Sulfate Local Coordination Environment in Schwertmannite

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Environmental Science & Technology

Sulfate Local Coordination Environment in Schwertmannite

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Xiaoming Wang†, ‡, Chunhao Gu†, Xionghan Feng‡, *, and Mengqiang Zhu†, *

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†

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82071

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‡

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Ministry of Agriculture, College of Resources and Environment, Huazhong Agricultural

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University, Wuhan 430070, China

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

Key Laboratory of Arable Land Conservation (Middle and Lower Reaches of Yangtze River),

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*Corresponding authors:

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Mengqiang Zhu, Tel: +1 307-766-5523; Email: [email protected]

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Xionghan Feng, Tel: +86 27-87280271; E-mail: [email protected]

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The total words: 4593 + 300 × 8 = 6993

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Submitted to Environmental Science and Technology

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1

ACS Paragon Plus Environment


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ABSTRACT

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Schwertmannite, a nano-crystalline ferric oxyhydroxy-sulfate mineral, plays an important role

22

in many environmental geochemical processes in acidic sulfate-rich environments. The sulfate

23

coordination environment in schwertmannite, however, remains unclear, hindering our

24

understanding of the structure, formation and environmental behavior of the mineral. In this study,

25

sulfur K-edge X-ray absorption near edge structure (XANES) and extended X-ray absorption fine

26

structure (EXAFS) spectroscopic analyses in combination with infrared spectroscopy were used to

27

determine the sulfate local atomic environment in wet and air-dried schwertmannite samples after

28

incubation at various pHs and ionic strengths. Results indicate that sulfate exists as both inner- and

29

outer-sphere complexes in schwertmannite. Regardless of the sample preparation conditions, the

30

EXAFS-determined

31

bidentate-binuclear sulfate inner-sphere complexes. XANES spectroscopy shows that the

32

proportion of the inner-sphere complexes decreases with increasing pH for both wet and dried

33

samples and that the dried samples contained much more inner-sphere complexes than the wet

34

ones at any given pH. Assuming that schwertmannite is a distorted akaganéite-like structure, the

35

sulfate inner-sphere complexation suggests that, the double chains of the edge-sharing Fe

36

octahedra, enclosing the tunnel, must contain defects, on which reactive singly-Fe coordinated

37

hydroxyl functional groups form for ligand exchange with sulfate. The drying effect suggests that

38

the tunnel contains readily-exchanged H2O molecules in addition to sulfate ions.

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INTRODUCTION

S-Fe

inter-atomic

distances

are

3.22

–

3.26

Å,

indicative

of

40

Schwertmannite is a nano-crystalline ferric oxyhydroxy-sulfate mineral with a variable

41

chemical composition, simplified as Fe8O8(OH)8-2x(SO4)x·nH2O where 1 ≤ x ≤ 1.75.1 It often 2


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forms in acidic sulfate-rich environments, such as acid mine drainage,2, 3 in which schwertmannite

43

plays important role in controlling the fate and transport of various contaminants.4-7 Although

44

some may disagree,8, 9 most researchers believe that schwertmannite is a structurally-distorted

45

akaganéite, and double chains of edge-sharing FeO6 octahedra form 2 × 2 tunnels where sulfate

46

ions are presumably located.2, 10-12 As a structural component, sulfate is indispensable for the

47

formation and stabilization of schwertmannite.10 Removal of sulfate destabilizes the structure that

48

subsequently transforms to goethite,6, 13 with the extent and rate of the transformation depending

49

on SO42- content in schwertmannite.5, 14 Sulfate also affects the environmentally-relevant reactivity

50

of schwertmannite. Sulfate enhances adsorption of toxic metal cations on schwertmannite through

51

formation of ternary sulfate-metal complexes.15, 16 Toxic oxyanion adsorption on schwertmannite

52

can occur via ion- and/or ligand exchange with tunnel sulfate in addition to with hydroxyl

53

functional groups.17

54

Discovering the sulfate coordination environment, therefore, has important implications for

55

understanding the role of sulfate in schwertmannite formation, structural stabilization and

56

reactivity. Based on Infrared (IR) spectroscopy, Bigham et al.10 proposed that sulfate form

57

inner-sphere complexes with Fe(III) in the schwertmannite bulk, whereas Boily et al.18 suggested

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that sulfate ions exist as both inner- and outer-sphere complexes and that the later could be

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hydrogen-bonded species or free protonated sulfate. Based on atomic pair distribution function

60

(PDF) analysis, X-ray diffraction (XRD), and density functional theory, Fernandez-Martinez et

61

al.11 concluded similarly as Boily et al.18 that two sulfate ions are present per unit cell in the tunnel

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structure with one as an outer-sphere complex and the other as an inner-sphere complex. However,

63

none of these studies was able to unambiguously determine the local atomic environment of 3



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sulfate.19

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Recently, S K-edge extended X-ray absorption fine structure (EXAFS) spectroscopy was used

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to directly characterize the structure of sulfate surface complexes on ferrihydrite.19 In the present

67

study, we used S K-edge XANES and EXAFS spectroscopy to address whether sulfate

68

inner-sphere complexes exist in the schwertmannite structure and if they do, how they are

69

coordinated to Fe(III). To investigate the impacts of environmental factors on sulfate complexation,

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both wet and air-dried samples were examined after incubation at various pHs and ionic strength.

71

Examination of the impacts of these factors also assists in unveiling the sulfate complexation

72

mechanism. Schwertmannite solid phase changes by incubation were characterized using XRD, Fe

73

K-edge EXAFS and PDF analyses.

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MATERIALS AND METHODS

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Synthesis of Schwertmannite. One liter deionized water was preheated to 60 oC in an oil

77

bath, followed with a quick addition of 5.4 g FeCl3·6H2O and 1.5 g of Na2SO4.10 Schwertmannite

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precipitated immediately and the obtained suspension was maintained at 60 oC for additional 12

79

min under stirring condition. After cooling to room temperature (22 oC), the suspension was

80

dialyzed for ~ 7 days with its final conductivity below 20 µs cm-1. The solid was collected by

81

centrifugation and air dried for two days under ambient conditions. The dried powder was ground

82

and stored at 4 oC in a refrigerator.

83

Incubation of Schwertmannite. A stock schwertmannite suspension (15 g L-1) was prepared

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by dispersing dry schwertmannite powder in 0.05 M or 0.5 M NaNO3 solution by ultrasonication

85

for 15 min followed by overnight stirring. The pH of the obtained suspension was ~ 3.2. Ten mL 4


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of the suspension was transferred to a 15-mL polyethylene centrifuge tube with its pH adjusted to

87

2, 3.2, 4, 5, 6, 7, or 8. The tubes were shaken on an orbital rotator for 24 h, during which, the pHs

88

of the suspensions were maintained at the pre-set values by adding 0.1 or 1 M NaOH or HNO3.

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After that, each suspension passed through a 0.2 µm membrane filter mounted on a vacuum

90

apparatus. The membrane loaded with the solid was divided into two parts with one kept wet and

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the other air dried. Both parts were subject to S K-edge XANES and EXAFS spectroscopic

92

analyses and attenuated total reﬂectance Fourier transform infrared (ATR-FTIR) spectroscopic

93

measurements as well. The air-dried samples were further analyzed by XRD, Fe K-edge EXAFS

94

and X-ray atomic PDF analysis. The filtrates were measured for dissolved S and Fe concentration

95

using inductively coupled plasma optical emission spectroscopy (ICP-OES) (Perkin Elmer). The S

96

and Fe contents in the air-dried solids were measured after dissolving ~ 10 mg powder of each

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sample in 10 mL solution consisting of 0.2 M (NH4)2C2O4 and 0.2 M H2C2O4 (pH 3).

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High Energy X-ray Scattering (HEXC) and Fe K-edge EXAFS Spectroscopy. To examine

99

potential solid phase changes after incubation, HEXC data were collected from the selected dried

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samples using X-rays of 58.6491 keV (λ = 0.2114 Å) for XRD and PDF analyses at beamline

101

11-ID-B at the Advanced Photon Source (APS), Argonne National Laboratory. The procedures for

102

data collection and processing are as described in Zhu et al.20 and SI-1. Fe EXAFS data were

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collected at beamline 13-BM-D at APS. The detailed procedure for data collection was provided in

104

the supporting information (SI-2).

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ATR-FTIR Spectroscopy. ATR-FTIR spectra were collected with a PerkinElmer

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spectrometer using the single bounce model with a diamond internal reflection element (IRE) at

107

ambient temperature. The wet and dried samples were coated directly on the diamond crystal 5



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surface. The spectra were collected against the air background. Thirty two scans with a resolution

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of 4 cm-1 were collected and averaged for each sample.

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S K-edge XANES and EXAFS Spectroscopy. Sulfur X-ray absorption spectra were

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collected in fluorescence mode at beamline 4-3 at the Stanford Synchrotron Radiation

112

Light-source (SSRL). The monochromator, a Si(111) crystal, was calibrated by setting the first

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absorption maxima of sodium thiosulfate at 2472.02 eV. S K-edge XANES spectra were collected

114

using a Passivated Implanted Planar Silicon (PIPS) detector with fine step sizes and two scans

115

were collected for each sample. S K-edge EXAFS spectra were collected using a four-element

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silicon drift detector (Vortex-ME4) that discriminates the fluorescence signal of S from that of Cl,

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avoiding the Cl interference with the S EXAFS data collection. Six to ten scans were collected and

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averaged for each sample. The XANES portion associated with the EXAFS data was collected

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with coarse step sizes and shorter dwelling time to save time for data collection of the EXAFS

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portion. The XANES data collected using the PIPS detector were those used for the XANES

121

analysis.

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For the data collection, the air-dried powders were scattered on the adhesive side of a

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polyester film that was mounted on a cell made of acrylic plate. Each wet sample was pasted on

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the adhesive side and the other side was sealed with polyethylene film window to prevent

125

dehydration of the samples during data collection. Both the polyester and polyethylene films were

126

S and Cl free. The samples were placed in a measurement chamber flowed with dry helium gas so

127

that beam attenuation caused by air and water absorption was minimized. At the completion of the

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data collection of each wet sample, the sample was immediately taken out of the chamber and

129

visually inspected for whether the sample was dried out. All wet samples were confirmed to be 6


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wet after data collection.

131

The program SixPack21 was used to average the XAS spectra, whereas background removal

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and post-edge normalization of the averaged spectra were performed with the program Athena.22

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The post-edge was normalized over 50 -700 eV and the parameter Rbkg was set at 0.9. The

134

obtained k3-weighted EXAFS spectra were Fourier transformed using the Kaiser-Bessel window

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over 4 - 13 Å-1. The EXAFS spectra were fitted using SixPack to determine sulfate local atomic

136

parameters according to Zhu et al.19 The amplitude reduction factor (S02) was fixed at 0.96 for all

137

samples.19

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For the S XANES spectra, similar data deduction procedures were used except that the

139

post-edge normalization was conducted over 50 - 330 eV. A linear combination fitting (LCF)

140

analysis of the XANES spectra was performed over 2462 - 2512 eV using the program Athena.

141 142

RESULTS

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High Energy X-ray Scattering. The XRD patterns and atomic PDFs within 1 – 5 Å for the

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dried schwertmannite samples are depicted in Figure 1. The XRD patterns show the eight major

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characteristic diffraction peaks of schwertmannite (Figure 1a),1, 11 and (110) peak of goethite. The

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goethite peak is very weak, indicating only minor goethite in the solid. This peak becomes more

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pronounced with increasing incubation pH, indicative of schwertmannite transformation to

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goethite at elevated pHs, in good agreement with Boily et al.18, which can be understood from the

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following reaction:

150

Fe8O8(OH)6SO4 + 2 OH- ⇔ 8 α-FeOOH + SO42-

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A closer inspection of the XRD peaks allows for examining subtle structural alteration of 7



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schwertmannite by incubation prior to its transformation to goethite. With increasing pH, the

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intensities of (200) and (301) peaks slightly decrease; meanwhile, (301), (225), and (040) peaks

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slightly shift to lower d-spacings and (200) peak to a higher d-spacing, which can be seen more

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clearly in an enlarged figure (Figure SI-3). These variations should not be caused by the presence

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of minor goethite because its contributions at these peak positions are negligible (Figure 1a). The

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peak variations likely suggest a schwertmannite structural distortion caused by the loss of sulfate

158

(see below) and the equilibrium at high pH during incubation. The incubation conditions favor

159

schwertmannite transformation to more stable goethite

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schwertmannite could be the precursor to goethite 9. These XRD analyses coincide with the PDF

161

modeling results showing monotonic variation trends in unit cell parameters (Table SI-1). But it

162

should be cautious to interpret the PDF results because of the potential interference from the

163

co-existing goethite as the PDF analysis utilizes the entire HEXS patterns. Additional Fe K-edge

164

EXAFS analyses of the schwertmannite samples is provided in the supporting information (SI-2).

165

Note that the contribution of the minor goethite to sulfate adsorption is negligible and should not

166

affect the following analyses of the sulfate coordination environment in schwertmannite.

5, 14

. Thus, the structurally distorted

167

Chemical Composition and Solution Analysis. The S content in as-synthesized

168

schwertmannite is 1.404 mmol g-1. With increasing incubation pH from 2 to 8, 2% to 51% of the

169

initial sulfate is released from the solid in 0.05 M NaNO3 solution (Figure 2). Increasing the ionic

170

strength to 0.5 M releases slightly more sulfate, suggesting that the structure contains

171

exchangeable sulfate outer-sphere complexes.23, 24 Schwertmannite dissolves slightly at pH 2 as

172

suggested by the detection of dissolved Fe (Figure SI-4a).

173

S K-edge XANES Spectra. The XANES spectra of the pH 5.5 sulfate solution, jarosite and 8


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selected schwertmannite samples, as well as the difference pre-edges with respect to the sulfate

175

solution spectra, are given in Figure 3 and SI-5. The pre-edge (peak A) is present for jarosite and

176

all schwertmannite samples (Figure 3a, 3b) but not for the sulfate solution.19 The pre-edge is

177

caused by the electronic transition from S 1s orbital to 3p orbital that is hybridized with Fe 3d

178

orbitals.25 Its presence indicates schwertmannite contains sulfate inner-sphere complexes. The

179

pre-edge intensity slightly decreases with increasing pH, suggesting less inner-sphere complexes

180

at higher pHs. At a given pH, the pre-edges of the dried samples are much stronger than those of

181

the wet samples (Figure 3b), indicative of much more inner-sphere complexes in the dried samples.

182

That is, drying converts outer-sphere to inner-sphere complexes. The conversion is reversible upon

183

re-wetting as the XANES spectra of the re-wet sample is almost the same as that of the initial wet

184

sample (Figure SI-6).

185

The difference pre-edges are obtained by subtracting the solution pre-edge from those jarosite

186

and schwertmannite.2 The difference pre-edge of jarosite has two peaks (Figure 3d), consistent

187

with the study by Majzlan and Myneni.2 Compared to jarosite in which each sulfate is coordinated

188

to three Fe atoms, the peaks of schwertmannite are located at slightly lower energies (Figure 3d),

189

suggesting less than three Fe atoms coordinated to each schwertmannite sulfate.2 Little variation in

190

the peak positions with pH and hydration degree indicates the same type of the sulfate

191

inner-sphere complexes existing in the schwertmannite samples regardless of the incubation

192

conditions.

193

The XANES spectrum of the pH 5.5 Na2SO4 solution shows a sharp white line peak ( peak B)

194

(Figure 3c), assigned to the 1s → t2* transition of sulfur.26 With decreasing solution pH from 5.5

195

to 0.7, peak B broadens and its position shifts to a higher energy by ~ 0.22 eV (Figure SI-7), 9



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which is ascribed to the formation of HSO4- with increased molecular asymmetry compared to

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SO42-.2 A broad peak (peak D, Figure 3c) ~15 eV above the absorption edge of the sulfate solution

198

corresponds to the 1s → continuum-state transition.26 When sulfate complexes with Fe in

199

schwertmannite, both peaks significantly shift to higher energies (Figure 3c). The peak positions

200

further vary with incubation pH and sample drying status. Drying and decreasing pH shift both

201

peaks to higher energies (Table SI-8), whereas the ionic strength impact is not obvious (Table

202

SI-8). In addition, all schwertmannite samples have peak C (Figure 3c) which is absent in the

203

solution spectra. Peak C of the wet samples are much stronger and located at lower energies than

204

those of the dried ones. With increasing pH, the peak grows with its position shifting to a lower

205

energy. As the pre-edges, the variations of peak B, C, and D probably also result from the

206

proportional changes of the inner- and outer-sphere complexes.

207

XANES LCF analysis was used to semi-quantify the proportions of these complexes. Based

208

on the above trend, we assume that the two end members for inner- and outer-sphere complexes

209

for the LCF analysis, respectively, are the air-dried sample after incubation at pH 2 in 0.5 M

210

NaNO3 solution and the wet sample after incubation at pH 8 in 0.05 M NaNO3 solution. The end

211

members are approximate as both contain a mixture of inner- and outer-sphere complexes

212

although one type is dominant. Figure SI-9 shows the comparison of the fits to the data and Figure

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4 gives the obtained proportions as a function of pH and ionic strength. The high goodness of fit

214

values suggest that two species only, i.e., inner- and outer-sphere complexes, are sufficient to

215

account for sulfate speciation changes in schwertmannite under various incubation conditions. As

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pH increases from 2 to 8, the inner-sphere complex proportion decreases from 100% to ~ 49% for

217

the dried samples and from ~ 60% at pH 2 to 0% at pH ≥ 6 for the wet ones (Figure 4c, 4d). The 10


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218

ionic strength effect on the proportions is insignificant, probably due to the LCF quantification

219

that is not accurate enough to capture the small proportional changes.

220

S K-edge EXAFS Spectroscopy. Compared to the XANES analyses on the structure of

221

sulfate inner-sphere complexes, the following EXAFS fitting provides more direct structural

222

information. Figure 5 shows the EXAFS spectra and their fits, and the obtained parameters are

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listed in Table 1. Jarosite spectra have additional oscillations in k space, e.g., at ~ 8.5 Å-1 and ~ 10

224

Å-1, compared to the sulfate solution spectra, which result from the backscattering of the three

225

structural Fe atoms. The dried schwertmannite samples have similar but weak peaks in k space at

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these locations (Figure 5a, 5c). These peaks are even weaker for the wet samples due to the low

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proportion of the inner-sphere complexes, which also results in the similarity of the wet sample

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and the solution EXAFS spectra (Figure SI-10). The EXAFS fitting results show that the S-O

229

bond lengths are 1.480 - 1.490 Å for all wet and dried samples. There are 1.0 – 1.2 Fe atoms

230

located at 3.22 – 3.24 Å from S for the dried samples and 0.7 – 0.9 Fe atoms at 3.23 – 3.26 Å for

231

the wet ones (Table 1). The fewer Fe atoms for the wet samples than for the dried samples are

232

consistent with the lower peak intensity at ~ 2.9 Å (R + ∆R) in R space (Figure 5b, 5d), where the

233

Fe atomic shell is located.

234

ATR-FTIR Spectroscopy. For both dried and wet samples, the sulfate ATR-FTIR spectra

235

consist of a broad triply degenerate asymmetric stretching (ν3) band at ~1105 cm-1 with two

236

shoulder bands at ~1040 cm-1 and 1170 cm-1, a ν1 fundamental of the symmetric sulfate stretching

237

at 980 cm-1, and a ν4 bending band at ~608 cm-1 (Figure 6), consistent with previous studies.1, 18

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These bands result from sulfate inner-sphere complexes. The shape of these bands are significantly

239

affected by sample incubation pH and hydration status. With increasing pH, the ν3 band splits to a 11



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240

lower degree for both dry and wet samples, suggesting less inner-sphere complexes.27 The

241

decreased intensities of these bands are ascribed to the increasing sulfate loss at higher pHs.

242

Moreover, the ν3 band for the dried samples splits more than that for the wet samples at a given pH,

243

suggesting that dried samples contain more inner-sphere complexes. The ATR-FTIR spectra of

244

schwertmannite samples incubated in 0.05 M vs 0.5 M NaNO3 solution (Figure SI-11) are almost

245

the same, indicating insignificant ionic strength effect. These results are consistent with the

246

XANES analyses.

247

Changes are also observed for other IR bands. The spectra (Figure 6a) of the dried samples

248

show the schwertmannite characteristic OH deformation (δOH) (~ 840 cm-1) and Fe-O stretch

249

bands (685 cm-1)10. The intensity of the δOH band decreases with increasing pH; meanwhile, it

250

splits into goethite characteristic δ-OH and γ-OH bands (~ 892 cm-1 and ~ 795 cm-1,

251

respectively18). The shape of the Fe-O stretch band remains almost unchanged, but its intensity

252

decreases with increasing pH, which may be due to increased schwertmannite structural disorder.

253

The δOH and Fe-O stretch bands of the wet samples (Figure 6b) exhibit significant differences

254

from those of the dried samples. The δOH and the Fe-O stretch bands shift to lower wavenumbers

255

(~786 cm-1 and ~ 665 cm-1, respectively), which may be caused by more extensive hydrogen

256

bonding in the schwertmannite structural frame at the wet state than at the drying state.

257 258

DISCUSSION

259

Tunnel versus Surface-adsorbed Sulfate. Sulfate associated with schwertmannite is

260

adsorbed either on the external surface or inside tunnels, and the former is presumably more prone

261

to desorption. Thus, the tunnel sulfate can become more dominant with increasing pH because of 12


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262

the preferential desorption of the surface sulfate. Given that surface adsorbed sulfate is negligible

263

at pH 8 on ferrihydrite,15, 23 it is plausible to infer that the sulfate remained in the schwertmannite

264

at this pH is mainly located in the tunnel, which would account for ~ 50 % of total sulfate initially

265

present (Figure 2). These tunnel sulfate ions are mainly outer-sphere complexed based on the

266

XANES analysis (Figure 3b) and not readily exchangeable under the experimental conditions of

267

this study.

268

The schwertmannite incubated at lower pHs than 8 must contain greater than 50% tunnel

269

sulfate as the tunnel sulfate desorbs less at lower pHs. The dominance of the tunnel sulfate can

270

also be inferred from the high sulfate content of schwertmannite. The S to Fe molar ratio (SSA10:

271

~ 250 m2·g-1) is ~ 0.17 for schwertmannite at pH 4. At the same pH, the ratios are ~ 0.09 and 0.01,

272

respectively, for sulfate-adsorbed ferrihydrite (SSA: 600 m2·g-1)15 and goethite (SSA: 82 m2·g-1)28

273

at the maximum sulfate adsorption loadings. When normalized to their respective specific surface

274

areas, schwertmannite contains 5.31 µmol sulfate·m-2, almost four times higher than those of

275

ferrihydrite (1.42 µmol sulfate·m-2) and goethite (1.41 µmol sulfate·m-2). This is possible only

276

when the majority of the sulfate are located in the tunnels rather than on the external surface.

277

Regardless of the proportion of the surface sulfate, the same S-Fe inter-atomic distance is

278

observed, indicating that surface sulfate has similar local atomic environment as the tunnel sulfate.

279

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

280

dried samples are 3.22 – 3.24 Å (Table 1), close to the distances of sulfate adsorption complexes

281

on ferrihydrite (3.22 ± 0.05 Å).19 Quantum chemical calculations predict that the S-Fe distances

282

are 3.25 – 3.29 Å for a bidentate-binuclear complex and 3.46 – 3.49 Å for a

283

monodentate-mononuclear complex.29,

30

Our observed S-Fe distance is also too long for 13



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bidentate-mononuclear and tridentate complexes 31. Therefore, the sulfate inner-sphere complex of

285

the dried samples is likely bidentate-binuclear, although a bent monodentate H-bonded complex

286

might have a similar distance.32 This conclusion is consistent with the selenate complexation in

287

selenate-exchanged schwertmannite.33 This is not surprising as sulfate and selenate have very

288

similar chemical nature. The smaller CNs (1.0 – 1.2) of the Fe shell than the theoretical value of 2

289

for a bidentate-binuclear complex can be ascribed to the co-existence of the outer-sphere

290

complexes.19

291

The wet samples have slightly longer S-Fe distances (3.23 – 3.26 Å) than the dried samples

292

(Table 1). But we believe that the inner-sphere complexes have the same structure as those of the

293

dried samples as the XANES spectra of all wet samples are well fitted with one of the dry sample

294

XANES spectra (i.e., the end member). The observed longer distances are likely due to the larger

295

uncertainties in EXAFS fitting (Table 1) because the wet samples contain much less inner-sphere

296

complexes compared to the dried samples.

297

The Structure of Outer-sphere Complexes. Our analyses clearly show that sulfate

298

outer-sphere complexes exist in the structure of schwertmannite. It is likely that only one chemical

299

form of outer-sphere complexes exists in schwertmannite regardless of the sample incubation pH

300

and ionic strength. XANES spectral profiles are very sensitive to changes in local environment

301

and electronic structure of the absorbing atom. If more than one type of outer-sphere complexes

302

present, three components at least are required for XANES LCF analysis and the fits would be

303

poor with only two end members, which contradicts to the obtained high goodness of fit values for

304

our samples. They outer-sphere complex must distinct from the free sulfate ion in solution based

305

on the fact that the peak C, observed for the wet schwertmannite samples, is not present in the 14


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306

XANES spectra of the sulfate solution (Figure 3c and Figure SI-5d). The outer-sphere complex

307

may not be a free protonated sulfate; otherwise, the XANES spectra would display the

308

characteristic feature of bisulfate, as observed for the pH 0.7 sulfate solution (Figure SI-7). More

309

detailed information on the chemical environment of the outer-sphere complexes needs to be

310

revealed to interpret the origin of peak C in future studies.

311

Structure and Formation of Schwertmannite. Our results on the sulfate atomic

312

coordination environment provide important insights into the structure and formation of

313

schwertmannite. Each O atom in the tunnel structure of akaganéite is coordinated to three Fe

314

atoms, which is not reactive for ligand exchange reactions,34, 35 such as sulfate adsorption36. If

315

schwertmannite has a similar tunnel structure as akaganéite does10, 11, sulfate is not able to form

316

inner-sphere complexes inside the tunnel. To enable the formation, the double chains of the

317

edge-sharing Fe octahedra, enclosing the tunnel, must have some structural defects, with respect to

318

akaganéite, on which singly-Fe coordinated functional groups (i.e., ≡Fe-OH-1/2 and ≡Fe-OH2+1/2)

319

form for ligand exchange with sulfate (Figure 7). The defects in schwertmannite could also

320

contribute to the distortion and enlargement of the tunnels so that the tunnels are big enough to

321

host sulfate ions that are larger than Cl-.

322

The conversion of sulfate outer-sphere complexes to inner-sphere complexes, driven by

323

sample dehydration suggests that the wet schwertmannite samples contain water molecules inside

324

the tunnel in addition to sulfate ions. The reversibility of the conversion indicates that the tunnel

325

readily undergoes hydration and dehydration. The tunnel water molecules may form H-bonds with

326

the schwertmannite structural framework and the tunnel sulfate as well. The loss of the tunnel

327

water molecules breaks some of the H-bonds, which could contribute to the peak position shifts of 15



328 329

Page 16 of 28

both δOH and Fe-O stretching IR bands upon drying (Figure 6 and Figure SI-11) Our results show that considerable sulfate inner-sphere complexes form at low pHs (e.g., ~ 50%

330

at pH 3.2) for the wet samples (Figure 4c). It is reasonable to infer that a significant proportion of

331

sulfate inner-sphere complexes also form on the intermediate FeO6 molecular clusters during the

332

formation of schwertmannite, directing the nucleation and growth of the mineral (Figure 7). But

333

this does not exclude possible contribution from outer-sphere complexation to schwertmannite

334

formation.2

335

Environmental Implications

336

Schwertmannite can effectively sequester toxic oxyanions, such as arsenic, selenium and

337

chromate,4, 6, 7 and metal ions5 through adsorption or structural incorporation. Structural sulfate of

338

schwertmannite inevitably interferes with these reactions, such as enhancing metal sorption by

339

forming metal-sulfate ternary complexes.15 The structure of the ternary complexes needs to be

340

implemented into surface complexation modeling to accurately predict toxic metal fate and

341

transport in acid mine drainage environment.15,

342

inner-sphere complexes provides constrains on the possible structures of the ternary complexes.

343

This study also has important implications for determining the crystal structure of schwertmannite.

344

In a schwertmannite structural model, one needs to know where and how sulfate ions are bound in

345

the structure, which, however, is hard to determine using X-ray scattering analyses.11 Discovering

346

the dominance of bidentate-binuclear sulfate complexes in dried schwertmannite in this study is

347

one of the critical steps towards the final resolution of the schwertmannite structure. Overall, this

348

study provides numerous important insights into the structure, formation and reactivity of

349

schwertmannite, an abundant and highly reactive nanomineral widely occurring in natural acidic

16

Our finding of sulfate bidentate-binuclear

16


Page 17 of 28

350


environments.

351 352

Acknowledgments

353

This work was partially supported by the Wyoming Agricultural Experimental Station

354

Competitive Grants Program. XW is grateful to the Chinese Scholarship Council fellowship. We

355

thank Dr. Alejandro Fernandez-Martinez at the University of Grenoble and CNRS and Dr. Glenn

356

A. Waychunas at the Lawrence Berkeley National Laboratory for stimulating discussions. Use of

357

the Stanford Synchrotron Radiation Light source is supported U.S. DOE-BES under Contract No.

358

DE-AC02-76SF00515. Use of the Advanced Photon Source, Argonne National Laboratory, is

359

supported by U.S. DOE-BES under Contract DE-AC02-06CH11357.

360

Supporting Information

361

The supporting information is available free of charge via the internet at http://pubs.acs.org.,

362

including (1) PDF and Fe EXAFS analyses, (2) S XANES spectra of schwertmannite incubated in

363

0.5 M NaNO3 solution and of sulfate solutions, (3) the peak positions of the XANES spectra, (4)

364

the S XANES LCF results, (5) the comparison of XANES spectra for solutions and wet and dried

365

samples, (6) ATR-FTIR spectra of schwertmannite incubated in 0.5 M NaNO3, and (7) the Fe

366

content in the solution and solid phases.

367 368

References

369 370 371 372 373 374

1.

Bigham, J. M.; Carlson, L.; Murad, E., Schwertmannite, a new iron oxyhydroxysulphate from pyhasalmi, finland,

and other localities. Mineral. Mag. 1994, 58, (393), 641-648. 2.

Majzlan, J.; Myneni, S. C. B., Speciation of iron and sulfate in acid waters: aqueous clusters to mineral

precipitates. Environ. Sci. Technol. 2005, 39, (1), 188-194. 3.

Peretyazhko, T.; Zachara, J. M.; Boily, J. F.; Xia, Y.; Gassman, P. L.; Arey, B. W.; Burgos, W. D., Mineralogical

transformations controlling acid mine drainage chemistry. Chem. Geol. 2009, 262, (3-4), 169-178. 17



375 376 377 378 379 380 381 382 383 384 385 386 387 388 389 390 391 392 393 394 395 396 397 398 399 400 401 402 403 404 405 406 407 408 409 410 411 412 413 414 415 416 417 418

4.

Page 18 of 28

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. 5.

Jonsson, J.; Persson, P.; Sjoberg, S.; Lovgren, L., Schwertmannite precipitated from acid mine drainage: phase

transformation, sulphate release and surface properties. Appl. Geochem. 2005, 20, (1), 179-191. 6.

Regenspurg, S.; Peiffer, S., Arsenate and chromate incorporation in schwertmannite. Appl. Geochem. 2005, 20,

(6), 1226-1239. 7.

3-

2-

Waychunas, G. A.; Xu, N.; Fuller, C. C.; Davis, J. A.; Bigham, J. M., XAS study of AsO4 and SeO4 substituted

schwertmannites. Physica B 1995, 208, (1-4), 481-483. 8.

Loan, M.; Cowley, J. M.; Hart, R.; Parkinson, G. M., Evidence on the structure of synthetic schwertmannite. Am.

Mineral. 2004, 89, (11-12), 1735-1742. 9.

French, R. A.; Caraballo, M. A.; Kim, B.; Rimstidt, J. D.; Murayama, M.; Hochella, M. F., The enigmatic iron

oxyhydroxysulfate nanomineral schwertmannite: Morphology, structure, and composition. Am. Mineral. 2012, 97, (8-9), 1469-1482. 10. Bigham, J. M.; Schwertmann, U.; Carlson, L.; Murad, E., A poorly crystallized oxyhydroxysulfate of iron formed by bacterial oxidation of Fe(II) in acid-mine waters. Geochim. Cosmochim. Acta 1990, 54, (10), 2743-2758. 11. Fernandez-Martinez, A.; Timon, V.; Roman-Ross, G.; Cuello, G. J.; Daniels, J. E.; Ayora, C., The structure of schwertmannite, a nanocrystalline iron oxyhydroxysulfate. Am. Mineral. 2010, 95, (8-9), 1312-1322. 12. Knorr, K. H.; Blodau, C., Controls on schwertmannite transformation rates and products. Appl. Geochem. 2007, 22, (9), 2006-2015. 13. Paikaray, S.; Peiffer, S., Abiotic schwertmannite transformation kinetics and the role of sorbed As(III). Appl. Geochem. 2012, 27, (3), 590-597. 14. Burton, E. D.; Johnston, S. G.; Kraal, P.; Bush, R. T.; Claff, S., Sulfate availability drives divergent evolution of arsenic speciation during microbially mediated reductive transformation of schwertmannite. Environ. Sci. Technol. 2013, 47, (5), 2221-2229. 15. Swedlund, P. J.; Webster, J. G., Cu and Zn ternary surface complex formation with SO4 on ferrihydrite and schwertmannite. Appl. Geochem. 2001, 16, (5), 503-511. 16. 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. 17. Antelo, J.; Fiol, S.; Gondar, D.; Perez, C.; Lopez, R.; Arce, F., Cu(II) incorporation to schwertmannite: Effect on stability and reactivity under AMD conditions. Geochim. Cosmochim. Acta 2013, 119, 149-163. 18. 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. 19. Zhu, M.; Northrup, P.; Shi, C.; Billinge, S. J. L.; Sparks, D. L.; Waychunas, G., Structure of sulfate adsorption complexes on ferrihydrite. Environ. Sci. Technol. Lett. 2014, 1, 97-101. 20. Zhu, M. Q.; Farrow, C. L.; Post, J. E.; Livi, K. J. T.; Billinge, S. J. L.; Ginder-Vogel, M.; Sparks, D. L., Structural study of biotic and abiotic poorly-crystalline manganese oxides using atomic pair distribution function analysis. Geochim. Cosmochim. Acta 2012, 81, 39-55. 21. Webb, S. M., SIXpack: a graphical user interface for XAS analysis using IFEFFIT. Phys. Scripta 2005, T115, 1011-1014. 22. Ravel, B.; Newville, M., ATHENA, ARTEMIS, HEPHAESTUS: data analysis for X-ray absorption spectroscopy using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537-541. 23. 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. 18


Page 19 of 28

419 420 421 422 423 424 425 426 427 428 429 430 431 432 433 434 435 436 437 438 439 440 441 442 443 444 445 446


24. Rietra, R. P. J. J.; Hiemstra, T.; van Riemsdijk, W. H., Comparison of selenate and sulfate adsorption on goethite. J. Colloid Interf. Sci. 2001, 240, (2), 384-390. 25. Okude, N.; Nagoshi, M.; Noro, H.; Baba, Y.; Yamamoto, H.; Sasaki, T. A., P and S K-edge XANES of transition-metal phosphates and sulfates. J. Electron Spectrosc. 1999, 101, 607-610. 26. Myneni, S. C. B., X-ray and vibrational spectroscopy of sulfate in earth materials. Rev. Mineral Geochem. 2000, 40, 113-172. 27. Peak, D.; Ford, R. G.; Sparks, D. L., An in situ ATR-FTIR investigation of sulfate bonding mechanisms on goethite. J. Colloid Interf. Sci. 1999, 218, (1), 289-299. 28. Liu, F.; He, J. Z.; Colombo, C.; Violante, A., Competitive adsorption of sulfate and oxalate on goethite in the absence orpresence of phosphate. Soil Sci. 1999, 164, (3), 180-189. 29. 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. 30. 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, (4), 978-988. 31. Lafferty, B. J.; Ginder-Vogel, M.; Zhu, M.; Livi, K. J. T.; Sparks, D. L., Arsenite oxidation by a poorly crystalline manganese-oxide. 2. Results from X-ray absorption spectroscopy and X-ray diffraction. Environ. Sci. Technol. 2010, 44, (22), 8467-8472. 32. Loring, J. S.; Sandstrom, M. H.; Noren, K.; Persson, P., Rethinking arsenate coordination at the surface of goethite. Chem. Eur. J. 2009, 15, (20), 5063-5072. 33. Peak, D.; Sparks, D. L., Mechanisms of selenate adsorption on iron oxides and hydroxides. Environ. Sci. Technol. 2002, 36, (7), 1460-1466. 34. Hiemstra, T.; VanRiemsdijk, W. H., A surface structural approach to ion adsorption: the charge distribution (CD) model. J. Colloid Interf. Sci. 1996, 179, (2), 488-508. 35. Rietra, R. P. J. J.; Hiemstra, T.; van Riemsdijk, W. H., The relationship between molecular structure and ion adsorption on variable charge minerals. Geochim. Cosmochim. Acta 1999, 63, (19-20), 3009-3015. 36. Kozin, P. A.; Boily, J. F., Proton binding and ion exchange at the akaganeite/water interface. J. Phys. Chem. C 2013, 117, (12), 6409-6419.

447 448

19



Page 20 of 28

449

Table 1. The structural parameters of sulfate in schwertmannite and references determined from S

450

K-edge EXAFS fitting. S-O Samples

S-Fe

pH d (Å)

CN a

σ2 (Å)

d (Å)

CN

σ2 b

∆E (eV)

R

Solution

5.5

1.493 (6)

4

0.0003 (3)

--

--

--

14 (3)

0.030

Jarosite

--

1.485 (7)

4

0.0004 (3)

3.24 (3)

2.7 (9)

0.006

11 (3)

0.027

Dried, I = 0.05 M

3.2

1.483 (4)

4

0.0015 (2)

3.22 (2)

1.2 (4)

0.006

11 (2)

0.008

5

1.485 (5)

4

0.0011 (2)

3.22 (3)

1.1 (5)

0.006

12 (2)

0.012

8

1.488 (5)

4

0.0013 (3)

3.22 (3)

1.2 (6)

0.006

13 (2)

0.014

3.2

1.486 (4)

4

0.0017 (2)

3.22 (2)

1.2 (5)

0.006

12 (2)

0.010

5

1.485 (5)

4

0.0012 (2)

3.23 (3)

1.0 (5)

0.006

12 (2)

0.012

8

1.488 (5)

4

0.0015 (3)

3.24 (4)

1.0 (6)

0.006

13 (2)

0.015

3.2

1.480 (7)

4

0.0009 (3)

3.24 (6)

0.9 (9)

0.006

13 (3)

0.027

5

1.485 (7)

4

0.0007 (3)

3.24 (7)

0.7 (9)

0.006

12 (3)

0.026

8

1.487 (7)

4

0.0005 (3)

3.26 (7)

0.9 (9)

0.006

13 (3)

0.028

3.2

1.480 (7)

4

0.0005 (3)

3.24 (6)

0.9 (9)

0.006

10 (3)

0.029

5

1.488 (7)

4

0.0010 (4)

3.26 (8)

0.7 (9)

0.006

13 (3)

0.028

8

1.490 (7)

4

0.0006 (4)

3.23 (9)

0.7 (9)

0.006

13 (3)

0.029

Dried, I = 0.5 M

Wet, I = 0.05 M

Wet, I = 0.5 M

a

The coordination number (CN) of S-O for all samples are fixed at 4.

b

The obtained Debye-Waller factor (σ2) from the fitting are round 0.006 for the dried samples and below

0.001 for the wet samples if floated during the fitting. To compare the CNs for different samples, we fixed σ2 at 0.006 for the fitting of all samples. 451 452

20


Page 21 of 28


pH 8 pH 7 pH 5 pH 3.2 pH 2 (301) (200) 3.35 5.27 Å G (110) 4.18

(b) 12

2.54 (221)

Fe-O

pH 8 pH 7 pH 5 pH 3.2 pH 2

Fe-Fe (CS)

9 (040) (042) (225) 1.51 1.46 1.65

Fe-Fe (ES) -2

(023) 2.28 (421) 1.95

G (r) (Å )

Intensity (a.u.)

(a)

6 3 0

S-O

-3 Goethite

6 5

453

4

3

2

-6 1.0

1.5

2.0

2.5

3.0

3.5

4.0

4.5

5.0

r (Å)

d spacing (Å)

454

Figure 1. X-ray diffraction patterns (G = goethite, goethite is included as a reference) (a) and

455

atomic pair distribution functions within 1 – 5 Å (b) for the air-dried schwertmannite samples after

456

incubated at various pHs in 0.05 M NaNO3 solution. The PDF peak located at ~ 1.48 Å

457

corresponds to the S-O atomic pair in sulfate and its intensity decreases with increasing pH is due

458

to the loss of more sulfate from schwertmannite at higher pHs.

459 460 461

21



-1

1.2

0.6 1.0

Solid

0.4

Sulfate in the solid (mmol g )

1.4

Circle (Ο): 0.5 M NaNO3

0.8

-1

Soluble sulfate (mmol g )

Square (): 0.05 M NaNO3

Page 22 of 28

Solution 0.8

0.2 0.6

0.0 2

3

4

5

6

7

8

462

pH

463

Figure 2. The amount of the sulfate released to the solution and of the sulfate remained in the solid

464

per gram schwertmannite after schwertmannite (15 g L-1) was incubated at various pHs in 0.05 M

465

or 0.5 M NaNO3 solutions.

466 467

22


Page 23 of 28


10

B

Normalized m(E)

8

6

D

Jarosite

A 0.4

0.2

Solution

C

A 0 2475

0.8

0.6

4

2

2480

2485

2490

2495

2500

2505

0.0 2476

2510

2477

E (eV)

468 10

1.6 1.4

B

(d)

D

6

2480

2481

0.2

Solution

4 2 0 2480

Jarosite 2481

2482

2483

2484

2485

E (eV)

1.2

C 1.0

Solution

0.8 2480

2479

Jarosite

8

Normalized m(E)

Normalized m(E)

1.8

2478

E (eV)

2.0 Normalized m(E)

(c)

469

(b)

pH 3.2_dry pH 5_dry pH 8_dry pH 3.2_wet pH 5_wet pH 8_wet Jarosite Sulfate solution_pH 5.5

Normalized m(E)

(a)

0.0

-0.2

Jarosite 2485

2490

2495

2500

2505

2510

2476

2477

E (eV)

2478

2479

2480

2481

E (eV)

470

Figure 3. The S K-edge XANES spectra (a), and their expanded pre-edge (b), white line and

471

post-edge region (c) for selected wet and air-dried schwertmannite samples after incubated at

472

various pHs in 0.05 M NaNO3 solution. The pre-edge regions of the difference spectra with

473

respect to pH 5.5 sulfate solution are given in (d). The spectra of jarosite and a Na2SO4 solution of

474

pH 5.5 were included as references.

475 476

23



(b)

8

Data Fit Differentiate Inner-sphere Outer-sphere

6

Normalization m(E)

Normalization m(E)

(a)

4

2

Page 24 of 28

8

6

4

2

0

0 2470

2480

2490

2500

2470

2510

2480

2490

Inner-sphere percentage (%)

(c) 100

0.05 M dried 0.5 M dried 0.05 M wet 0.5 M wet

2510

80 60 40 20 0

(d) 100 Outer-sphere percentage (%)

477

2500

E (eV)

E (eV)

80 60 40 20 0

2

3

4

5

6

7

8

2

3

pH

478

4

5

6

7

8

pH

479

Figure 4. The XANES spectra and their linear combination fits for representative air-dried (a) and

480

wet (b) schwertmannite samples incubated at pH 5 in 0.5 M NaNO3 solution, and the obtained

481

molar percentages of the sulfate inner- (c) and outer-sphere complexes (d) as a function of pH and

482

ionic strength.

483 484

24


Page 25 of 28


(a)

pH 8_0.5 M

(b)

pH 5_0.5 M

pH 5_0.5 M pH 3.2_0.5 M

-4

pH 8_0.05 M

pH 8_0.5 M

|χ(R)| (Å ) (a.u.)

K χ(k) (a.u.)

pH 3.2_0.5 M

pH 5_0.05 M

3

pH 3.2_0.05 M

pH 8_0.05 M pH 5_0.05 M

Jarosite

pH 3.2_0.05 M Jarosite

4

6

8

10

12

14

16

0

-1

1

K (Å )

485

2

3

4

5

6

R+ ∆R (Å)

(c)

Sulfate solution

(d)

pH 8_0.5 M

pH 8_0.5 M pH 5_0.5 M

-4

pH 3.2_0.5 M

Sulfate solution

|χ(R)| (Å ) (a.u.)

K χ(k) (a.u.)

pH 5_0.5 M

3

pH 8_0.05 M pH 5_0.05 M

pH 3.2_0.5 M pH 8_0.05 M pH 5_0.05 M pH 3.2_0.05 M

pH 3.2_0.05 M

4

486

6

8

10

12 -1

14

16

0

1

K (Å )

2

3

4

5

6

R+ ∆R (Å)

487

Figure 5. Sulfur K-edge EXAFS spectra and their fits for selected air-dried (a, b) and wet (c, d)

488

schwertmannite samples and references.

489 490

25



1105

Fe-O 685

Dried

ν3

1040

ν3

ν1

1170

980

ν3

(b)

ν3

610

δOH δOH γOH

Fe-O 665 ν4 608

Wet

1105

ν4

1040

892 843 795

pH 8

Abs (a.u.)

pH 7 pH 6 pH 5

δOH

ν3

1170

786

ν1

Abs (a.u.)

ν3

(a)

Page 26 of 28

pH 8

980

pH 7 pH 6 pH 5

pH 4

pH 4

pH 3.2

pH 3.2 pH 2

pH 2

1200

491

1100

1000

900

800

700

600

1200

1100

1000

900

800

700

600

-1

-1

Wavenumber (cm )

Wavenumber (cm )

492

Figure 6. ATR-FTIR spectra of the air-dried (a) and the wet (b) schwertmannite samples after

493

incubation at various pHs in 0.05 M NaNO3 solution.

494 495

26


Page 27 of 28


496 497

Figure 7. A schematic diagram of a potential formation route and the possible structures of wet and

498

dried schwertmannite. The structural framework is based on the structural model (a distorted 2×2

499

tunnel) developed by Fernandez-Martinez et al.11.

500 501

27



Page 28 of 28

502 503

TOC figure

28


Sulfate Local Coordination Environment in Schwertmannite

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