of 32


280

negligible in this system.

281

Sb(III) oxidation by pyrite in the presence of light

282

Pyrite is a semiconductor with a small band gap of 0.95 eV, and it can be activated

283

by light.36,

37

284

sunlight on Sb(III) oxidation in pyrite suspension at pH 6 (Figure 5). A small amount

285

of Sb(III) oxidation was observed in the absence of pyrite under light. Remarkably,

286

Sb(III) was completely transformed to Sb(V) after 2 h of reaction in the presence of

287

pyrite under light. However, only 35.5% of Sb(III) was oxidized after a 2 h reaction in

288

dark. The oxidation efficiency of Sb(III) under light (Kobs = 0.0142 min-1) was

289

4.3-fold the rate under dark (Kobs = 0.0033 min-1). The efficiency of Sb(III) oxidation

290

by pyrite was greatly increased with exposure to simulated sunlight.

Experiments were performed to investigate the effect of simulated

291

When pyrite is illuminated, a photon can be captured and promote an electron from

292

a non-bonding iron d-level to the conduction band of pyrite, and hole-electron pairs

293

(e-, h+) in valence band are excited on the surface of pyrite. A hole (h+) migrates to the

294

pyrite surface, and adsorbed H2O dissociates to form ·OH in pyrite suspension (Equ.

295

22, 23).38 This process was confirmed by ESR (Figure S10), which showed an

296

increasing concentration of ·OH under the illumination of light in pyrite suspension.

297

Hence, the formation of ·OH via the photochemical process can contribute to the

298

promotion of the oxidation efficiency of Sb(III).

299

The oxidation efficiency of Sb(III) was altered less by the addition of TBA under

300

light than in the dark, indicating that •OH is not the only oxidant generated by light.

301

·OH could combine to form H2O2 (Equ. 25). A previous study demonstrated 15



302

increasing H2O2 production under the illumination of light in pyrite suspension.39

303

Proposed reaction mechanisms

304 305

Based on the above results, possible mechanisms of Sb(III) oxidation in the presence of pyrite were proposed as follows (Figure 6):

306

❶ ≡FeII induced ·OH and H2O2 on pyrite surface. ≡FeII can be oxidized by O2 to

307

generate H2O2 and simultaneously provide ≡FeIII at a sulfur-deficient defect site, and

308

then, the adsorbed H2O reacts with ≡FeIII, generating adsorbed ·OH and H2O2.

309

❷ FeII induced ·OH, H2O2 and Fe(IV) in pyrite solution. The dissolved FeII can

310

react with O2 to generate H2O2. In neutral and alkaline solutions, Fe(IV) can form

311

through the reaction of FeIIOH+ and H2O2. However, in acidic solutions, more ·OH

312

forms through the Fenton reaction.

313 314 315 316

➌ Light-induced ·OH and H2O2. More ·OH and H2O2 form in the illumination of light in pyrite suspension. The pyrite-induced ·OH, H2O2 and Fe(IV) can oxidize Sb(III) to Sb(V) (Equ.1-3).

Environmental implications

317

When redox speciation determinations of Sb are performed in aquatic systems, the

318

results have commonly shown that Sb(V) predominates and that Sb(III) is only found

319

at low concentrations in oxic conditions.1 Significant concentrations of Sb(V) are

320

found in anoxic waters,40 which contrasts with thermodynamic equilibrium

321

predictions. In oxic environments, Sb(V) may come from the catalytic oxidation of

322

Sb(III) by natural Fe oxyhydroxides11, 14 and humic acid16 in the presence of dissolved

323

oxygen or light, or from direct oxidation of Sb(III) by Mn oxyhydroxides13. Whereas 16


Page 16 of 32

Page 17 of 32


324

this study indicated that Sb(III) can also be catalytically oxidized by pyrite induced

325

hydroxyl radicals and hydrogen peroxide. The process can greatly increase the

326

mobility and reduce the toxicity of Sb. The oxidation efficiency of Sb(III) is

327

conspicuously enhanced under visible light, and it is 4.3-fold faster than in the dark.

328

This oxidation process may be more important in surface soils or waters. In anoxic

329

environments, pyrite-induced oxidation of Sb(III) can also occur, which could partly

330

explain the existence of Sb(V) in anoxic environments. Therefore, pyrite can greatly

331

influence the fate of Sb(III) and play an important role in the biogeochemical cycle of

332

antimony.

333

ACKNOWLEDGEMENTS

334

This work was supported by the National Science Foundation for Innovative Research

335

Group (Grant No. 51421065), the National Natural Science Foundation of China (21177011)

336

and the Major Science and Technology Program for Water Pollution Control and Treatment

337

(2012ZX07202002). We are grateful to the editors and the anonymous reviewers for their

338

numerous valuable comments and suggestions on our paper.

339

Supporting Information Available

340

Details of SEM, EDS and XRD analysis of pyrite and SOP, photochemical reactor,

341

Sb(III)/TSb changes in solution in the absence of pyrite, XPS analysis of Sb speciation, ESR

342

spectra in dark and light generated by pyrite, Sb(III) oxidation by H2O2, Fe2+ and Fe3+ and

343

XPS analysis of Fe 2p on the surface of SOP. The Supporting information is available free of 17



344

charge via the Internet at http://pubs.acs.org/.

345 346 347 348 349 350 351 352 353

REFERENCES (1)

Filella, M.; Belzile, N.; Chen, Y. W. Antimony in the environment: a review

focused on natural waters I. Occurrence. Earth-Sci. Rev., 2002, 57 (1-2), 125-176. (2)

He, M.; Wang, X.; Wu, F.; Fu, Z. Antimony pollution in China. Sci. Total Environ.,

2012, 421-422, 41-50. (3)

He, M. Distribution and phytoavailability of antimony at an antimony mining and

smelting area, Hunan, China. Environ. Geochem. Hlth, 2007, 29 (3), 209-219. (4)

Scheinost, A. C.; Rossberg, A.; Vantelon, D.; Xifra, I.; Kretzschmar, R.; Leuz, A.;

354

Funke, H.; Johnson, C. A. Quantitative antimony speciation in shooting-range soils by

355

EXAFS spectroscopy. Geochim. Cosmochim. Ac., 2006, 70 (13), 3299-3312.

356 357 358

(5)

Belzile, N.; Chen, Y. W.; Filella, M. Human exposure to antimony: I. Sources and

intake. Crit. Rev. Env. Sci. Tec., 2011, 41 (14), 1309-1373. (6)

Filella, M.; Belzile, N.; Chen, Y. W. Human exposure to antimony. II. Contents in

359

some human tissues often used in biomonitoring (hair, nails, teeth). Crit. Rev. Env. Sci. Tec.,

360

2012, 42 (10), 1058-1115.

361 362 363

(7)

Filella, M.; Belzile, N.; Chen, Y. W. Antimony in the environment: a review

focused on natural waters II. Relevant solution chemistry. Earth-Sci. Rev., 2002, 59, 265-285. (8)

Quentel, F.; Filella, M.; Elleouet, C.; Madec, C. L. Kinetic studies on Sb(III)

364

oxidation by hydrogen peroxide in aqueous solution. Environ. Sci. Technol., 2004, 38 (10),

365

2843-2848.

366

(9)

Leuz, A.; Johnson, C. A. Oxidation of Sb(III) to Sb(V) by O2 and H2O2 in aqueous 18


Page 18 of 32

Page 19 of 32


367 368 369 370 371 372 373 374

solutions. Geochim. Cosmochim. Ac., 2005, 69 (5), 1165-1172. (10)

Quentel, F.; Filella, M.; Elleouet, C.; Madec, C. Sb(III) oxidation by iodate in

seawater: A cautionary tale. Sci. Total Environ., 2006, 355 (1-3), 259-263. (11)

Fan, J.; Wang, Y.; Fan, T.; Cui, X.; Zhou, D. Photo-induced oxidation of Sb(III) on

goethite. Chemosphere, 2014, 95, 295-300. (12)

Xi, J.; He, M.; Wang, K.; Zhang, G. Adsorption of antimony(III) on goethite in the

presence of competitive anions. J. Geochem. Explor., 2013, 132, 201-208. (13)

Wang, X.; He, M.; Lin, C.; Gao, Y.; Zheng, L. Antimony(III) oxidation and

375

antimony(V) adsorption reactions on synthetic manganite. Chem. Erde-Geochem., 2012, 724,

376

41-47.

377

(14)

Leuz, A. K.; Monch, H.; Johnson, C. A. Sorption of Sb(III) and Sb(V) to goethite:

378

Influence on Sb(III) oxidation and mobilization. Environ. Sci. Technol., 2006, 40 (23),

379

7277-7282.

380 381 382 383 384 385 386

(15)

Belzile, N.; Chen, Y.; Wang, Z. Oxidation of antimony (III) by amorphous iron and

manganese oxyhydroxides. Chem. Geol., 2001, 174 (4), 379-387. (16)

Buschmann, J.; Canonica, S.; Sigg, L. Photoinduced oxidation of antimony(III) in

the presence of humic acid. Environ. Sci. Technol., 2005, 39 (14), 5335-5341. (17)

Bae, S.; Kim, D.; Lee, W. Degradation of diclofenac by pyrite catalyzed Fenton

oxidation. Appl. Catal. B-Environ., 2013, 134, 93-102. (18)

Benetoli, L.; Cadorin, B. M.; Baldissarelli, V. Z.; Geremias, R.; de Souza, I. G.;

387

Debacher, N. A. Pyrite-enhanced methylene blue degradation in non-thermal plasma water

388

treatment reactor. J. Hazard. Mater., 2012, 237, 55-62. 19



389 390 391

(19)

Che, H.; Bae, S.; Lee, W. Degradation of trichloroethylene by Fenton reaction in

pyrite suspension. J. Hazard. Mater., 2011, 185 (2-3), 1355-1361. (20)

Cohn, C. A.; Borda, M. J.; Schoonen, M. A. RNA decomposition by pyrite-induced

392

radicals and possible role of lipids during the emergence of life. Earth Planet. Sc. Lett., 2004,

393

225 (3-4), 271-278.

394

(21)

Cohn, C. A.; Laffers, R.; Simon, S. R.; O'Riordan, T.; Schoonen, M. A. A. Role of

395

pyrite in formation of hydroxyl radicals in coal: possible implications for human health. Part.

396

Fibre Toxicol., 2006, 3, 16-16.

397

(22)

Wang, W.; Qu, Y. P.; Yang, B.; Liu, X. Y.; Su, W. H. Lactate oxidation in pyrite

398

suspension: A Fenton-like process in situ generating H2O2. Chemosphere, 2012, 86 (4),

399

376-382.

400 401 402 403 404 405 406

(23)

Weerasooriya, R.; Dharmasena, B. Pyrite-assisted degradation of trichloroethene

(TCE). Chemosphere, 2001, 42 (4), 389-396. (24)

Zhang, Y. L.; Zhang, K.; Dai, C. M.; Zhou, X. F. Performance and mechanism of

pyrite for nitrobenzene removal in aqueous solution. Chem. Eng. Sci., 2014, 111, 135-141. (25)

Zuo, Y. Kinetics of photochemical/chemical cycling of iron coupled with organic

substances in cloud and fog droplets. Geochim. Cosmochim. Ac., 1995, 59 (15), 3123-3130. (26)

Xu, J.; Li, J.; Wu, F.; Zhang, Y. Rapid Photooxidation of As(III) through Surface

407

Complexation with Nascent Colloidal Ferric Hydroxide. Environ. Sci. Technol., 2014, 48 (1),

408

272-278.

409 410

(27)

Cohn, C. A.; Mueller, S.; Wimmer, E.; Leifer, N.; Greenbaum, S.; Strongin, D. R.;

Schoonen, M. A. A. Pyrite-induced hydroxyl radical formation and its effect on nucleic acids. 20


Page 20 of 32

Page 21 of 32


411 412

Geochem. T., 2006, 7 (3). (28)

Borda, M. J.; Elsetinow, A. R.; Strongin, D. R.; Schoonen, M. A. A mechanism for

413

the production of hydroxyl radical at surface defect sites on pyrite. Geochim. Cosmochim. Ac.,

414

2003, 67 (5), 935-939.

415

(29)

Guevremont, J. M.; Strongin, D. R.; Schoonen, M. Photoemission of adsorbed

416

xenon, X-ray photoelectron spectroscopy, and temperature-programmed desorption studies of

417

H2O on FeS2(100). Langmuir, 1998, 14 (6), 1361-1366.

418

(30)

Guevremont, J. M.; Strongin, D. R.; Schoonen, M. Effects of surface imperfections

419

on the binding of CH3OH and H2O on FeS2(100): using adsorbed Xe as a probe of mineral

420

surface structure. Surf. Sci., 1997, 391 (1-3), 109-124.

421

(31)

Wu, F.; Deng, N. S. Photochemistry of hydrolytic iron (III) species and

422

photoinduced degradation of organic compounds. A minireview. Chemosphere, 2000, 41 (8),

423

1137-1147.

424

(32)

Schoonen, M. A. A.; Harrington, A. D.; Laffers, R.; Strongin, D. R. Role of

425

hydrogen peroxide and hydroxyl radical in pyrite oxidation by molecular oxygen. Geochim.

426

Cosmochim. Ac., 2010, 74 (17), 4971-4987.

427

(33)

Todd, E. C.; Sherman, D. M.; Purton, J. A. Surface oxidation of pyrite under

428

ambient atmospheric and aqueous (pH=2 to 10) conditions: Electronic structure and

429

mineralogy from X-ray absorption spectroscopy. Geochim. Cosmochim. Ac., 2003, 67 (5),

430

881-893.

431 432

(34)

Chandra, A. P.; Gerson, A. R. The mechanisms of pyrite oxidation and leaching: A

fundamental perspective. Surf. Sci. Rep., 2010, 65 (9), 293-315. 21



433

(35)

Guo, X.; Wu, Z.; He, M.; Meng, X.; Jin, X.; Qiu, N.; Zhang, J. Adsorption of

434

antimony onto iron oxyhydroxides: Adsorption behavior and surface structure. J. Hazard.

435

Mater., 2014, 276, 339-345.

436

(36)

Ferrer, I. J.; Nevskaia, D. M.; Delasheras, C.; Sanchez, C. About the band-gap

437

nature of FeS2 as determined from optical and photoelectrochemical measurements. Solid

438

State Commun., 1990, 74 (9), 913-916.

439 440 441 442 443

(37)

Liu, S. L.; Li, M. M.; Li, S.; Li, H. L.; Yan, L. Synthesis and

adsorption/photocatalysis performance of pyrite FeS2. Appl. Surf. Sci., 2013, 268, 213-217. (38)

Schoonen, M.; Elsetinow, A.; Borda, M.; Strongin, D. Effect of temperature and

illumination on pyrite oxidation between pH 2 and 6. Geochem. T., 2000, 1 (4), 23. (39)

Borda, M. J.; Elsetinow, A. R.; Schoonen, M. A.; Strongin, D. R. Pyrite-induced

444

hydrogen peroxide formation as a driving force in the evolution of photosynthetic organisms

445

on an early earth. Astrobiology, 2001, 1 (3), 283-8.

446

(40)

Chen, Y.; Deng, T.; Filella, M.; Belzile, N. Distribution and early diagenesis of

447

antimony species in sediments and porewaters of freshwater lakes. Environ. Sci. Technol.,

448

2003, 37 (6), 1163-1168.

449

(41)

Leuz, A. K.; Hug, S. J.; Wehrli, B.; Johnson, C. A. Iron-mediated oxidation of

450

antimony(III) by oxygen and hydrogen peroxide compared to arsenic(III) oxidation. Environ.

451

Sci. Technol., 2006, 40 (8), 2565-2571.

452

(42)

McKibben, M. A.; Barnes, H. L. Oxidation of pyrite in low temperature acidic

453

solutions: Rate laws and surface textures. Geochim. Cosmochim. Ac., 1986, 50 (7),

454

1509-1520. 22


Page 22 of 32

Page 23 of 32


455 456 457

(43)

King, D. W. Role of carbonate speciation on the oxidation rate of Fe(II) in aquatic

systems. Environ. Sci. Technol., 1998, 32 (19), 2997-3003. (44)

Rush, J. D.; Bielski, B. Pulse radiolytic studies of the reactions of HO2/O2- with Fe

458

(II)/Fe (III) ions. The reactivity of HO2/O2- with ferric ions and its implication on the

459

occurrence of the Haber-Weiss reaction. J. Phys. Chem., 1985, 89 (23), 5062-5066.

460 461 462

(45)

Sutglik, Z.; Zagorski, Z. P. Pulse-radiolysis of neutral iron(II) solutions-oxidation of

ferrous-ions by .OH radicals. Radiat. Phys. Chem., 1981, 17 (4), 229-233. (46)

Hug, S. J.; Leupin, O. Iron-catalyzed oxidation of arsenic(III) by oxygen and by

463

hydrogen peroxide: pH-dependent formation of oxidants in the fenton reaction. Environ. Sci.

464

Technol., 2003, 37 (12), 2734-2742.

465

(47)

Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical-review of

466

rate constants for reactions of hydrated electrons, hydrogen-atoms and hydroxyl radicals

467

(.OH/.O-) in aqueous-solution. J. Phys. Chem. Ref. Data, 1988, 17 (2), 513-886.

468 469

(48)

Anipsitakis, G. P.; Dionysiou, D. D. Radical generation by the interaction of

transition metals with common oxidants. Environ. Sci. Technol., 2004, 38 (13), 3705-3712.

470

23



471

TABLE CAPTIONS

472

Table 1. Chemical Reactions and Corresponding Rate Constants in Pyrite Suspension

473

Table 2. The initial and normalized oxidation rate of Sb(III) by pyrite and SOP

474

FIGURES CAPTIONS

475

Fig. 1 Concentration of TSb and Sb(V) in solution (a) and the oxidation efficiency of

476

Sb(III)/TSb in solution by pyrite (b) under oxic and dark conditions at pH 3, 5, 7 and 9.

477

Inset: pseudo-first-order model for the determination of the observed rate constant Kobs.

478

[FeS2]=0.25 g L-1, [Sb(III)]=20 µM.

479

Fig. 2 Quenching effect of TBA and catalase on Sb(III) oxidation in pyrite suspension under

480

oxic and dark conditions at pH 3 (a), 5 (b) and 9 (c). [FeS2]=0.25 g L-1, [Sb(III)]=20 µM,

481

[TBA]=200 mM, [catalase]=0.075 g L-1.

482

Fig. 3 Sb(III)/TSb ratios under anoxic and oxic conditions in dark at pH 5. Inset:

483

pseudo-first-order model for the determination of the observed rate constant Kobs.

484

[FeS2]=0.25 g L-1, [Sb(III)]=20 µM.

485 486

Fig. 4 Comparison of Sb(III)/TSb ratios in solution between pyrite and SOP at pH 3, 5, 7 and

9 under oxic and dark conditions. [FeS2]=0.25 g L-1, [SOP]=0.25 g L-1, [Sb(III)]=20 µM.

487

Fig. 5 Sb(III)/TSb ratios during oxidation under dark and light conditions in oxic solutions at

488

pH 6. Inset: pseudo-first-order model for the determination of the observed rate constant

489

Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 µM, [TBA]=200 mM.

490

Fig. 6 Possible mechanisms of Sb(III) oxidation in pyrite suspension ❶ FeIII(pyrite)-induced

491

oxidation of Sb(III). ❷ FeII(aq)-induced oxidation of Sb(III). ➌ Light-induced ·OH and H2O2.

492

24


Page 24 of 32

Page 25 of 32


TABLE 1. Chemical Reactions and Corresponding Rate Constants in Pyrite Suspension Rate coefficients

Equ.

(s-1 or M-1 s-1 M-1 s-1)

ref.

Antimony Oxidations Sb(III)+·OH →Sb(IV)

2

Sb(III) + H2 O2 → Sb(IV)

3

Sb(III) +Fe(IV)→ Sb(IV)+Fe3+

9 41

●

4

41

8 × 109

1

Sb(IV) + O2 → Sb(V) + O2

1.1 × 109

41

Pyrite Oxidations 2+

4SO24

38

+ 4H+

5

2FeS2 + 7O2 + 2H2 O → 2Fe

6

+ FeS2 + 14Fe3+ + 8H2 O →15Fe2+ + 2SO24 + 16H

38

7

+ 2FeS2 + 15H2 O2 → 2Fe3+ + 4SO24 + 2H +14H2 O

42

8

FeII(pyrite) FeII(pyrite) FeIII (pyrite)

+ O2 →

FeII(aq) FeII(aq) FeII(aq) FeII(aq)

O2 → O2 ● O2 + 2H → FeIII aq + H2 O2 ● + III HO2 + H → Feaq + H2 O2 ·OH → FeIII aq + OH

+

Reactions on Pyrite Sulfur-Deficient Defect Sites

9 10

FeIII pyrite

●

+

O2

+

+ 2H →

27

●

+ O2

III Fepyrite

27

+ H2 O2

28

+ H2 O → FeII(pyrite) + ·OH FeII Oxidations

11 12 13 14

+ + + +

43

●

FeIII aq + +

44

1 × 107 1.2 × 106

44

8

45

3.2 × 10 Fenton Reactions

2+

3+

15

Fe

16

Fe OH + H2 O2 → FeIII OH2+ + ·OH + OH

+ H2 O2 → Fe

+ ·OH + OH

+

II

Modified Fenton Reactions 2+

46

17

Fe

18

FeII OH+ + H2 O2 → INT-OH

46

19

INT-OH + H+ → INT

46

20

INT → FeIII OH2+ + ·OH

46

21

INT-OH → Fe(IV)

+ H2 O2 → INT

46

Reactions on Pyrite under Illumination of Light →e

+

22

FeS2 + hv

23

H2 O + ℎ → H+ + ·OH

24

O2 + e- → O2

25

·OH + ·OH → H2 O2

+h

●

Reactive Oxygen Radical Reactions

26

●

HO2 + O2 + H+ → H2 O2 + O2 ●

27

●

●

HO2 + HO2 → H2 O2 + O2

5.2 × 109

47

9.7 × 107

44

8.3 × 105

44

9

6.6 × 10

47

3 × 107

47

●

28

HO2 + ·OH → O2 + H2 O ●

29

H2 O2 + ·OH → HO2 + H2 O

30

O2 + H+ ↔ HO2

31

·OH + TBA → O2 + products

●

1 × 1012/1.58 × 107

●

44

(pKa 4.8) ●-

(3.8~7.6)×108

493 25


48


Page 26 of 32

494 495 496

Table 2. The initial and normalized oxidation rate of Sb(III) for pseudo-first-order model by pyrite and SOP Pyrite

SOP

Initial rate Normalized rate

Initial rate Normalized rate

(min-1)

(min-1 m-2)

(min-1)

(min-1 m-2)

pH 3

0.0003

0.0016

0.0007

0.0025

pH 5

0.0012

0.0057

0.0048

0.0171

pH 7

0.0015

0.0068

0.0050

0.0178

pH 9

0.0254

0.1186

0.0575

0.2046

497 498

26


Page 27 of 32


Figure 1. Concentration of TSb and Sb(V) in solution (a) and the oxidation efficiency of Sb(III)/TSb in solution by pyrite (b) under oxic and dark conditions at pH 3, 5, 7 and 9. Inset: pseudo-first-order model for the determination of the observed rate constant Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM.



Figure 2. Quenching effect of TBA and catalase on Sb(III) oxidation in pyrite suspension under oxic and dark conditions at pH 3 (a), 5 (b) and 9 (c). [FeS2]=0.25 g L-1, [Sb(III)]=20 μM, [TBA]=200 mM, [catalase]=0.075 g L-1.


Page 28 of 32

Page 29 of 32


Figure 3. Sb(III)/TSb ratios under anoxic and oxic conditions in dark at pH 5. Inset: pseudofirst-order model for the determination of the observed rate constant K obs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM.



Figure 4. Comparison of Sb(III)/TSb ratios in solution between pyrite and SOP at pH 3, 5, 7 and 9 under oxic and dark conditions. [FeS2]=0.25 g L-1, [SOP]=0.25 g L-1, [Sb(III)]=20 μM.


Page 30 of 32

Page 31 of 32


Figure 5. Sb(III)/TSb ratios during oxidation under dark and light conditions in oxic solutions at pH 6. Inset: pseudo-first-order model for the determination of the observed rate constant Kobs. [FeS2]=0.25 g L-1, [Sb(III)]=20 μM, [TBA]=200 mM.



Figure 6. Possible mechanisms of Sb(III) oxidation in pyrite suspension ❶ ≡FeIII induced oxidation of Sb(III). ❷ FeII induced oxidation of Sb(III). ➌ Light-induced ∙OH and H2O2.


Page 32 of 32

Mechanisms of Sb(III) Oxidation by Pyrite-Induced Hydroxyl Radicals

Recommend Documents