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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

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Deposition Kinetics of Colloidal Manganese Dioxide onto

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Representative Surfaces in Aquatic Environments: The Role

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of Humic Acid and Biomacromolecules

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

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

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† Key

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

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

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University, 400044, China

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

10



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

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China

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

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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

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21 22

ABSTRACT

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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environmental surfaces and have significant implications for predicting the transport

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

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1. INTRODUCTION 2

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

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

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waters through various processes, including mineral weathering and dissolved

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

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

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

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

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

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surface

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

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

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

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

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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,

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Colloidal MnO2 aggregation kinetics have been previously investigated by dynamic

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

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

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

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

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

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

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

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

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

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

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

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are not available to date.

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Deposition kinetics have been widely investigated by employing different

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

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

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

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

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

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

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interpreting experimental data of deposition and proposing the deposition mechanisms

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

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

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

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

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

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

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nanoparticles.24,28

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Moreover, macromolecular organic matter widely distributed in natural aquatic

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

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

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

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

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

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collector and nanoparticle surfaces.28,

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

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

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

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

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

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

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protein commonly employed for examining the influence of protein on environmental

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

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

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

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deposition kinetics under the influence of BSA in aqueous dispersions.

33

Humic substances, polysaccharides, and

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

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

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

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

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

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

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

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

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colloidal MnO2 deposition kinetics for these representative surfaces were also

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

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

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

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explored as well.

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2. MATERIALS AND METHODS

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2.1. Materials

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

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

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

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

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

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

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41

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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

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

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temperature conditions (Model Multi3100, Jana, Germany).

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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

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

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

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

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

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

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

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and deposition experiment. The concentrations of MnO2 colloidal suspensions were

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

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

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

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

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

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

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

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

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

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elaborated in our previous publication.11

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2.3. Deposition Study Employing a Quartz Crystal Microbalance with a

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Dissipation Monitoring System

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

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

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

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

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

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crystal sensors were cleaned following protocols modified based on those

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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equation:

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Δ𝑚 = ―

𝐶Δ𝑓𝑛

(1)

𝑛

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

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

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in the crystal’s dissipation unit (D):

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𝐸𝑑𝑖𝑠𝑠𝑖𝑝𝑎𝑡𝑖𝑜𝑛

𝐷=

(2)

2𝜋𝐸𝑠𝑡𝑜𝑟𝑒𝑑

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

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energy in the oscillator.

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

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

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

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

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of MnO2 colloids:

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𝑟𝑓 =

|( ) | 𝑑𝛥𝑓(3) 𝑑𝑡

(3)

𝑡→0

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The deposition kinetics at each electrolyte concentration was quantified by the

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

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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

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concentration, while the denominator represents the corresponding deposition rate

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under favorable conditions obtained in the same electrolyte solutions.

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

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

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

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

200

Information (Text S2, Figures S2-S5).

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2.4. Calculation of Interaction Energies between Colloidal MnO2 and Surfaces

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DLVO and EDLVO theories were employed to better interpret the mechanisms

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

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

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

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

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

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

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

210

Information (Text S4).

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3. RESULTS AND DISCUSSIONS

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

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HA and Biomacromolecules

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Representative normalized frequency shifts at the third overtone when colloidal MnO2

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

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

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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

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

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

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

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

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

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

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

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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

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

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

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

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

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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

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

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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

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

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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

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

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

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

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

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

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

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

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

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

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

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

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

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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

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

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

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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

Deposition Rate (Hz/min)

103

(b) 10

1

NaNO3 Concentration (mM)

Al2O3 102

101

100

(c) 10-1

257

Fe3O4

101

NaNO3 Concentration (mM)

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.

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

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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

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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

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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

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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

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

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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

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48

The observed

of

HA

and

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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.

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

291

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

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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.

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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

NaNO3 Concentration (mM)

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

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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

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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

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NaNO3 NaNO3+HA

Deposition Rate (Hz/min)

103

NaNO3+Alginate

SiO2

NaNO3+BSA 102

101

100

10-1

(a) 1

10

NaNO3 Concentration (mM)

Deposition Rate (Hz/min)

103

Fe3O4

102

101

100

10-1

(b) 1

10

NaNO3 Concentration (mM)

Deposition Rate (Hz/min)

103

Al2O3

102

101

100

10-1

(c) 1

10

NaNO3 Concentration (mM)

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

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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

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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

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and

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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

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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

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NaNO3 NaNO3+HA NaNO3+Alginate

Attachment Efficiency

101

NaNO3+BSA

SiO2

100

10-1

10-2

10-3

101

(a) 10

NaNO3 Concentration (mM)

Attachment Efficiency

Fe3O4 100

10-1

10-2

10-3

(b) 10

1

NaNO3 Concentration (mM) 101

Attachment Efficiency

Al2O3 100

10-1

10-2

10-3

(c) 10

1

NaNO3 Concentration (mM)

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

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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

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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.

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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

NaNO3 Concentration (mM)

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

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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

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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

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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

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 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

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