Fe-Impregnated Mineral Colloids for Peroxide Activation: Effects of

Subscriber access provided by UNIV OF CAMBRIDGE

Article

Fe-impregnated Mineral Colloids for Peroxide Activation: Effects of Mineral Substrate and Fe Precursor Yue Li, Libor Machala, and Weile Yan Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.5b03970 • Publication Date (Web): 29 Dec 2015 Downloaded from http://pubs.acs.org on January 2, 2016

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

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

Page 1 of 36

Environmental Science & Technology

1

Fe-impregnated Mineral Colloids for Peroxide Activation: Effects of Mineral

2

Substrate and Fe Precursor

3

Yue Li1, Libor Machala2, Weile Yan1,*

4 5

1

Department of Civil and Environmental Engineering, Texas Tech University, Lubbock, Texas,

6 7 8

United States 2

Regional Centre of Advanced Technologies and Materials, Department of Experimental Physics, Palacký University in Olomouc, Olomouc, Czech Republic

9 10 11 12 13 14 15 16 17 18 19 20 21 22

*Corresponding author. Tel: (806) 834 3478; Fax: (806) 742-3449

23

Email address: [email protected]

24

ACS Paragon Plus Environment


25

ABSTRACT

26

Heterogeneous iron species at the mineral/water interface are important catalysts for the

27

generation of reactive oxygen species at circumneutral pH. One significant pathway leading to

28

the formation of such species arises from deposition of dissolved iron onto mineral colloids due

29

to changes in redox conditions. This study investigates the catalytic properties of Fe-impregnated

30

on silica, alumina, and titania nanoparticles (as prototypical mineral colloids). Fe impregnation

31

was carried out by immersing the mineral nanoparticles in dilute Fe(II) or Fe(III) solutions at pH

32

6 and 3, respectively, in an aerobic environment. The uptake of iron per unit surface area

33

follows the order of nTiO2 > nAl2O3 > nSiO2 for both types of Fe precursors. Impregnation of

34

mineral particles in Fe(II) solutions results in predominantly Fe(III) species due to efficient

35

surface-mediated oxidation. The catalytic activity of the impregnated solids to produce OH∙

36

from H2O2 decomposition was evaluated using benzoic acid as a probe compound under dark

37

conditions. Invariably, the rates of benzoic acid oxidation with different Fe-laden particles

38

increase with the surface density of Fe until a critical density above which the catalytic activity

39

approaches a plateau, suggesting active Fe species are formed predominantly at low surface

40

loadings. The critical surface density of Fe varies with the mineral substrate as well as the

41

aqueous Fe precursor. Fe-impregnated on TiO2 exhibits markedly higher activity than their

42

Al2O3 and SiO2 counterparts. The speciation of interfacial Fe is analyzed with diffuse

43

reflectance UV-Vis analysis and interpretation of the data in the context of benzoic oxidation

44

rates suggests that the surface activity of the solids for OH∙ generation correlates strongly with

45

the isolated (i.e., mononuclear) Fe species. Therefore, iron dispersed on mineral colloids is a

46

significant form of reactive iron surfaces in the aquatic environment.


Page 2 of 36

Page 3 of 36


47

INTRODUCTION

48

Environmental redox reactions catalyzed by mineral surfaces play an important role in

49

contaminant transformations, cycling of nutrients and trace elements, and mineral formation and

50

dissolution [1,2]. Among many catalytically active metals, iron is of special interest to

51

researchers because of its ubiquitous presence in the aquatic environment [3,4] and its distinct

52

chemistry that includes facile cycling between Fe(II) and Fe(III) valence states within

53

ecologically important redox conditions, strong affinity for oxygen-containing ligands and

54

mineral surfaces, and the presence of an atomic exchange mechanism enabling efficient material

55

exchange across the solid/water interface [5,6].

56

Iron plays an essential role in catalyzing the production of reactive oxygen species (ROS) and

57

the decomposition of organic carbons in natural waters. Photochemical reactions involving

58

dissolved iron, mainly in complexes with small organic ligands, account for an important source

59

of H2O2 in weakly acidic water such as rain or fog water [7,8]. At environmentally relevant pH

60

(4 – 9), however, iron is predominantly associated with the solid phase in an oxic environment.

61

Hence, there is a substantial interest in exploring the use of iron-containing minerals (e.g.,

62

amorphous iron oxides, ferrihydrite, goethite, and hematite) as heterogeneous catalysts for

63

organic carbon degradation with naturally occurring or artificially injected oxidants

64

[9,10,11,12]. The latter process, known as in situ chemical oxidation (ISCO), is considered a

65

viable alternative to the more energy and resource-intensive ex situ treatment methods for

66

groundwater and soil decontamination.

67

Many studies have been performed to assess the catalytic activity of various forms of iron-

68

containing minerals. Lin and Gurol evaluated the decomposition of H2O2 on goethite and

69

observed the kinetics to conform to the Langmuir-Hinshelwood model [9]. As there exist



70

competing pathways leading to H2O2 decomposition [13,14], the efficiency of the solid phase to

71

produce hydroxyl radical (OH• ) with respect to unproductive decomposition of H2O2 to O2 and

72

H2O cannot be assessed from the bulk H2O2 consumption rates. Using probe compounds that

73

react rapidly with OH•, the rates of OH• production on different surfaces have been measured

74

[10,15,16]. It was found that considerably large variations exist of OH• production rates among

75

different forms of iron oxides [10], and studies using native aquifer materials have observed

76

rapid H2O2 decay without quantitative oxidation of contaminants [11,17], implying oxidant yield

77

is strongly dependent on the nature of the minerals and their surface or bulk characteristics. In

78

spite of many studies, the key attributes affecting the reactivity of mineral surface are not well

79

understood, and attempts to predict the rates of reactive oxidant formation have achieved limited

80

success.

81

With an aim to understand the nature of iron-bearing minerals, our focus in this work is placed

82

on a type of heterogeneous surfaces consisting of iron residing on the surface of a background

83

mineral substrate. Although extensive studies have been performed to understand interactions at

84

the surface of single-phase iron minerals (e.g., goethite, ferrihydrite, and hematite), the

85

environmental chemistry of iron impregnated on common mineral solids such as silica and

86

aluminosilica is less explored. Iron-coated mineral solids are a common occurrence in both

87

natural and engineered treatment systems in the surface or near-surface environment [18,19]. In

88

groundwater, oxidation of iron ores (e.g., pyrite) or reduction of ferric oxides contribute to a

89

significant level of dissolved iron (mainly as Fe(II)). When the water approaches oxic zones,

90

rapid Fe(II) oxidation is catalyzed by mineral surfaces, while homogenous Fe(II) oxidation in the

91

aqueous phase is a much slower process at acidic or neutral pH [2,20]. Therefore, in the

92

presence of abundant solid particles, impregnation of iron onto background colloids could be a


Page 4 of 36

Page 5 of 36


93

more preferable pathway for iron deposition than homogeneous oxidation and nucleation. The

94

resultant interfacial iron species may exist as well-dispersed surface complexes, hydrolyzed

95

Fe(III) polymers, or discrete ferric (oxyhydr)oxide precipitates [21]. One or more of these species

96

may exist simultaneously and their distribution is sensitive to the solution chemistry as well as

97

the nature of the underlying substrate [19,22,23]. Although not much has been learned about the

98

chemistry of these interfacial iron species, in the context of synthetic iron catalysts, incorporation

99

of iron onto a separate solid phase is expected to have a profound impact on the quantity (i.e.,

100

surface density) and the activity of reactive sites [24,25,26], and such effect may be attributed to a

101

suite of factors such as improved catalyst dispersion as well as catalyst-support interactions that

102

alter the structure, acid/base properties, and charge density of the catalytically active species [14,

103

24 27

104

different from their natural counterparts, limited inference can be made about the properties of

105

iron-impregnated minerals arising in dilute aqueous systems of environmental interest.

106

The objective of this study is to examine the morphology, surface chemistry, and catalytic

107

activity of iron anchored on three colloid-sized mineral particles, namely silica, alumina and

108

titania. Iron impregnation took place by exposing the pristine mineral particles to dilute aqueous

109

solutions of ferrous and ferric ions. The activity of these iron-impregnated solids was evaluated

110

using peroxide oxidation with benzoic acid as the probe compound. Additionally, spectroscopic

111

characterization of the mineral particles was performed to determine the different forms of Fe

112

species present at the surface in an effort to understand the relationship between Fe speciation

113

and catalytic activity.

, ]. As the reaction conditions and precursors leading to synthetic iron catalysts are vastly



114

EXPERIMENTAL SECTION

115

Materials. A complete list of reagents used in this study is available in the Supporting

116

Information (SI). Nanoparticles of silica (catalogue #718483), aluminum oxide (#718475), and

117

titanium(IV) oxide (#718467) were purchased from Sigma-Aldrich. They are denoted as nSiO2,

118

nAl2O3, and nTiO2, respectively, in this work. The particles of nSiO2 and nAl2O3 were

119

synthesized using aerogel preparation routes, and nTiO2 is a commercial equivalent of the widely

120

used P25 as confirmed by the presence of anatase (~ 80%) and rutile (~ 20%) with X-ray

121

diffraction analysis (Table 1 and Figure S1). The nanoparticles were used as-received.

122

Stock solutions of Fe(III) at 20 mM and Fe(II) at 5 mM were prepared from ferric chloride and

123

ferrous sulfate, respectively, on the day of experiments and were used without acidification.

124

Buffer solutions of 2-(N-Morpholino)ethanesulfonic acid (MES) and 3-Morpholinopropane-1-

125

sulfonic acid (MOPS) were prepared and adjusted to desired pH values with dilute NaOH or

126

HCl.

127

Preparation of Iron-Impregnated Mineral Colloids. Iron-impregnated mineral solids were

128

prepared by immersing pristine nanoparticles in a series of Fe(II) or Fe(III) aqueous solutions

129

with Fe concentrations varying between 0.05 and 0.8 m M and the particle loading fixed at 0.6

130

g/L. All impregnation steps involved in this study were performed under atmospheric conditions.

131

A separate study investigating the deposition of Fe(II) in anoxic environment is underway. Two

132

types of aqueous iron precursors (i.e., ferrous and ferric ions) were used. Specifically, an Fe(II)

133

impregnation solution was prepared by adding a small aliquot of Fe(II) stock solution to a

134

background electrolyte solution containing 5 mM NaCl and 5 mM MES. The pH of the

135

electrolyte solution was adjusted and allowed to equilibrate for 2 h prior to Fe(II) addition, such

136

that the mixture had an initial pH of 6.0 +/- 0.1. Immediately after the addition of Fe(II) stock,


Page 6 of 36

Page 7 of 36


137

pristine mineral particles were added and the suspension was mixed on a wrist-action mechanical

138

shaker at 250 rpm for 24 h. The resultant solids were collected by centrifugation at 1000 rpm for

139

10 min (Beckman Coulter Avanti J-E Centrifuge), dried in air for 48 h, and stored at 4 ̊C prior to

140

use. No rinsing was performed during solid collection to prevent changes in surface chemistry or

141

iron leaching. Impregnation in Fe(III) solutions followed the same procedure as the above except

142

that the Fe(III) solution contains a background electrolyte of 10 mM NaCl and the initial pH was

143

acidified as needed to 3.0 to avoid precipitation of ferric hydroxide.

144

Samples were withdrawn from the solutions prior to and at the end of the impregnation and were

145

analyzed for total iron concentration to determine the loading of Fe on different solids. All

146

solution samples were filtered through 0.2 μm syringe filter (Millipore SLLG025SS) and

147

acidified before analysis. Acid digestion of the filtrate suggests the separation efficiency of the

148

filtration step is greater than 98% for all solids.

149

Aqueous Oxidation Experiments. Catalytic oxidation reactions were carried out at room

150

temperature (20 oC) in 40-mL glass vials wrapped with aluminum foil to prevent potential

151

photochemical reactions. Benzoic acid (or benzoate, with pKa = 4.2) was used as a probe to

152

determine the rate of oxidant formation during decomposition of H2O2 on iron-impregnated

153

mineral oxides. The initial concentration of benzoic acid was 6.9 x 10-5 M and iron-impregnated

154

solids were dosed at 1.5 g/L. Upon introduction of the solids to a benzoic acid solution, the

155

suspension was sonicated for 3 min to break up loose aggregates and was amended immediately

156

with an aliquot of H2O2 solution (Co = 11 mM) to initiate the oxidation reaction. The reactors

157

were mixed on a shaker at 250 rpm for up to 48 h. Except for a subset of experiments to assess

158

the effect of pH, no buffer was used in the experiments due to concerns over their consumption

159

of reactive oxidants. For the unbuffered systems, the final pH was approximately 4 and the



Page 8 of 36

160

variation of pH was found to be less than one unit throughout the reaction periods. Experiments

161

at higher pH were conducted using 5 mM of MES (pH 5) and MOPs (pH 7), respectively.

162

During the oxidation experiments, samples were withdrawn at predetermined intervals and

163

filtered through 0.2 μm syringe filters (Millipore SLLG025SS) with the filtrate collected in 2-mL

164

amber HPLC vials for benzoate analysis. A small amount (32 μL) of pure methanol was spiked

165

into vials to quench further reactant oxidation by residue H2O2. Descriptions of Analytical

166

Methods are available in the SI.

167

Surface Characterizations. Details on BET, TEM and XRD analysis were available in the SI.

168

UV-Vis diffuse reflectance spectra (DRS) were acquired on a Cary 5000 UV-Vis-NIR

169

spectrophotometer equipped with a Praying MantisTM diffuse reflectance sample holder (Harrick

170

Scientific Products). A small amount of dry solids was loaded into a sample cup and the surface

171

of the sample was gently pressed against a glass slide. The process was repeated several times to

172

ensure the particles were densely packed. Spectrum acquisition was performed in the range of

173

200 nm to 800 nm at a step size of 1 nm and an averaging time of 1 second. The spectra of iron-

174

impregnated solids were taken using the respective pristine solid (i.e., unamend nSiO2, nAl2O3,

175

nTiO2) as the reference standard. For a sample of infinite thickness (or in practical terms, a

176

sample thickness greater than a few hundreds of microns), the diffuse reflectance of the sample

177

(𝑅∞ ) is related to the ratio of the apparent absorption (K) and apparent scattering (S) coefficients

178

through the Schuster-Kubelka-Munk (SKM) function (F(𝑅∞ )) [28,29]:

179

𝐹(𝑅∞ ) =

(1−𝑅∞ )2 2𝑅∞

=

𝐾

(1)

𝑆

180

It is commonly assumed that S varies little within the spectrum window [28], thus 𝐹(𝑅∞ ) is a

181

direct measure of K, and the latter increases linearly with the concentration of the absorbing

182

species within a limited range of 𝐹(𝑅∞ ) (0 < 𝐹(𝑅∞ ) < 0.5). For more strongly absorbing


Page 9 of 36


183

samples, dilution of solids with a non-absorbing reference is often used, but we found this

184

method tends to create inhomogeneous mixture that results in poor spectral reproducibility.

185

Theoretical calculations indicate that 𝐹(𝑅∞ ) deviates from a linear dependence on the true

186

absorption coefficient by less than 10% when 𝐹(𝑅∞ ) increases from 0.5 to 3.5 [30], therefore, the

187

samples were analyzed as is without dilution.

188

RESULTS AND DISCUSSION

189

Characterization of Mineral Particles

190

TEM images of the fresh nSiO2, nAl2O3, and nTiO2 and those immersed in 0.2 mM of Fe(II) or

191

Fe(III) solutions were shown in Figure 1. Pristine nSiO2 appears amorphous (Figure 1a), whereas

192

the pristine nAl2O3 and nTiO2 exhibit clear lattice fringes and surface facets (insets of Figure 1b

193

and c). X-ray diffraction (XRD) analysis confirms that the dominant crystalline phases in the

194

alumina particles are ε-Al2O3 and δ-Al2O3, and nTiO2 is comprised of a mixture of anatase and

195

rutile (Table 1 and Figure S1). No distinct crystalline phase can be identified through XRD for

196

nSiO2. The average size of the primary particles estimated through TEM images agrees on a

197

large part with the size info provided by the manufacturers (Table 1). All particles show

198

considerable aggregation, and in the case of nSiO2, fusing of particles into porous foam-like

199

structure is evident, consistent with its aerogel synthesis history. Images of particles exposed to

200

0.2 mM Fe(II) solutions for 24 h suggest there is no significant morphological changes. In

201

contrast, nSiO2 immersed in an equivalent concentration of Fe(III) solution develops small

202

surface deposits that appear as dark speckles on the substrate. The speckles measure approx. 3 - 5

203

nm in size in the high resolution image (inset of Figure 1a third image). No iron precipitates

204

could be discerned on nAl2O3 and nTiO2 impregnated in ferric solutions. The measured BET

205

surface areas of the pristine and Fe-impregnated solids are listed in Table 1. Slight changes in



206

surface area were noted with iron deposition, however the trend is not consistent among different

207

types of minerals and the changes are relatively small (< 6% of the original surface area). Given

208

this, the surface areas of the pristine particles are used to calculate surface normalized iron

209

loading in the subsequent discussion.

210

Uptake of Aqueous Iron Species by Mineral Particles

211

A series of iron-coated mineral particles were prepared by mixing the pristine particles with

212

varying concentrations of Fe(II) or Fe(III) solutions (0.05 – 0.8 mM) for 24 h. All experiments

213

were performed in the atmospheric environment. Since surface-catalyzed Fe(II) oxidation

214

proceeds at a considerable rate at circumneutral pH [20,31], it is expected that the Fe(II) initially

215

adsorbed on the mineral particles was readily converted to Fe(III) during the impregnation period

216

(24 h). Hence, regardless of the initial valence state of the Fe precursor in the aqueous solutions,

217

we expect it to be predominantly retained as Fe(III) on the solid surface at the end of the

218

impregnation step. This notion is validated by measuring Fe(II) and total Fe content in the solids

219

that had been immersed in Fe(II) solutions following the method by Amonette and Templeton

220

[32]. Total Fe recovery was between 70 to 93%, and Fe(II) was undetected in all solids (Table

221

S1).

222

The amounts of iron immobilized onto the three types of mineral particles as a function of the

223

final dissolved Fe concentration are plotted in Figure 2. With Fe(II) solutions, the final uptake of

224

iron per unit surface area at a given initial Fe(II) concentration follows the order of nTiO2 >

225

nAl2O3 > nSiO2 (Figure 2a). A vertical dashed line was used for nTiO2 because the final aqueous

226

Fe concentration was consistently below the instrument detection limit of 2 μM across the range

227

of the initial Fe concentration evaluated, implying that all ferrous ions (at up to 0.8 mM) were


Page 10 of 36

Page 11 of 36


228

effectively immobilized onto the nTiO2. The highest surface Fe density obtained for nTiO2 is 16

229

Fe atoms/nm2, a value significantly higher than the maximum cation adsorption sites on rutile

230

and anatase (~12.5 and 11 sites/nm2, respectively) [33,34]. Compared to that of nTiO2, the uptake

231

curves of nAl2O3 and nSiO2 show sign of saturation with increasing aqueous Fe(II)

232

concentration, however, neither curve can be fit adequately to the classical monolayer adsorption

233

isotherms. The highest iron loading was observed to be 6.7 and 3.3 atoms/nm2 on nAl2O3 and

234

nSiO2, respectively, and as a reference, the calculated adsorption site density of alumina (as γ-

235

Al2O3) and amorphous silica are approximately 8 and 4.6 sites/nm2 [35,36]. The experimentally

236

observed adsorption capacity is typically substantially smaller than the calculated site density

237

due to steric and electrostatic repulsion and the formation of multi-dentate complexes, therefore,

238

the results of Fe(II) uptake on the three types of mineral colloids suggest that multilayer

239

deposition or Fe(III) hydroxide precipitates are likely present at high Fe loadings. Deposition of

240

multiple layers of Fe(III) species on surfaces is plausible under aerobic conditions as the

241

oxidation of adsorbed Fe(II) followed by hydrolysis of the resultant Fe(III) species can create

242

additional sites for Fe(II) uptake [2]. Subsequent oxidation of the Fe(II) will propagate the

243

adsorption-oxidation cycle. As impregnation was conducted under similar conditions for all three

244

mineral particles, the trends in Figure 2a indicate that nTiO2 possesses the most favorable surface

245

for Fe(II) oxygenation among the mineral nanoparticles studied.

246

When impregnation took place in Fe(III) solutions (Figure 2b), the surface loading of iron

247

follows the same order as that in Fe(II) impregnation experiments with nTiO2 attracting the

248

greatest amount of Fe(III) from the aqueous phase, followed by nAl2O3 and nSiO2. It should be

249

noted that Fe(II) impregnation was carried out at pH 6.0, whereas that of Fe(III) was conducted

250

in acidic solutions (pH 3.0) due to limited solubility of Fe(III) at neutral pH. At pH 3, the



251

dominant aqueous Fe(III) species are Fe3+, Fe(OH)2+ and Fe(OH)2+, with polynuclear

252

hydroxocomplexes (e.g., Fe2(OH)24+) being relatively unimportant (Figure S2) [37]. The point of

253

zero charge (pHpzc) values of various mineral phases reported in the literature fall generally in the

254

range of 8 – 9 for alumina, 6 – 7 for P25 TiO2, and < 3 for silica materials [38,39]. Therefore,

255

adsorption of Fe(III) on nSiO2 is electrostatically favorable, whereas nAl2O3 and nTiO2 present

256

varying degrees of electrostatic repulsion against Fe(III) attachment. Evidently, the observed

257

impregnation behavior cannot be rationalized on the basis of surface charge, and we attribute the

258

affinity of nAl2O3 and nTiO2 for Fe(III) to strong chemical interaction involving inner-sphere

259

complex formation. Notably, all curves depict a rising tail at a high Fe(III) concentration, which

260

we interpret as surface-initiated Fe(III) precipitation on the ground that homogeneous nucleation

261

is very slow at this acidic pH. With heterogeneous precipitation being involved, attributes such

262

as density of surface functional groups, surface crystallographic orientation, and porosity of the

263

particles can all affect the structure of the Fe(III) formation. Delineating the dynamic process of

264

surface precipitation is not the main interest of this study, rather, we aim to relate the reactivity

265

of the impregnated iron species to their measurable structural properties.

266

Aqueous Catalytic Oxidation of Benzoic Acid

267

The reactivity of various solids in activating peroxide was evaluated using benzoic acid as a

268

probe compound. Compared with other probes (e.g., methanol or isopropanol), Benzoic acid is

269

known to react more selectively towards OH• among several forms of reactive oxygen species

270

generated during decomposition of H2O2 on iron surface (e.g., high-valent iron species,

271

superoxide or peroxyl radicals) [40,41,42]. Figure 3 shows the changes in benzoate acid

272

concentration as a function of reaction time. Each plot contains results of a particular type of

273

mineral particles impregnated in either Fe(II) or Fe(III) solutions. Legends indicate the initial


Page 12 of 36

Page 13 of 36


274

concentration of the iron precursors, and for ease of visual comparison, only two solids

275

(representing those impregnated in solutions of a relatively low or high Fe content) were shown

276

in each plot.

277

Heterogeneous catalytic reactions are inherently more complex than reactions in homogeneous

278

systems. In the present study, two factors complicate the analysis of solid phase activity: 1)

279

dissolution of iron species into the solution phase and activation of H2O2 by dissolved Fe via the

280

classical Fenton reaction, and 2) apparent loss of the probe compound through surface sorption.

281

We expect leaching of surface-anchored iron to be insignificant for the reasons that the Fe

282

precursor was essentially retained as Fe(III) on the particle surface and ferric species has

283

sparingly low aqueous solubility at the pH of the oxidation experiments. Analysis of dissolved

284

iron concentration at the end of the oxidation experiments detected negligible iron concentration

285

for all solids except for nSiO2 impregnated in Fe(III) solutions, of which a discussion will be

286

offered later. Marginal contribution of homogeneous H2O2 activation is further supported by

287

parallel experiments carried out in benzoic acid solutions in the presence of H2O2 under different

288

settings (Figure S3).. Insignificant change in benzoic acid concentration upon removal of the Fe-

289

impregnated solids serve as evidence that benzoic acid oxidation relies critically on surface-

290

catalyzed reactions. This point has been recognized in other peroxide oxidation systems

291

employing iron-bearing minerals at near neutral pH [14,10,9]. Our second concern related to

292

non-reactive sorption of benzoic acid was investigated by experiments mixing benzoate acid with

293

various solids without adding H2O2. pH of these experiments was adjusted to the same level as

294

the corresponding oxidation experiments and the results were shown as dashed lines in Figure 3

295

for solids exposed to the highest concentration of Fe(II) or Fe(III) solutions.



296

Several observations were notable in Figure 3. Adsorption of benzoic acid is insignificant onto

297

surface impregnated nSiO2. There is appreciable adsorption, however, onto the Fe-laden nAl2O3

298

and nTiO2. Furthermore, adsorption was considerably more rapid than benzoic acid oxidation

299

and the equilibrium of the former process was established within 2 h of catalyst addition.

300

Examining benzoic acid concentration beyond the initial 2 h reveals significant differences in its

301

rates among different forms of mineral substrates. Catalysts derived from nSiO2 and nAl2O3 are

302

significantly less reactive compared to those of nTiO2. Control experiments using pristine

303

particles suggest that only nTiO2 carries moderate catalytic activity (Figures 3c and 3f), while

304

bare nSiO2 and nAl2O3 were unable to activate H2O2 (Figure S4). Importantly, the reactivity of

305

nTiO2 increases considerably with increasing amount of iron deposited on the surface. Benzoic

306

acid degradation on nTiO2 impregnated in Fe(III) solutions was so rapid that the time frames of

307

oxidation and adsorption of the probe compound became comparable. For this reason, a reduced

308

dose (0.5 g/L) was used for nTiO2 impregnated in Fe(III) solutions. This does not affect our

309

subsequent discussion as surface activity of all solids is compared on a surface-area-normalized

310

basis. With nSiO2 and nAl2O3, we noted that iron precursor exerts a different effect on the

311

benzoic acid-solid interactions. Specifically, the silica and alumina pre-exposed to Fe(III)

312

solutions have greater adsorptive affinity for benzoic acid, but they seem to bestow lower

313

oxidation activity than solids derived from Fe(II) solutions.

314

More insights into the catalytic performance of the various particles can be gleaned by plotting

315

the surface activity as a function of iron loading (Figure 4). Surface activity is represented by the

316

surface area-normalized average benzoate degradation rate during the reaction time of 2 to 8 h to

317

exclude benzoic acid loss due to sorption. A strong and consistent substrate-dependent effect was

318

noticed across the three mineral series: namely, impregnated nTiO2 presents the most active


Page 14 of 36

Page 15 of 36


319

surface, followed by nAl2O3 and nSiO2 when compared at the same Fe loading. Another

320

common behavior recognizable from all curves in Figure 4 is that the reactivity of all iron-laden

321

particles approaches a plateau beyond a critical surface loading. Consequently, the catalytic

322

activity of these minerals does not scale with the total amount of iron impregnated on the

323

surface; rather, it is strongly influenced by a subset of iron species formed predominantly at a

324

low Fe coverage. The critical iron loading where surface activity saturates varies considerably

325

with the type of the minerals and the impregnation media. In the case of solids immersed in

326

Fe(II) solutions (Figure 4a), saturation occurs at a surface density of 0.4 Fe/nm2 on nSiO2 and at

327

~ 4 Fe/nm2 on nAl2O3 and nTiO2. nSiO2 and nAl2 O3 exposed to Fe(III) (Figure 4b) were

328

considerably less reactive than their Fe(II) counterparts in Figure 4a, with benzoic oxidation rate

329

peaks at as low as 0.4 Fe/nm2. On the contrary, nTiO2 demonstrates a steep rise in surface

330

activity with increasing Fe loading, although the trend tapers off beyond ~ 4 Fe atom/nm2.

331

Comparison of the results with Figure 2 suggest there is a reasonable agreement between the

332

critical Fe density and the range of Fe loading where the iron uptake transits from simple

333

adsorption to multilayer deposition or surface precipitation.

334

We further assessed the activity of impregnated nTiO2 at high pH (Figure 5), where 5 mM of

335

MES and MOPS were used to maintain the solution pH at 5.2 and 6.9, respectively. As expected,

336

lower rates of benzoic acid oxidation were recorded with increasing pH due to changes in Fe(III)

337

ligand environment (i.e., increasing number of hydroxo groups in the coordination sphere) [ 42],

338

scavenging effect of the buffering molecules, and higher rate of H2O2 decomposition at elevated

339

pH [15]. Nonetheless, at a constant pH, the data features consistently a surface saturation pattern

340

with the critical density occurring at a similar Fe loading. This close resemblance suggests that



341

the same form of Fe(III) species is likely responsible for OH∙ production, and the proportion of

342

the active species among total surface Fe seems to be primarily controlled by the solids

343

impregnation history and is not strongly affected by the pH of H2O2 reactions in the

344

circumneutral range.

345

Finally, an important consideration in using peroxide-based chemicals for advanced oxidation is

346

the efficiency of H2O2 consumption [17]. Catalysts of high activity may offer low selectivity

347

resulting in rapid decay of H2O2 in the reaction media. We examined the H2O2 utilization

348

efficiency of mineral particles impregnated in Fe(II) solutions (Figure S5). The efficiency is

349

computed as the molar ratio of benzoic acid degraded versus the amount of H2O2 consumed

350

during catalytic oxidation of benzoic acid (between 2 to 8 h). The impregnated nTiO2 has an

351

efficiency of approximately 1% across a broad range of Fe surface loading. In comparison,

352

impregnated nAl2O3 and nSiO2 display greater peroxide efficiency, but there is a marked decline

353

in efficiency at a higher Fe loading for both types of particles. We did not attempt to measure the

354

oxidant utilization efficiency of solids impregnated in Fe(III) solutions due to insignificant

355

changes in H2O2 concentration during the reaction period that prohibit accurate quantification.

356

Solid Phase Characterization by DR-UV-Vis

357

DR-UV-Vis is a useful technique to interrogate the nature and speciation of transition metal ions

358

on inorganic substrates. The absorption bands of Fe3+ in the UV-Vis region originate from two

359

types of electronic transitions within the 3d5 shell of Fe3+, namely, transitions due to the ligand

360

field of Fe3+ (i.e., d-d transitions) and ligand-to-metal charge-transfer (LMCT) [29,43]. d-d

361

transitions of Fe3+ ions are symmetry and spin-forbidden for isolated Fe3+ species and are

362

therefore very week for dispersed iron samples [29,44]. O  Fe3+ charge transfer occurs generally

363

below 300 nm for isolated Fe3+ species. Formation of dimerized or polymerized Fe(III)


Page 16 of 36

Page 17 of 36


364

complexes on the surface enables magnetic coupling of neighboring Fe3+ ions, thereby reducing

365

the energy required for LMCT transitions and shifting the corresponding absorption bands to

366

longer wavelengths [45].

367

Figure 6a presents the DR-UV-Vis spectra of nSiO2 impregnated in Fe(II) solutions. The

368

calculated Fe surface loading for each material was indicated in the legend. At the lowest Fe

369

loading, only two major bands at ~ 225 nm and ~ 265 nm can be discerned. The positions agree

370

with the two charge transfer transitions of isolated Fe3+ in an octahedral coordination

371

environment [45,46]. Increasing surface deposition by exposing the silica to a more concentrated

372

Fe(II) solution (0.2 mM) led to emergence of two shoulders at ~ 300 nm and ~ 350 nm. The 350

373

nm band has been assigned to trimeric ferric species based on UV-Vis and magnetic

374

characterizations [45]. The shoulder at 300 nm is tentatively attributed to dimeric iron. At higher

375

wavelength, a shoulder with a wide tailing appears at 480 nm. The spectra feature is

376

characteristic of ferric hydroxide nanoparticles and its broad shape represents superposition of

377

bands from particles of varying sizes and geometry [44,47,48]. Further increasing the Fe loading

378

results in uprising of all major bands and broadening of the low energy shoulder.

379

The spectra of silica nanoparticles exposed to Fe(III) solutions suggests polymeric or

380

nanoparticle were formed even at a very low Fe surface coverage (Figure 6b). Interestingly,

381

nSiO2 exposed to 0.2 mM Fe(III) has more intense UV-Vis absorption than that in contact with

382

0.8 mM Fe(III). This anomaly behavior coincides with a decrease in benzoic acid oxidation rate

383

with nSiO2 exposed to higher concentration of Fe(III) (Figure 4). It is also noted that the

384

suspension of nSiO2 impregnated in elevated concentrations of Fe(III) solutions have an

385

appreciable level of ‘dissolved’ Fe after reactions with benzoic acid. These observations can be

386

reasonably accounted for if loose, unattached iron oxide nanoparticles were formed in the



387

suspensions. One clue that favors this postulation is our observation of discrete particulate matter

388

on nSiO2 exposed to a ferric solution in the TEM image (Figure 1a, last image). It is possible that

389

these nanoparticles originate from surface precipitation and are prone to dislodgment from the

390

mineral support during subsequent experiments. Nano-precipitates may also arise from structural

391

re-arrangement of the initially adsorbed species since benzoic acid oxidation was conducted at a

392

higher pH than the impregnation procedure. Further investigations would be required to elucidate

393

their formation, and based on the present data, this phenomenon is restricted to nSiO 2 exposed to

394

Fe(III) solutions only.

395

The spectra of nAl2O3, shown in Figures 6c and 6d, reflect concomitant increases of single,

396

polymerized, and nanoparticulate species with increasing iron deposition irrespective of its initial

397

valence state. The concurrent emergence of isolated as well as clustered Fe species on nAl2O3

398

implies that surface nucleation and precipitation may set off at a relatively low surface coverage

399

prior to exhaustion of surface adsorption sites. Whether isolated Fe predominates at low loadings

400

(as in the case of nSiO2) or multiple species arise in parallel would depend on the energies of the

401

respective processes (adsorption vs. nucleation and precipitation) and the situation should vary

402

with the nature of the substrates and solution conditions [22].

403

Given the complex nature of iron deposition on mineral surfaces, our ability to differentiate iron

404

speciation on the surface becomes all the more relevant. Unfortunately, exact quantification of

405

species in a UV-Vis spectrum is challenging due to their different molar extinction coefficients.

406

It is permissible, however, to determine the relative abundance of a particular species in a series

407

of solids based on the area of the respective spectral peak(s). The spectra were deconvoluted

408

using a set of Gaussian bands at the positions indicated in Figure 6. Signals in the regions of 350

409

– 450 nm and > 450 nm, representing polymeric and discrete ferric hydroxide nanoparticles,


Page 18 of 36

Page 19 of 36


410

respectively, were fitted using one or two broad Gaussian components since each region contains

411

a multitude of transitions that cannot be spectrally resolved [44,48]. The deconvoluted spectra are

412

shown in Figure S6. Plotting the rates of benzoic oxidation against the signals of different

413

interfacial iron species reveals a strong dependence of reaction rates on the abundance of surface

414

Fe3+ monomers (Figure 6e), while correlation with other forms of Fe3+ produces substantial

415

scatter, particularly for nSiO2. The slope of the correlation lines is similar between nSiO 2 and

416

nAl2O3 and is independent of the iron precursor used. The data thus presents a strong inference

417

to the dominant role of isolated Fe3+ species in catalyzing the production of OH• from H2O2

418

decomposition. We are unable to examine iron speciation on nTiO2 because of very strong

419

absorbance by pristine TiO2 in the UV region (𝜆 < 350 nm). Nonetheless, the trend of TiO2

420

catalytic activity where surface reactivity shows signs of saturation at higher Fe loadings is in

421

line with our notion that surface activity arises predominantly from the dispersed iron species.

422

Environmental Implications

423

Interaction of aqueous iron with background minerals is an important process in generating

424

reactive surfaces in the aquatic environment. With our interest primarily focused on the activity

425

of iron-enriched surfaces in converting H2O2 to reactive oxidants, we showed that exposing clean

426

minerals to dissolved iron can generate surfaces of considerably different catalytic properties,

427

and such difference is strongly tied to the structure and quantities of surface iron species.

428

Collectively, results of iron impregnation, benzoic acid oxidation, and spectroscopic analysis

429

seem to suggest that the dispersed iron species are substantially more reactive than aggregated

430

iron on the surface. Although the reactivity of iron oxides were not evaluated here, data by Kwan

431

and Voelker using formic acid as a probe compound suggests that the surface-area-normalized

432

formic acid oxidation rates on ferrihydrite, goethite, and hematite are on the order of 1 x 10-11 -



433

10-8 mole/h-m2 at pH 4 [10]. The range sits at the lower end of the reactivity spectrum observed

434

in this study, even after correcting for the difference in the reaction rates of the two probe

435

compounds at the experiment pH, and this prompt us to underline the importance of interfacial

436

iron species instead of iron oxide particles in contributing to catalytic oxidation of contaminants

437

at the solid-water interface.

438

Among numerous environmental factors that may exert an impact on the iron deposition process,

439

the effect of the initial valence state of dissolved iron was examined here owing to the distinct

440

acid-base and coordination chemistry of Fe(II) and Fe(III). Lower catalytic activity was

441

observed with nSiO2 and nAl2O3 immersed in Fe(III) solutions, which we attributed mainly to a

442

greater tendency of ferric ions to hydrolyze and initiate surface precipitation than ferrous species.

443 444

Impregnated TiO2 emerged consistently as the most reactive surfaces for reactive oxidant

445

generation among the mineral particles examined. Furthermore, Fe(III) precursor offers

446

moderately increased activity with respect to Fe(II) at equivalent surface loading. The rutile or

447

anatase phases are known to form strong inner-sphere complexes with transition metal cations

448

over a broad range of pH [49,50]. The strong substrate-iron interactions would favor the deposition

449

of iron as a monolayer of surface complexes. Consequently, nTiO2 particles are able to

450

accommodate a large number of reactive sites per unit surface before saturation of catalytic

451

activity and the strong interaction between nTiO2 and Fe(III) prevents the formation of

452

aggregated species that are catalytically less active. However, we argue that this surface

453

templating effect alone is unable to account for their high catalytic activity, as benzoic acid

454

oxidation rates generated by nTiO2 far exceed those of alumina and silica at similar surface iron

455

loading. Strong catalyst-support interactions and coupling of charge transfer process of the


Page 20 of 36

Page 21 of 36


456

surface Fe with that of titania may have played a role in enhancing the redox cycling of surface

457

reactive sites. This is consistent with several recent studies which reported a strong chemical

458

synergy between surface deposited Fe(III) and TiO2 nanoparticles resulting in composite

459

materials with reduced band gap energies and remarkable increases in photocatalytic activity

460

[51,52,53]. Enhanced reactivity may also stem from facile interfacial charge transfer between

461

Fe(III) and the TiO2 phase [52], which is not available for the Fe(III) impregnated on non-

462

reducible mineral substrates such as alumina and silica. Although our studies were performed

463

exclusively in the dark, the prospect of nTiO2 being a favorable substrate for iron under diverse

464

conditions (dark or illuminated; with exposure to either ferrous or ferric precursors) makes it a

465

viable candidate to be incorporated in remediation or treatment systems.

466 467

Acknowledgement

468

WLY acknowledges the funding provided by the U.S. National Science Foundation (CHE-

469

1308726). LM acknowledges the financial support provided by the project LO1305 of the

470

Ministry of Education, Youth and Sports of the Czech Republic. The authors thank Dr. Kamil

471

Klier for suggestions on DRS analysis and Drs. Klara Cepe, Moira Ridley and Juliusz

472

Warzywoda for assistance in HR-TEM, BET, and XRD analysis, respectively.

473 474

Supporting Information

475

Information about the materials, methods related to solid characterization and Fe(II) extraction,

476

results of Fe(II) extraction, XRD spectra of the solids, Fe(III) speciation diagram, auxiliary

477

experiments of benzoic acid oxidation, and fitting of DR-UV-Vis data are available in the



478

Supporting Information. This material is available free of charge via the Internet at

479

http://pubs.acs.org.

480


Page 22 of 36

Page 23 of 36


481

REFERENCES

482 483 484 485 486 487 488 489 490 491 492 493 494 495 496 497 498 499 500 501 502 503 504 505 506 507 508 509 510 511 512 513 514 515 516 517 518 519

1. Borch, T.; Kretzschmar, R.; Kappler, A.; Van Cappellen, P.; Ginder-Vogel, M.; Voegelin, A.; Campbell, K., Biogeochemical Redox Processes and their Impact on Contaminant Dynamics. Environmental Science & Technology 2010, 44 (1), 15-23. 2. Wehrli, B.; Sulzberger, B.; Stumm, W., Redox processes catalyzed by hydrous oxide surfaces. Chemical Geology 1989, 78 (3-4), 167-179. 3. Taylor, K. G.; Konhauser, K. O., Iron in Earth Surface Systems: A Major Player in Chemical and Biological Processes. Elements 2011, 7 (2). 4. Muller, D. B.; Wang, T.; Duval, B.; Graedel, T. E., Exploring the engine of anthropogenic iron cycles. Proceedings of the National Academy of Sciences of the United States of America 2006, 103 (44), 16111-16116. 5. Yanina, S. V.; Rosso, K. M., Linked reactivity at mineral-water interfaces through bulk crystal conduction. Science 2008, 320 (5873), 218-222. 6. Handler, R. M.; Beard, B. L.; Johnson, C. M.; Scherer, M. M., Atom Exchange between Aqueous Fe(II) and Goethite: An Fe Isotope Tracer Study. Environmental Science & Technology 2009, 43 (4), 1102-1107. 7. Zuo, Y. G.; Hoigne, J., Photochemical decomposition of oxalic, glyoxalic, and pyruvic acid catalyzed by iron in atmospheric waters. Atmospheric Environment 1994, 28 (7), 1231-1239. 8. Weschler, C. J.; Mandich, M. L.; Graedel, T. E., Speciation, photosensitivity, and reactions of transition metal ions in atmospheric droplets. Journal of Geophysical Research-Atmospheres 1986, 91 (D4), 5189-5204. 9. Lin, S.-S.; Gurol, M. D., Catalytic decomposition of hydrogen peroxide on iron oxide: kinetics, mechanism, and implications. Environmental Science & Technology 1998, 32 (10), 1417-1423. 10. Kwan, W. P.; Voelker, B. M., Rates of hydroxyl radical generation and organic compound oxidation in mineral-catalyzed Fenton-like systems. Environmental science & technology 2003, 37 (6), 1150-1158. 11. Miller, C. M.; Valentine, R. L., Mechanistic studies of surface catalyzed H2O2 decomposition and contaminant degradation in the presence of sand. Water Research 1999, 33 (12), 2805-2816. 12. Kong, S. H.; Watts, R. J.; Choi, J. H., Treatment of petroleum-contaminated soils using iron mineral catalyzed hydrogen peroxide. Chemosphere 1998, 37 (8), 1473-1482. 13. Pignatello, J. J.; Oliveros, E.; MacKay, A., Advanced oxidation processes for organic contaminant destruction based on the Fenton reaction and related chemistry. Critical Reviews in Environmental Science and Technology 2006, 36 (1), 1-84. 14. Pham, A. L.-T.; Lee, C.; Doyle, F. M.; Sedlak, D. L., A Silica-Supported Iron Oxide Catalyst Capable of Activating Hydrogen Peroxide at Neutral pH Values. Environmental Science & Technology 2009, 43 (23).



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

15. Watts, R. J.; Foget, M. K.; Kong, S. H.; Teel, A. L., Hydrogen peroxide decomposition in model subsurface systems. Journal of Hazardous Materials 1999, 69 (2), 229-243. 16. Huling, S. G.; Arnold, R. G.; Sierka, R. A.; Miller, M. R., Measurement of hydroxyl radical activity in a soil slurry using the spin trap alpha-(4-pyridyl-1-oxide)-N-tertbutylnitrone. Environmental Science & Technology 1998, 32 (21), 3436-3441. 17. Pham, A. L. T.; Doyle, F. M.; Sedlak, D. L., Kinetics and efficiency of H2O2 activation by iron-containing minerals and aquifer materials. Water Research 2012, 46 (19), 64546462. 18. Poulton, S. W.; Raiswell, R., Chemical and physical characteristics of iron oxides in riverine and glacial meltwater sediments. Chemical Geology 2005, 218 (3-4), 203-221. 19. Wolthoorn, A.; Temminghoff, E. J. M.; Weng, L. P.; van Riemsdijk, W. H., Colloid formation in groundwater: effect of phosphate, manganese, silicate and dissolved organic matter on the dynamic heterogeneous oxidation of ferrous iron. Applied Geochemistry 2004, 19 (4), 611-622. 20. Sung, W.; Morgan, J. J., Kinetics and Product of Ferrous Iron Oxygenation in Aqueous Systems. Environmental Science & Technology 1980, 14 (5), 561-568. 21. Henry, M.; Jolivet, J. P.; Livage, J., Aqueous chemistry of metal cations: Hydrolysis, condensation and complexation. Chemistry, Spectroscopy and Applications of Sol-Gel Glasses Structure and Bonding 1992, 77, 153-206. 22. Hu, Y.; Neil, C.; Lee, B.; Jun, Y.-S., Control of Heterogeneous Fe(III) (Hydr)oxide Nucleation and Growth by Interfacial Energies and Local Saturations. Environmental Science & Technology 2013, 47 (16), 9198-9206. 23. Jun, Y.-S.; Lee, B.; Waychunas, G. A., In situ observations of nanoparticle early development kinetics at mineral− water interfaces. Environmental science & technology 2010, 44 (21), 8182-8189. 24. Balu, A. M.; Pineda, A.; Yoshida, K.; Campelo, J. M.; Gai, P. L.; Luque, R.; Romero, A. A., Fe/Al synergy in Fe2O3 nanoparticles supported on porous aluminosilicate materials: excelling activities in oxidation reactions. Chemical Communications 2010, 46 (41), 78257827. 25. Lim, H.; Lee, J.; Jin, S.; Kim, J.; Yoon, J.; Hyeon, T., Highly active heterogeneous Fenton catalyst using iron oxide nanoparticles immobilized in alumina coated mesoporous silica. Chemical Communications 2006, (4), 463-465. 26. Shi, F.; Tse, M. K.; Pohl, M.-M.; Brueckner, A.; Zhang, S.; Beller, M., Tuning catalytic activity between homogeneous and heterogeneous catalysis: Improved activity and selectivity of free nano-Fe2O3 in selective oxidations. Angewandte Chemie-International Edition 2007, 46 (46), 8866-8868. 27. Hermanek, M.; Zboril, R.; Medrik, N.; Pechousek, J.; Gregor, C., Catalytic efficiency of iron(III) oxides in decomposition of hydrogen peroxide: Competition between the surface area and crystallinity of nanoparticles. Journal of the American Chemical Society 2007, 129 (35), 10929-10936.


Page 24 of 36

Page 25 of 36

560 561 562 563 564 565 566 567 568 569 570 571 572 573 574 575 576 577 578 579 580 581 582 583 584 585 586 587 588 589 590 591 592 593 594 595 596 597 598 599


28. Kortum, G., Reflectance Spectroscopy Principles, Methods, Applications SpringerVerlag: New York, 1969. 29. Torrent, J.; Barron, V., Diffuse reflectance spectroscopy of iron oxides. In Encylopedia of surface and colloid science Marcel Dekker 2002; pp 1438-1446. 30. Klier, K., Adsorption and scattering in plane parallel turbid media. Journal of the Optical Society of America 1971, 62 (7), 882-885. 31. Tamura H.; Kawamura, S.; Hagayama, M., Acceleration of the oxidation of Fe2+ ions by Fe(III)-oxyhydroxides Corrosion Science 1980, 20, 963. 32. Amonette, J. E.; Templeton, J. C., Improvements to the quantitative assay of nonrefractory minerals for Fe(II) and total Fe using 1,10-phenanthroline. Clays and Clay Minerals 1998, 46 (1), 51-62. 33. Livi, K. J. T.; Schaffer, B.; Azzolini, D.; Seabourne, C. R.; Hardcastle, T. P.; Scott, A. J.; Hazen, R. M.; Erlebacher, J. D.; Brydson, R.; Sverjensky, D. A., Atomic-Scale Surface Roughness of Rutile and Implications for Organic Molecule Adsorption. Langmuir 2013, 29 (23), 6876-6883. 34. Koretsky, C. M.; Sverjensky, D. A.; Sahai, N., A model of surface site types on oxide and silicate minerals based on crystal chemistry: Implications for site types and densities, multi-site adsorption, surface infrared spectroscopy, and dissolution kinetics. American Journal of Science 1998, 298 (5), 349-438. 35. Sverjensky, D. A., Prediction of the speciation of alkaline earths adsorbed on mineral surfaces in salt solutions. Geochimica Et Cosmochimica Acta 2006, 70 (10), 24272453. 36. Hiemstra, T.; Yong, H.; Van Riemsdijk, W. H., Interfacial charging phenomena of aluminum (hydr)oxides. Langmuir 1999, 15 (18), 5942-5955. 37. Stumm, W.; Morgan, J. J., Aquatic Chemistry Chemical Equilibria and Rates in Natural Waters. 3rd ed.; Wiley Interscience: 1995. 38. Kosmulski, M., pH-dependent surface charging and points of zero charge: III. Update. Journal of Colloid and Interface Science 2006, 298 (2), 730-741. 39. Kosmulski, M., The pH-dependent surface charging and points of zero charge V. Update. Journal of Colloid and Interface Science 2011, 353 (1), 1-15. 40. Zhou, X. L.; Mopper, K., Determination of photochemically produced hydroxyl radicals in seawater and freshwater Marine Chemistry 1990, 30 (1-3), 71-88. 41. Lee, C.; Sedlak, D. L., Enhanced Formation of Oxidants from Bimetallic Nickel-Iron Nanoparticles in the Presence of Oxygen. Environmental Science & Technology 2008, 42 (22), 8528-8533. 42. Lee, H.; Lee, H.-J.; Sedlak, D. L.; Lee, C., pH-Dependent reactivity of oxidants formed by iron and copper-catalyzed decomposition of hydrogen peroxide. Chemosphere 2013, 92 (6), 652-658. 43. Sherman, D. M.; Waite, T. D., Electronic spectra of Fe3+ oxides and oxide hydroxide in the near IR to near UV American Mineralogist 1985, 70 (11-12), 1262-1269.



600 601 602 603 604 605 606 607 608 609 610 611 612 613 614 615 616 617 618 619 620 621 622 623 624 625 626 627 628 629 630 631 632 633 634 635

44. Perez-Ramirez, J.; Groen, J. C.; Bruckner, A.; Kumar, M. S.; Bentrup, U.; Debbagh, M. N.; Villaescusa, L. A., Evolution of isomorphously substituted iron zeolites during activation: comparison of Fe-beta and Fe-ZSM-5. Journal of Catalysis 2005, 232 (2), 318334. 45. Pirngruber, G. D.; Roy, P. K.; Prins, R., On determining the nuclearity of iron sites in Fe-ZSM-5 - a critical evaluation. Phys. Chem. Chem. Phys. 2006, 8 (34), 3939-3950. 46. Schwidder, M.; Kumar, M. S.; Klementiev, K.; Pohl, M. M.; Brückner, A.; Grünert, W., Selective reduction of NO with Fe-ZSM-5 catalysts of low Fe content: I. Relations between active site structure and catalytic performance. Journal of Catalysis 2005, 231 (2), 314-330. 47. Schwidder, M.; Kumar, M. S.; Klementiev, K.; Pohl, M. M.; Bruckner, A.; Grunert, W., Selective reduction of NO with Fe-ZSM-5 catalysts of low Fe content - I. Relations between active site structure and catalytic performance. Journal of Catalysis 2005, 231 (2), 314-330. 48. Kumar, M. S.; Schwidder, M.; Grunert, W.; Bruckner, A., On the nature of different iron sites and their catalytic role in Fe-ZSM-5 DeNO(x) catalysts: new insights by a combined EPR and UV/VIS spectroscopic approach. Journal of Catalysis 2004, 227 (2), 384-397. 49. Ridley, M. K.; Hiemstra, T.; van Riemsdijk, W. H.; Machesky, M. L., Inner-sphere complexation of cations at the rutile-water interface: A concise surface structural interpretation with the CD and MUSIC model. Geochimica Et Cosmochimica Acta 2009, 73 (7), 1841-1856. 50. Ludwig, C.; Schindler, P. W., Surface Complexation on TiO2 1. Adsorption of H+ and Cu2+ ions onto TiO2 (anatase) Journal of Colloid and Interface Science 1995, 169 (2), 284-290. 51. Tada, H.; Jin, Q.; Nishijima, H.; Yamamoto, H.; Fujishima, M.; Okuoka, S.-i.; Hattori, T.; Sumida, Y.; Kobayashi, H., Titanium(IV) Dioxide Surface-Modified with Iron Oxide as a Visible Light Photocatalyst. Angewandte Chemie-International Edition 2011, 50 (15), 3501-3505. 52. Liu, M.; Qiu, X.; Miyauchi, M.; Hashimoto, K., Energy-Level Matching of Fe(III) Ions Grafted at Surface and Doped in Bulk for Efficient Visible-Light Photocatalysts. Journal of the American Chemical Society 2013, 135 (27), 10064-10072. 53. Nolan, M., Electronic coupling in iron oxide-modified TiO2 leads to a reduced band gap and charge separation for visible light active photocatalysis. Phys. Chem. Chem. Phys. 2011, 13 (40), 18194-18199.

636


Page 26 of 36

Page 27 of 36

637

638 639 640 641


Table 1. Properties of mineral nanoparticles examined in this study. Ave. Primary Particle Diametera

Crystallographic phaseb

nSiO2

12 nm

amorphous

nAl2O3

13 nm

ε-Al2O3, δ-Al2O3

BET surface area, m2/g Pristine Feimpregnatedc 201 +/- 0.1 190 +/- 0.3 99 +/- 0.6

102 +/- 0.1

21 nm anatase, rutile 51 +/- 0.2 53.5 +/- 0.1 nTiO2 a Particle size info was provided by vendor and was verified qualitatively with TEM characterization. b Identified with X-ray diffraction analysis. c Particles immersed in 0.2 mM Fe(III) for nSiO2 and nAl2O3 and 0.8 mM Fe(III) for nTiO2 for 24 h.



642 643 644

Figure 1. TEM micrographs of as-received and impregnated (a) nSiO2, (b) nAl2 O3, and (c)

645

nTiO2. Impregnation was performed by immersing 0.6 g/L of nanoparticles in 0.2 mM ferrous

646

(pH 6.0) or ferric solutions (pH 3.0) for 24 h.

647


Page 28 of 36

Page 29 of 36


(a)

648

649 650

Figure 2. Uptake of dissolved iron by nanoparticles after immersion in (a) ferrous or (b) ferric

651

solutions for 24 h. Particle loading in all experiments was 0.6 g/L. pH of Fe(II) solutions was

652

controlled at 6.0 with 5 mM MES. Initial pH of Fe(III) solutions was adjusted to 3.0 using dilute

653

HCl and no buffering agent was used.



654


Page 30 of 36

Page 31 of 36


655

Figure 3. Catalytic oxidation of benzoic acid in the presence of impregnated silica (a, d),

656

alumina (b, e), and titania (c, f) nanoparticles. The initial concentration and valence state of iron

657

solutions were denoted in the legends. Dashed lines represent the amount of benzoic acid

658

adsorbed by particles with the highest loading of iron in the series. ([benzoic acid] 0 = 0.069 mM ;

659

[H2O2]0 = 11 mM; particle dose = 1.5 g/L for all studies except for (f) which was reduced to 0.5

660

g/L due to very rapid reactions)

661



662

663 664

Figure 4. Benzoic acid degradation rates as a function of apparent surface density of iron on

665

different mineral nanoparticles. (a) nanoparticles impregnated in Fe(II) solutions; and b)

666

nanoparticles impregnated in Fe(III) solutions. Dashed lines represent benzoic acid degradation

667

in the presence of pristine nTiO2. Insets in (b) are enlarged views of data of nSiO2 and nAl2O3.


Page 32 of 36

Page 33 of 36


668 669

Figure 5. Effect of pH on benzoic acid degradation by nTiO2 impregnated in Fe(III) solutions.

670

Experiments at pH 5.2 and 6.9 were buffered using 5 mM of MES and MOPS, respectively.

671

Other conditions follow those of the unbuffered experiments.



672

673 674


Page 34 of 36

Page 35 of 36


675

Figure 6. Diffuse Reflectance UV-Vis spectra of various iron-impregnated nSiO2 (a, b) and

676

nAl2O3 (c, d). Legends denote the initial concentration and valence state of iron impregnation

677

solutions and values in [ ] are surface iron loadings upon impregnation. (e) Correlation between

678

benzoic acid oxidation rates and the peak areas of isolated Fe3+ species.

679



Page 36 of 36


Fe-Impregnated Mineral Colloids for Peroxide Activation: Effects of

Recommend Documents