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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
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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
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
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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
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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
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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
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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
99
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
112
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
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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
185
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
190
The deposition kinetics at each electrolyte concentration was quantified by the
191
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
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
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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).
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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
213
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
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
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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
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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
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
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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
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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
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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
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Deposition Rate (Hz/min)
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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.
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
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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
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
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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.
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.
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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
18
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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
21
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of
HA
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
23
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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.
26
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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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