Deposition Kinetics of Colloidal Manganese Dioxide onto

Subscriber access provided by TULANE UNIVERSITY

Environmental Processes

Deposition Kinetics of Colloidal Manganese Dioxide onto Representative Surfaces in Aquatic Environments: The Role of Humic Acid and Biomacromolecules Xiaoliu Huangfu, Chengxue Ma, Ruixing Huang, Qiang He, Caihong Liu, Jian Zhou, Jin Jiang, Jun Ma, Yinying Zhu, and Muhua Huang Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.8b04274 • Publication Date (Web): 30 Nov 2018 Downloaded from http://pubs.acs.org on December 2, 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.

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 42

Environmental Science & Technology

1

Deposition Kinetics of Colloidal Manganese Dioxide onto

2

Representative Surfaces in Aquatic Environments: The Role

3

of Humic Acid and Biomacromolecules

4

Xiaoliu Huangfu†*, Chengxue Ma†, Ruixing Huang†, Qiang He†*, Caihong Liu†, Jian

5

Zhou†, Jin Jiang‡, Jun Ma‡, Yinying Zhu†, Muhua Huang†

6

† Key

7

of Education, Faculty of Urban Construction and Environmental Engineering, National

8

Centre for International Research of Low-carbon and Green Buildings, Chongqing

9

University, 400044, China

Laboratory of Eco-environments in the Three Gorges Reservoir Region, Ministry

10

‡

11

Municipal and Environmental Engineering, Harbin Institute of Technology, 150090,

12

China

State Key Laboratory of Urban Water Resource and Environment, School of

13

Colloidal MnO 2

Alginate

HA

BS A

Enhanced Deposition

Hindered Deposition

Repulsive

14

Steric Interaction DLVO Interactions

Attractive Sites

DLVO Interactions

Steric Interaction DLVO Interactions

15 16 17 18 19 20 1

ACS Paragon Plus Environment


21 22

ABSTRACT

23

The initial deposition kinetics of colloidal MnO2 on three representative surfaces in

24

aquatic systems (i.e., silica, magnetite, and alumina) in NaNO3 solution were

25

investigated in the presence of model constituents, including humic acid (HA), a

26

polysaccharide (alginate), and a protein (bovine serum albumin (BSA), using laboratory

27

quartz crystal microbalance with dissipation monitoring equipment (QCM-D). The

28

results indicated that the deposition behaviors of MnO2 colloids on three surfaces were

29

in good agreement with classical Derjaguin-Landau-Verwey-Overbeek (DLVO) theory.

30

Critical deposition concentrations (CDC) were determined to be 15.5 mM NaNO3 and

31

9.0 mM NaNO3 when colloidal MnO2 was deposited onto silica and magnetite,

32

respectively. Both HA and alginate could largely retard the deposition of MnO2 colloids

33

onto three selected surfaces due to steric repulsion, and HA was more effective in

34

decreasing the deposition rate relative to alginate. However, the presence of BSA can

35

provide more attractive deposition site and thus lead to greater deposition behavior of

36

MnO2 colloids onto surfaces. The dissipative properties of the deposited layer were also

37

influenced by surface type, electrolyte concentration, and organic matter characteristics.

38

Overall, these results provide insights into the deposition behavior of MnO2 colloids on

39

environmental surfaces and have significant implications for predicting the transport

40

potential of common MnO2 colloids in natural environments and engineered systems.

41

1. INTRODUCTION 2


Page 2 of 42

Page 3 of 42


42

The occurrence of manganese dioxide (MnO2) colloids has been widely reported in both

43

natural and engineered water environments.1-3 They are commonly formed in natural

44

waters through various processes, including mineral weathering and dissolved

45

manganese (Mn(II)) oxidation, etc.4,

46

engineered systems when permanganate is utilized as an oxidant or when conducting

47

Mn(II) oxidation during water treatment processes,6 such as methanogenic propionate

48

and butyrate degradation in anaerobic ecosystems.7 The formed MnO2 colloids may

49

impact the transport and fate of both natural and synthetic contaminants due to its high

50

surface

51

adsorption/desorption, and catalysis, with relevant contaminants in both natural and

52

engineered aquatic environments.3, 8-10 However, MnO2 colloids could also undergo

53

aggregation and deposition in aquatic environments, resulting in changes in their

54

transport behavior,11 and therefore affect the corresponding interactions and the fate of

55

MnO2 to a certain extent.12-17

activity,

which

might

5

Moreover, they are normal solid products in

lead

to

various

interactions,

e.g.,

redox,

56

Colloidal MnO2 aggregation kinetics have been previously investigated by dynamic

57

light scattering (DLS) technology, and the influence of humic substance (HS),

58

biomacromolecules, and redox processes on the aggregation kinetics of MnO2 colloids

59

has been reported as well.11, 18, 19 The deposition process of MnO2 colloids onto surfaces

60

is also critical for understanding their mobility and persistence and the associated

61

contaminants in both natural and engineered environments. It has been reported that the

62

addition of MnO2 nanoparticles can efficiently enhance the removal of trace thallium

63

(Tl) in quartz sand filtration systems by adsorption of Tl onto nanosized MnO2 and 3



64

subsequent deposition of MnO2 particles onto quartz sand surfaces.20 In the work

65

published by Ya Cheng and co-authors, NH3 oxidation in a manganese oxide coated

66

quartz sand filter had been reported for water treatment.21 However, the fundamental

67

data on the deposition kinetics of colloidal MnO2 on environmentally relevant surfaces

68

are not available to date.

69

Deposition kinetics have been widely investigated by employing different

70

approaches (e.g., theoretical methods, packed column experiments, and quartz crystal

71

microbalance with dissipation monitoring (QCM-D) technology) for various

72

nanomaterials in synthetic and/or natural waters, such as silver nanoparticles (AgNPs),

73

fullerene (C60), quantum dots (QDs), carbon nanotubes, graphene oxide (GO) NPs,

74

TiO2 NPs and CeO2 NPs.22-30 Classical Derjaguin-Landau-Verwey-Overbeek (DLVO)

75

and extended DLVO (EDLVO) theory have been considered as the basic theories for

76

interpreting experimental data of deposition and proposing the deposition mechanisms

77

of NPs in environmental conditions.24-26, 28 According to classical DLVO theory, the

78

deposition of colloids on surfaces is strongly dependent upon their energy barrier,

79

which is affected by colloidal properties and water chemistry (e.g., pH, ionic

80

strength).24-28, 31 The characteristics of surfaces in aquatic systems (e.g., the ubiquitous

81

minerals of silicon oxide (SiO2); iron oxide (Fe3O4) and aluminum oxide (Al2O3) etc.)

82

is also one of the critical conditions could impact the fate and transport of

83

nanoparticles.24,28

84

Moreover, macromolecular organic matter widely distributed in natural aquatic

85

environments could affect the properties of both colloids and the surface to a large 4


Page 4 of 42

Page 5 of 42


86

extent in aquatic systems; thus, more complicated interactions between colloids and

87

surfaces should be considered.25, 26, 32 The extent to which these macromolecules could

88

impact the deposition of colloids highly depends on the combined effect of the

89

electrostatic, steric, and bridging interactions induced by their adsorption on both

90

collector and nanoparticle surfaces.28,

91

proteins are the most important components of natural organic molecules in surface

92

waters.34 It is widely reported that humic acid (HA) is able to enhance the stability and

93

mobility of NPs (e.g., fullerene, ZnO NPs, and QDs) by inducing electrostatic and/or

94

steric repulsive energies.33, 35, 36 Similar enhancement has been observed in the case of

95

alginate (a representative polysaccharide commonly found in natural waters) by

96

increasing steric repulsion.28 Bovine Serum Albumin (BSA) is an important model

97

protein commonly employed for examining the influence of protein on environmental

98

behaviors of nanoparticles due to its high structural stability. 11, 37, 38 It is hypothesized

99

that the impacts of proteins on the mobility of nanoparticles are different from humic

100

sustances,37 however, limited experimental works have been conducted on the colloid

101

deposition kinetics under the influence of BSA in aqueous dispersions.

33

Humic substances, polysaccharides, and

102

In the work described herein, the first experimental data were obtained for

103

estimating the initial deposition kinetics of colloidal MnO2 in dilute NaNO3 solutions

104

by employing QCM-D. Since metal oxide surfaces existing as coating patches on

105

natural sands can significantly influence the colloid transport in the environment, 39 The

106

sensors coated by silicon oxide (SiO2); iron oxide(Fe3O4) and aluminum oxide (Al2O3)

107

were applied as representative surfaces to quantify the effect of the mineral components 5



108

on the deposition kinetics of MnO2 colloids. The effects of model constituents of humic

109

acid (HA) and biomacromolecules (i.e., alginate and bovine serum albumin (BSA)) on

110

colloidal MnO2 deposition kinetics for these representative surfaces were also

111

elucidated and discussed. Moreover, a combination of DLVO and EDLVO calculations

112

was performed for a better understanding of the mechanisms controlling the deposition

113

behavior of MnO2 colloids. Finally, the dissipative property of the deposited layer was

114

explored as well.

115

2. MATERIALS AND METHODS

116

2.1. Materials

117

Chemicals, including KMnO4, Na2S2O3, NaOH, HNO3, and NaNO3, were obtained

118

from Sinopharm Chemical Reagent Co., Ltd. Poly-L-lysine hydrobromide (PLL,

119

molecular weight 70000-150000 by viscosity, P-1247), sodium alginate (no. 180947),

120

and BSA (no. V900933) were purchased from Sigma Aldrich, St. Louis, MO. These

121

chemicals were all used as received. HA (Fluka no. 53680) was purchased from Sigma

122

Aldrich, Inc. (Milwaukee, WI), for which the characteristics were reported previously.40,

123

41

124

mixing in double deionized (DDI) water (>18.2 MΩ/cm), being filtrated by a cellulose

125

acetate filter (Whatman ME 24, 0.2 μm) and stored at 4°C. The total organic carbon

126

(TOC) concentrations were determined by the oxidation method under high

127

temperature conditions (Model Multi3100, Jana, Germany).

128

2.2 Preparation and Characterization of MnO2 Colloidal Suspensions

The stock solutions (500 mg/L, pH 6.0) of alginate and BSA were prepared by first

6


Page 6 of 42

Page 7 of 42


129

MnO2 colloids were synthesized by the method utilized in our previously published

130

papers.10, 11 Briefly speaking, a stoichiometric amount of a Na2S2O3 solution was added

131

dropwise into a rapidly stirred KMnO4 solution with a magnetic stirrer, and the N2 was

132

purged to maintain an anaerobic environment during the synthesis of MnO2 colloids.

133

Dark brown MnO2 colloids were finally formed. The colloidal stock suspension was

134

continuously stirred overnight and stored in the dark at 4°C prior to the measurement

135

and deposition experiment. The concentrations of MnO2 colloidal suspensions were

136

determined by ICP-MS (NexION 300Q, PerkinElmer Cop, IS). The average particle

137

size of MnO2 colloids was approximately 59.1 ± 0.3 nm (n = 30) from dynamic light

138

scattering (DLS) measurements using a Nano ZetaSizer (Nano ZS90, Malvern, UK).

139

The diameters of freshly prepared MnO2 colloids determined by TEM were shown in

140

Figure S1. The average oxidation state of Mn in MnO2 nanoparticles was determined

141

by an iodimetric method (4.03 ± 0.04 (n=5)), accorded with the measurement of XPS

142

spectra (PHI 5700 ESCA, US). The absolute zeta potential (ζ potential) of MnO2

143

colloids was negative over the range of 1-20 mM NaNO3 at pH 6.0 (Zetasizer Nano

144

ZS90, Malvern, UK). Detailed colloidal properties of synthesized MnO2 colloids were

145

elaborated in our previous publication.11

146

2.3. Deposition Study Employing a Quartz Crystal Microbalance with a

147

Dissipation Monitoring System

148

Colloidal MnO2 deposition on selected surfaces was investigated by utilizing a QCM-

149

D instrument (E4, Q-sense, Biolin Scientific, Sweden), which could simultaneously

150

monitor the shifts in frequency (Δf) and energy dissipation (ΔD). Crystal sensors coated 7



151

with silica (SiO2, QSX303), magnetite (Fe3O4, QSX326) or alumina (Al2O3, QSX306)

152

surfaces were employed in the deposition experiments. Prior to each determination,

153

crystal sensors were cleaned following protocols modified based on those

154

recommended by Q-Sense: i.e., UV/ozone pretreatment, followed by rinsing with DDI

155

water and a drying process with ultrapure nitrogen. The detailed procedures of crystal

156

cleaning are provided in the Supporting Information (Text S1). The flow of all

157

suspensions was kept constant at a rate of 0.15 mL/min in the module, and the

158

temperature was maintained at 25 ± 0.2°C throughout the experiment. Before starting

159

the deposition experiments, the sensors were first equilibrated with DDI water to reach

160

initial stabilization by monitoring frequency (Δf(3)) and dissipation (ΔD(3)) signal

161

overtones. Stabilization of the system was achieved when the frequency (Δf(3)) shift was

162

not more than 0.3 Hz in a period of 10 min.33 Then, a particle-free electrolyte solution

163

was used to rinse the crystal surfaces until stabilized. The deposition experiment was

164

started once a stable baseline was observed. The suspension containing 1 mM (0.087

165

mg/L) MnO2 colloids at desired background electrolyte was injected into the crystal

166

chamber. For the experiments in the presence of macromolecular organic matter, the

167

collector surface was rinsed with electrolyte solution of interest, followed by

168

macromolecules solution with the same ionic strength for ~30 min to establish the

169

baseline. The mixture of a premeasured volume of diluted MnO2 colloidal suspension

170

and macromolecules at the TOC of 5.0 mg/L was vortexed in the desired electrolyte

171

solution for 10 s, and was then introduced into the chamber and monitored over a time

172

period of 20-60 min until a sufficient frequency shift was obtained. 8


Page 8 of 42

Page 9 of 42


173

With deposition of MnO2 colloids on surfaces, the mass of the crystal sensor leads to

174

a negative shift in the overtone frequencies (Δfn). The direct relationship between the

175

deposition mass (Δm) and the shift in frequencies could be described by the Sauerbrey

176

equation:

177

Δ𝑚 = ―

𝐶Δ𝑓𝑛

(1)

𝑛

178

where n is the overtone number, n=1, 3, 5, 7…., and C is the crystal sensitivity constant,

179

17.7 ng/(Hz•cm2). The deposition of colloids on surfaces can also result in an increase

180

in the crystal’s dissipation unit (D):

181

𝐸𝑑𝑖𝑠𝑠𝑖𝑝𝑎𝑡𝑖𝑜𝑛

𝐷=

(2)

2𝜋𝐸𝑠𝑡𝑜𝑟𝑒𝑑

182

where Edissipation is the dissipated energy in one oscillation cycle, and Estored is the stored

183

energy in the oscillator.

184

Since the deposited mass of MnO2 colloids onto collector surfaces is linearly related

185

to the frequency change, the deposition rates can be indicated by the rates of frequency

186

shift. 26, 27, 29 Normalized frequency shifts monitored by QCM-D at the third overtone

187

during early stage (~1-3 min) were used to calculate the initial deposition rates (i.e., rf)

188

of MnO2 colloids:

189

𝑟𝑓 =

|( ) | 𝑑𝛥𝑓(3) 𝑑𝑡

(3)

𝑡→0

190

The deposition kinetics at each electrolyte concentration was quantified by the

191

attachment efficiency αD, and could be calculated from rf :42

192

193

𝑟𝑓

𝛼𝐷 = (𝑟𝑓)

𝑓𝑎𝑣

=

|( |(

𝑑𝛥𝑓(3) 𝑑𝑡

𝑑𝛥𝑓(3) 𝑑𝑡

)

)

𝑡→0

| |

(4)

𝑓𝑎𝑣, 𝑡→0

In eq. 4, the numerator is the rate of shift in the normalized f(3) at the tested electrolyte 9



194

concentration, while the denominator represents the corresponding deposition rate

195

under favorable conditions obtained in the same electrolyte solutions.

196

For deposition experiments under favorable conditions, cationic Poly-L-lysine

197

hydrobromide (PLL) was used to modify the silica, magnetite and alumina surfaces.

198

Details on the protocols of surfaces modification with PLL and relevant discussion of

199

the favorable deposition kinetics of nanoparticles can be found in the Supporting

200

Information (Text S2, Figures S2-S5).

201

2.4. Calculation of Interaction Energies between Colloidal MnO2 and Surfaces

202

DLVO and EDLVO theories were employed to better interpret the mechanisms

203

determining the deposition behavior of MnO2 colloids. The interaction energies for

204

MnO2 colloids approaching tested surfaces were calculated in the absence and presence

205

of macromolecules. According to classical DLVO theory, the total interaction energy

206

equals the sum of the van der Waals (VDW) energy and the electrical double layer

207

(EDL) energy, as reported in previous works.29, 43 The total interaction energy for the

208

EDLVO model was modified by the incorporation of a steric repulsion, in addition to

209

the VDW and EDL energies. Details of the calculation are provided in the Supporting

210

Information (Text S4).

211

3. RESULTS AND DISCUSSIONS

212

3.1. Deposition of Colloidal MnO2 on Environmental Surfaces in the Absence of

213

HA and Biomacromolecules

214

Representative normalized frequency shifts at the third overtone when colloidal MnO2

215

was deposited on tested surfaces (i.e., SiO2, Fe3O4 and Al2O3) were presented in Figure 10


Page 10 of 42

Page 11 of 42


216

S6, S7, and S8. In the initial stages of the experiment, the crystal sensor was first rinsed

217

with DDI water and then a predetermined volume of electrolyte solution to obtain a

218

stable Δf(n) response. When deposition took place, significant decreases in Δf at all

219

overtones were observed. The profiles of frequency shifts at the third overtone shown

220

in Figure S6 indicated that efficient deposition of MnO2 colloids took place in the

221

presence of 9, 12, 13, 15, 17, 18, and 20 mM NaNO3 on the SiO2 surface, while no

222

change and only a negligible shift in frequency were observed at 1 mM NaNO3 and 5

223

mM NaNO3, respectively, suggesting that MnO2 colloids might not be deposited on the

224

SiO2 surface under these low electrolyte conditions. The values of 𝑟𝑓 of MnO2 colloids

225

onto silica were presented as a function of the NaNO3 concentration in Figure 1a. When

226

the NaNO3 concentration increased from 7 mM to 15 mM, 𝑟𝑓 dramatically increased

227

from 0.5 Hz/min to 42.8 Hz/min. Similar deposition behavior had also been previously

228

reported for C60 NPs deposition onto a silica surface in the presence of NaCl.33 The

229

observed increase in the deposition rate might be attributed to the progressive

230

compression of the electrical double-layer resulting from more effective charge

231

screening under a higher Na+ concentration. Nevertheless, the further increase in the

232

NaNO3 concentration from 15 mM to 20 mM resulted in a notable decrease of 𝑟𝑓 from

233

42.8 Hz/min to 0.8 Hz/min. This was consistent with the time-resolved DLS

234

measurements conducted at the same electrolyte conditions (Table S1), showing that

235

the initial aggregation rate of MnO2 colloids dramatically increased as the Na+

236

concentration increased from 15 mM to 20 mM due to the decrease of electrostatic

237

energy barrier between the negatively charged colloids. The formation of large 11



238

aggregates of MnO2 colloids under this higher Na+ condition decreased their diffusion

239

coefficient to silica and therefore reduced the absolute deposition rate, despite the fact

240

that the changes in water chemistry were beneficial to deposition. A decrease in 𝑟𝑓 has

241

also been reported under higher concentrations of the electrolyte for various NPs, and

242

similar mechanisms have been proposed for the corresponding deposition behavior, e.g.,

243

GO, ZnO, iron oxide etc. 24, 44, 45

244

In the case of the magnetite (i.e., Fe3O4) surface, the deposition rate presented in

245

Figure 1b described a similar deposition behavior of MnO2 as that for the silica surface.

246

The lack of deposition in 1 mM NaNO3 (Figure S7) and an increase in deposition rates

247

with an increasing Na+ concentration at the low electrolyte level were observed. The

248

value of 𝑟𝑓 reached the maximum of 185.1 Hz/min at 10 mM NaNO3 and significantly

249

decreased when IS exceeded 15 mM. It should be noted that higher values of 𝑟𝑓

250

compared to those on silica were obtained at the same electrolyte concentration (Figure

251

1b vs. Figure 1a), indicating that the magnetite surface has a higher affinity than the

252

silica surface for the deposition of MnO2 colloids. The surface charges for silica and

253

magnetite surfaces were both negative for the pH in the present study, i.e., pH 6.0 (Table

254

S2);46 thus, the greater deposition might be owing to the weaker electrostatic repulsion

255

between the colloids and less negative Fe3O4 surface. This finding was consistent with

256

a previous report of the higher attachment of AgNPs onto hematite than silica.47

12


Page 12 of 42

Page 13 of 42


Deposition Rate (Hz/min)

103

SiO2 102

101

100

(a) 10-1

101

NaNO3 Concentration (mM) Deposition Rate (Hz/min)

103

102

101

100

10-1


103

(b) 10

1

NaNO3 Concentration (mM)

Al2O3 102

101

100

(c) 10-1

257

Fe3O4

101


258

Figure 1. Deposition rate of colloidal MnO2 on surfaces: (a) SiO2, (b) Fe3O4, and (c)

259

Al2O3 as a function of the NaNO3 concentration at pH 6.0.

260

For deposition on alumina surface, the frequency profiles presented in Figure S8

261

shows that a noticeable shift in f(3) can be observed in the presence of 1 mM NaNO3,

262

which was contrary to the observation of insignificant deposition behavior on silica or

263

magnetite under this condition. This could be attributed to the fact that the Al2O3 surface 13



Page 14 of 42

264

was positively charged at the tested conditions in the present study and could provide

265

favorable electrostatic conditions for MnO2 deposition, while the SiO2 and Fe3O4

266

surfaces were negative and deemed repulsive (Table S2). The deposition rates of MnO2

267

onto the alumina surface are shown in Figure 1c. The values of 𝑟𝑓 remained relatively

268

constant in low electrolyte solutions (≤10 mM), which was similar to the deposition

269

behavior observed on the positive PLL surface (Figure S5), suggesting that the

270

deposition of colloids onto alumina was favorable. As the electrolyte concentration

271

increased from 10 mM to 15mM, 𝑟𝑓 slightly decreased, as similarly observed for

272

MWNTs and TiO2 nanoparticles under favorable conditions.29,

273

decrease in 𝑟𝑓 can be attributed to the attenuated electrostatic attraction between the

274

colloid and surface due to more effective counterion screening.29 This behavior was in

275

agreement with the classic DLVO theory where attractive electrostatic forces

276

predominated (see below).48, 49 The further decrease in 𝑟𝑓 at higher Na+ concentrations

277

(>15 mM) was consistent with those of MnO2 deposition on silica and magnetite

278

surfaces (Figure 1a,1b), which can be explained by the aggregation of MnO2 colloids

279

resulting in diffusion limited transport to alumina surface.50

280

3.2

281

Biomacromolecules

282

To normalize the colloidal deposition rates in Figure 1 by the favorable deposition rates

283

under the model PLL-coated surfaces (Figure S5, Text S2), the attachment efficiencies,

284

αD, over the range of Na+ concentrations were derived for all three tested surfaces

285

(Figure 2).

Deposition

Attachment

Efficiency

in

the

14


Absence

48

The observed

of

HA

and

Page 15 of 42


286

In Figure 2a, the deposition kinetics of MnO2 colloids onto silica could be recognized

287

as DLVO-type behaviors.51 The αD values for deposited MnO2 colloids increased

288

significantly from 0.003 to 1 when the NaNO3 concentration increased from 7 mM to

289

15 mM, then a plateau was achieved after a further increase of the NaNO3 concentration.

290

Thus, the critical deposition concentration (CDC), i.e., the minimum electrolyte

291

concentration that allows fast deposition to take place, for MnO2 colloidal deposition

292

onto silica was approximately 15.5 mM NaNO3. It has been reported previously that

293

two regimes of deposition kinetics were observed for the deposition of nanoparticles

294

onto the silica surface (i.e., C60, MWCNs), and CDC on silica were obtained as 32.1

295

mM and 39.3 mM NaCl.26, 52 For the magnetite surface, a similar deposition behavior

296

as that on the silica surface was observed (Figure 2b). The deposition kinetics for MnO2

297

colloid deposition have also been observed as two regimes: slow deposition (i.e., from

298

5 mM to 9 mM NaNO3) and fast deposition (i.e., >10 mM NaNO3). Thus, the CDC for

299

the deposition of colloidal MnO2 onto magnetite was obtained as 9 mM NaNO3. Profiles

300

in Figure 2c also revealed that the αD values for MnO2 colloid deposition onto alumina

301

were near 1.0 and independent of the NaNO3 concentration, indicative of the fast

302

deposition of MnO2 colloids on alumina.

15



Attachment Efficiency, D

101

SiO2

100

10-1

10-2

(a) 10-3

101

NaNO3 Concentration (mM) Attachment Efficiency, D

101

Fe3O4 100

10-1

10-2

(b) 10-3

101

NaNO3 Concentration (mM) Attachment Efficiency, D

101

Al2O3 100

10-1

10-2

(c) 10-3

303

101


304

Figure 2. Attachment efficiencies of colloidal MnO2 onto surfaces (i.e., (a)the SiO2

305

surface, (b) the Fe3O4 surface, and (c) the Al2O3 surface) as a function of the NaNO3

306

concentration at pH 6.0. The respective CDC values were determined from the

307

intersections of the extrapolations (dashed lines) of two deposition regimes of 15.5 mM

308

NaNO3 for the SiO2 surface and 9.0 mM NaNO3 for the Fe3O4 surface.

309

To further understand the mechanisms controlling the observed MnO2 colloidal

310

deposition behavior onto silica, magnetite, and alumina surfaces, the DLVO interaction 16


Page 16 of 42

Page 17 of 42


311

energy profiles at representative Na+ concentrations were calculated and are presented

312

in Figure S10, S11 and S12. As can be seen, the electrolyte concentration played a

313

critical role in total colloidal interaction energies, thus controlling MnO2 colloidal

314

deposition behavior. Significantly high repulsive energies were observed when a MnO2

315

nanoparticle approached the silica surface at both 1 and 5 mM NaNO3, in agreement

316

with the fact that no or little deposition had been observed on silica under these

317

conditions where strong repulsive electrostatic interactions existed. Further increases

318

of the electrolyte concentration lead to lower repulsive energy, and MnO2 colloidal

319

deposition was thus observed. Similarly, MnO2 colloids could not be deposited on the

320

Fe3O4 surface in the presence of 1 mM NaNO3, while deposition was observed at 5 mM

321

NaNO3. This result might be attributed to the lower repulsive energy between particles

322

and the Fe3O4 surface (30 kT at 1 mM NaNO3). Higher

323

concentrations of NaNO3 also resulted in an increase in αD due to a further decrease in

324

repulsive energies when approaching magnetite surfaces. The overall lower repulsive

325

energy barriers for magnetite relative to silica at the same electrolyte concentration

326

explained the higher affinity for the magnetite surface observed before.

327

In contrast, the interaction energy profiles in Figure S12 suggested that no energy

328

barrier was present in the case of the Al2O3 surface, validating the completely favorable

329

conditions for deposition of negatively charged MnO2 colloids onto positively charged

330

Al2O3 surface (Figure 2c). However, the profiles showed that the separation distances

331

of the occurrence of attractive energy well, namely, the separation distance where MnO2

332

colloids experienced attractive force, were larger at low electrolyte strength (i.e., 1 mM 17



333

and 5 mM) relative to those at high electrolyte strength (>10 mM). Previously,

334

researchers have proposed that the decreased separation distance of the energy well

335

with the increasing ionic strength could lead to decreased deposition rate of

336

nanoparticles,29,

337

deposition when NaNO3 concentrations were higher than 10 mM (Figure 1c).

338

3.3 Deposition of Colloidal MnO2 in the Presence of HA and Biomacromolecules

44

which was consistent with the observation of the attenuated

339

The effects of HA and biomacromolecules on the deposition rate of MnO2 colloids

340

onto tested surfaces as a function of Na+ were examined and are presented in Figure 3.

341

Generally, the overall trend for 𝑟𝑓 shifts suggested that the presence of HA and alginate

342

significantly retarded deposition in the same NaNO3 concentrations on all examined

343

surfaces, indicating that HA and alginate could dramatically strengthen the mobility of

344

MnO2 colloids in aquatic environments. The slower deposition rate of MnO2 onto

345

collector surfaces in the presence of HA and alginate might result from the electrosteric

346

repulsion that originated from the adsorption on MnO2 colloidal surfaces, as proposed

347

for the deposition of other NPs.26, 32, 33, 48 It also should be noticed that the deposition

348

rate of MnO2 colloids in the presence of alginate were higher than those in the presence

349

of HA on all selected surfaces (Figure 3).The higher values of 𝑟𝑓 were possibly due to

350

the rougher surface induced by the extended conformation of the larger alginate

351

macromolecules compared to HA.50,

352

deposition for alginate could be interpreted by the EDLVO theory, as discussed in the

353

subsequent subsection.

53

Other evidence supporting this greater

18


Page 18 of 42

Page 19 of 42


NaNO3 NaNO3+HA


103

NaNO3+Alginate

SiO2

NaNO3+BSA 102

101

100

10-1

(a) 1

10



103

Fe3O4

102

101

100

10-1

(b) 1

10



103

Al2O3

102

101

100

10-1

(c) 1

10


354 355

Figure 3. Deposition rate of colloidal MnO2 onto surfaces (i.e., (a) the SiO2 surface, (b)

356

the Fe3O4 surface, and (c) the Al2O3 surface) as a function of the NaNO3 concentration

357

in the presence of HA, alginate and BSA at pH 6.0. The concentration of HA, alginate

358

and BSA was maintained at 5 mg/L of TOC. The error range shows the standard

359

deviation. 19



Page 20 of 42

360

However, the 𝑟𝑓 values of MnO2 colloids in the presence of BSA were unexpectedly

361

higher than those in the presence of HA and alginate or with no background

362

macromolecules. The observed greater deposition was independent of the collector

363

surface type (i.e., negative silica, magnetite and positive alumina in present study),

364

revealing a more effective enhancement of colloidal MnO2 mobility by BSA when

365

deposited onto environmental surfaces. This finding was consistent with the previous

366

observation of enhanced colloid deposition by absorbed BSA.54, 55, 56 To deepen the

367

understanding of this deposition behavior, the representative values of 𝑟𝑓 for BSA,

368

MnO2 colloids, and MnO2 colloids in the presence of BSA deposited onto the three

369

surfaces at different electrolyte concentrations are presented in Table S4. The values of

370

𝑟𝑓 for BSA were obtained, indicating that BSA could be absorbed onto the three

371

surfaces, consistent with results reported before.57, 58, 59, 60 However, the simultaneous

372

deposition of BSA was not the origin of the greater depositional behavior of MnO2

373

colloids, supported by the higher value of |rf, MnO2 colloids+BSA| than the value of |rf,BSA|+|rf,

374

MnO2 colloids|

375

colloidal deposition in the presence of BSA to the effects of: (1) the attractive

376

interactions generated from - the adsorption of BSA on hydrophilic collector surfaces,57,

377

61

378

aggregation due to the high steric repulsion induced by BSA molecules.11,

379

Jeyachandran et.al showed that the relatively hydrophilic BSA can strongly adsorbed

380

on hydrophilic collector surfaces (i.e., silica, magnetite and alumina) used in present

381

study.63, 64 The absorption of BSA on the surfaces approaching/reaching saturation can

when IS > 10 mM. We attributed this observed enhancement of MnO2

and (2) more effective colloid in the deposition system originated from less

20


62

Page 21 of 42


382

generate attractive deposition sites for colloids,56 and therefore increased the deposition

383

rate. Previous publications have proposed that the specific conformation of absorbed

384

BSA on particulate colloids may alter the nature and magnitude of surface interaction

385

forces exerted by a BSA molecule54,

386

particulate colloid deposition.56 In addition, BSA proteins have the strongest effect with

387

respect to retarding the MnO2 aggregation rate in the Na+ solutions relative to HA and

388

alginate due to the steric repulsion (Table S1). The strong steric repulsive forces

389

imparted by the adsorbed BSA layers can reduce the impact of limited diffusion

390

resulting from the formation of aggregates on decreasing deposition more efficiently,

391

and this mechanism also explained that the dramatic decrease in MnO2 colloidal

392

deposition rate was not observed until Na+ concentration was higher than 20 mM in the

393

presence of BSA (Figure 3, Table S4).

394

3.4

395

Biomacromolecules

396

MnO2 colloidal deposition kinetics in the presence of humic acid and

397

biomacromolecules on tested surfaces in terms of attachment efficiency (αD) are

398

presented in Figure 4. When the nanoparticles were deposited on silica and magnetite

399

surfaces, the αD values increased with increasing ionic strength in the presence of

400

macromolecules, which was similar to the deposition kinetics of MnO2 colloids with no

401

background macromolecules (Figure 4a, b). However, unlike the classical colloidal

402

deposition behaviors of MnO2 with two distinct deposition regimes (fast and slow)

403

observed in the absence of HA and biomacromolecules, the deposition kinetics of MnO2

Deposition

Attachment

65

and also exposed its attractive regions for

Efficiency

in

the

21


Presence

of

HA

and


404

in the presence of these macromolecules did not reach the fast deposition regime (αD=1)

405

in the examined electrolyte concentrations. This demonstrated that classic DLVO forces

406

were not the predominant mechanisms for the deposition of MnO2 colloids onto silica

407

and magnetite surfaces in the presence of background macromolecules.

408

To further elucidate the mechanisms for the deposition of MnO2 colloids in the

409

presence of macromolecules, the EDLVO interaction energy profiles incorporated by a

410

quantitative steric repulsive energy for the interacting MnO2 colloids with silica and

411

magnetite surfaces in the presence of macromolecules are presented in Figure S13 and

412

S14. As observed, the increase in electrolytes lead to a lower repulsive energy barrier,

413

consistent with the increased αD with increasing electrolyte concentrations. The

414

calculations of EDLVO energy also showed that the repulsive energy barriers for MnO2

415

colloid deposition onto negative silica and magnetite were higher at all conditions in

416

the presence of HA and alginate compared to the energy profiles in the absence of the

417

two macromolecules (Figure S13a, b, S14a, b vs. Figure S10, S11), implying more

418

unfavorable conditions for nMnO2 deposition, which was in accordance with the

419

observed decrease in αD (Figure 4a, b). Furthermore, the higher interaction energy

420

barriers for HA (120 kT-240 kT and 35 kT-80 kT for silica and magnetite, respectively)

421

relative to alginate (50 kT-200 kT and 10 kT-50 kT for silica and magnetite,

422

respectively) in all water chemistry conditions suggested a more repulsive steric

423

interaction originated from the HA layer than from alginate, and this could explain the

424

previously observed deposition rates for MnO2 (i.e., 𝑟𝑓Alginate > 𝑟𝑓HA; Figure 3a, b) onto

425

silica and magnetite surfaces. 22


Page 22 of 42

Page 23 of 42


426

In the presence of BSA, the EDLVO calculations (Figure S13c, 14c) for the

427

interaction energy of nMnO2 on silica and magnetite surfaces revealed that addition of

428

BSA molecules made the energy barrier of MnO2 colloid to surface decreased to 10 kT

429

-125 kT and 1 kT -26 kT for silica and magnetite, respectively. The theoretic energy

430

profiles calculated from the EDLVO model suggested a more favorable condition for

431

deposition of MnO2 in the presence of BSA, agreeing with the enhanced attachment of

432

MnO2 colloids observed before (Figure 4a, b). BSA is an amphiphilic molecule with an

433

isoelectric point of approximately pH 4.7,66 thus BSA is negatively charged under the

434

experimental condition (pH 6.0). However, there are large amount of positively charged

435

Lys residues on its surface,67, 68 and these positively charged domain of BSA could bind

436

with negatively charged surfaces (i.e., silica and magnetite) via electrostatic attraction,

437

as discussed before. Besides, the EPM experiments also indicated that the absorption

438

of BSA can reduce the absolute EPM value of MnO2 (Table S5), thus decreased the

439

electrostatic repulsion between negatively colloids and surfaces and leading to greater

440

deposition. The relative higher deposition attachment efficiencies of engineered MnO2

441

colloids onto surfaces in the presence of BSA compared to other two macromolecules

442

indicated that the transport and fate of these nanoparticles may be more greatly

443

governed by protein in the environment.

444

23



NaNO3 NaNO3+HA NaNO3+Alginate

Attachment Efficiency

101

NaNO3+BSA

SiO2

100

10-1

10-2

10-3

101

(a) 10



Fe3O4 100

10-1

10-2

10-3

(b) 10

1

NaNO3 Concentration (mM) 101


Al2O3 100

10-1

10-2

10-3

(c) 10

1


445 446

Figure 4. Attachment efficiencies (𝛼𝐷) of colloidal MnO2 onto surfaces (i.e., (a) the

447

SiO2 surface, (b) the Fe3O4 surface, and (c) the Al2O3 surface) as a function of the

448

NaNO3 concentration in the presence of alginate and BSA at pH 6.0. The concentration

449

of humic acid, alginate and BSA was maintained at 5 mg/L of TOC. The error range

450

shows the standard deviation.

451

For alumina, the overall deposition behavior in the presence of HA and two 24


Page 24 of 42

Page 25 of 42


452

biomacromolecules was similar to the trend observed for silica and magnetite surfaces

453

(Figure 4c). However, given that silica and magnetite surfaces were negative and

454

unfavorable for deposition of MnO2 colloids while the alumina surface was positive

455

and exposed attractive interactions for deposition in our study, this observation was not

456

in agreement with the EDLVO prediction for the alumina surface presented in Figure

457

S15. The occurrence of force barriers for HA and BSA and the decreased separation

458

distances for alginate with increasing electrolyte strength expected a reduction of αD as

459

the Na+ concentration increased, while increased values of αD in the presence of

460

macromolecules on alumina were observed. The absolute zeta potential of HA and the

461

two biomacromolecules were reported to be negative at pH 6.0 in the examined

462

electrolyte range.69 Antonius et al.70 found that the adsorption of negatively dissolved

463

organic matter on positively charged surfaces was irreversible and rigid, and their

464

adsorption imparted more negative charges and even reversed the surface charge from

465

positive to negative.34, 64, 71 Hence, the adsorption of these negative macromolecules

466

onto the oppositely charged alumina surface may lead to the charge reversal of the

467

positive alumina surface, and similar deposition behavior of MnO2 colloids onto

468

alumina as that for silica and magnetite surfaces was therefore observed.

469

3.5 Understanding the Dissipative Properties of Deposited Layers Using

470

|∆D(3)/∆f(3)| Values

471

In addition to the frequency shift, the deposition of colloidal mass onto crystal

472

sensors could also lead to energy dissipation,29 which has been employed for the

473

deposition of colloids,26, 72-74 bacteria,75 and viruses,76 as well as for the MnO2 colloids 25



474

described herein. Moreover, ΔD combined with Δf provides insights not only on

475

information the coupled mass, but also on the changes in dissipative properties and the

476

layer structure of deposited colloids. 77 A lower value of |ΔD/Δf| is an indication of the

477

rigidity of deposited layers, whereas an elevated slope suggests the formation of a

478

dissipative layer.78 The changes in the viscoelastic properties of a deposited layer at a

479

solid liquid interface are of considerable interest for their use as controllable surfaces.79

480

The |∆D(3)/∆f(3)| values of MnO2 colloidal deposition are shown as function of NaNO3

481

concentration in Figure 5. Generally, the values of |∆D(3)/∆f(3)| increased with the

482

electrolyte concentration, indicating the formation of a dissipative deposited layer at

483

higher electrolyte concentrations. The aggregation of MnO2 colloids might be the cause

484

for this reduction in the rigidity of the deposited layer. In the presence of a low

485

electrolyte concentration, MnO2 colloids (or macromolecules coated colloids) existed

486

in aqueous solution individually and thus were deposited onto surfaces individually. In

487

the case of high electrolyte concentration levels, MnO2 colloids aggregated into large

488

clusters, which partially associated with collectors. Therefore, the deposited layer

489

became more loosely attached to collector surface, enhancing the crystal’s ability to

490

dissipate, as a result of frictional losses in the deposited layer.

26


Page 26 of 42

Page 27 of 42


SiO2-NaNO3 SiO2-NaNO3+Alginate

|D(3)/f(3)| (10-6/Hz)

1.5

(a)

SiO2-NaNO3+BSA SiO2-NaNO3+HA SiO2-PLL-NaNO3

1.0

0.5

0.0 Fe3O4-NaNO3 Fe3O4-NaNO3+Alginate

1.5

(b)

|D(3)/f(3)| (10-6/Hz)

Fe3O4-NaNO3+BSA Fe3O4-NaNO3+HA Fe3O4-PLL-NaNO3

1.0

0.5

0.0 Al2O3-NaNO3

1.5

|D(3)/f(3)| (10-6/Hz)

(c)

Al2O3-NaNO3+Alginate Al2O3-NaNO3+BSA Al2O3-NaNO3+HA Al2O3-PLL-NaNO3

1.0

0.5

0.0 0

491

20

40

60


492

Figure 5. |∆D(3)/∆f(3)| values as functions of electrolyte concentrations for colloidal

493

MnO2 deposited onto selected surfaces: (a) SiO2, (b) Fe3O4, and (c) Al2O3 as a function

494

the NaNO3 concentration in the absence and presence of HA, alginate, BSA and PLL

495

at pH 6.0. The error range shows the standard deviation.

496

Moreover, the values of the ratio |∆D(3)/∆f(3)| obtained when MnO2 colloids were 27



497

deposited onto silica surface in the absence and presence of macromolecules (0.12-1.29

498

×10-6 Hz−1) are generally larger than the values obtained during when these colloids

499

were deposited onto magnetite and alumina (0.10-0.72 ×10-6 Hz−1 for magnetite and

500

0.08-0.53 × 10-6 Hz−1 for alumina, respectively) (Figure 5b, c), indicating that the

501

“softest” deposited MnO2 colloidal layer was formed on the silica in the same water

502

solution condition, and a more rigid layer formed from the deposition of MnO2 colloids

503

onto magnetite and alumina. Since the energy barrier between MnO2 colloids and silica

504

was higher than that with magnetite in both the absence and presence of

505

macromolecules, as shown in the calculations of total interaction energy (Figure S10,

506

S11 and S13, S14), the negative deposited colloid might partially stick out into the bulk

507

solution, originating from the more repulsive silica. Correspondingly, lower repulsion

508

might result in lower flexibility of the deposited colloidal MnO2 onto magnetite. In

509

contrast, the positive alumina surface could draw deposited MnO2 colloids due to

510

electrostatic attraction (Figure S12, S15) between oppositely charged colloids and the

511

surface, and thus, a more rigid layer was observed. Likewise, the existence of

512

electrostatic attraction between the positive PLL coating and MnO2 colloids lead to a

513

lowest value of |∆D(3)/∆f(3)| (0.10-0.40×10-6 Hz−1), implying the formation of the most

514

rigid deposited layer. As presented in Figure 5a, 5b and 5c, |∆D(3)/∆f(3)| values of MnO2

515

colloids obtained in the presence of HA and alginate were slightly higher than those in

516

the absence of macromolecules for all tested surfaces. This slight enhancement could

517

be interpreted as MnO2 colloids that were coated by adsorption layer of HA or alginate

518

were less fully coupled to surfaces than naked colloids. Interestingly, the presence of 28


Page 28 of 42

Page 29 of 42


519

BSA led to lower values of |∆D(3)/∆f(3)|, indicating the formation of a rigid deposited

520

layer in the presence of BSA. It was previously reported that BSA had good affinity for

521

the negative silica surface and formed monolayers.57, 58, 59, 60 Moreover, BSA has the

522

strongest effect on retarding the aggregation rate of MnO2 colloids relative HA and

523

alginate, thus leading to enhanced individual deposition of these colloids onto surfaces.

524

Consequently, the strong association of BSA-coated MnO2 colloids with surfaces might

525

be responsible for the higher rigidity of the deposited layer. In general, the deposited

526

colloidal MnO2 layers in the presence of HA and alginate exhibited higher dissipation

527

energy, whereas a more rigid deposited layer formed in the presence of BSA.

528

4. ENVIRONMENTAL IMPLICATIONS

529

Manganese dioxide colloids are one of the most abundant Mn species, and their

530

retention on environmental surfaces is highly influenced by their interaction with

531

ubiquitous macromolecular organic matter (i.e., humic substance, polysaccharide and

532

protein) in aquatic systems. The data obtained herein imply that electrostatic surface

533

properties are the most critical surface characteristics controlling MnO2 colloidal

534

deposition in monovalent sodium solutions. Higher deposition kinetics of colloids may

535

indicate their lower mobility in the water system containing the related surfaces (e.g.,

536

silica, iron oxides, or/and aluminum oxides, etc.) in the presence of organic

537

macromolecules. Moreover, a further extended DLVO calculation verifies that the

538

deposition of MnO2 colloids could be hindered to a large extent in the presence of HA

539

and alginate due to the existence of steric repulsion. While BSA can decrease the energy

540

barrier for MnO2 colloids deposition on surfaces and thus increase the affinity of 29



541

colloids onto surfaces via electrostatic attraction. The enhanced deposition observed for

542

BSA indicates that BSA molecules play a more significant role compared to the other

543

two molecules with respect to the transport and retention of MnO2 colloids on

544

environmental surfaces.

545

This mechanistic investigation of the role of humic acid and biomacromolecules

546

on the deposition behavior of MnO2 colloids on environmental surfaces also has

547

significant implications for predicting the interactions with their associated

548

contaminants in aquatic systems. Chao and coauthors found that the oxidation of

549

CrxFe1–x(OH)3 solids by MnO2 are controlled by the diffusion of Cr(III)aq ion to

550

MnO2.80 Thus, Higher mobility of MnO2 in the presence of HA and alginate may

551

increase the proximity between Cr(III)aq to MnO2 oxidant, and thus might benefit to the

552

oxidation. Simultaneously, Mn could also be detected in the effluent due to the

553

hindrance effect of HA and alginate for the deposition of MnO2 colloids onto quartz

554

sand column. Besides, for removal of Tl from surface water by MnO2 colloids enhanced

555

quartz sand filtration process, the presence of BSA may efficiently enhance the removal

556

of Tl due to the enhanced deposition of MnO2 particles onto sand surfaces.

557

 ASSOCIATED CONTENT

558

Supporting Information

559

QCM-D crystal sensor cleaning methods. Representative TEM images of the MnO2

560

colloids. Aggregation measurement of colloidal MnO2. Deposition profiles of MnO2

561

colloids onto environmental surfaces and PLL-coated surfaces. Estimation of DLVO

562

and EDLVO interaction energy. Rate of BSA in relevant deposition processes. 30


Page 30 of 42

Page 31 of 42


563

 AUTHOR INFORMATION

564

Corresponding Author

565

Phone: 86-23-65120980; fax: 86-23-65120980; email: [email protected] (X. H.);

566

[email protected] (Q. H.)

567

Notes

568

The authors declare no competing financial interests.

569

 ACKNOWLEDGEMENTS

570

The present work has been financially supported by the National Natural Science

571

Foundation of China (51608067, 51878092), the Graduate Research and Innovation

572

Foundation of Chongqing, China (Grant CYS18029), the Scientific and Technological

573

Innovation

574

(cstc2015shmsztzx0053), the China Postdoctoral Science Foundation (Grant

575

2016M592643), the Chongqing Postdoctoral Science Foundation (Grant Xm2016059),

576

and the Program for Innovation Team Building at Institutions of Higher Education in

577

Chongqing. The authors thank Dr. Shihong Lin for the constructive advices in revising

578

the manuscript.

579

 REFERENCES

580

1.

581

riverine and estuarine suspended particles from the Ouse-Trent system, UK. Colloid.

582

Surface A 1997, 120, (1-3), 183-198.

583

2.

584

and manganese in aquatic systems: chemical and microscopic evidence. Geochim.

Special

Program

of

Social

Livelihood

of

Chongqing

Ferreira, J. R.; Lawlor, A. J.; Bates, J. M.; Clarke, K. J.; Tipping, E., Chemistry of

Lienemann, C. P.; Taillefert, M.; Perret, D.; Gaillard, J. F., Association of cobalt

31



585

Cosmochim. Acta 1997, 61, (7), 1437-1446.

586

3.

587

glutathione by soluble polymeric MnO2. Environ. Sci. Technol. 2003, 37, (15), 3332-

588

3338.

589

4.

590

nanoparticles. J Air Wast Mange. 2007, 9, (12), 1306-16.

591

5.

592

1. structure and behavior of colloidal material. Environ. Sci. Technol. 1995, 29, (9),

593

2169-75.

594

6.

595

preoxidation-destruction of precursors. J Water Supply Res. T.-Aqua 1996, 45, (6), 308-

596

315.

597

7.

598

propionate and butyrate degradation in anaerobic digestion. Journal of Hazardous

599

Materials 2018.

600

8.

601

biogenic manganese oxide. Environ. Sci. Technol. 2005, 39, (2), 569-576.

602

9.

603

inorganic oxides. Langmuir 2009, 25, (6), 3571-3576.

604

10. Jiang, J.; Pang, S. Y.; Ma, J., Oxidation of triclosan by permanganate (Mn(VII)):

605

importance of ligands and in situ formed manganese oxides. Environ. Sci. Technol.

606

2009, 43, (21), 8326-31.

Herszage, J.; dos Santos Afonso, M.; Luther, G. W., Oxidation of cysteine and

Wigginton, N. S.; Haus, K. L.; Hochella, M. F., Jr., Aquatic environmental

Buffle, J.; Leppard, G. G., Characterization of aquatic colloids and macromolecules.

Ma, J.; Graham, N., Controlling the formation of chloroform by permanganate

Tian, T.; Qiao, S.; Yu, C.; Zhou, J., Effects of nano-sized MnO2 on methanogenic

Villalobos, M.; Bargar, J.; Sposito, G., Mechanisms of Pb(II) sorption on a

Yang, K.; Lin, D. H.; Xing, B. S., Interactions of humic acid with nanosized

32


Page 32 of 42

Page 33 of 42


607

11. Huangfu, X.; Jiang, J.; Ma, J.; Liu, Y.; Yang, J., Aggregation kinetics of manganese

608

dioxide colloids in aqueous solution: influence of humic substances and

609

biomacromolecules. Environ. Sci. Technol. 2013, 47, (18), 10285-92.

610

12. Wang, C. Y.; Groenzin, H.; Shultz, M. J., Comparative study of acetic acid,

611

methanol, and water adsorbed on anatase TiO2 probed by sum frequency generation

612

spectroscopy. J. Am. Chem. Soc. 2005, 127, (27), 9736-9744.

613

13. Zhang, H. Z.; Penn, R. L.; Hamers, R. J.; Banfield, J. F., Enhanced adsorption of

614

molecules on surfaces of nanocrystalline particles. J. Phys. Chem. B 1999, 103, (22),

615

4656-4662.

616

14. Tseng, Y. H.; Lin, H. Y.; Kuo, C. S.; Li, Y. Y.; Huang, C. P., Thermostability of

617

nano-TiO2 and its photocatalytic activity. React. Kinet. Catal. L. 2006, 89, (1), 63-69.

618

15. Lu, X. X.; Huangfu, X. L.; Ma, J., Removal of trace mercury(II) from aqueous

619

solution by in situ formed Mn-Fe (hydr)oxides. J. Hazard. Mater. 2014, 280, 71-78.

620

16. Huangfu, X.; Ma, C.; Ma, J.; He, Q.; Yang, C.; Jiang, J.; Wang, Y.; Wu, Z.,

621

Significantly improving trace thallium removal from surface waters during coagulation

622

enhanced by nanosized manganese dioxide. Chemosphere 2017, 168, 264–271.

623

17. Huangfu, X.; Jiang, J.; Lu, X.; Wang, Y.; Liu, Y.; Pang, S.-Y.; Cheng, H.; Zhang,

624

X.; Ma, J., Adsorption and Oxidation of Thallium(I) by a Nanosized Manganese

625

Dioxide. Water, Air, & Soil Pollution 2014, 226, (1), 2272.

626

18. Huangfu, X.; Jiang, J.; Ma, J.; Wang, Y.; Liu, Y.; Lu, X.; Zhang, X.; Cheng, H.,

627

Reduction-induced aggregation and/or dissolution of MnO2 colloids by organics.

628

Colloid. Surface A 2015, 482, (2015), 485-490. 33



629

19. Huangfu, X.; Jiang, J.; Wang, Y.; Liu, Y.; Pang, S.-Y.; Lu, X.; Zhang, X.; Cheng,

630

H.; Ma, J., Reduction-induced aggregation of manganese dioxide colloids by guaiacol.

631

Colloid. Surface A 2015, 465, (2015), 106-112.

632

20. X. Huangfu, C. M., J. Ma, Q. He, C. Yang, J. Jiang, Y. Wang and J. Zhou, Effective

633

removal of trace thallium from surface water by nanosized manganese dioxide

634

enhanced quartz sand filtration. Chemosphere 2017, 189, 1-9.

635

21. Cheng, Y.; Huang, T. L.; Sun, Y. K.; Shi, X. X., Catalytic oxidation removal of

636

ammonium from groundwater by manganese oxides filter: Performance and

637

mechanisms. Chemical Engineering Journal 2017, 322, 82-89.

638

22. Mcnew, C. R.; LeBoeuf, E. J., nC(60) deposition kinetics: the complex contribution

639

of humic acid, ion concentration, and valence. J. Colloid Interf. Sci. 2016, 473, 132-

640

140.

641

23. Kubiak, K.; Adamczyk, Z.; Ocwieja, M., Kinetics of silver nanoparticle deposition

642

at PAH mono layers: reference QCM results. Langmuir 2015, 31, (10), 2988-2996.

643

24. Chowdhury, I.; Duch, M. C.; Mansukhani, N. D.; Hersam, M. C.; Bouchard, D.,

644

Deposition and release of graphene oxide nanomaterials using a quartz crystal

645

microbalance. Environ Sci Technol 2014, 48, (2), 961-9.

646

25. Quevedo, I. R.; Olsson, A. L. J.; Tufenkji, N., Deposition kinetics of quantum dots

647

and polystyrene latex nanoparticles onto alumina: Role of water chemistry and particle

648

coating. Environ. Sci. Technol. 2013, 47, (5), 2212-2220.

649

26. Chang, X. J.; Bouchard, D. C., Multiwalled carbon nanotube deposition on model

650

environmental surfaces. Environ. Sci. Technol. 2013, 47, (18), 10372-10380. 34


Page 34 of 42

Page 35 of 42


651

27. Qu, X. L.; Alvarez, P. J. J.; Li, Q. L., Impact of sunlight and humic acid on the

652

deposition kinetics of aqueous fullerene nanoparticles (nC(60)). Environ. Sci. Technol.

653

2012, 46, (24), 13455-13462.

654

28. Liu, X. Y.; Chen, G. X.; Su, C. M., Influence of collector surface composition and

655

water chemistry on the deposition of cerium dioxide nanoparticles: QCM-D and column

656

experiment approaches. Environ. Sci. Technol. 2012, 46, (12), 6681-6688.

657

29. Fatisson, J.; Domingos, R. F.; Wilkinson, K. J.; Tufenkji, N., Deposition of TiO2

658

nanoparticles onto silica measured using a quartz crystal microbalance with dissipation

659

monitoring. Langmuir 2009, 25, (11), 6062-6069.

660

30. Chen, K. L.; Elimelech, M., Aggregation and deposition kinetics of fullerene (C60)

661

nanoparticles. Langmuir 2006, 22, (26), 10994-1001.

662

31. Yi, P.; Chen, K. L., Influence of surface oxidation on the aggregation and

663

deposition kinetics of multiwalled carbon nanotubes in monovalent and divalent

664

electrolytes. Langmuir 2011, 27, (7), 3588-3599.

665

32. Furman, O.; Usenko, S.; Lau, B. L. T., Relative importance of the humic and fulvic

666

fractions of natural organic matter in the aggregation and deposition of silver

667

nanoparticles. Environ. Sci. Technol. 2013, 47, (3), 1349-1356.

668

33. Chen, K. L.; Elimelech, M., Interaction of fullerene (C-60) nanoparticles with

669

humic acid and alginate coated silica surfaces: Measurements, mechanisms, and

670

environmental implications. Environ. Sci. Technol. 2008, 42, (20), 7607-7614.

671

34. Philippe, A.; Schaumann, G. E., Interactions of dissolved organic matter with

672

natural and engineered inorganic colloids: a review. Environ Sci Technol 2014, 48, (16), 35



673

8946-62.

674

35. Jiang, X. J.; Tong, M. P.; Li, H. Y.; Yang, K., Deposition kinetics of zinc oxide

675

nanoparticles on natural organic matter coated silica surfaces. J. Colloid Interf. Sci.

676

2010, 350, (2), 427-434.

677

36. Quevedo, I. R.; Tufenkji, N., Influence of solution chemistry on the deposition and

678

detachment kinetics of a CdTe quantum dot examined using a quartz crystal

679

microbalance. Environ. Sci. Technol. 2009, 43, (9), 3176-3182.

680

37. Sun, B.; Zhang, Y.; Chen, W.; Wang, K.; Zhu, L., Concentration Dependent Effects

681

of Bovine Serum Albumin on Graphene Oxide Colloidal Stability in Aquatic

682

Environment. Environ Sci Technol 2018, 52, (13), 7212-7219.

683

38. Sheng, A.; Liu, F.; Xie, N.; Liu, J., Impact of Proteins on Aggregation Kinetics and

684

Adsorption Ability of Hematite Nanoparticles in Aqueous Dispersions. Environ Sci

685

Technol 2016, 50, (5), 2228-35.

686

39. Tamura, H.; Katayama, N.; Furuichi, R., Modeling of ion-exchange reactions on

687

metal oxides with the frumkin isotherm. 1. Acid−base and charge characteristics of

688

MnO2, TiO2, Fe3O4, and Al2O3 surfaces and adsorption affinity of alkali metal ions.

689

Environ Sci Technol 1996, 30, (4), 1198-1204.

690

40. Fetsch, D.; Havel, J., Capillary zone electrophoresis for the separation and

691

characterization of humic acids. J. Chromatogr. A 1998, 802, (1), 189-202.

692

41. Rigol, A.; Vidal, G. R., Ultrafiltration–capillary zone electrophoresis for the

693

determination of humic acid fractions. J. Chromatogr. A 1998, 807, (2), 275–284.

694

42. Chen, Q.; Xu, S.; Liu, Q.; Masliyah, J.; Xu, Z., QCM-D study of nanoparticle 36


Page 36 of 42

Page 37 of 42


695

interactions. Adv Colloid Interface Sci 2016, 233, 94-114.

696

43. Petosa, A. R.; Jaisi, D. P.; Quevedo, I. R.; Elimelech, M.; Tufenkji, N., Aggregation

697

and deposition of engineered nanomaterials in aquatic environments: Role of

698

physicochemical interactions. Environ. Sci. Technol. 2010, 44, (17), 6532-6549.

699

44. Jiang, X.; Tong, M.; Li, H.; Yang, K., Deposition kinetics of zinc oxide

700

nanoparticles on natural organic matter coated silica surfaces. J Colloid Interface Sci

701

2010, 350, (2), 427-434.

702

45. Li, W.; Liu, D.; Wu, J.; Kim, C.; Fortner, J. D., Aqueous Aggregation and Surface

703

Deposition Processes of Engineered Superparamagnetic Iron Oxide Nanoparticles for

704

Environmental Applications. Environ Sci Technol 2014, 48, (20), 11892-11900.

705

46. Erdemoğlu, M.; Sarıkaya, M., Effects of heavy metals and oxalate on the zeta

706

potential of magnetite. J Colloid Interf Sci 2006, 300, 795–804.

707

47. Lin, S.; Cheng, Y.; Bobcombe, Y.; K, L. J.; Liu, J.; Wiesner, M. R., Deposition of

708

silver nanoparticles in geochemically heterogeneous porous media: predicting affinity

709

from surface composition analysis. Environ Sci Technol 2011, 45, (12), 5209-5215.

710

48. Chang, X.; Bouchard, D. C., Multiwalled Carbon Nanotube Deposition on Model

711

Environmental Surfaces. Environ Sci Technol 2013, 47, (18), 10372-10380.

712

49. Quevedo, I. R.; Olsson, A. L.; Tufenkji, N., Deposition kinetics of quantum dots

713

and polystyrene latex nanoparticles onto alumina: role of water chemistry and particle

714

coating. Environ Sci Technol 2013, 47, (5), 2212-2220.

715

50. Chowdhury, I.; Duch, M. C.; Mansukhani, N. D.; Hersam, M. C.; Bouchard, D.,

716

Interactions of Graphene Oxide Nanomaterials with Natural Organic Matter and Metal 37



717

Oxide Surfaces. Environ Sci Technol 2014, 48, (16), 9382-9390.

718

51. Elimelech, M.; Gregory, J.; Jia, X.; Williams, R. A., Particle Deposition and

719

Aggregation: Measurement, Modeling, and Simulation. Butterworth-Heinemann:

720

Oxford, England 1995.

721

52. Kai, L. C.; Elimelech, M., Interaction of Fullerene (C60) Nanoparticles with Humic

722

Acid and Alginate Coated Silica Surfaces: Measurements, Mechanisms, and

723

Environmental Implications. Environ Sci Technol 2008, 42, (20), 7607-7614.

724

53. Tong, M.; Zhu, P.; Jiang, X.; Kim, H., Influence of natural organic matter on the

725

deposition kinetics of extracellular polymeric substances (EPS) on silica. Colloids and

726

Surfaces B: Biointerfaces 2011, 87, (1), 151-158.

727

54. Yang, X.; Flynn, R.; von der Kammer, F.; Hofmann, T., Modeling colloid

728

deposition on a protein layer adsorbed to iron-oxide-coated sand. J Contam Hydrol

729

2012, 142-143, 50-62.

730

55. Xiao, Y.; Wiesner, M. R., Transport and Retention of Selected Engineered

731

Nanoparticles by Porous Media in the Presence of a Biofilm. Environ Sci Technol 2013,

732

47, (5), 2246-2253.

733

56. Flynn, R. M.; Yang, X.; Hofmann, T.; von der Kammer, F., Bovine serum albumin

734

adsorption to iron-oxide coated sands can change microsphere deposition mechanisms.

735

Environ Sci Technol 2012, 46, (5), 2583-91.

736

57. Kubiak-Ossowska, K.; Tokarczyk, K.; Jachimska, B.; Mulheran, P. A., Bovine

737

serum albumin adsorption at a silica surface explored by simulation and experiment. J.

738

Phys. Chem. B 2017, 121, (16), 3975–3986. 38


Page 38 of 42

Page 39 of 42


739

58. Park, J. H.; Sut, T. N.; Jackman, J. A.; Ferhan, A. R.; Yoonab, B. K.; Cho, N.-J.,

740

Controlling adsorption and passivation properties of bovine serum albumin on silica

741

surfaces by ionic strength modulation and cross-linking. Phys. Chem. Chem. Phys. 2017,

742

19, 8854-8865

743

59. Borah, B. M.; Saha, B.; Dey, S. K.; Das, G., Surface-modification-directed

744

controlled adsorption of serum albumin onto magnetite nanocuboids synthesized in a

745

gel-diffusion technique. J. Colloid Interf. Sci. 2010, 349, (1), 114-121.

746

60. Kawashita, M.; Hayashi, J.; Li, Z. X.; Miyazaki, T.; Hashimoto, M.; Hihara, H.;

747

Kanetaka, H., Adsorption characteristics of bovine serum albumin onto alumina with a

748

specific crystalline structure. J. Mater. Sci.-Mater. M. 2014, 25, (2), 453-459.

749

61. Shim, M.; Kam, N. W. S.; Chen, R. J.; Li, Y. M.; Dai, H. J., Functionalization of

750

carbon nanotubes for biocompatibility and biomolecular recognition. Nano Lett 2002,

751

2, (4), 285-288.

752

62. Bazkin, A.; Norde, W., Physical Chemistry of Biological Interfaces. 1999, 115-135.

753

63. Jeyachandran, Y. L.; Mielczarski, E.; Rai, B.; Mielczarski, J. A., Quantitative and

754

qualitative evaluation of adsorption/desorption of bovine serum albumin on hydrophilic

755

and hydrophobic surfaces. Langmuir 2009, 25, (19), 11614-20.

756

64. Li, W.; Liao, P.; Oldham, T.; Jiang, Y.; Pan, C.; Yuan, S.; Fortner, J. D., Real-time

757

evaluation of natural organic matter deposition processes onto model environmental

758

surfaces. Water Res 2018, 129, 231-239.

759

65. LC Xu, B. L., Interaction Forces between Colloids and Protein-Coated Surfaces

760

Measured Using an Atomic Force Microscope. Environ Sci Technol 2005, 39, (10), 39



761

3592-3600.

762

66. Brewer, S. H. G., W. R.; Johnson, M. C.; Knaq, M. K.; Franzen, S. , Probing BSA

763

Binding to Citrate-Coated Gold Nanoparticles and Surfaces. Langmuir 2005, 21, (20),

764

9303−9307.

765

67. Kubiak-Ossowska, K.; Tokarczyk, K.; Jachimska, B.; Mulheran, P. A., Bovine

766

Serum Albumin Adsorption at a Silica Surface Explored by Simulation and Experiment.

767

J Phys Chem B 2017, 121 (16), 3975–3986.

768

68. Kubiak-Ossowska, K.; Jachimska, B.; Mulheran, P. A., How Negatively Charged

769

Proteins Adsorb to Negatively Charged Surfaces: A Molecular Dynamics Study of BSA

770

Adsorption on Silica. J Phys Chem B 2016, 120, 10463−10468.

771

69. Contreras, A. E.; Kim, A.; Li, Q., Combined fouling of nanofiltration membranes:

772

Mechanisms and effect of organic matter. Journal of Membrane Science 2009, 327, (1-

773

2), 87-95.

774

70. Armanious, A.; Aeppli, M.; Sander, M., Dissolved organic matter adsorption to

775

model surfaces: adlayer formation, properties, and dynamics at the nanoscale. Environ

776

Sci Technol 2014, 48, (16), 9420-9.

777

71. Rezwan, K.; Meier, L. P.; Rezwan, M.; Vörös, J.; Textor, M.; Gauckler, L. J.,

778

Bovine Serum Albumin Adsorption onto Colloidal Al2O3 Particles: A New Model

779

Based on Zeta Potential and UV−Vis Measurements. Langmuir 2004, 20, (23), 10055-

780

10061.

781

72. Fatisson, J.; Ghoshal, S.; Tufenkji, N., Deposition of carboxymethylcellulose-

782

coated zero-valent iron nanoparticles onto silica: roles of solution chemistry and 40


Page 40 of 42

Page 41 of 42


783

organic molecules. Langmuir 2010, 26, (15), 12832-12840.

784

73. Xu, D.; Hodges, C.; Ding, Y. L.; Biggs, S.; Brooker, A.; York, D., A QCM study

785

on the adsorption of colloidal laponite at the solid/liquid interface. Langmuir 2010, 26,

786

(11), 8366-8372.

787

74. Sirk, K. M.; Saleh, N. B.; Phenrat, T.; Kim, H. J.; Dufour, B.; Ok, J.; Golas, P. L.;

788

Matyjaszewski, K.; Lowry, G. V.; Tilton, R. D., Effect of adsorbed polyelectrolytes on

789

nanoscale zero valent iron particle attachment to soil surface models. Environ. Sci.

790

Technol. 2009, 43, (10), 3803-3808.

791

75. Poitras, C.; Tufenkji, N., A QCM-D-based biosensor for E. coli O157:H7

792

highlighting the relevance of the dissipation slope as a transduction signal. Biosens.

793

Bioelectron. 2009, 24, (7), 2137-2142.

794

76. Steinmetz, N. F.; Bock, E.; Richter, R. P.; Spatz, J. P.; Lomonossoff, G. P.; Evans,

795

D. J., Assembly of multilayer arrays of viral nanoparticles via biospecific recognition:

796

A quartz crystal microbalance with dissipation monitoring study. Biomacromolecules

797

2008, 9, (2), 456-462.

798

77. Rodahl, M.; Hook, F.; Fredriksson, C.; Keller, C. A.; Krozer, A.; Brzezinski, P.;

799

Voinova, M.; Kasemo, B., Simultaneous frequency and dissipation factor QCM

800

measurements of biomolecular adsorption and cell adhesion. Faraday Discuss. 1997,

801

107, 229-246.

802

78. Gall, I.; Herzberg, M.; Oren, Y., The effect of electric fields on bacterial attachment

803

to conductive surfaces. Soft Matter 2013, 9, (8), 2443-2452.

804

79. Styrkas, D. A.; Bütün, V.; Lu, J. R.; Keddie, J. L.; Armes, S. P., pH-Controlled 41



805

Adsorption of Polyelectrolyte Diblock Copolymers at the Solid/Liquid Interface.

806

Langmuir 2000, 16, (14), 5980-5986.

807

80. Pan, C.; Liu, H.; Catalano, J. G.; Qian, A.; Wang, Z. M.; Giammar, D. E., Rates of

808

Cr(VI) Generation from CrxFe1-x(OH)(3) Solids upon Reaction with Manganese

809

Oxide. Environmental Science & Technology 2017, 51, (21), 12416-12423.

810

42


Page 42 of 42

Deposition Kinetics of Colloidal Manganese Dioxide onto

Recommend Documents