Reactions at particle edges versus vacancy sites - ACS Publications

Subscriber access provided by UNIV OF DURHAM

Article

Diffusion- and pH-dependent reactivity of layer-type MnO: Reactions at particle edges versus vacancy sites 2

Yuheng Wang, Sassi Benkaddour, Francesco Femi Marafatto, and Jasquelin Pena Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.7b05820 • Publication Date (Web): 12 Feb 2018 Downloaded from http://pubs.acs.org on February 13, 2018

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 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 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 29

Environmental Science & Technology

1

Diffusion- and pH-dependent reactivity of layer-type

2

MnO2: Reactions at particle edges versus vacancy sites

3

Yuheng Wang1, Sassi Benkaddour1, Francesco Femi Marafatto1, Jasquelin Peña1*

4

5 6

1

Institute of Earth Surface Dynamics, University of Lausanne, CH-1015 Lausanne, Switzerland

7 8 9 10

*Corresponding author: [email protected]

11 12 13

1 Environment ACS Paragon Plus


14

Page 2 of 29

TOC Art Co(II)

edge species

time

10 min

0 min

Co(III)

pH

vacancy species

vacancy species

12 h edge species

δ-MnO2

15 16

4

2 ACS Paragon Plus Environment

8

Page 3 of 29

17


Abstract

18

Layer-type manganese oxides are among the strongest solid-phase oxidants in surface

19

environments and readily oxidize a range of chemical species. However, knowledge of the

20

role played by different surface sites in contaminant oxidation is scarce. In this study, we

21

investigate the reactivity of particle edges versus vacancy sites in δ-MnO2 by combining Co

22

sorption kinetic experiments with quick X-ray absorption spectroscopy. During the fast

23

kinetic phase (t < 10 min), Co sorption and oxidation occurred dominantly at edge sites at pH

24

8; at pH 6 and pH 4, reactions also occurred at vacancy sites but were limited in extent. At

25

longer reaction times (t > 10 min), continuous accumulation of Co at vacancy sites was

26

observed, while the amount of Co at particle edges decreased or remained constant depending

27

on the absence or presence of aqueous Co(II), respectively. These data are consistent with the

28

diffusion-limited transport of metal cations to vacancy sites. In addition, at higher pH values,

29

the kinetics and extent of reaction at particle edges are greater than at pH 4 – 6. These results

30

suggest that although particle edges will be the first to react, layer vacancies will serve as the

31

long-term sorption and oxidation sites for contaminant metals in MnO2-rich systems.

32



33 34

Page 4 of 29

Introduction Manganese oxide (nominally MnO2) is among the most strongly oxidizing solid-phase

35

species commonly found in Earth surface environments.1,

2

36

birnessite and the nanoscale and turbostratic vernadite, and δ-MnO2, the synthetic analog for

37

vernadite3, typically display hexagonal sheet symmetry4 and under-coordinated oxygen atoms

38

both at layer vacancy sites5-7 and particle edges.8, 9 The high redox reactivity of these phases

39

depends on the i) amount of Mn(III) in the mineral, which can range from 0% to 35% and can

40

be located either in layer or interlayer positions, ii) different oxidation capacity of Mn(III) and

41

Mn(IV) in Mn oxides,10-14 iii) crystallinity of Mn oxide particles,15 and iv) different

42

adsorption reactivity of singly- and doubly-coordinated oxygen atoms at vacancy and edge

43

sites towards cations including protons and trace metals.11, 16, 17

Layer-type MnO2, such as

44

Manganese oxides are known to degrade a wide variety of organic contaminants,

45

including pesticides,18 antibiotics18 and organoarsenic species,19 and oxidize the reduced

46

forms of trace and toxicant elements such as divalent cobalt (CoII),10, 20-23 trivalent arsenic

47

(AsIII),8, 14, 15, 24 trivalent chromium (CrIII)7, 13, 25-27 and tetravalent uranium (UIV).28, 29 Previous

48

spectroscopic studies have relied primarily on batch experiments to elucidate the mechanism

49

of contaminant oxidation by birnessite. However, even with advanced characterization of the

50

solid phase, the relevant surface site(s) (edge vs. vacancy) and oxidant(s) (MnIII vs. MnIV) can

51

only be identified from site-specific information of reaction kinetics and surface speciation of

52

a probe compound.

53

Cobalt(II) is an ideal probe compound to investigate the site-specific oxidation

54

reactivity of birnessite due to its well-established coordination chemistry and speciation on

55

Mn oxides.10, 30 Kinetic studies of Co(II) oxidation by Mn oxides have shown that Co sorption

56

takes place in two kinetic steps: an initial rapid sorption step followed by a slow sorption

57

phase that extends over several days.31, 32 For biogenic Mn oxide, Co(II) was oxidized to


Page 5 of 29


58

Co(III) primarily during the second phase.33 Recently, Simanova and Peña provided a

59

mechanistic explanation for the fast and slow kinetic steps by investigating Co(II) oxidation

60

by Mn(III)-rich δ-MnO2 (hereafter referred to as “δ-MnIV,IIIO2”).34 The authors found that

61

Mn(IV,III) at edges sites were mainly responsible for Co(II) oxidation in the initial fast

62

kinetic step (t < 10 min), whereas interlayer Mn(III) reacted with Co(II) during the slow

63

sorption step (t > 10 min) but only to a limited extent. Their study34 provided evidence for the

64

low redox reactivity of interlayer Mn(III) and its passivation of vacancy sites with respect to

65

cation sorption. However, the reaction kinetics at vacancy sites not occupied by Mn(III)

66

relative to that of edge sites have not been assessed. Moreover, the pH dependence of these

67

site-specific reactions is not known. The adsorption of metals by δ-MnO2 depends on

68

suspension pH relative to the point-of-zero-charge (PZC) of the reactive surface oxygen

69

atoms, which increases with decreasing oxygen coordination.17, 24, 35, 36 Therefore, we expect

70

that singly-coordinated oxygen atoms at the particle edges will show strong pH dependence

71

relative to vacancy sites. Finally, even though Mn(IV) has been suggested to be stronger

72

oxidant than Mn(III) located in the MnO2 layer,13-15, 37 no mechanistic study has confirmed

73

this hypothesis to date.

74

In this study, we combined batch kinetic experiments with quick X-ray absorption

75

spectroscopy (QXAS) to study the site-specific sorption and oxidation of Co(II) by a fully

76

oxidized δ-MnO2 at suspension pH values of 4 – 8. To better discriminate the reactivity of

77

edge sites versus vacancy sites, we worked under conditions where Co(II) was either fully

78

removed from solution within the first few minutes or present throughout the 24 hr reaction.

79

These results were compared with those from Simanova and Peña34 to evaluate the difference

80

in reactivity between Mn(III)-poor and Mn(III)-rich δ-MnO2. Our work provides mechanistic

81

insight regarding the reactivity of particle edges and vacancy sites in δ-MnO2 over a range of

82

environmentally relevant pH values. 5 Environment ACS Paragon Plus


83 84

Materials and Methods

85

Synthesis of δ-MnO2. The δ-MnO2 used in this study was prepared according to the redox

86

method.38, 39 Briefly, we added 300 ml of 0.3 M MnCl2 at a rate of 72 mL min-1 into 300 ml of

87

a solution 0.2 M KMnO4, and 340 ml of 0.4 M NaOH. After 1 hr, suspension was washed and

88

centrifuged five times. The resulting wet paste was then freeze-dried and the δ-MnO2 powder

89

was stored at -20 °C. The solid-phase was characterized by a crystallite size of 7.2 ± 0.5 nm

90

(ab plane), a specific surface area of 259 ± 13 m2/g, a Na:Mn ratio of 15 % (mol mol-1) and

91

AMON (average Mn oxidation number) of 4.03 ± 0.01. Details for the synthesis are presented

92

in the Supporting Information (SI), whereas detailed chemical and structural characterization

93

of the δ-MnO2 used in this study are provided elsewhere.39

94 95

Kinetic Experiments. The δ-MnO2 powder was ground in an agate mortar and dispersed in

96

ultrapure water using an ultrasonic bath to obtain a 4 g L-1 stock suspension. Subsequently, 50

97

mL of the δ-MnO2 stock suspension, 20 mL of 100 mM NaCl and 120 mL ultrapure water

98

were added to a 250 mL beaker on a magnetic stirrer. The pH of the suspension was adjusted

99

to 4.0, 6.0 or 8.0 before adding Co and was kept constant using a Metrohm 718 STAT Titrino.

100

To initiate an experiment, 1 or 4 mL of 100 mM CoCl2·6H2O solution was added to the

101

suspension and the final volume of the suspension was adjusted to 200 mL using ultrapure

102

water. Four experimental conditions were tested, which are identified throughout the text by

103

the suspension pH value and maximum achievable Co loading in units of mol%: pH6-Co8%,

104

pH6-Co40%, pH8-Co32% and pH4-Co37%.

105

The time at which Co(II) was added to the δ-MnO2 suspension was defined as the start

106

of the reaction (t0). At defined time intervals, 10 mL aliquots were subsampled and filtered

107

onto a 0.2 micron Millipore membrane. The filter membranes were then analyzed


Page 6 of 29

Page 7 of 29


108

immediately by QXAS. About 1 min was required between the subsampling and the start of

109

data acquisition. The filtrates were retained for chemical analysis. The reaction time (t) for

110

each subsample was defined as the time elapsed between the beginning of QXAS data

111

acquisition and t0. Due to sampling constraints, there was ~1 min discrepancy between the

112

extent of Co sorption, as captured by wet chemical analysis, and Co surface speciation, as

113

determined from QXAS. At the end of each experiment, 1 mL of the remaining suspension

114

was sampled for analysis of total Mn and Co concentrations. Further details are provided in

115

the SI.

116

Aqueous (cMn and cCo) and total metal concentrations (cMnTOT and cCoTOT) were

117

determined by inductively coupled plasma optical emission spectrometry (ICP-OES) on a

118

PerkinElmer Optima 8300 spectrometer (Table S1). All analyses were done in triplicate using

119

scandium as an internal standard and reading three emission lines for each element. For

120

analysis of cMnTOT and cCoTOT, 1 mL of the Co-bearing δ-MnO2 suspensions were digested in a

121

10 mL solution of 3.5 wt.% HNO3 and 10 mM oxalic acid (H2C2O4). Cobalt loadings on δ-

122

MnO2 (qCo) were calculated according to (cCoTOT − cCo)/(cMnTOT − cMn).

123 124

Quick X-ray Absorption Spectroscopy. QXAS experiments were performed at beamline

125

SuperXAS-X10DA of the Swiss Light Source using a quick-scanning monochromator with a

126

Si(111) channel-cut crystal, which oscillated between 7.4 and 8.5 keV at a frequency of 2 Hz

127

(2 scans s-1). The monochromator energy was calibrated using a Co foil (7709 eV) and

128

monitored continuously in transmission mode during data acquisition. Sample spectra were

129

collected from the wet pastes at room temperature in transmission mode. Data acquisition

130

lasted 2 min for sub-samples collected at t < 10 h and 4 min for those at t > 10 h, respectively.

131

The scans collected in these time intervals were averaged after confirming that there was no

132

significant difference between the first and last spectra in a batch.



133

The software JAQ40 was used to convert the raw scan files from the angular encoder to

134

an energy scale in units of electron volts (eV), rebin the raw data containing about 600 000

135

data points into individual Co K-edge XAS scans with 4 000 data points and average the

136

scans. Averaged QXAS spectra were aligned in energy, normalized and background

137

subtracted in Athena41 using the following parameters: E0 = 7721 eV, k-weight = 3, Rbkg = 1

138

Å, dk= 1 Å−1, no clamps for the spline function and using fitting a linear function for the pre-

139

edge (−125 to −30 eV below E0) and a quadratic function for the post-edge (150 to 800 eV

140

above E0).

141

Normalized Co K-edge quick X-ray absorption near edge structure (QXANES) spectra

142

were analyzed in Athena using the linear combination fitting (LCF) module. Spectra were

143

fitted from 7710 to 7755 eV using CoCl2(aq) and CoIIIOOH(s) as references for Co(II) and

144

Co(III), respectively,34 and allowing the fractions of Co(II) and Co(III) to float between 0 and

145

1. The fit values were then normalized to the sum of the two components in order to obtain a

146

component sum of 1.00 (Table S2). The scaled fractional parameters, fCoII and fCoIII, were

147

used to calculate Co(II) and Co(III) loadings (qCoII and qCoIII) according to qCoII = qCo * fCoII

148

and qCoIII = qCo * fCoIII, respectively (Table S3).

149

Background-subtracted and k3-weighted quick extended X-ray absorption fine structure

150

(QEXAFS) spectra were Fourier-transformed (FT). The FT QEXAFS spectra were analyzed

151

by shell-by-shell fitting in R-space in the range of 0 − 3.5 Å using SixPack42 and the

152

structural model developed by Simanova et al.,34 which includes the four species shown in

153

Figure 1. Two of these species are located at vacancy sites, triple corner-sharing Co(II)

154

(Co(II)-TCS) and incorporated Co(III) (Co(III)-INC) complexes, and two at edge sites,

155

double corner-sharing Co(II) (Co(II)-DCS) and double edge-sharing Co(III) (Co(III)-DES)

156

complexes. In addition, Co(II) is assumed to form only corner-sharing complexes (CS, i.e.


Page 8 of 29

Page 9 of 29


157

Co(II)-TCS and Co(II)-DCS), whereas Co(III) is assumed to form only edge-sharing

158

complexes (ES, i.e. Co(III)-INC and Co(III)-DES).

159

Based on geometric constraints, four single-scattering paths were used in shell-by-shell

160

fits: Co(II)-O, Co(III)-O, Co(II)-MnCS, and Co(III)-MnES. Interatomic distances (R) and

161

Debye-Waller factors (σ2) were free to float for all paths, except that the σ2 of the two Co-

162

Mn paths were co-varied. The difference between the user-defined threshold energy and the

163

experimentally determined one (ΔE0) was co-varied for the four paths. The coordination

164

numbers (CNs) of the Co(II) and Co(III) shells were scaled by the fraction of Co(II) and

165

Co(III) in the samples as determined from analysis of the QXANES spectra according to:

166

CN·fCoII for Co(II)-O and Co(II)-MnCS and CN·fCoIII for Co(III)-O and Co(III)-MnES. Finally,

167

the CNs for the Co(II)-O and Co(III)-O shells were set to their structural model value (i.e.,

168

CN = 6), while the CNs for the Co(II)-MnCS and Co(III)-MnES shells were allowed to float.

169

This fitting strategy yielded good fits in terms of R, σ2, ΔE0, and their uncertainties (Figures

170

S2-S5, Tables S4-S7). Fitting uncertainties for the CNs of the Co-Mn shells correlated with

171

the CN, where CNs > 4 typically had uncertainties between 8% and 20% and CNs < 3

172

typically had uncertainties between 15% and 40%. Out of the 64 CNs fitted across all samples,

173

31 uncertainties were ≤ 20 % and 5 were > 40 %. In general, CNs cannot be determined to

174

accuracy greater than 10% and uncertainty in CN as high as 40%-50% can be commonly

175

found even with satisfactory structure models.43 The fit results are supported by the R-factor

176

values, which also provide a measure of goodness of fit. All the R-factors of our fits fall

177

within the range of 0.01 – 0.04 (Table S4-S7), which are lower than the threshold for a

178

reliable fit at 0.05.43 The reliability of the structural model and fit results are also supported

179

by the close agreement between the experimental and calculated Fourier-transformed EXAFS

180

and back Fourier-transformed signals (Figures S2-S5).

181



Page 10 of 29

182

Estimates of Species-Specific Surface Loadings. The distribution of Co between the four

183

species shown in Figure 1 (i.e., Co(II)-TCS, Co(III)-INC, Co(II)-DCS and Co(III)-DES) was

184

estimated from the EXAFS-derived CNs and QXANES-derived Co(II) and Co(III) loadings.

185

By assuming ideal values of 6 and 2 for the CNs of species formed at vacancy sites (i.e.,

186

Co(II)-TCS and Co(III)-INC) and edge sites (i.e., Co(II)-DCS and Co(III)-DES), we could

187

estimate the fractions of Co(III) at vacancy sites (x) and edge sites (1 - x) and those of Co(II)

188

at vacancy sites (y) and edge sites (1 - y) using Eq. 1 and Eq. 2.:

189

CN(CoIII-MnES) = 6 · x + 2 · (1 - x)

Eq. 1

190

CN(CoII-MnCS) = 6 · y + 2 · (1 - y)

Eq. 2

191

where CN(CoIII-MnES) and CN(CoII-MnCS) were obtained directly from shell-by-shell fits

192

(Table S4-S7). The surface loading of each species was then estimated according to:

193

qCoIII-INC = qCoIII · x

Eq. 3

194

qCoIII-DES = qCoIII · (1 – x)

Eq. 4

195

qCoII-TCS = qCoII · y

Eq. 5

196

qCoII-DCS = qCoII · (1 – y)

Eq. 6

197

The estimates of species-specific surface loadings are reported in Table S3. While uncertainty

198

may be large for about half of the sub-samples (> 20 % based on fitting uncertainties in CN

199

obtained from shell-by-shell fitting of QEXAFS data), the trends in surface speciation are

200

supported unequivocally by the evolution of amplitudes of the FT transform peaks shown in

201

Figure S6.

202 203

Characterization of Reacted δ-MnO2. All Mn characterization measurements are described

204

in the SI. Briefly, to assess changes in AMON and average local coordination environment of

205

Mn in δ-MnO2, Mn K-edge XANES and EXAFS spectra were collected at the end of each


Page 11 of 29


206

experiment (t ~ 24 h). In addition, trivalent Mn (Mn(III)) accumulated in δ-MnO2 was

207

quantified directly in separate experiments by extraction with sodium pyrophosphate (PP).44

208 209

Results

210

Kinetics of Co Sorption and Mn Solubilization. The kinetics and extent of Co sorption by

211

δ-MnO2 varied strongly with suspension pH (Figure 2a). Inspection of the data showed up to

212

three kinetic regimes: rapid (< 3.5 min), moderate (3.5 – 11 min), and minimal or static

213

sorption (11 min to 24 h). The low Co loading experiment, pH6-Co8%, transitioned sharply

214

from the rapid sorption phase to the static phase, whereas the high loading experiments, i.e.,

215

pH8-Co32%, pH6-Co40% and pH4-Co37%, showed all three kinetic phases. As reported in

216

Table 1, the rates of Co(II) sorption onto δ-MnO2 during the fast kinetic phase increased with

217

increasing pH. However, the transition and static kinetic phases showed no clear trends with

218

pH.

219

Figure 2b shows the kinetics of Mn release from δ-MnO2. Overall, the sorption and

220

oxidation of Co(II) led to minimal accumulation of Mn in solution: at most 2% of the Mn

221

initially in the solid phase as Mn(IV), which corresponds to 90 – 120 µM Mn, was released to

222

solution in the pH6-Co40% and pH4-Co37% experiments; aqueous Mn was negligible in the

223

pH8-Co32% and pH6-Co8% experiments (Table S1). These observations indicate that Mn

224

reduced from Mn(IV, III) to Mn(III, II) upon Co(II) oxidation to Co(III) must largely remain

225

associated with the solid phase. In addition, because no ligand was present in solution to

226

stabilize aqueous Mn(III), we assume that any aqueous Mn measured was Mn(II) produced

227

either from the transfer of two electrons to one Mn(IV) ion or from the disproportionation of

228

2 Mn(III) to form Mn(II) and Mn(IV).

229

After 24 hours of reaction, both the loading (qCo) and fraction of Co sorbed by δ-MnO2

230

[(cCoTOT − cCo)/cCoTOT] increased with increasing suspension pH (Table S1). Loadings of 0.23,



231

0.30 and 0.32 Co mol Mn mol-1 were achieved in the pH4-Co37% pH6-Co40% and pH8-

232

Co32% experiments, respectively (Table S1). For the pH8-Co32% experiment, less than 3 %

233

of cCoTOT or 54 µM remained in solution after 11 min of reaction. In contrast, about 27% and

234

39 % of cCoTOT or 774 and 1079 µM remained in solution at the end of experiment in the pH6-

235

Co40% and pH4-Co37% experiments, respectively. Similar to the pH 8 experiment, the pH6-

236

Co8% experiment showed complete Co sorption, approaching its maximum achievable qCo

237

value of 0.08 after 3.5 min. In contrast, the high concentrations of aqueous Co(II) present

238

throughout the reaction in the pH6-Co40% and pH4-Co37% experiments provided a large

239

pool of aqueous Co(II) that could be resupplied to the δ-MnO2 surface upon oxidative loss of

240

adsorbed Co(II) species.

241 242

Kinetics of Co(II) Oxidation. The Co K-edge QXANES spectra in Figure S2 show the

243

increase in Co(II) oxidation to form Co(III) over time, as indicated by the decrease in

244

absorbance at 7725 eV and increase in absorbance at 7730 eV. The LCF-derived fractions of

245

Co(II) and Co(III) in the samples (fCoII and fCoIII) (Table S2) returned fitting uncertainties of

246

less than 10 % and the sum of the two components before scaling to 1.00 ranged from 1.02 –

247

1.09 for all samples, indicating good fit results. In all experiments, except for the pH 6-Co8%

248

experiment, 30 – 50 % of the sorbed Co remained as Co(II) after 24 hours of reaction.

249

Time-course plots of Co(II) and Co(III) loadings on δ-MnO2 are presented in Figure 3

250

and Table S3. These loadings were used to determine the rate of Co(II) oxidation in each of

251

the three kinetic phases, as reported in Table 1. The initial rate of Co(II) oxidation (10±1 –

252

45±4 × 10-3 mol Co mol-1 Mn min-1) slowed by at least one order of magnitude in the 3.5 –

253

11 min phase (11±6 – 35±15 × 10-4 mol Co mol-1 Mn min-1) and again between 11 min and

254

24 hours (22±4 – 45±11 × 10-6 mol Co mol-1 Mn min-1). Similar to the rate of sorption, the

255

rate of Co(II) oxidation in the first 3.5 min increased with increasing pH (Table 1).


Page 12 of 29

Page 13 of 29

256


Trends in the rates of Co(II) oxidation differed between experiments where aqueous

257

Co(II) resupply was present or absent during the reaction. For the pH6-Co40% and pH4-Co37%

258

experiments (with Co(II) resupply), qCoII remained constant after 11 min (Table 1), which

259

indicates similar rates of Co(II) oxidation and Co(II) sorption since faster Co(II) oxidation

260

relative to sorption would have resulted in a decrease in the Co(II) surface loading. For the

261

pH8-Co32% and pH6-Co8% experiments (no resupply of aqueous Co(II)), qCoII decreased

262

and qCoIII increased continuously after 3.5 min and after 11 min, respectively, indicating

263

higher rate of Co(II) oxidation than that of Co(II) sorption.

264 265

Time-Resolved Surface Speciation of Co. All FT QEXAFS spectra showed an apparent

266

increase in the amplitude of the Co-O shell at R ≈ 2.0 Å and Co-MnES shell at 2.9 Å, and

267

modest decrease or no change in the amplitude of the Co-MnCS shell at R ≈ 3.5 Å with

268

increasing reaction times (Figures S2-S5). The FT amplitude of the Co-O shell at ∼2.0 Å

269

consists of destructive interference of the photoelectron scattering from Co(II)-O pairs at ∼2.1

270

Å and Co(III)-O pairs at ∼1.9 Å, whereas the Co-MnCS shell at ∼3.5 Å and Co-Mn shell at

271

∼2.9 Å indicate Co(II) corner-sharing (Co(II)-CS, i.e. Co(II)-TCS and Co(II)-DCS) and

272

Co(III) edge-sharing (Co(III)-ES, i.e. Co(III)-INC and Co(III)-DES) complexes. The increase

273

in the FT magnitude of Co(III)-MnES shell at R ≈ 2.9 Å and the modest decrease or no change

274

in that of the Co(II)-MnCS shell at R ≈ 3.5 Å with time are plotted in Figure S6. These trends

275

correspond to an increase in the amount of Co(III)-ES species with time, as corroborated by

276

the increasing CN of the Co(III)-Mn shell at ∼2.9 Å (Tables S4-S7). For example, the Co-

277

MnES CN increased from 2.9 ± 0.9 at 7.5 min to 4.8 ± 0.4 at 1405 min and from 3.0 ± 1.0 at

278

7.5 min to 5.8 ± 0.7 at 1379 min for the pH6_Co40% and pH6_Co8% experiments,

279

respectively.



280

The surface loading estimates for the four species considered in the EXAFS fitting (i.e.,

281

Co(II)-TCS, Co(III)-INC, Co(II)-DCS and Co(III)-DES) are reported in Table S3 and plotted

282

in Figure 4. For all four experiments, the QEXAFS spectra acquired at ~3.5 min were not of

283

sufficiently high quality for shell-by-shell fitting. Based on the EXAFS spectra acquired at ~7

284

min, we assume that the fast Co sorption kinetics observed for t < 3.5 min was due primarily

285

to the formation of Co(III)-DES and Co(II)-DCS species.

286

For the pH8-Co32% experiment, where 97 % sorption was achieved after 11 min, the q

287

values of Co(III)-INC or Co(II)-TCS at vacancy sites remained low or insignificant for t < 31

288

min (Table S3). The high sorption capacity and the fast oxidation kinetics observed in the fast

289

and moderate kinetic phases at pH 8 were due to reaction at edge sites, which can

290

accommodate Co(III)-DES and Co(II)-DCS species (Figure 4a). Between 11 min and 27 h,

291

the surface loading of Co(III)-INC increased at the expense of Co(III)-DES and Co(II)-DCS

292

species (Table S3, Figure 4a). The redistribution of Co from edge sites to vacancy sites was

293

also observed in the pH6-Co8% experiment (Figure 4d), where greater than 99% sorption

294

was achieved after 3.5 min.

295

For the pH6-Co40% and pH4-Co37% experiments, Co(II)-DCS and Co(III)-DES

296

complexes at edge sites were also the dominant species at short reaction times. For t > 11 min,

297

the amount of Co(II)-DCS, Co(III)-DES and Co(II)-TCS species remained approximately

298

constant, indicating that the increase in the Co loading was due mainly to the formation of

299

Co(III)-INC species (Figure 4b and c). In addition, the constant loading for Co(II)-TCS

300

species suggests similar rates of Co(II) adsorption and oxidation at vacancy sites. Finally, the

301

combination of all edge species at all time points were generally lower at pH 4 and 6 than at

302

pH8. By contrast, the loadings of the Co(III)-INC and Co(II)-TCS species at vacancy sites

303

were higher at pH6 and pH4 than at pH8 (Figure 4a-c, Table S3). Thus, these results show


Page 14 of 29

Page 15 of 29


304

that increasing pH enhances the reactivity of edge sites, whereas it leads to a moderate

305

decrease in the extent of sorption at vacancy sites only for the pH 8 experiment.

306 307

Mn valence in Reacted δ-MnO2. Linear combination of the Mn K-edge XANES reference

308

spectra reproduced well the experimental spectra as shown in Figure S7 and indicated by the

309

sum of the fractional parameters for Mn(IV), Mn(III) and Mn(II), which ranged from 0.95 –

310

0.96 for all samples before scaling to 1.00. After 24 h of reaction, Mn(IV) accounted for 0.76

311

± 0.04 to 0.92 ± 0.04 of the total Mn in the solid phase (Table S8). The AMON values

312

calculated from LCF for the δ-MnO2 decreased from 3.87 ± 0.04 at pH 4 to 3.70 ± 0.04 at pH

313

8 after reaction with Co(II) (Table S9). This decrease in AMON was directly proportional to

314

the increase in Co(III) (R2 = 0.93). In addition, estimates of Mn(III) content obtained by

315

extraction with PP in low loading samples prepared at pH 6 and pH 4 showed equal

316

proportions of Co(III) and Mn(III) (Table S10).

317

A small fraction of adsorbed Mn(II) (0.06 ± 0.04) was found in the pH8-Co32% and

318

pH6-Co40% experiments. It is difficult to determine if in fact the samples contain Mn(II).

319

The presence of a small fraction of aqueous Mn(II) at pH 6 suggest that the samples may

320

indeed contain a small but measurable fraction of Mn(II). The presence of Mn(II) is consistent

321

with favorable conditions for cation adsorption at near-neutral and alkaline pH values.

322

However, adsorbed Mn(II) is expected to comproportionate with Mn(IV) in δ-MnO2 over the

323

timescale of our measurements.45 These result suggest a decrease in the oxidizing capacity of

324

the solid phase after reaction with Co, which leads the accumulation of 0.22 and 0.18 mol

325

Co(III) Mn mol-1 and 0.018 and 0.13 mol Mn(III) mol-1 Mn at pH 8 and pH 6, respectively

326

(Table S3).

327



328

Local Structure of Reacted δ-MnO2. Manganese K-edge EXAFS spectra (Figure S7)

329

obtained at t ~24 h showed similar features to the unreacted sample. Therefore, no major

330

changes in average Mn coordination environment occurred within the spatial resolution of

331

EXAFS spectroscopy (R < 6 Å) upon: 1) adjusting the pH of the mineral suspension to values

332

between 4 and 8; 2) increasing the surface excess of Co to values as high as 0.3 mol Co mol-1

333

Mn; or 3) decreasing the AMON of δ-MnO2 from 4.0 to 3.7. An overlay plot of the Mn K-

334

edge FT EXAFS spectra (Figure S7) reveal higher amplitudes in the Mn-O and Mn-Mn

335

shells near 1.9 and 2.9 Å in the samples with greater amounts of Co incorporation into MnO2

336

layers (Table S3). These differences are consistent with an increasing ordered local structure

337

as the extent of Co(III) incorporation increases.

338 339

Discussion

340

Diffusion-Limited Reactions and Formation of Stable Co(III) Species at Vacancy Sites.

341

Cobalt(II) adsorption and oxidation by δ-MnO2 showed an initial fast kinetic phase followed

342

by one (for low Co experiment pH6-8%) or two (for all the three high Co experiments) slower

343

kinetic phases, consistent with previous work.31, 32, 34 Analysis of the QEXAFS data identified

344

the particle edges in δ-MnO2 as the sites for initial fast Co(II) adsorption and oxidation,

345

indicating that these surface sites were readily accessible to aqueous species. However,

346

significant amounts of Co remained as Co(II) at the edge sites in all experiments, suggesting

347

that the oxidation capacity of edge sites was exceeded. As the reaction proceeded (t > 10 min),

348

increasing amounts of Co(III) were incorporated at vacancy sites (Figure 4). The kinetically-

349

limited reaction at vacancy sites suggests that sorption of aqueous Co(II) as Co(II)-TCS and

350

subsequent oxidation to Co(III)-INC was constrained by its ability to diffuse into the

351

interlayer and to available vacancy sites in δ-MnO2. The hypothesis that reactions at vacancy

352

sites are diffusion-limited is especially supported by the pH6-Co40% and pH4-Co37%


Page 16 of 29

Page 17 of 29


353

experiments (i.e., Co(II)aqueous >> 0), where the availability of aqueous Co(II) did not limit

354

reactions at vacancy sites. Diffusion-limited sorption by minerals surfaces46,47 has been

355

invoked to explain adsorption kinetics of Cu and Pb on ferrihydrite48 as well as Zn on MnO2-

356

coated montmorillonite.49

357

In the pH8-Co32% and pH6-Co8% experiments, aqueous Co(II) was largely consumed

358

within the first few minutes of the reaction (Table S1). Therefore, the subsequent formation

359

of Co(III)-INC species required the net transfer of Co from edge sites to vacancy sites over

360

time. Since we assume minimal transfer of Co(III) from edge to vacancy sites due to its low

361

solubility relative to Co(II),50 the most plausible pathway would be the desorption of Co(II)-

362

DCS species enabled by the dynamic equilibrium maintained between Co(II)-DCS species

363

and aqueous Co(II), and subsequent diffusion of aqueous Co(II) to vacancy sites. Another

364

process that could lead to the formation of Co(III)-INC species is the re-crystallization of

365

Mn(II,III) around Co(III)-DES species. While we cannot rule out re-crystallization or crystal

366

growth processes that may be favored under acidic conditions,39, 51, 52 they are unlikely to be

367

significant because this mechanism requires significant Mn(II,III) re-oxidation at particles

368

edges but the oxidation capacity of the edge sites is largely depleted by Co(II) within the first

369

few minutes of the reaction.

370

Notwithstanding the slower reaction kinetics at vacancy sites relative to edge sites,

371

vacancy-bound Co species appear to be more thermodynamically stable than surface

372

complexes particles edges. Specifically, Co(III)-INC species form at the expense of Co(II)

373

edge species. The higher stability of Co(III)-INC species in δ-MnO2 can be explained by i)

374

the greater number of shared O atoms between CoO6 octahedra and neighboring MnO6 ones,

375

and ii) the similar ionic radii between the low-spin octahedral Co(III)30 and Mn(IV) and thus

376

high compatibility of CoO6 octahedra within MnO6 layers.

377



Page 18 of 29

378

Reactivity of particle edges and vacancy sites as a function of suspension pH. The edge

379

sites of δ-MnO2 become increasingly reactive with increasing pH as shown by the kinetics

380

and extent of sorption (Table 1, Figure 4). Previous research correlating the point-of-zero-

381

charge (PZC) and the saturation degree of singly-coordinated O atoms for various minerals as

382

well as trends in oxyanion sorption by δ-MnO216, 35 has suggested a pKa value between 6 and

383

8 for singly-coordinated O atoms on the edge sites of MnO2. The greater degree of

384

deprotonation, and thus under-saturation, of singly-coordinated O atoms on δ-MnO2 at higher

385

pH makes these sites stronger Lewis bases that readily share lone pair electrons and bind

386

aqueous Co(II) more favorably. On the other hand, the greater incorporation of Co(III) at

387

vacancy sites at pH 6 than at pH 8 (Figure 4) corroborates the enhanced reactivity of edge

388

sites at higher pH values and their ability to outcompete vacancy sites. Moreover, if the

389

kinetics of Co sorption at vacancy sites are diffusion-limited, the decrease in Co solubility

390

with increasing pH would also decrease the rate of Co transfer from edge sites to vacancy

391

sites.

392 393

Electron Balance between Mn and Co Species. The large amounts of Co(III) that

394

accumulate in δ-MnO2, the low amount of Mn(II) accumulated in solution, and similarity in

395

Mn EXAFS spectra before and after reaction with Co(II) raise the question of whether or not

396

electron balance between reduced Mn and oxidized Co is supported by the data. If Mn(IV)

397

and Mn(III) are the only oxidants of Co(II), then each mole of Mn(III) or Mn(II) quantified

398

should correspond to one or two moles of e- captured by δ-MnO2, respectively, such that the

399

moles of e- equivalents acquired per mole of δ-MnO2 (qe-) can be calculated according to

400

qe- = fMnIII + 2(fMnII + fMnIIaq)


eq. 8

Page 19 of 29


401

The magnitude of qe- can also be computed as the fraction of Mn(IV) lost upon reaction with

402

Co (qe-= ΔfMnIV = 1 - fMnIV). If we then assume that Co(II) is the only reductant in the system,

403

qe- should be equal to the Co(III) surface loading.

404

Equation 8 did not provide a reliable measure of electron balance for two reasons:

405

there is a high uncertainty in qe- (± 0.09 as calculated by error propagation) and low fMnII

406

values (0.02 ± 0.04 to 0.06 ± 0.04), which are multiplied by 2 to calculate qe-. Although

407

values of qe- (Table S9) and qCoIII (Table S3) were within error of each other, qe- was

408

systematically higher than qCoIII. We assume that the values obtained with Eq. 8 are biased for

409

the aforementioned reasons. However, our measurements showed similar values for ΔfMnIV

410

and qCoIII for all experiments. The balance between Mn(IV) consumed and Co(III) generated

411

is well supported at low loadings (~0.05 mol Co mol-1 Mn) where Mn(III) quantification using

412

PP showed equal Co(III) and Mn(III) amounts (Table S10). Thus, our data suggest that

413

electron balance between Mn and Co species is likely achieved and that no other oxidant, e.g.

414

dissolved oxygen, nor other reductant, e.g. water molecules, alter significantly the AMON of

415

δ-MnO2 upon reaction with Co(II).

416 417

Redox Reactivity of Mn(III)-free versus Mn(III)-rich δ-MnO2. The comparison between

418

the reactivity of δ-MnIVO2 and δ-MnIV,IIIO2 (Mn(III) ~ 34 %)34 is only valid under conditions

419

where the Mn(III) content in the reacted δ-MnO2 is low. For δ-MnIVO2, this condition was

420

only met in the first few minutes of the reaction since the accumulation of Co(III) in δ-MnO2

421

leads to the accumulation of Mn(III). At pH of 6.0 and 6.5 and high Co loadings (Figure S8,

422

Co_0.020 and pH6_Co40%), the amount of Co species at vacancy sites was much higher in δ-

423

MnIVO2 (0.051 mol Co mol−1 Mn) than in δ-MnIV,IIIO2 (~0). A similar trend was observed for

424

the low Co-loading samples (Figure S8). These data indicate that vacancy sites were free to

425

react with Co(II) in δ-MnIVO2 but not in δ-MnIV,IIIO2. In the latter material, the vacancy sites



Page 20 of 29

426

were passivated by Mn(III),6 which prevented the adsorption of Co34. This phenomenon also

427

has been observed for Ni and Cu reacted with Mn(III)-rich δ-MnO253, 54. Consequently Co(II)

428

at those sites can only be oxidized slowly through outer-sphere electron transfer, which is

429

kinetically hindered relative to inner-sphere electron transfer.34,

430

characterized by a Co loading that was about 1.5 times higher for δ-MnIVO2 than for δ-

431

MnIV,IIIO2 at similar steady-state concentrations of aqueous Co(II). Thus, the edge sites are

432

also more reactive towards Co in δ-MnIVO2 than in δ-MnIV,IIIO2. This finding is consistent

433

with previous studies which propose higher sorption11 and oxidation capacity13,

434

Mn(IV)-rich Mn oxides compared to Mn(III)-rich Mn oxides.

55, 56

The edge sites were

14, 57

for

435 436

Environmental Implications. The QXAS analyses presented here allowed us to investigate

437

multiple adsorption and redox processes between Co and δ-MnIVO2 nanoparticles at a range

438

of environmentally relevant pH values. Remarkably high Co loadings on Mn(III)-free δ-

439

MnO2 were observed, as well as significant but incomplete oxidation of Co(II) in experiments

440

where the Co loading exceeded the vacancy content. The surface site-specific and pH-

441

dependent reactivity of δ-MnO2 is consistent with diffusion-limited adsorption processes that

442

slow reactions at vacancy sites, higher thermodynamic stability of Co incorporated at vacancy

443

sites that drives the transfer of Co(II) from edge sites to vacancy sites, and a pKa value for

444

singly-coordinated O atoms at the edges of δ-MnO2 situated between pH 6 and 8, which

445

enhances the reactivity of edge sites under alkaline conditions. These results also demonstrate

446

that suspension pH, through its effect on surface charge and particle size, exerts major control

447

over the reactivity of surface sites at particle edges versus vacancy sites. Finally, the

448

comparison of the reactivity of Mn(III)-free δ-MnO2 to that of Mn(III)-rich δ-MnO234

449

suggests that Mn(III) located at vacancy sites passivates these sites with respect to cation


Page 21 of 29


450

adsorption53

54

451

preceded by the formation of an inner-sphere surface complex.

and oxidation, presumably because electron transfer from Mn to Co is

452

This work has implications for systems rich in Mn oxides, where edge sites would be

453

the first to react with metal contaminants, but unoccupied vacancy sites would act as the long-

454

term contaminant-scavengers if they are not occupied by Mn(III). This is especially true as

455

natural birnessite particles should show more ordered stacking along the c-axis relative to

456

synthetic δ-MnO2 and the former should suffer more from diffusion limitation that results in

457

slower access to vacancy sites. Moreover, we propose that the density and accessibility to

458

vacancy sites is more important to consider than the surface area of Mn oxides when studying

459

long-term equilibration between metal contaminants and Mn oxides in the environment.7 On

460

the other hand, in biogenic or biofilm-associated Mn oxides in natural or technical systems,

461

sorption of organic molecules to the mineral particles may reduce the reactivity of the edge

462

sites.58, 59 Enhanced reactivity in biogenic systems may be facilitated by the ability of bacteria

463

to re-oxidize reduced Mn(II, III) and precipitate Mn oxides with a high abundance of vacancy

464

sites.29, 60 Finally, this work may inform the mechanism through which Co enhances the

465

reactivity of Co-rich birnessite used as an oxygen evolution catalyst.61

466 467

ASSOCIATED CONTENT

468

Supporting Information

469

Supporting Information, including additional details on the methods used in this study, all

470

QXANES and QEXAFS spectra and corresponding fit results, and comparison in site-specific

471

speciation for Mn(III)-poor and Mn(III)-rich δ-MnO2, is available free of charge on the ACS

472

Publications website at DOI: xxx

473 474

AUTHOR INFORMATION



475

Corresponding author

476

*E-mail: [email protected]

477

Notes

478

The authors declare no competing financial interest.

479 480

ACKNOWLEDGMENTS

481

Funding for this work was provided by the Swiss National Science Foundation (200020-

482

162825) and awards to JP from the Sandoz Family foundation and the BCV Foundation. We

483

thank Anna Simanova for pyrophosphate extraction experiments and initial data collection on

484

Mn(III)-rich δ-MnO2, Maarten Nachtegaal for technical assistance during the QXAS beam

485

time at the Swiss Lightsource, and Laetitia Monbaron for laboratory support.

486 487

References

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

1. Borch, T.; Kretzschmar, R.; Kappler, A.; Cappellen, P. V.; Ginder-Vogel, M.; Voegelin, A.; Campbell, K. Biogeochemical redox processes and their Impact on contaminant dynamics. Environmental Science & Technology 2010, 44, 15-23. 2. Manceau, A.; Marcus, M. A.; Tamura, N.; Fenter, P. A.; Rivers, M. L.; Sturchio, N. C.; Sutton, S. R., Applications of Synchrotron Radiation in Low-Temperature Geochemistry and Environmental Science. 2002; p 341. 3. Grangeon, S.; Lanson, B.; Lanson, M.; Manceau, A. Crystal structure of Ni-sorbed synthetic vernadite: a powder X-ray diffraction study. Mineralogical Magazine 2008, 72, 1279-1291. 4. Post, J. E. Manganese oxide minerals: Crystal structures and economic and environmental significance. Proceedings of the National Academy of Sciences 1999, 96, 3447-3454. 5. Drits, V. A.; Silvester, E.; Gorshkov, A. I.; Manceau, A. Structure of synthetic monoclinic Na-rich birnessite and hexagonal birnessite .1. Results from X-ray diffraction and selected-area electron diffraction. American Mineralogist 1997, 82, 946-961. 6. Silvester, E.; Manceau, A.; Drits, V. A. Structure of synthetic monoclinic Na-rich birnessite and hexagonal birnessite .2. Results from chemical studies and EXAFS spectroscopy. American Mineralogist 1997, 82, 962-978. 7. Manceau, A.; Charlet, L. X-ray absorption spectroscopic study of the sorption of Cr(III) at the oxidewater interface: I. Molecular mechanism of Cr(III) oxidation on Mn oxides. Journal of Colloid and Interface Science 1992, 148, 425-442. 8. Manning, B. A.; Fendorf, S. E.; Bostick, B.; Suarez, D. L. Arsenic(III) Oxidation and Arsenic(V) Adsorption Reactions on Synthetic Birnessite. Environmental Science & Technology 2002, 36, 976-981. 9. Foster, A. L.; Brown, G. E.; Parks, G. A. X-ray absorption fine structure study of As(V) and Se(IV) sorption complexes on hydrous Mn oxides. Geochimica et Cosmochimica Acta 2003, 67, 1937-1953. 10. Manceau, A.; Drits, V. A.; Silvester, E.; Bartoli, C.; Lanson, B. Structural mechanism of Co2+ oxidation by the phyllomanganate buserite. American Mineralogist 1997, 82, 1150-1175. 11. Zhao, W.; Cui, H.; Liu, F.; Tan, W.; Feng, X. Relationship between Pb2+ adsorption and average Mn oxidation state in synthetic birnessites. Clays and Clay Minerals 2009, 57, 513-520. 12. Zhu, M.; Ginder-Vogel, M.; Parikh, S. J.; Feng, X.-H.; Sparks, D. L. Cation Effects on the Layer Structure of Biogenic Mn-Oxides. Environmental Science & Technology 2010, 44, 4465-4471.


Page 22 of 29

Page 23 of 29

517 518 519 520 521 522 523 524 525 526 527 528 529 530 531 532 533 534 535 536 537 538 539 540 541 542 543 544 545 546 547 548 549 550 551 552 553 554 555 556 557 558 559 560 561 562 563 564 565 566 567 568 569 570 571 572 573 574 575 576


13. Landrot, G.; Ginder-Vogel, M.; Livi, K.; Fitts, J. P.; Sparks, D. L. Chromium(III) Oxidation by Three Poorly-Crystalline Manganese(IV) Oxides. 1. Chromium(III)-Oxidizing Capacity. Environmental Science & Technology 2012, 46, 11594-11600. 14. 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. Environmental Science & Technology 2010, 44, 8467-8472. 15. Tournassat, C.; Charlet, L.; Bosbach, D.; Manceau, A. Arsenic(III) Oxidation by Birnessite and Precipitation of Manganese(II) Arsenate. Environmental Science & Technology 2002, 36, 493-500. 16. Tonkin, J. W.; Balistrieri, L. S.; Murray, J. W. Modeling sorption of divalent metal cations on hydrous manganese oxide using the diffuse double layer model. Applied Geochemistry 2004, 19, 29-53. 17. Villalobos, M., The Role of Surface Edge Sites in Metal(loid) Sorption to Poorly-Crystalline Birnessites. In Acs Sym Ser, American Chemical Society: 2015; Vol. 1197, pp 65-87. 18. Remucal, C. K.; Ginder-Vogel, M. A critical review of the reactivity of manganese oxides with organic contaminants. Environmental Science: Processes & Impacts 2014, 16, 1247-1266. 19. Joshi, T. P.; Zhang, G.; Jefferson, W. A.; Perfilev, A. V.; Liu, R.; Liu, H.; Qu, J. Adsorption of aromatic organoarsenic compounds by ferric and manganese binary oxide and description of the associated mechanism. Chemical Engineering Journal 2017, 309, 577-587. 20. Yin, H.; Li, H.; Wang, Y.; Ginder-Vogel, M.; Qiu, G.; Feng, X.; Zheng, L.; Liu, F. Effects of Co and Ni co-doping on the structure and reactivity of hexagonal birnessite. Chemical Geology 2014, 381, 10-20. 21. Yu, Q.; Sasaki, K.; Tanaka, K.; Ohnuki, T.; Hirajima, T. Structural factors of biogenic birnessite produced by fungus Paraconiothyrium sp. WL-2 strain affecting sorption of Co2+. Chemical Geology 2012, 310–311, 106-113. 22. Yin, H.; Liu, F.; Feng, X.; Liu, M.; Tan, W.; Qiu, G. Co2+-exchange mechanism of birnessite and its application for the removal of Pb2+ and As(III). Journal of Hazardous Materials 2011, 196, 318-326. 23. Kwon, K. D.; Refson, K.; Sposito, G. Understanding the trends in transition metal sorption by vacancy sites in birnessite. Geochimica et Cosmochimica Acta 2013, 101, 222-232. 24. Villalobos, M.; Escobar-Quiroz, I. N.; Salazar-Camacho, C. The influence of particle size and structure on the sorption and oxidation behavior of birnessite: I. Adsorption of As(V) and oxidation of As(III). Geochimica et Cosmochimica Acta 2014, 125, 564-581. 25. Landrot, G.; Ginder-Vogel, M.; Sparks, D. L. Kinetics of Chromium(III) Oxidation by Manganese(IV) Oxides Using Quick Scanning X-ray Absorption Fine Structure Spectroscopy (Q-XAFS). Environmental Science & Technology 2010, 44, 143-149. 26. Tang, Y.; Webb, S. M.; Estes, E. R.; Hansel, C. M. Chromium(iii) oxidation by biogenic manganese oxides with varying structural ripening. Environmental Science: Processes & Impacts 2014, 16, 2127-2136. 27. Fendorf, S. E.; Zasoski, R. J. Chromium(III) oxidation by δ-manganese oxide (MnO2). 1. Characterization. Environmental Science & Technology 1992, 26, 79-85. 28. Wang, Z.; Lee, S.-W.; Kapoor, P.; Tebo, B. M.; Giammar, D. E. Uraninite oxidation and dissolution induced by manganese oxide: A redox reaction between two insoluble minerals. Geochimica et Cosmochimica Acta 2013, 100, 24-40. 29. Wang, Z. M.; Giammar, D. E. Metal Contaminant Oxidation Mediated by Manganese Redox Cycling in Subsurface Environment. Acs Sym Ser 2015, 1197, 29-50. 30. Burns, R. G. The uptake of cobalt into ferromanganese nodules, soils, and synthetic manganese (IV) oxides. Geochimica et Cosmochimica Acta 1976, 40, 95-102. 31. McKenzie, R. The reaction of cobalt with manganese dioxide minerals. Soil Research 1970, 8, 97-106. 32. Murray, J. W. The interaction of cobalt with hydrous manganese dioxide. Geochimica et Cosmochimica Acta 1975, 39, 635-647. 33. Tanaka, K.; Yu, Q.; Sasaki, K.; Ohnuki, T. Cobalt(II) Oxidation by Biogenic Mn Oxide Produced by Pseudomonas sp. Strain NGY-1. Geomicrobiology Journal 2013, 30, 874-885. 34. Simanova, A. A.; Peña, J. Time-Resolved Investigation of Cobalt Oxidation by Mn(III)-Rich δ-MnO2 Using Quick X-ray Absorption Spectroscopy. Environmental Science & Technology 2015, 49, 10867-10876. 35. van Genuchten, C. M.; Pena, J. Sorption selectivity of birnessite particle edges: a d-PDF analysis of Cd(II) and Pb(II) sorption by δ-MnO2 and ferrihydrite. Environmental Science: Processes & Impacts 2016, 18, 1030-1041. 36. Morgan, J. J.; Stumm, W. Colloid-chemical properties of manganese dioxide. Journal of Colloid Science 1964, 19, 347-359. 37. 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. Environmental Science & Technology 2009, 43, 6655-6661. 38. Villalobos, M.; Toner, B.; Bargar, J.; Sposito, G. Characterization of the manganese oxide produced by pseudomonas putida strain MnB1. Geochimica et Cosmochimica Acta 2003, 67, 2649-2662.



577 578 579 580 581 582 583 584 585 586 587 588 589 590 591 592 593 594 595 596 597 598 599 600 601 602 603 604 605 606 607 608 609 610 611 612 613 614 615 616 617 618 619 620 621 622 623 624 625 626 627 628 629 630

39. Marafatto, F. F.; Lanson, B.; Pena, J. Crystal growth and aggregation in suspensions of δ-MnO2 nanoparticles: implications for surface reactivity. Environmental Science: Nano 2018, DOI: 10.1039/C7EN00817A. 40. Mueller, O. JAQ Analyzes QEXAFS, 3.3.46; Bergische Universität Wuppertal, 2015. 41. Ravel, B.; Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: data analysis for X-ray absorption spectroscopy using IFEFFIT. Journal of Synchrotron Radiation 2005, 12, 537-541. 42. Webb, S. M. SIXpack: a graphical user interface for XAS analysis using IFEFFIT. Physica Scripta 2005, T115, 1011-1014. 43. Kelly, S. D.; Hesterberg, D.; Ravel, B., Analysis of Soils and Minerals Using X-ray Absorption Spectroscopy. In Methods of Soil Analysis Part 5—Mineralogical Methods, Ulery, A. L.; Richard Drees, L., Eds. Soil Science Society of America: Madison, WI, 2008; pp 387-463. 44. Wang, Y.; Stone, A. T. Phosphonate- and Carboxylate-Based Chelating Agents that Solubilize (Hydr)oxide-Bound MnIII. Environmental Science & Technology 2008, 42, 4397-4403. 45. Elzinga, E. J.; Kustka, A. B. A Mn-54 Radiotracer Study of Mn Isotope Solid–Liquid Exchange during Reductive Transformation of Vernadite (δ-MnO2) by Aqueous Mn(II). Environmental Science & Technology 2015, 49, 4310-4316. 46. Seri-Levy, A.; Avnir, D. Kinetics of diffusion-limited adsorption on fractal surfaces. The Journal of Physical Chemistry 1993, 97, 10380-10384. 47. Alvarez, N. J.; Walker, L. M.; Anna, S. L. Diffusion-limited adsorption to a spherical geometry: The impact of curvature and competitive time scales. Physical Review E 2010, 82, 011604. 48. Scheinost, A. C.; Abend, S.; Pandya, K. I.; Sparks, D. L. Kinetic Controls on Cu and Pb Sorption by Ferrihydrite. Environmental Science & Technology 2001, 35, 1090-1096. 49. Boonfueng, T.; Axe, L.; Xu, Y.; Tyson, T. A. The impact of Mn oxide coatings on Zn distribution. Journal of Colloid and Interface Science 2006, 298, 615-623. 50. Lide, D. R., CRC Handbook of Chemistry and Physics. 87th ed.; Taylor and Francis: Boca Raton, Florida, USA, 2007. 51. Liang, X. R.; Zhao, Z. X.; Zhu, M. Q.; Liu, F.; Wang, L. J.; Yin, H.; Qiu, G. H.; Cao, F. F.; Liu, X. Q.; Feng, X. H. Self-assembly of birnessite nanoflowers by staged three-dimensional oriented attachment. Environmental Science: Nano 2017, 4, 1656-1669. 52. Raju, M.; van Duin, A. C. T.; Fichthorn, K. A. Mechanisms of Oriented Attachment of TiO2 Nanocrystals in Vacuum and Humid Environments: Reactive Molecular Dynamics. Nano Letters 2014, 14, 1836-1842. 53. Simanova, A. A.; Kwon, K. D.; Bone, S. E.; Bargar, J. R.; Refson, K.; Sposito, G.; Peña, J. Probing the sorption reactivity of the edge surfaces in birnessite nanoparticles using nickel(II). Geochimica et Cosmochimica Acta 2015, 164, 191-204. 54. Peña, J.; Bargar, J. R.; Sposito, G. Copper sorption by the edge surfaces of synthetic birnessite nanoparticles. Chemical Geology 2015, 396, 196-207. 55. Rosso, K. M.; Morgan, J. J. Outer-sphere electron transfer kinetics of metal ion oxidation by molecular oxygen. Geochimica et Cosmochimica Acta 2002, 66, 4223-4233. 56. Silvester, E.; Charlet, L.; Manceau, A. Mechanism of chromium(III) oxidation by Na-buserite. The Journal of Physical Chemistry 1995, 99, 16662-16669. 57. Lanson, B.; Drits, V. A.; Gaillot, A. C.; Silvester, E.; Plancon, A.; Manceau, A. Structure of heavymetal sorbed birnessite: Part 1. Results from X-ray diffraction. American Mineralogist 2002, 87, 1631-1645. 58. Johnson, K.; Purvis, G.; Lopez-Capel, E.; Peacock, C.; Gray, N.; Wagner, T.; März, C.; Bowen, L.; Ojeda, J.; Finlay, N.; Robertson, S.; Worrall, F.; Greenwell, C. Towards a mechanistic understanding of carbon stabilization in manganese oxides. Nature Communications 2015, 6, 7628. 59. Simanova, A.; Kroll, A.; Pena, J. In Organo-mineral interactions in Pseudomonas putida-birnessite assemblages: Impact on mineral reactivity, EGU General Assembly Conference, Vienna, 2016. 60. Tebo, B. M.; Bargar, J. R.; Clement, B. G.; Dick, G. J.; Murray, K. J.; Parker, D.; Verity, R.; Webb, S. M. BIOGENIC MANGANESE OXIDES: Properties and Mechanisms of Formation. Annual Review of Earth and Planetary Sciences 2004, 32, 287-328. 61. Thenuwara, A. C.; Shumlas, S. L.; Attanayake, N. H.; Aulin, Y. V.; McKendry, I. G.; Qiao, Q.; Zhu, Y.; Borguet, E.; Zdilla, M. J.; Strongin, D. R. Intercalation of Cobalt into the Interlayer of Birnessite Improves Oxygen Evolution Catalysis. ACS Catalysis 2016, 6, 7739-7743.

631


Page 24 of 29

Page 25 of 29

632 633 634 635 636 637 638 639 640 641 642 643 644 645 646 647 648 649 650 651 652 653 654 655 656


a

b

c

d

Figure 1. Structural model for the four considered Co surface species on δ-MnO2: (a) Co(III)INC and (b) Co(II)-TCS associated with vacancy sites and (c) Co(III)-DES and (d) Co(II)DCS associated with edge sites. Gray octahedra: Mn; red spheres: O; red octahedra: Co(III); green spheres: Co(II). The color pattern at the bottom right corner of each panel represents the color code used for each species in Figure 4.



0.40

a

0.04 fCo

= 1.00

fCo

= 0.73

fCo

= 0.61

0.03

ad

0.10

0.02

0.01 fCo

ad

657 658 659 660 661 662 663

aq

0.20

Mn /Mn

ad

0.00

b

total

0.30

-1

q (mol Co mol Mn)

ad

0

= 1.00

10 20 30 400 800 1200 1600 Reaction Time (min)

0.00

0

10 20 30 400 800 1200 1600 Reaction Time (min)

Figure 2. (a) Surface loading of Co (qCo) on δ-MnO2 and (b) the fraction of aqueous Mn released from δ-MnO2: red: pH8-Co32%; green: pH6-Co40%; blue: pH4-Co37%; light green: pH6-Co8%. The fraction of adsorbed Co, fCoad, reported in Figure 2a, corresponds to the fraction of Co adsorbed on δ-MnO2 relative to total Co added in each experiment (Table S1). Notice the break and the change of scale on the x-axis after 35 min.


Page 26 of 29

Page 27 of 29


a: pH 8

0.20

0.10

0

10 20 30

664

0.30

c: pH 4

0.10

0

0.20

q Co(II) q Co(III)

10 20 30

400 800 1200 1600 Time (min)

d: pH 6

q Co(II) q C0(III)

0.15

-1

q (mol Co mol Mn)

0.30

0.20

0.00

400 800 1200 1600 Time (min)

-1

q (mol Co mol Mn)

0.40

0.20

0.10

0.00

666 667 668 669 670 671 672

q Co(II) q Co(III)

b: pH 6

-1

q (mol Co mol Mn)

0.30

0.00

665

0.40

q Co(II) q Co(III)

-1

q (mol Co mol Mn)

0.40

0

10 20 30

400 800 1200 1600 Time (min)

0.10

0.05

0.00

0

10 20 30

400 800 1200 1600 Time (min)

Figure 3. Loading of Co(II) (qCoII) and Co(III) (qCoIII) on δ-MnO2: (a) red: pH8-Co32%; (b) green: pH6-Co40%; (c) blue: pH4-Co37%; (d) light green: pH6-Co8%. The data used in figures are reported in Table S3 and are calculated from the qCo values (Table S1) and results from LCF analysis of the QXANES spectra (Table S2). Note the break and change of scale on the x-axis after 35 min for all panels. Note the different scale of y-axis in the panel d relative to the other panels.



0.40

0.40

q (mol Co mol Mn)

b: pH 6 0.30

-1

0.30

-1

q (mol Co mol Mn)

a: pH 8

0.20 0.10 0.00

673

7

0.20 0.10 0.00

11 16.7 31.2 62 179 817 1612 Time (min)

0.40

q (mol Co mol Mn)

29 61 568 804 1405 Time (min)

d: pH 6 0.30

-1

0.30

-1

q (mol Co mol Mn)

7.5 11.7 17

0.40

c: pH 4

0.20 0.10 0.00

674 675 676 677 678 679 680 681 682 683 684

Page 28 of 29

6.5

11 20.5 29 61.5 285 692 1500 Time (min)

q Co(III)-INC on vacancy sites q Co(II)-TCS on vacancy sites q Co(III)-DES on edge sites q Co(II)-DCS on edge sites

0.20 0.10 0.00

7.5 11.5 23.5 34.5 59.2 275 473 1379 Time (min)

Figure 4. Site-specific surface speciation of Co: (a) pH8-Co32%; (b) pH6-Co40%; (c) pH4Co37%; (d) pH6-Co8%. The values used in the figures are reported in Table S3 and are estimated from the surface loadings (Table S1), fCoII and fCoIII values obtained from LCF of the QXANES spectra (Table S2), and the distribution of Co(II) between vacancy sites (y) and edge sites (1 - y) and that of Co(III) between vacancy sites (x) and edge sites (1 - x) as obtained from shell-by-shell fitting of the QEXAFS data (Table S4-S7). Uncertainty for sitespecific surface loadings is lower than 20 % for about half of the sub-samples; the trends in surface speciation are consistent with the evolution of amplitudes of the FT transform peaks shown in Figure S6. Uncertainties are not shown in the figures for clarity.


Page 29 of 29

685 686 687


Table 1. Rates of Co(II) sorption and oxidation on δ-MnO2. Values within brackets indicate uncertainty in the last reported digit of Co(II) oxidation rates propagated from the LCF analysis of XANES spectra. Rate (Δq/Δt, mol Co mol-1 Mn min-1) Experiment Process 0 – 3.5 min 3.5 – 11 min 11 min – end of exp. -3 -4 pH8-Co32% Co(II) sorption 82 × 10 33 × 10 6 × 10-6 Co(II) oxidation 45 (4) × 10-3 21 (20)× 10-4 31 (9) × 10-6 pH6-Co40%

Co(II) sorption Co(II) oxidation

56 × 10-3 24 (4) × 10-3

55 × 10-4 35 (15) × 10-4

41 × 10-6 45 (11) × 10-6

pH4-Co37%


47 × 10-3 17 (3) × 10-3

21 × 10-4 17 (15) × 10-4

34 × 10-6 36 (8) × 10-6

pH6-Co8%


23 × 10-3 10 (1) × 10-3

0 11 (6) × 10-4

0 22 (4) × 10-6

688


Reactions at particle edges versus vacancy sites - ACS Publications

Recommend Documents