X-ray Absorption Spectroscopic Quantification and ... - ACS Publications

Subscriber access provided by The University of British Columbia Library

Article

X-ray Absorption Spectroscopic Quantification and Speciation Modeling of Sulfate Adsorption on Ferrihydrite Surfaces Chunhao Gu, Zimeng Wang, James David Kubicki, Xiaoming Wang, and Mengqiang Zhu Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b00753 • Publication Date (Web): 05 Jul 2016 Downloaded from http://pubs.acs.org on July 6, 2016

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

X-ray Absorption Spectroscopic Quantification and Speciation Modeling of Sulfate

1

Adsorption on Ferrihydrite Surfaces

2 3

Chunhao Gu, † Zimeng Wang, ‡1 James D. Kubicki, §

4

Xiaoming Wang, † and Mengqiang Zhu†, *

5 6 7

†

Department of Ecosystem Science and Management, University of Wyoming, Laramie, WY 82071

8

‡

Department of Civil and Environmental Engineering, Stanford University, Stanford, CA 94305

9

§

Department of Geological Sciences, University of Texas, El Paso, TX 79968

10

11

*Corresponding author: Mengqiang Zhu

12

Tel: +1 307-766-5523

13

Email: [email protected]

14

1

15

University, Baton Rouge, LA 70803

Present Address: Department of Civil and Environmental Engineering, Louisiana State

16 17

Words of text: 5,242

18

Tables: 2

19

Figures: 5

20 21

Revision submitted to Environmental Science and Technology

22

June 25, 2016

23 24

1 ACS Paragon Plus Environment


Page 2 of 32

Table of Content Graph

25 26 27 28

XANES Linear Combination Fitting

Surface Complexation Modeling

Normalized µ(E)

Outer-sphere Inner-sphere

% Sulfate adsorption

100 80 60 40 20 0 3

4

5

6

7

8

pH

Quantum Chemical Calculation

Data Fit 2-

SO4

2480

2490

2500

E(ev)


Page 3 of 32

29


ABSTRACT

30

Sulfate adsorption on mineral surfaces is an important environmental chemical process,

31

but the structures and respective contribution of different adsorption complexes under various

32

environmental conditions are unclear. By combining sulfur K-edge XANES and EXAFS

33

spectroscopy, quantum chemical calculations, and surface complexation modeling (SCM), we

34

have shown that sulfate forms both outer-sphere complexes and bidentate-binuclear inner-sphere

35

complexes on ferrihydrite surfaces. The relative fractions of the complexes vary with pH, ionic

36

strength (I) and sample hydration degree (wet versus air-dried), but their structures remained the

37

same. The inner-sphere complex adsorption loading decreases with increasing pH while

38

remaining unchanged with I. At both I = 0.02 and 0.1 M, the outer-sphere complex loading

39

reaches maximum at pH ~ 5 and then decreases with pH, whereas it monotonically decreases

40

with pH at I = 0.5 M. These observations result from a combination of the ionic strength effect,

41

the pH dependence of anion adsorption, and the competition between inner- and outer-sphere

42

complexation. Air-drying drastically converts the outer-sphere complexes to the inner-sphere

43

complexes. The respective contributions to the overall adsorption loading of the two complexes

44

were directly modeled with the Extended Triple Layer SCM by implementing the bidentate-

45

binuclear inner-sphere complexation identified in the present study. These findings improve our

46

understanding of sulfate adsorption and its effects on other environmental chemical processes,

47

and have important implications for generalizing adsorption behavior of anions forming both

48

inner- and outer-sphere complexes on mineral surfaces.



49

Page 4 of 32

INTRODUCTION

50

Sulfate is a common anion in soil, water and atmospheric environments, and its

51

adsorption on mineral surfaces in these environments plays an important role in many

52

environmental chemical processes. The adsorption controls plant-available sulfur (S)

53

concentration 1, affects fate and transport of metals

54

weakens ice nucleation abilities of aerosol minerals

55

responsible for sulfate adsorption on Fe(III) oxides is also critical to the formation of

56

environmentally-relevant Fe(III) oxyhydroxy-sulfate minerals, such as schwertmannite and

57

jarosite 15-20.

2-6

, oxyanions, and organics 13, 14

3, 7-12

, and

. Fe(III)-sulfate complexation

58

Knowledge of the molecular structure of sulfate surface complexes is of fundamental

59

importance to understand the mechanisms of these critical mineral surface-controlled

60

environmental processes. Sulfate adsorption complexes and mechanisms on iron (Fe) oxide

61

surfaces have been characterized using infrared (IR) spectroscopy

62

calculations

63

spectroscopy

64

forms both inner- and outer-sphere surface complexes, and the relative proportion of the

65

inner-sphere complexes decreases with increasing pH and decreasing I

66

found that dried samples contained more sulfate inner-sphere surface complexes than

67

corresponding wet ones, and that drying may result in bisulfate formation on surfaces 22, 33, 36.

68

25-28

, surface complexation modeling (SCM)

31, 32

29, 30

21-24

, quantum chemical

, and S K-edge X-ray absorption

combined with macroscopic experiments. These studies showed that sulfate

21, 30, 33-35

. It was also

The molecular structure of sulfate inner-sphere complexes, however, had been 10,

21,

22,

30,

37-39

69

controversial

, until recently it was convincingly determined to be

70

bidentate-binuclear on ferrihydrite (Fhy) surfaces using S K-edge extended X-ray absorption fine

71

structure (EXAFS) spectroscopy and differential atomic pair distribution function analysis

31

.


Page 5 of 32


72

Nevertheless, limited experimental conditions, i.e., pH 4 and air-dried samples, were examined

73

in that study

74

loadings on the structure of sulfate surface complexes remain yet unknown. A recent study

75

quantified respective contributions of co-existing inner- and outer-sphere complexes to overall

76

sulfate adsorption using multivariate curve resolution (MCR) analysis of IR spectra which is,

77

however, an indirect approach with arbitrary assignments of the two types of complexes. A direct

78

approach is desired to determine the respective contributions to improve quantitative

79

understanding of the sulfate adsorption mechanism. For example, the experimentally-determined

80

fraction of each sulfate surface species provides additional constraints for surface complexation

81

models that predict the fractions, so that the models become more physically meaningful. In

82

addition, the previous SCM studies

83

sphere complexation which however, has not been proved.

31

. The effects of pH, I, sample hydration status (wet versus dried) and surface

29, 30

24

implemented monodentate-mononuclear (MM) inner-

84

In the present study, we characterized the structure of sulfate inner-sphere surface

85

complexes on Fhy as a function of pH, I, surface loading and hydration status (wet versus dried)

86

using S K-edge EXAFS spectroscopy, complemented by quantum chemical calculations.

87

Respective adsorption loadings of inner- and outer-sphere complexes were determined using S

88

K-edge XANES linear combination fitting (LCF) analysis. Subsequently, the EXAFS-

89

determined inner-sphere structure was implemented into the Extended Triple Layer SCM to

90

directly simulate the XANES-derived respective adsorption loadings, i.e., the sulfate adsorption

91

envelopes at three different ionic strengths. The combination of these approaches provides new

92

insights into adsorption on mineral surfaces of sulfate and other oxyanions that form both inner-

93

and outer-sphere complexes.

94

MATERIALS AND METHODS 5 ACS Paragon Plus Environment


Page 6 of 32

95

All glassware and plastic containers used in this study were soaked in 1 M HCl overnight

96

before use. All chemicals were of reagent grade. The salts used were nitrate rather than chloride

97

to prevent the interference of Cl with S K-edge EXAFS data collection as the Cl absorption edge

98

is located in the middle of the S EXAFS region 32.

99

Sulfate Adsorption Envelopes. The synthesis and characterization of 2-line Fhy is

100

provided in the supporting information (SI-1). Three sulfate adsorption envelopes were obtained

101

with 1.2 mM sulfate and 2.45 g/L Fhy at three ionic strengths controlled by 0.02, 0.1 and 0.5 M

102

NaNO3, respectively, at room temperature (20 ± 0.5 oC). Each suspension with sulfate loaded

103

was shaken on a tube rotator for 24 hours with pH maintained at 3 - 8 by adding small volumes

104

of 0.1 - 1 M HNO3 or NaOH. Ferrihydrite dissolution at pH 3 is negligible (SI-1). Atmospheric

105

CO2 was not excluded from the reaction system. Thus, dissolved CO2 at pH 8 might compete

106

with sulfate, which, however, is not expected to significantly affect sulfate adsorption loading

107

because of the high dissolved sulfate concentration at pH 8. The pH measurement in 0.5 M

108

adsorption experiments was not corrected for the high ionic strength effect, which may shift the

109

envelope to the left by ~ 0.15 pH unit as compared to the corrected one

110

samples were prepared using 5 mM sulfate at pH 3 and 6 to determine the effects of sulfate

111

loadings on the structure of sulfate inner-sphere complexes. At the termination of each

112

experiment, the solid was collected by vacuum filtration onto filter membranes with a 0.2 µm

113

pore size. The water associated with the solid was minimized by the filtration. The collected wet

114

solid of each sample was split into two portions for XAS analysis, with one kept wet and the

115

other air dried for ≥ 48 hours under ambient conditions. Sulfate concentration in the filtrates was

116

measured by inductively coupled plasma atomic emission spectroscopy. Sulfate adsorption

40

. A few additional


Page 7 of 32


117

loadings were calculated based on sulfate concentration difference to obtain the adsorption

118

envelopes.

119

X-ray Absorption Spectroscopy. Sulfur K-edge XANES spectroscopy was used to

120

qualitatively and quantitatively determine the relative proportions of the inner- and outer-sphere

121

complexes while S EXAFS spectroscopy for the structural characterization of the inner-sphere

122

complexes

123

adsorption samples as well as from 0.1 M Na2SO4 solutions of pH 0.7, 2 and 5.5 at beamline 4-3

124

at the Stanford Synchrotron Radiation Laboratory (SSRL), Menlo Park, CA. The program

125

Athena

126

for the XANES LCF analysis to determine the relative proportions of the two types of complexes.

127

Details about the data collection and analysis are provided in SI-2.

41

31, 32

. Both XANES and EXAFS spectra were collected from the wet and dried

was used for background removal, normalization, and extraction of EXAFS data, and

128

Surface Complexation Modeling. A surface complexation model was developed for

129

sulfate adsorption on Fhy using the formulations of the Extended Triple Layer Model (ETLM)

130

proposed by Fukushi and Sverjensky 30 for the same system. In ETLM, the involvement of water

131

molecules in a surface reaction (i.e., water dipole reaching or leaving a charged surface) was

132

considered for correcting the energetics of the equilibrium. The calculation was implemented in

133

MINEQL + 4.6.42 This model incorporated the surface acid-base reactions of Fhy, outer-sphere

134

adsorption of the electrolyte ions (Na+ and NO3-), sulfate adsorption (both inner- and outer-

135

sphere) on Fhy and the relevant aqueous speciation reactions of sulfate. The surface acid-base

136

equilibrium constants, the site density, the specific surface area, two capacitance values, and the

137

electrolyte ion outer-sphere adsorption equilibrium constants were directly taken from Fukushi

138

and Sverjensky 30 and kept intact. These intrinsic parameters of Fhy are reliable as indicated by



Page 8 of 32

139

their successful application in modeling numerous independent data sets of sulfate adsorption on

140

Fhy 29, 30.

141

Our present model quantitatively incorporated the new finding of the EXAFS-determined

142

sulfate BB structure whereas the selected outer-sphere surface complex was referenced to

143

Fukushi and Sverjensky

144

reactions were identified by systematically performing multiple forward calculations of the

145

model to directly model the XANES-determined respective inner-sphere and outer-sphere

146

complex adsorption loadings (in % of the total sulfate) while minimizing the residual sum of

147

squared errors. A homemade Excel Macro based tool (MINFIT)43 was used to implement the

148

optimization of the fitting parameters.

149

30

. The optimal values of the equilibrium constants for the two surface

Density Functional Theory (DFT). A Fhy nanoparticle structural model was 44, 45

150

constructed based on the structure proposed by Michel et al.

for geometric optimization of

151

sulfate surface complexes. The model was trimmed down to ~ 1.6 nm in diameter with a

152

stoichiometry of Fe38O104H962+ + SO42-, Fe38O104H95+ + HSO4-, or Fe38O104H962+ + SO42-·14(H2O)

153

using Materials Studio 7 (Accelrys, San Diego, CA) 46. The three initial structures were a sulfate

154

BB complex on a dry surface (SO4_Dry), a bisulfate BB complex on a dry surface (HSO4_Dry)

155

and a sulfate BB complex on a wet surface (SO4_Wet), respectively. The starting configurations

156

for these complexes were taken from previous simulations of CrO42- on the same surface 47. An

157

energy minimization was performed using the Vienna Ab-initio Simulation Package

158

starting configurations to allow all atoms to relax. Energy minimizations were carried out with

159

the periodic box surrounding the nanoparticle constrained to 20 × 20 × 20 Å3. This allowed for a

160

gap of ~ 6 Å between periodic images of the nanoparticles which should sufficiently minimize

48

on these


Page 9 of 32


161

interactions of the particle with itself. Other details for the geometric optimization are described

162

in SI-3.

163

RESULTS

164

Adsorption Envelopes. The sulfate adsorption envelopes on Fhy at different I are shown

165

in Figure 1A. With increasing pH from 3 to 8, the sulfate adsorption loading decreases

166

monotonically, which is expected as the pKa2 of H2SO4 < 3

167

drastically with increasing I (Figure 1A), indicating the presence of a considerable fraction of

168

sulfate outer-sphere complexes according to the conventional interpretation of the ionic strength

169

effects 49, 50.

29

. Sulfate adsorption decreases

170

EXAFS Spectroscopy. The EXAFS spectra of the three sulfate solutions are provided in

171

Figure 2 with a comparison of the fits to the spectra given in Figure SI-4C and 4D and obtained

172

parameters listed in Table 1. The solution spectra differ significantly in both k and R space

173

(Figure 2). With increasing solution pH from 0.7 to 5.5, the position of the O peak in R space

174

significantly shifts to the right with concomitant peak intensity increase (Figure 2B), suggesting

175

increased average S-O bond length (ds-o) of dissolved sulfate with a narrower length distribution.

176

EXAFS fitting results confirm this and show that as pH increases, ds-o increases from 1.47 ± 0.01

177

Å to 1.49 ± 0.01 Å with decreasing σ2 (Table 1). These changes result from decreasing sulfate

178

protonation degree (HSO4- SO42-) with elevated pH that increases the repulsion among the

179

four O2- of SO42-, elongating ds-o and narrowing its distribution.

180

EXAFS spectra and their fits for selected adsorption samples are compared in Figure SI-4

181

and the parameters are listed in Table 1. For the air-dried samples (Figure SI-4A and 4B),

182

EXAFS fitting identifies an Fe atomic shell at 3.22 – 3.25 Å, besides the O shell located at 1.47 –



Page 10 of 32

31, 32

183

1.49 Å, indicating that dried samples contain sulfate BB inner-sphere complexes

. The wet

184

samples and the solutions have similar spectra (Figure SI-4C and 4D), but a comparison of their

185

pre-edges (see the pre-edge analysis below) clearly shows that the former contain inner-sphere

186

complexes but with lower proportions than the corresponding dried samples. EXAFS fitting

187

shows that both ds-o and ds-Fe of the wet samples are similar to the distances of the dried ones

188

(Table 1). It is hard to tell whether the CNs of the Fe shell are significantly lower for the wet

189

samples (0.60 - 0.93) than for the dried ones (0.90 -1.87) due to the large fitting uncertainties.

190

The S-Fe distances, hence the type of sulfate inner-sphere complexes, remain essentially the

191

same with pH, I, and sulfate loadings (Table 1).

192

Pre-edges and XANES Spectra. The pre-edge is present only for inner-sphere

193

complexation due to the S-Fe orbital hybridization 51-53. The pre-edge intensity is proportional to

194

the fraction of the inner-sphere complexes

195

sample preparation conditions on the contribution of the inner- and outer-sphere complexation to

196

overall sulfate adsorption. Figure 3 shows the pre-edges of the adsorption samples prepared

197

under selected conditions. The pre-edges for all experimental conditions are provided in Figure

198

SI-5.

32

, and can thus be used to determine the effects of

199

For both wet and dried samples, the pre-edge peak generally becomes weaker with

200

increasing pH from 3 to 7 at each I (Figure 3A and SI-5A); at a given pH, the pre-edge intensity

201

increases with increasing I (Figure 3B and SI-5B); and the dried samples have much stronger

202

pre-edge peaks than the wet ones (Figure 3A and SI-5). These results indicate that lower pH,

203

higher I, and drying all favor the formation of the inner-sphere complexes.

204

To quantify the effects of the environmental factors on sulfate complexation, we

205

determined the inner- (ƒinner) and outer-sphere (ƒouter) complex fractions using XANES LCF


Page 11 of 32


206

analysis 32. Based on the above qualitative analysis, ƒinner increases with increasing I and drying

207

degree, and decreasing pH. Thus, the air-dried sample prepared at pH 3 and I = 0.5 M and an

208

additional wet sample prepared at pH 8 and I = 0 M (i.e., in DI water) were used to approximate

209

the inner- and outer-sphere end-members, respectively. The XANES spectra of the end-members

210

have distinct whiteline positions and post-edge profiles in addition to the pre-edges (Figure 3C),

211

warranting the accuracy of the LCF analysis. It is hard to quantify the overall errors of LCF

212

analyses while it is generally up to 10%. A comparison of the fits to the data for selected samples

213

is illustrated in Figure SI-6B and the obtained ƒinner and ƒouter are provided in Figure 1D and 1E.

214

The high goodness of fit indicates that sulfate speciation under different conditions is well

215

represented by a combination of the two types of sulfate surface complexes.

216

For the wet samples (Figure 1D), ƒinner generally decreases while ƒouter increases with

217

increasing pH and decreasing I, consistent with the qualitative comparison of the pre-edge

218

intensities described above. The respective adsorption loadings (Figure 1B) of the two

219

complexes were determined by combining the fractions in Figure 1D with the adsorption

220

envelopes in Figure 1A. With increasing pH at each I, the inner-sphere adsorption loading

221

decreases monotonically. The outer-sphere complex loadings show a similar trend at I = 0.5;

222

however, at I = 0.02 and 0.1 M, it reaches maximum at pH ~ 5 and then decreases. At each pH,

223

the outer-sphere complex loadings decrease drastically with increasing I whereas the loadings for

224

the inner-sphere complexes remain essentially unchanged, consistent with the ionic strength

225

effects 49, 54.

226

For the air-dried samples, the fractions (Figure 1E) and the loadings (Figure 1C) for

227

inner-sphere complexes were drastically increased compared to the wet samples while the

228

outer-sphere complexation was suppressed. The changes of the fractions and the loadings with



Page 12 of 32

229

pH are similar to those of the wet samples. The fractions of the inner-sphere complexes at I = 0.5

230

M are the highest while their fractions at I = 0.02 M and 0.1 M are similar, suggesting that ionic

231

strength modulates the effects of drying on sulfate complexation.

232

Surface Complexation Modeling. The surface complexation model was able to simulate

233

the general trends of respective adsorption loadings of the inner-sphere and outer-sphere

234

complexes as a function pH and I, hence the overall sulfate adsorption (i.e., the adsorption

235

envelops) (Figure 4). The model successfully predicted the co-existence of both types of

236

complexes at low pH. As the pH increases, the model shows a faster diminishing of inner-sphere

237

complexes than that of outer-sphere complexes, and that the outer-sphere complex adsorption

238

reaches maximum at pH 5 – 6 for I = 0.02 and 0.1 M. The model-predicted outer-sphere complex

239

maximum for I = 0.5 M was not evident in our XANES LCF results. As I increases, the model

240

reproduces the shift of the overall sulfate adsorption envelope to a lower pH and also predicts the

241

drastic suppression of the outer-sphere complexation as opposed to the inner-sphere

242

complexation. All of those trends simulated by the model are generally consistent with the

243

results from both the macroscopic adsorption experiments and the spectroscopic analyses with a

244

remarkable resolution that exceeds most previous studies 29, 30.

245

DFT Calculations. The energy-minimized structures of the sulfate complexes are shown

246

in Figure 5. After the optimization, the SO4_dry remains as a BB complex whereas both

247

HSO4_dry and SO4_wet convert to MM complexes. For the optimized SO4_dry complex, the

248

predicted ds-o are 1.47 - 1.57 Å with an average value of 1.51 ± 0.05 Å (Table SI-3), slightly

249

longer than the EXAFS-determined bond lengths (1.47 - 1.49 Å) for the dried samples. The

250

predicted ds-Fe are 3.21 and 3.24 Å (Figure 5), in good agreement with the EXAFS-determined

251

values (3.22 – 3.24 Å in Table 1). As to the optimized structure of HSO4_Dry, the predicted ds-o


Page 13 of 32


252

(Table SI-3) are close to the experimental, but ds-Fe (3.41 and 3.61 Å) do not match with the

253

EXAFS values (Table 1). In addition, the SO4_dry complex is more stable than HSO4_dry by

254

about -125 kJ/mol. Thus, the HSO4- MM complexes on dried Fhy surfaces can be excluded. As

255

to the optimized structure of SO4_wet, the predicted ds-Fe are 3.28 Å and 3.58 Å, respectively,

256

suggesting a MM complex. The 3.28 Å distance is close to the experimental (3.22 Å to 3.25 Å in

257

Table 1) for the wet samples, but the distance of 3.58 Å is not experimentally observed.

258

DISCUSSION

259

Structure of Sulfate Inner-sphere Complexes. The EXAFS-determined S-Fe distances

260

(3.22 – 3.24 Å) for the dried samples are consistent with the DFT prediction of the BB

261

complexes on a dried surface. On the wet surface, the DFT prediction suggests that sulfate forms

262

MM complexes, but the longer ds-Fe (3.58 Å) is not detected in the EXAFS analysis. The

263

inconsistency could be due to the low fractions of the inner-sphere complexes in the wet samples,

264

not allowing the EXAFS fitting to resolve the long ds-Fe. Alternatively, it could be due to

265

inaccuracy of the DFT prediction for wet surfaces because of potentially insufficient handling of

266

H-bonding in the DFT method 55.

267

The XANES LCF analysis, however, strongly supports the presence of BB inner-sphere

268

complexes on the wet Fhy surfaces because the spectra of the wet samples are fitted well with

269

the spectra of a dried sample (i.e., the inner-sphere end member) that contains mainly BB

270

complexes. This conclusion is further supported by the similarity of the pre-edge peak positions

271

between the dried and wet samples, which is evident from the difference pre-edges (Figure 3C)

272

with respect to pH 5.5 sulfate solution. The peak position does not vary among samples prepared

273

under different conditions, suggesting that the numbers of Fe atoms coordinated to each sulfate

274

ion are the same, i.e., the same type of surface complexes

19

. Therefore, although the 13

ACS Paragon Plus Environment


Page 14 of 32

275

experimental conditions change its fraction, the type of inner-sphere complexes (BB) remains

276

unaffected.

277

Surface Complexation Modeling. This study presents the first sulfate surface

278

complexation model by implementing the convincing sulfate BB adsorption geometry obtained

279

from the EXAFS spectroscopy. Although there are some discrepancies between the modeling

280

results and the macroscopic data, which could be due to the errors in the XANES LCF

281

quantification or insufficiency in SCM, or both, our modeling predications capture most

282

experimental observations, both macroscopically and microscopically. Compared with previous

283

studies on sulfate adsorption on ferrihydrite

284

quantitative speciation data of the individual inner- and outer-sphere complexes at various pH

285

and I. Therefore, our model incorporates more physically validated information in the speciation

286

calculation compared the previous models which only consider the overall sulfate uptake results

287

as the fitting constraint.

288

29, 30

, our present model was built by fitting the

The model framework used in this study were the same as those in

Fukushi and

289

Sverjensky 30 except that the inner-sphere surface complex used was BB, i.e., (≡FeO)2SO2, while

290

they used MM, i.e., ≡FeOSO3−. The selection of the outer-sphere surface complex was less

291

constrained. The outer-sphere species (≡FeOH2+)2-SO42−, as suggested for sulfate adsorption on

292

goethite surfaces in Fukushi and Sverjensky

293

only emerged above pH 5 (data not shown), which contradicts to the coexistence of inner- and

294

outer-sphere surface complexes at pH 3 and 4 (Figure 1D). Fukushi and Sverjensky 29, 30 reported

295

that bisulfate outer-sphere surface complex (≡FeOH2+-HSO4−) was required for modeling sulfate

296

adsorption on ferrihydrite. The present model found that the combination of (≡FeO)2SO2 and

56

, gave very poor fits to the adsorption data and


Page 15 of 32


297

≡FeOH2+-HSO4− gives good fits to all results including the macroscopic adsorption percentage

298

and the respective fractions of each surface complex (Figure 4).

299

The bisulfate outer-sphere complex, however, may not exist on the surfaces although it

300

gives the best modeling results and was also used in previous modeling studies 6, 29, 30. At pH > 4,

301

dissolved sulfate is entirely SO42-, and its protonation to form HSO4- surface complexes on a

302

positively-charged surface is apparently disfavored. In addition, the XANES spectra of the high

303

pH samples (7 and 8) do not have any characteristics, such as the whiteline peak broadening, as

304

seen for the bisulfate solution (pH 0.7, Figure SI-6A), another evidence for the absence of

305

bisulfate surface species at high pH. The bisulfate species used in the SCM is more likely a

306

representation of some unknown H-bonded SO42- outer-sphere complexes that have

307

spectroscopically distinct nature compared to the free SO42- ion based on their different XANES

308

profiles (Figure SI-6A). The difference is also observed between sulfate solution and the

309

structural sulfate of schwertmannite

310

surface complexes cannot be reflected in the classic framework of the triple layer models.

32

. However, such delicate feature of the outer-sphere

311

Effects of Ionic Strength. According to Hayes et al. 49, 54, the macroscopic investigation

312

of ionic strength effects can be used to distinguish between inner- and outer-sphere adsorption in

313

the study of ion adsorption on metal oxides. Adsorption that decreases with increasing I indicates

314

outer-sphere surface complexation (not excluding the co-existence of inner-sphere complexes)

315

whereas inner-sphere adsorption is not affected or increased marginally by increasing I

316

suppression of the outer-sphere complexation could be understood by considering both

317

competitive adsorption from background electrolytes and the electric double layer (EDL)

318

contraction. Background electrolyte NO3- forms outer-sphere complexes and competes with

319

outer-spherically bound sulfate, suppressing sulfate outer-sphere complexation. A higher I

57

. The



Page 16 of 32

320

induces a more pronounced EDL contraction and thus a lower electrical potential (positive), at

321

the adsorption plane

322

experimental study spectroscopically validating the prevailing interpretation for the ionic

323

strength effect for an adsorption system with co-existing inner- and outer-sphere complexes of

324

the same type of ion.

325

58

, which can also decrease sulfate adsorption. Our work is the first

Effects of pH. At a given I, finner decreases and fouter increases with increasing pH (Figure 21, 33

326

1D and 1E), consistent with previous IR studies

. This indicates that sulfate inner-sphere

327

complexation is more vulnerable to increasing pH than outer-sphere complexation. As pH

328

increases, >FeOH and/or >FeO- became increasingly dominant, disfavoring the ligand exchange

329

(i.e., inner-sphere adsorption) between the surface groups and sulfate as the Fe-O bonds in these

330

two species are stronger and harder to break than that in >FeOH2+ 59. The less positively charged

331

surface with increasing pH impairs sulfate approaching the surface, further disfavoring sulfate

332

inner-sphere adsorption. However, outer-sphere complexation is disfavored only by the surface

333

charge changes, probably accounting for its more resistance than inner-sphere complexation to

334

increasing pH.

335

The adsorption loading of the outer-sphere complexes monotonically decreases at I = 0.5

336

M with increasing pH, but at I = 0.02 and 0.1 M, it reaches maximum at ~ pH 5. These can be

337

understood by considering both the surface charge changes and the changes in the amount of

338

available surface sites for outer-sphere complexation as a result of competition from inner-sphere

339

complexation. With increasing pH, more surface sites become available for outer-sphere

340

complexation because of the faster decrease of the inner-sphere complexation, favoring

341

formation of outer-sphere complexes; meanwhile, surfaces become less positively charged,

342

disfavoring the outer-sphere complexation. At low I (0.02 and 0.1 M), the influence of the two 16 ACS Paragon Plus Environment

Page 17 of 32


343

factors are likely comparable, resulting in the maxima for outer-sphere complexation with

344

increasing pH. At I = 0.5 M, however, outer-sphere complexation is much disfavored by the high

345

I, and therefore, the maximum of the outer-sphere complexation is absent with increasing pH.

346

The outer-sphere complexation maxima do not exist for an adsorption system with a high sulfate

347

loading (6.05 mM, data not shown), either. In this case, the majority of the surface sites are

348

occupied and the amount of the sites released by a rapid decrease of inner-sphere complexation

349

are inadequate to significantly boost the amount of available sites for outer-sphere complexation.

350

The absence of outer-sphere complexation maxima was also predicted in our SCM at 6.05 mM

351

total sulfate (data not shown). These underlying mechanisms described above for the observed

352

pH- or I-dependent changes of sulfate inner- and outer-sphere complexation are more or less

353

reflected in our ETL surface complexation model that roughly captures the adsorption data.

354

Effects of Hydration. An ion is hydrated in aqueous solution and its hydration sphere

355

has to be partially/completely removed prior to its adsorption as an inner-sphere complex. The

356

following inner-sphere complexation reaction, if neglecting the ligand exchange with surface

357

adsorbed water molecules 35,

358 359

SO42-(H2O)n + 2(>Fe-OH) ⇌ (>Fe-O)2(SO2)(H2O)m + (n – m)H2O + 2OH-

∆

(1)

can be split into two, i.e.,

360

SO42-(H2O)n ⇌ (SO42-)(H2O)m + (n-m)H2O

361

SO42-(H2O)m + 2(>Fe-OH) ⇌ (>Fe-O)2(SO2)(H2O)m + 2OH-

∆

(2)

∆

(3)

362

Thus, the overall Gibbs free energy change (∆) including the Coulombic interaction can be

363

written as

364

∆ = ∆ + ∆ = ∆ + ∆ + ∆ ,

(4) 17

ACS Paragon Plus Environment


Page 18 of 32

365

among which ∆ > 0, ∆ < 0, while ∆ can be either positive or negative. If

366

∆ > 0, it means that the amount of energy required for dehydration is larger than the net amount

367

of energy released from the chemical binding and the Coulombic interaction. In this case, the

368

hydration sphere likely remains and outer-sphere complexation is favored; otherwise, inner-

369

sphere complexes form. In another word, the proportions of inner-sphere and outer-sphere

370

complexes depend on the relative contribution of the dehydration energy to the overall

371

adsorption reaction energy change.

372

This general rule could be used to explain the different behavior of anion adsorption on

373

mineral surfaces. For phosphate, arsenate, selenite and silicate, they strongly bind to mineral

374

surfaces (i.e., ∆ is very negative) and ∆ contribution to the adsorption is insignificant,

375

thus forming mainly inner-sphere complexes with negligible outer-sphere complexation. Anions,

376

such as nitrate, perchlorate, chromate, and Cl-, do not or very weakly bind to mineral surfaces

377

(i.e., ∆ is small), and their∆ , albeit small, could dominate∆ , thus forming mainly

378

outer-sphere complexes. For anions, such as CO32-, sulfate, selenate and arsenite, their hydration

379

energies could be comparable to the net energy change of chemical binding and Coulombic

380

interactions (i.e., ∆ + ∆ ), resulting in coexisting inner- and outer-sphere

381

complexes

382

hydration (some are listed in Table SI-7) and chemical binding energy data. Calorimetry and

383

DFT calculations of adsorption energies would help 63. Note that the above mechanism does not

384

consider the mineral surface dehydration, which contributes to the energy landscape for ion

385

adsorption and could also alter the proportion of inner- and outer-sphere complexes formed on

386

the surface 35.

33, 60-62

. The validation of this interpretation depends on the availability of ion


Page 19 of 32


387

The tendency to form inner-sphere complexes is improved by sample drying as it

388

enhances dehydration of both the adsorbate ion and the mineral surfaces. Our XAS results

389

clearly indicate that drying drastically converted sulfate outer-sphere complexes to inner-sphere

390

complexes on Fhy surfaces (Figure 1C and 1E) while not affecting the type of the inner-sphere

391

complexes (i.e., BB), as also observed for sulfate in schwertmannite

392

previous IR and DFT studies although they did not give convincing evidence on the changes 21, 22,

393

34, 36, 37

394

32

. This is consistent with

. In addition, drying wet paste may concentrate the acid in the solution associated with the

395

solid, resulting in protonated sulfate (HSO4- )

396

their XANES spectra, with the former having a shoulder peak on the left side of the white line.

397

This peak, however, is not obvious in the spectra of the dried samples, even for those prepared at

398

pH 3, suggesting that HSO4- on the surface was negligible if any under the conditions of the

399

present study. This could be due to the vacuum filtration, leaving much less solution with the

400

solids than directly drying the suspension does, as shown in Hug

401

concentration, a small amount of NaNO3 could precipitate as a solid during drying of the

402

adsorption samples while not visually observed. Further studies might be needed to determine

403

whether and how precipitated salts affect sulfate surface speciation and other surface chemical

404

properties of dried adsorption samples.

405

ENVIRONMENTAL IMPLICATIONS

22, 37, 64

. HSO4- differs significantly from SO42- in

22

. As used with a high

406

A systematic study of the effects of environmental conditions on sulfate complexation on

407

Fe(III) oxide surfaces is of fundamental importance to the understanding of the fate and transport

408

of sulfate and toxic metal(loid)s, and the formation and occurrence of Fe(III)-sulfate minerals in

409

sulfate-rich environments. Sulfate imposes strong influence on metal complexation on mineral 19 ACS Paragon Plus Environment


Page 20 of 32

410

surfaces by forming metal-sulfate ternary complexes, and the revelation of the sulfate adsorption

411

mechanism helps develop structurally consistent surface complexation model for predicting

412

adsorption of both sulfate and the metal on Fe(III) oxides

413

sulfate and Fe oxide surfaces and its pH dependence can be used to infer reactions between

414

sulfate and soluble Fe3+ monomers and clusters, leading to formation of Fe(III)-sulfate minerals,

415

including schwertmannite and jarosite

416

surface speciation changes during natural wet-dry cycles. This study also has important

417

implications for generalizing adsorption behavior of toxic and/or environmentally abundant

418

anions with concurring inner- and outer-sphere complexation, such as selenate, arsenite,

419

carbonate, etc. 33, 60-62.

420

ACKNOWLEDGEMENT

15-20, 65

2-6

. The reaction mechanisms between

. The drying effects provide insights into sulfate

421

This work was funded by the Wyoming Agricultural Experimental Station Competitive

422

Research Grant. We are grateful to Dr. Sabine Goldberg at the U.S. Salinity Laboratory for

423

helpful discussion on surface complexation, and Dr. Tjisse Hiemstra at the Wageningen

424

University for providing his unpublished surface protonation scheme. Comments and

425

suggestions of three anonymous reviewers and Associate Editor David Waite significantly

426

improved the quality of an earlier version of the paper. Use of the Stanford Synchrotron

427

Radiation Lightsource, SLAC National Accelerator Laboratory, is supported by the U.S.

428

Department of Energy, Office of Science, Office of Basic Energy Sciences under Contract No.

429

DE-AC02-76SF00515.

430

CyberInfrastructure computational resources provided by The Institute for CyberScience at The

431

Pennsylvania State University (http://ics.psu.edu).

432

SUPPORTIING INFORMATION

Portions

of

this

research

were

conducted

with

Advanced


Page 21 of 32


433

The supporting information is available free of charge via the internet at

434

http://pubs.acs.org, including ferrihydrite synthesis, XAS data collection and analysis, the DFT

435

calculations and the results, pre-edges of adsorption samples, LCF analysis for selected samples,

436

and hydration energies for some anions.

437



438 439 440 441 442 443 444 445 446 447 448 449 450 451 452 453 454 455 456 457 458 459 460 461 462 463 464 465 466 467 468 469 470 471 472 473 474 475 476 477 478 479 480 481 482

Page 22 of 32

Literature 1. Wilhelm Scherer, H., Sulfur in soils. J. Plant Nutr. Soil Sci. 2009, 172, (3), 326-335. 2. Swedlund, P. J.; Webster, J. G.; Miskelly, G. M., Goethite adsorption of Cu(II), Pb(II), Cd(II), and Zn(II) in the presence of sulfate: properties of the ternary complex. Geochim. Cosmochim. Acta 2009, 73, (6), 1548-1562. 3. Ali, M. A.; Dzombak, D. A., Interactions of copper, organic acids, and sulfate in goethite suspensions. Geochim. Cosmochim. Acta 1996, 60, (24), 5045-5053. 4. Walter, M.; Arnold, T.; Reich, T.; Bernhard, G., Sorption of uranium(VI) onto ferric oxides in sulfate-rich acid waters. Environ. Sci. Technol. 2003, 37, (13), 2898-2904. 5. Hoins, U.; Charlet, L.; Sticher, H., Ligand effect on the adsorption of heavy metals: The sulfate — Cadmium — Goethite case. Water Air Soil Pollut. 1993, 68, (1-2), 241-255. 6. Hinkle, M. A. G.; Wang, Z.; Giammar, D. E.; Catalano, J. G., Interaction of Fe(II) with phosphate and sulfate on iron oxide surfaces. Geochim. Cosmochim. Acta 2015, 158, 130-146. 7. Geelhoed, J. S.; Hiemstra, T.; Van Riemsdijk, W. H., Phosphate and sulfate adsorption on goethite: single anion and competitive adsorption. Geochim. Cosmochim. Acta 1997, 61, (12), 2389-2396. 8. Jain, A.; Loeppert, R. H., Effect of competing anions on the adsorption of arsenate and arsenite by ferrihydrite. J. Environ. Qual. 2000, 29, (5), 1422-1430. 9. Wilkie, J. A.; Hering, J. G., Adsorption of arsenic onto hydrous ferric oxide: effects of adsorbate/adsorbent ratios and co-occurring solutes. Colloids Surf., A Physicochem. Eng. Asp. 1996, 107, 97-110. 10. Lefèvre, G.; Fédoroff, M., Sorption of sulfate ions onto hematite studied by attenuated total reflection-infrared spectroscopy: kinetics and competition with other ions. Phys. Chem. Earth., Pt A/B/C 2006, 31, (10–14), 499-504. 11. Heather J. Shipley, S. Y., Amy T. Kan, Mason B. Tomson, A sorption kinetics model for arsenic adsorption to magnetite nanoparticles. Environ. Sci. Pollut. Res. 2010, 17, 1053 – 1062. 12. Vaishya, R.; Gupta, S., Arsenic removal from groundwater by iron impregnated sand. J. Environ. Eng. 2002, 129, (1), 89-92. 13. Yang, Z.; Bertram, A. K.; Chou, K. C., Why do sulfuric acid coatings influence the ice nucleation properties of mineral dust particles in the atmosphere? J. Phys. Chem. Lett. 2011, 2, (11), 1232-1236. 14. Cziczo, D. J.; Froyd, K. D.; Gallavardin, S. J.; Moehler, O.; Benz, S.; Saathoff, H.; Murphy, D. M., Deactivation of ice nuclei due to atmospherically relevant surface coatings. Environ. Res. Lett. 2009, 4, (4), 044013. 15. Burton, E. D.; Bush, R. T.; Johnston, S. G.; Watling, K. M.; Hocking, R. K.; Sullivan, L. A.; Parker, G. K., Sorption of arsenic (V) and arsenic (III) to schwertmannite. Environ. Sci. Technol. 2009, 43, (24), 9202-9207. 16. Regenspurg, S.; Peiffer, S., Arsenate and chromate incorporation in schwertmannite. Appl Geochem 2005, 20, (6), 1226-1239. 17. Waychunas, G.; Xu, N.; Fuller, C.; Davis, J.; Bigham, J., XAS study of AsO43− and SeO42− substituted schwertmannites. Physica B: Condensed Matter 1995, 208, 481-483. 18. Jönsson, J.; Persson, P.; Sjöberg, S.; Lövgren, L., Schwertmannite precipitated from acid mine drainage: phase transformation, sulphate release and surface properties. Appl Geochem 2005, 20, (1), 179-191.


Page 23 of 32

483 484 485 486 487 488 489 490 491 492 493 494 495 496 497 498 499 500 501 502 503 504 505 506 507 508 509 510 511 512 513 514 515 516 517 518 519 520 521 522 523 524 525 526 527 528


19. Majzlan, J.; Myneni, S. C. B., Speciation of iron and sulfate in acid waters: Aqueous clusters to mineral precipitates. Environ. Sci. Technol. 2004, 39, (1), 188-194. 20. Zhu, M.; Legg, B.; Zhang, H.; Gilbert, B.; Ren, Y.; Banfield, J. F.; Waychunas, G. A., Early-stage formation of iron oxyhydroxides during neutralization of simulated acid mine drainage solutions. Environ. Sci. Technol. 2012, 46, (15), 8140-8147. 21. Peak, D.; Ford, R. G.; Sparks, D. L., An in situ ATR-FTIR investigation of sulfate bonding mechanisms on goethite. J. Colloid Interface Sci. 1999, 218, (1), 289-299. 22. Hug, S. J., In Situ Fourier transform infrared measurements of sulfate adsorption on hematite in aqueous solutions. J. Coll. Interf. Sci. 1997, 188, (2), 415-422. 23. Boily, J.-F.; Gassman, P. L.; Peretyazhko, T.; Szanyi, J.; Zachara, J. M., FTIR spectral components of schwertmannite. Environ. Sci. Technol. 2010, 44, (4), 1185-1190. 24. Johnston, C. P.; Chrysochoou, M., Mechanisms of chromate, selenate, and sulfate adsorption on Al-substituted ferrihydrite: Implications for ferrihydrite surface structure and reactivity. Environ. Sci. Technol. 2016, 50, (7), 3589-3596. 25. Paul, K. W.; Borda, M. J.; Kubicki, J. D.; Sparks, D. L., Effect of dehydration on sulfate coordination and speciation at the Fe−(Hydr)oxide−water interface: a molecular orbital/density functional theory and fourier transform infrared spectroscopic investigation. Langmuir 2005, 21, (24), 11071-11078. 26. Paul, k. W.; Kubicki, J. D.; Sparks, D. L., Quantum chemical calculations of sulfate adsorption at the Al- and Fe-(Hydr)oxide-H2O interfaces: estimation of Gibbs free energies. Environ. Sci. Technol. 2006, 40, 7717-7724. 27. Paul, K. W.; Kubicki, J. D.; Sparks, D. L., Sulphate adsorption at the Fe (hydr)oxide– H2O interface: comparison of cluster and periodic slab DFT predictions. Eur. J. Soil Sci. 2007, 58, 978-988. 28. Kubicki, J. D.; Kwon, K. D.; Paul, K. W.; Sparks, D. L., Surface complex structures modelled with quantum chemical calculations: carbonate, phosphate, sulphate, arsenate and arsenite. Eur. J. Soil Sci. 2007, 58, 932-944. 29. Fukushi, K.; Aoyama, K.; Yang, C.; Kitadai, N.; Nakashima, S., Surface complexation modeling for sulfate adsorption on ferrihydrite consistent with in situ infrared spectroscopic observations. Appl Geochem 2013, 36, 92-103. 30. Fukushi, K.; Sverjensky, D. A., A surface complexation model for sulfate and selenate on iron oxides consistent with spectroscopic and theoretical molecular evidence. Geochim. Cosmochim. Acta 2007, 71, (1), 1-24. 31. Zhu, M.; Northrup, P.; Shi, C.; Billinge, S. J.; Sparks, D. L.; Waychunas, G. A., Structure of sulfate adsorption complexes on ferrihydrite. Environ. Sci. Technol. Lett. 2014, 1, (1), 97-101. 32. Wang, X.; Gu, C.; Feng, X.; Zhu, M., Sulfate local coordination environment in schwertmannite. Environ. Sci. Technol. 2015, 49, (17), 10440 – 10448. 33. Peak, D.; Elzinga, E. J.; Sparks, D. L., Understanding sulfate adsorption mechanisms on iron (III) oxides and hydroxides: Results from ATR-FTIR spectroscopy. In Heavy Metals Release in Soils, Sparks, H. M. S. a. D. L., Ed. CRC Press: 2001; pp 167-190. 34. Wijnja, H.; Schulthess, C. P., Vibrational spectroscopy study of selenate and sulfate adsorption mechanisms on Fe and Al (Hydr)oxide surfaces. J. Colloid Interface Sci. 2000, 229, (1), 286-297. 35. Sverjensky, D. A.; Fukushi, K., Anion adsorption on oxide surfaces: inclusion of the water dipole in modeling the electrostatics of ligand exchange. Environ. Sci. Technol. 2006, 40, (1), 263-271. 23 ACS Paragon Plus Environment


529 530 531 532 533 534 535 536 537 538 539 540 541 542 543 544 545 546 547 548 549 550 551 552 553 554 555 556 557 558 559 560 561 562 563 564 565 566 567 568 569 570 571 572 573 574

Page 24 of 32

36. Eggleston, C. M.; Hug, S.; Stumm, W.; Sulzberger, B.; Dos Santos Afonso, M., Surface complexation of sulfate by hematite surfaces: FTIR and STM observations. Geochim. Cosmochim. Acta 1998, 62, (4), 585-593. 37. Paul, K. W.; Borda, M. J.; Kubicki, J. D.; Sparks, D. L., Effect of dehydration on sulfate coordination and speciation at the Fe-(hydr) oxide-water interface: A molecular orbital/density functional theory and Fourier transform infrared spectroscopic investigation. Langmuir 2005, 21, (24), 11071-11078. 38. Parfitt, R. L.; Smart, R. S. C., The mechanism of sulfate adsorption on iron oxides. Soil Sci. Soc. Am. J. 1978, 42, (1), 48-50. 39. Turner, L.; Kramer, J., Sulfate ion binding on goethite and hematite. Soil Sci. 1991, 152, (3), 226-230. 40. Wiesner, A. D.; Katz, L. E.; Chen, C.-C., The impact of ionic strength and background electrolyte on pH measurements in metal ion adsorption experiments. J. Colloid Interface Sci. 2006, 301, (1), 329-332. 41. Ravel, á.; Newville, M., ATHENA, ARTEMIS, HEPHAESTUS: data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, (4), 537-541. 42. Schecher, W. D.; McAvoy, D. C., MINEQL+: A chemical equilibrium modeling system, version 4.6. In Environmental Research Software, Hallowell, ME, 2007. 43. Wang, Z. MINFIT: Simply Use MINEQL+ to Fit Data. http://minfit.strikingly.com 44. Michel, F. M.; Barrón, V.; Torrent, J.; Morales, M. P.; Serna, C. J.; Boily, J.-F.; Liu, Q.; Ambrosini, A.; Cismasu, A. C.; Brown, G. E., Ordered ferrimagnetic form of ferrihydrite reveals links among structure, composition, and magnetism. Proc. Natl. Acad. Sci. 2010, 107, (7), 27872792. 45. Barrón, V.; Torrent, J.; Michel, F. M., Critical evaluation of the revised akdalaite model for ferrihydrite—Discussion. Am. Mineral. 2012, 97, (1), 253-254. 46. Materials Studio 7.0 Accelrys Inc., S. D., CA. (2014). 47. Kubicki J.D., C. E., Strongin D.R. , Density functional theory modeling of ferrihydrite nanoparticle charging and adsorption behavior. . In Abstracts of the American Chemical Society, in press, (2015) 48. Kresse, G.; Furthmüller, J., Efficient iterative schemes for ab initio total-energy calculations using a plane-wave basis set. Physical Review B 1996, 54, (16), 11169-11186. 49. Hayes, K. F.; Papelis, C.; Leckie, J. O., Modeling ionic strength effects on anion adsorption at hydrous oxide/solution interfaces. J. Colloid Interface Sci. 1988, 125, (2), 717-726. 50. Lützenkirchen, J., Ionic strength effects on cation sorption to oxides: macroscopic observations and their significance in microscopic interpretation. J. Colloid Interface Sci. 1997, 195, (1), 149-155. 51. Majzlan, J.; Myneni, S. C., Speciation of iron and sulfate in acid waters: aqueous clusters to mineral precipitates. Environ. Sci. Technol. 2005, 39, (1), 188-194. 52. Okude, N.; Nagoshi, M.; Noro, H.; Baba, Y.; Yamamoto, H.; Sasaki, T., P and S K-edge XANES of transition-metal phosphates and sulfates. J. Electron Spectros. Relat. Phenomena 1999, 101, 607-610. 53. Majzlan, J.; Alpers, C. N.; Koch, C. B.; McCleskey, R. B.; Myneni, S. C.; Neil, J. M., Vibrational, X-ray absorption, and Mössbauer spectra of sulfate minerals from the weathered massive sulfide deposit at Iron Mountain, California. Chem. Geol. 2011, 284, (3), 296-305. 54. Hayes, K.; Leckie, J., Modeling ionic strength effects on cation adsorption at hydrous oxide/solution interfaces. J. Colloid Interface Sci. 1987, 115, (2), 564-572. 24 ACS Paragon Plus Environment

Page 25 of 32

575 576 577 578 579 580 581 582 583 584 585 586 587 588 589 590 591 592 593 594 595 596 597 598 599 600 601 602 603 604 605 606 607 608 609 610 611 612 613 614 615 616 617


55. Paul, K.; Kubicki, J.; Sparks, D., Sulphate adsorption at the Fe (hydr) oxide–H2O interface: comparison of cluster and periodic slab DFT predictions. Eur. J. Soil Sci. 2007, 58, (4), 978-988. 56. Fukushi, K.; Sverjensky, D. A., A surface complexation model for sulfate and selenate on iron oxides consistent with spectroscopic and theoretical molecular evidence. Geochim. Cosmochim. Acta 2007, 71, (1), 1-24. 57. Goldberg, S.; Johnston, C. T., Mechanisms of arsenic adsorption on amorphous oxides evaluated using macroscopic measurements, vibrational spectroscopy, and surface complexation modeling. J. Colloid Interface Sci. 2001, 234, (1), 204-216. 58. Zhang, G. Y.; Brümmer, G. M.; Zhang, X. N., Effect of perchlorate, nitrate, chloride and pH on sulfate adsorption by variable-charge soils. Geoderma 1996, 73, (3–4), 217-229. 59. Zhu, M.; Paul, K. W.; Kubicki, J. D.; Sparks, D. L., Quantum chemical study of arsenic (III, V) adsorption on Mn-Oxides: Implications for arsenic(III) oxidation. Environ. Sci. Technol. 2009, 43, (17), 6655-6661. 60. Bargar, J. R.; Kubicki, J. D.; Reitmeyer, R.; Davis, J. A., ATR-FTIR spectroscopic characterization of coexisting carbonate surface complexes on hematite. Geochim. Cosmochim. Acta 2005, 69, (6), 1527-1542. 61. Arai, Y.; Elzinga, E. J.; Sparks, D. L., X-ray absorption spectroscopic investigation of arsenite and arsenate adsorption at the aluminum oxide–water interface. J. Colloid Interface Sci. 2001, 235, (1), 80-88. 62. Peak, D.; Sparks, D. L., Mechanisms of selenate adsorption on iron oxides and hydroxides. Environ. Sci. Technol. 2002, 36, (7), 1460-1466. 63. Kubicki, J. D.; Paul, K. W.; Kabalan, L.; Zhu, Q.; Mrozik, M. K.; Aryanpour, M.; PierreLouis, A.-M.; Strongin, D. R., ATR–FTIR and Density Functional Theory Study of the Structures, Energetics, and Vibrational Spectra of Phosphate Adsorbed onto Goethite. Langmuir 2012, 28, (41), 14573-14587. 64. Kubicki, J.; Kwon, K.; Paul, K.; Sparks, D., Surface complex structures modelled with quantum chemical calculations: carbonate, phosphate, sulphate, arsenate and arsenite. Eur. J. Soil Sci. 2007, 58, (4), 932-944. 65. Zhu, M.; Frandsen, C.; Wallas, A. F.; Legg, B.; Khalid, S.; Zhang, H.; Mørup, S.; Banfield, J. F.; Waychunas, G. A., Precipitation pathways for ferrihydrite formation in acidic solutions. Geochim Cosmochim Acta 2016, 172, 247 - 264. 66. Fukushi, K.; Aoyama, K.; Yang, C.; Kitadai, N.; Nakashima, S., Surface complexation modeling for sulfate adsorption on ferrihydrite consistent with in situ infrared spectroscopic observations. Appl. Geochem. 2013, 36, 92-103. 67. Sverjensky, D. A., Standard states for the activities of mineral surface sites and species. Geochim. Cosmochim. Acta 2003, 67, (1), 17-28. 68. Wang, Z.; Giammar, D. E., Mass action expressions for bidentate adsorption in surface complexation modeling: theory and practice. Environ Sci Technol 2013, 47, (9), 3982-96. 69. Sverjensky, D. A., Prediction of the speciation of alkaline earths adsorbed on mineral surfaces in salt solutions. Geochim. Cosmochim. Acta 2006, 70, (10), 2427-2453. 70. Davis, J. A.; Leckie, J. O., Surface ionization and complexation at the oxide/water interface. 3. Adsorption of anions. J. Colloid Interface Sci. 1980, 74, (1), 32-43.

618 619 25 ACS Paragon Plus Environment


620 621

Table 1. EXAFS-determined structural parameters of sulfate adsorbed on ferrihydrite surfaces and of sulfate in solutions of different pH. Error bars of the parameters are given in parentheses. S-O d (Å)

622 623 624 625

Page 26 of 32

S-Fe

CNa σ2 (Å2)

d (Å)

CN

σ2 (Å2)a

∆E (eV)

R

Sol_pH 0.7

1.47 (1)

4

0.0021(6)

---

6 (2)

0.0119

Sol_pH 2.0

1.49 (1)

4

0.0015 (6)

---

11 (3)

0.0164

Sol_pH 5.5

1.49 (1)

4

0.0008 (6)

---

11 (2)

0.0125

D_0.02_3

1.48 (1)

4

0.0009 (3)

3.22 (0.04)

1.3 (0.8)

0.006

11 (3)

0.0219

D_0.02_4

1.48 (1)

4

0.0004 (5)

3.24 (0.04)

1.9 (1.0)

0.006

11 (3)

0.0186

D_0.02_6

1.47 (2)

4

X-ray Absorption Spectroscopic Quantification and ... - ACS Publications

Recommend Documents