Enhanced Colloidal Stability of CeO2 Nanoparticles by Ferrous Ions

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Enhanced Colloidal Stability of CeO2 Nanoparticles by Ferrous Ions: Adsorption, Redox Reaction, and Surface Precipitation Xuyang Liu, Jessica Renee Ray, Chelsea W. Neil, Qingyun Li, and Young-Shin Jun Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/es506363x • Publication Date (Web): 07 Apr 2015 Downloaded from http://pubs.acs.org on April 12, 2015

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Environmental Science & Technology

Enhanced Colloidal Stability of CeO2 Nanoparticles by Ferrous Ions: Adsorption, Redox Reaction, and Surface Precipitation

Xuyang Liu †, Jessica R. Ray †, Chelsea W. Neil †, Qingyun Li, and Young-Shin Jun* Department of Energy, Environmental and Chemical Engineering, Washington University in St. Louis, St. Louis, MO 63130, United States

E-mail: [email protected] http://encl.engineering.wustl.edu/ Submitted: December 2014 Revised: March 2015

Environmental Science &Technology

†These

authors contributed equally to the current work.

*To whom correspondence should be addressed.

ACS Paragon Plus Environment


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ABSTRACT

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Due to the toxicity of cerium oxide (CeO2) nanoparticles (NPs), a better understanding of the redox

3

reaction-induced surface property changes of CeO2 NPs and their transport in natural and

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engineered aqueous systems is needed. This study investigates the impact of redox reactions with

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ferrous ions (Fe2+) on the colloidal stability of CeO2 NPs. We demonstrated that under anaerobic

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conditions suspended CeO2 NPs in a 3 mM FeCl2 solution at pH = 4.8 were much more stable

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against sedimentation than those in the absence of Fe2+. Redox reactions between CeO2 NPs and

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Fe2+ lead to the formation of 6-line ferrihydrite on the CeO2 surfaces, which enhanced the colloidal

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stability by increasing the zeta potential and hydrophilicity of CeO2 NPs. These redox reactions

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can affect the toxicity of CeO2 NPs by increasing cerium dissolution and by creating new Fe(III)

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(hydr)oxide reactive surface layers. Thus, these findings have significant implications for

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elucidating the phase transformation and transport of redox reactive NPs in the environment.

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INTRODUCTION

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Cerium oxide (CeO2) nanoparticles (NPs) have been enumerated in the priority list of

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engineered nanomaterials for risk evaluation by the Organization for Economic Co-operation and

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Development (OECD) due to their wide application in industry and daily life.1 CeO2 NPs are ideal

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in catalysis applications, such as diesel engine catalytic converters,2 due to their oxygen storage

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capabilities3 and their ability to readily participate in Ce4+/Ce3+ redox processes. The use and

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production of CeO2 NPs will inevitably result in increased concentrations in natural or manmade

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aqueous environments, such as wastewater.4 Therefore, in order to better predict the life cycle of

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these engineered nanoparticles, it is vital to improve our understanding of their fate and transport,

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particularly in aqueous environments.

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Based on a 2001 report on human health risks of cerium from diesel fuels, the average

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worldwide estimated level of cerium in soils was 20–60 ppm.5 Recent studies have also shown that

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nanometer sized CeO2 particles are found in automobile exhaust.6,

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warrants immediate attention to prevent harmful effects to the biosphere. The bioavailability and

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toxicity of CeO2 NPs are largely determined by their fate and transport in the environment, which

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is in turn affected by their surface charge and aggregation state.8, 9 While recent investigations in

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this field have focused on the influence of solution chemistry and organic matter on the fate and

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transport of CeO2 NPs,2, 10-16 few studies have considered the effect of redox reactions on CeO2

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NP surface properties in the presence of redox-active ions in diverse aqueous environments.17, 18

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This high concentration

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When released in the environment, CeO2 NPs can coexist and interact with redox sensitive

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elements. One important element to consider is ferrous iron (Fe2+), which is widely distributed in

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natural aqueous systems (e.g., acid mine drainage),19 as well as in engineered systems for odor and

2



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corrosion control, precipitating hydrogen sulfide, and phosphorus removal.20, 21 The concentration

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of ferrous ions used in these engineered systems can be as high as 9.85 mM for phosphorus

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removal,22 and 1.5–3.0 mM to remove cyanide from industrial wastewater.23 In addition, in a study,

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Fe3O4/CeO2-impregnated NPs were synthesized and used as Fenton-like catalysts for hydroxyl

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radical generation and 4-chlorophenol degradation with the addition of H2O2.24 Ce4+ has also been

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used in the production of silicate glasses to rid the glass of Fe2+ ions.25 Therefore, anthropogenic

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CeO2 NPs originated from industrial wastewater can interact with ferrous ions during their

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industrial applications, as well as during wastewater treatment processes.26, 27

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According to the difference in the standard redox potential of Fe3+/Fe2+ (0.77V) and

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Ce4+/Ce3+ (1.44 V), redox reactions between Fe2+ and CeO2 can occur when they coexist in

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solution.28 However, it is largely unknown how this reaction with aqueous Fe2+ will affect the

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surface properties and colloidal stability of CeO2 NPs in aqueous environments. In this study, we

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systematically investigated changes in CeO2 NP surface properties when aqueous Fe2+ is present.

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Knowledge obtained in this study can help improve our understanding of the fate and transport of

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redox active nanomaterials during their lifetime, which will in turn give insight into the expected

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bioavailability of these anthropogenic NPs when released into the environment.

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EXPERIMENTAL SECTION

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Preparation of Nanoparticle Dispersions for Aggregation and Sedimentation Tests

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All solution preparations and the following wet experiments were conducted in an

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anaerobic Coy chamber to prevent the influence of dissolved oxygen (DO). By distinguishing Fe2+

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oxidation by CeO2 from oxidation by molecular oxygen, we can provide a better understanding of

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the redox reactions and mechanisms in aqueous Fe2+–CeO2 NP systems. Depleted oxygen can

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occur under specific conditions, such as in anoxic wastewater treatment processes, underground,

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or at the bottom of stratified lakes.29 Furthermore, in acidic environment, the Fe2+ oxidation rate

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by dissolved O2(aq) equilibrated at 0.2 atm is quite slow. Thus, our anaerobic experimental

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conditions can still be applicable in wastewater systems where Fe2+ is present and the local pH is

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acidic. To ensure anoxic conditions, deionized (DI) water was boiled prior to use in these

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experiments to remove DO, and then cooled to room temperature in the anaerobic chamber.

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We applied commercially available CeO2 NPs (Sigma-Aldrich, St. Louis, MO), so that our

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results have closer connections to real engineered NP–environmental systems. CeO2 NP

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dispersions of 50 mg/L were created in 50 mL test tubes inside the chamber, and the test tubes

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were indirectly ultrasonicated for 1 h before reaction using a Fisher Scientific ultrasonic cleaner

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(model no. FS6) with a frequency of 50/60 kHz and power of 30W. For all experiments, the ionic

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strength (IS) was maintained at 10 mM. In the reaction system, 3 mM FeCl2 and 1 mM NaCl were

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added to the CeO2 dispersions to have an ionic strength (IS) of 10 mM and a pH of 4.8 ± 0.2. The

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final CeO2 concentration of the mixed solution was 45 mg/L. A weakly acidic pH can occur in

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acid mine drainage environments.30, 31 In addition, when applying iron in wastewater treatment

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plants, the hydrolysis of iron could lead to pHs of 5 or less if no pH control is conducted., 21, 32 At

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pH 4.8, initial ferrous iron remains largely in the ferrous state, which could have greater

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toxicological effects than other iron species on organisms, e.g., the feeding activity and motility of

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the mayfly larvae.33, 34 This is also the pH of the system solution after FeCl2 addition, so errors in

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the ionic strength through pH adjustment can be minimized. The concentration of 3 mM Fe2+ used

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in our experimental system is commonly found in natural environments (e.g. 2–4 mM in an anoxic

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lake in Massachusetts) and wastewater treatment processes that use additive ferrous iron.22, 23, 35 In

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the control experiment,10 mM NaCl (with no FeCl2) was added to the CeO2 dispersion and the pH 4



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was adjusted to pH 4.8 with dilute HCl solution to match that of the Fe2+-containing system and

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maintain the final CeO2 concentration of the mixed solution as 45 mg/L. The pH change during

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the reaction period was monitored (Figure S6 in the SI). In addition, these conditions give valuable

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insight into the interactions that occur between Fe2+ and CeO2, and can be used to bolster further

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studies of CeO2 NP stability in more complex aqueous systems.

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During sedimentation experiments, aliquots were taken from the supernatant at elapsed

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times using a pipette with minimal disturbance to the suspension. The concentrations of CeO2 NPs

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were measured by UV-vis spectroscopy (Varian Inc., Cary 50 Bio UV-Vis Spectrophotometer,

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Palo Alto, California) at a wavelength of 305 nm, where the highest absorbance was obtained

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(Figure S1 in the Supporting Information). Sedimentation experiments were run for 93 hours in

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triplicate. To verify the changes in CeO2 surface properties following redox reactions, after

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reaction the zeta potentials and particle sizes (Dynamic Light Scattering, DLS) of CeO2

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nanoparticle aggregates in the Fe2+ and control systems were measured using a Zetasizer (Nano

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ZS, Malvern Instruments Ltd., Westborough, MA). The zeta potential was derived from the

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original electrophoretic mobility using the Smoluchowski equation. We have characterized the

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particle size using DLS (Table S1) and TEM (Figure S7). The hydrodynamic particle sizes by DLS

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were collected for at least three measurements (with each measurement taken over 10 s) and

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triplicate or more experiments were conducted for each sample. The hydrodynamic diameter is

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defined as the diameter of the aggregate plus that of the hydration layer.

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CeO2 NP Dissolution Experiments

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Due to the low CeO2 solubility,36 CeO2 suspensions of 250 mg/L and 15 mM FeCl2 were

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used in dissolution experiments in order to achieve aqueous Ce concentrations above the 25 μg/L

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detection limit of the inductively coupled plasma-optical emission spectrometer (ICP-OES).

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Similar or even higher concentrations of CeO2 for dissolution experiments have been also

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investigated by other researchers.37, 38 Because the Fe2+ concentration scales up proportionally, we

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expect that this elevated concentration system is relevant to lower suspended CeO2

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concentrations.37, 39, 40 Thus, the concentration we used for our dissolution tests reasonably allows

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us to compare our results with previous CeO2 NP studies.

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Reaction and control solutions were prepared in the anaerobic chamber as described for

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sedimentation experiments. Triplicate batches of the reaction and control systems were separated

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into 5 mL test tubes and put into test tube rotators in the chamber.39, 41, 42 At 1, 2, 3, 4, 5, and 6

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hour time points, triplicate samples were taken from the two systems. Samples were placed in

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ultracentrifuge tubes (PC Oak Ridge Tubes, Fisher Scientific) in the anaerobic chamber, capped,

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and removed from the chamber. To separate the supernatant from the CeO2 NPs, samples were

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then ultracentrifuged using a Thermo Scientific Sorvall WX Ultra Series Centrifuge with a T-865

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Fixed Angle Rotor at 40,000 rpm (or 115,861 x g) for 30 minutes. In addition to centrifugation,

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the samples were placed back inside the anaerobic chamber and filtered by a 0.2 µm filter

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(Millipore syringe filter) to ensure the removal of all bulk CeO2 NPs and aggregates. We have

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verified the efficacy of combining ultracentrifugation and filtration to separate CeO2 NPs in

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preliminary tests, and this method has also been commonly used in CeO2 NP dissolution and

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separation studies.39, 41, 42 The filtrate was collected and acidified to 1% v/v nitric acid for ICP-

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OES measurements.

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Phase Identification of Fe(III) (Hydr)oxides on CeO2 NP Surfaces

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We conducted replicate experiments for X-ray absorption spectroscopy (XAS) analysis.

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After reacting for 6 hours, solutions were transferred to centrifugation tubes and capped in the

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anaerobic chamber, then ultracentrifuged at 40,000 rpm for 30 minutes. Once centrifuged, the

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supernatant was poured off, leaving the CeO2 NPs fixed at the bottom of the tube. The supernatant

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was then replaced with deoxygenated DI water and the tubes were then capped and removed from

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the chamber for an additional 30 min of ultracentrifugation to remove excess salt. After the second

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ultracentrifugation, the DI water was poured off in the anaerobic chamber, and the solids were

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allowed to dry in the chamber overnight to prevent any oxidation. The cerium spectra were

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measured in transmission mode, and iron spectra were measured in fluorescence mode.

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Experiments were conducted at Beamline 13BM-D at the Advanced Photon Source (APS),

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Argonne National Laboratory. This station utilized a Si(111) monochromator, giving it a focused

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beam size of 10 m by 30 m and a resolution of 1  10-4 E/E. The energy flux was 1  109 at 10

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keV. The energy range for this station was 4.5–70 keV. The iron XANES edge was measured at

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7.119 keV and the cerium edge was measured at 40.444 keV.

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Secondary Mineral Phase Precipitation on CeO2-Sputtered Substrates by Physical Vapor

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Deposition (PVD)

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To identify the properties of reaction products on the surface of NPs, CeO2 substrates were

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created by sputtering CeO2 NPs on clean Si wafers using a PVD process (Kurt Lesker PVD 75,

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Livermore, CA). The DC (diode) mode was applied under 1 mTorr argon pressure, with 100 watts

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power input for 2000 seconds. The deposition was monitored in situ via a built-in quartz crystal

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microbalance. Because CeO2 nanoparticles were stable, solid phase analyses were conducted under

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atmospheric conditions. The thickness of the CeO2–sputtered wafer was quantified by alpha-SE

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ellipsometry (Lincoln, NE) and atomic force microscopy (AFM, Veeco, Nanoscope V) after the

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PVD sputtering process. During experiments, the CeO2 substrate was exposed to the same solution

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chemistry as the aqueous CeO2 NP experiments, containing either NaCl only or NaCl and Fe2+

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ions (Figure S2). A clean Si wafer was also used as a control to compare precipitation in the system

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containing NaCl and Fe2+ ions in the absence of CeO2. The precipitates on the CeO2-sputtered

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substrate and Si wafer were analyzed by AFM using tapping mode. We collected height, amplitude,

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and phase contrast information simultaneously for 5 µm × 5 µm areas, and analyzed the images

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using the Nanoscope 7.20 software. Experimental details on the AFM setup have been reported in

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our former studies.43, 44

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Grazing Incidence Small-Angle X-ray Scattering (GISAXS) and SAXS Measurements

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The surface properties and precipitates on the CeO2-sputtered substrates were also

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characterized in situ by GISAXS. In this experiment, the solution was prepared under the same

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solution chemistry as the CeO2 NP sedimentation tests, i.e., 3 mM FeCl2 and 1 mM NaCl at pH =

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4.8. For the control system, a 10 mM NaCl solution was used. The CeO2 substrates were placed

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flat at the bottom of a specially designed GISAXS cell. The solutions were injected at the top of

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the cell, and then the cell was capped. The substrates were reacted for 1 hour. During GISAXS

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measurements, incident X-ray beams at 18 keV were passed through the cell, where they interacted

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with particles precipitating on the substrate surface (GISAXS). The scattered X-ray beams were

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collected by a 2D detector. X-ray scattering data was processed by cutting along the Yoneda wing.

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All data reduction was conducted using the GISAXS shop macro, a software package available at

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APS beamline 12-ID. More detailed data fitting procedure descriptions are available in our

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previous publications.43-45 Prior to size fitting, the background image of the CeO2 sputtered wafer 8



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in DI water was subtracted from all 2D GISAXS scattering images. Thus, observed changes in

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particle size were from the precipitation of Fe(III) (hydr)oxides. During the particle size fitting,

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the shapes of the particles were assumed to be spherical, thus a form factor for polydisperse

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spherical particles. A Schultz size distribution was also assumed for fitting.

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RESULTS & DISCUSSION

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Enhanced Colloidal Stability of CeO2 NPs in Fe2+ Solutions

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First, we determined how the presence of Fe2+ affects the aggregation and subsequent

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sedimentation of CeO2 NPs. During the 4 day monitoring period, the concentration of CeO2 NPs

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in the supernatant was higher in the system with added Fe2+ than in the system without aqueous

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Fe2+ (Figure 1A), indicating that the presence of 3 mM Fe2+ promoted the colloidal stability of

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CeO2 NPs. For example, only 46% of the initial CeO2 NP’s concentration remained in the

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supernatant in the control system while 94% of the initial CeO2 NPs remained stable after 16 h in

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the presence of 3 mM FeCl2. In addition, the significant difference in NP sedimentation due to the

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presence of Fe2+ can be observed visually in the solutions after stirring overnight (Figure S3,

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Supporting Information). It is noted that the effect of double layer compression was similar in the

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solutions because they had same 10 mM ionic strength and same anion identity (i.e., Cl-), which

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will dominate double layer compression effects due to the positive charge of CeO2 NPs.

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Analyses of particle sizes and zeta potentials in the two experimental systems also help to

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illustrate the observed sedimentation trends. The zeta potential of CeO2 was 28.6 ± 1.7 mV in the

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presence of Fe2+, while in the absence of Fe2+, the zeta potential was 8.7 ± 0.8 mV (Table S1,

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Supporting Information). The isoelectric point, pHiep, of CeO2 NPs in 10 mM NaCl was measured

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to be pH 4.95 (Figure S4A), which is within the range of 3.0–7.6 found in the literature. The wide

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variation in this range can be due to the varying surface properties of the nanoparticles produced

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through different methods.46-48 In our experimental system, the unreacted CeO2 NPs should be

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positively charged under the test conditions.

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The hydrodynamic size for stable suspensions (Table S1) and XRD spectra (Figure S4B)

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for the unreacted CeO2 NPs are given in the SI. The hydrodynamic size was 137.7 ± 0.7 nm for

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stable dispersion of unreacted CeO2 in DI water, while the particle sizes measured by TEM ranged

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from 5–30 nm (Figure S7), consistent with the manufacture's nominal size. After 16 hours reaction,

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the hydrodynamic diameter (DH) of CeO2 NPs in the 3 mM Fe2+ solution was 154 ±1 nm, while

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the DH of CeO2 NPs in the 10 mM NaCl system was 3829 ± 358 nm (Table S1). The smaller size

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of CeO2 aggregates when Fe2+ is present is in accordance with the higher suspended concentrations

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of CeO2 NPs compared to the system without Fe2+, because the smaller aggregates will be more

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stable in solution. In contrast, larger aggregates tend to settle out because gravitational forces are

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predominant over Brownian motion.

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Identification of the Reaction Products and CeO2 Dissolution

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We hypothesize that the Fe2+-promoted stability of CeO2 NPs results from redox

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reactions28 between Fe2+ and CeO2 NPs and the formation of Fe(III)products at the CeO2 NP

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surfaces (eqn. (1)). Throughout the current manuscript, we expressed the oxidation state of solid

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and sorbed species as Roman numerals and the oxidation state of aqueous species as Arabic

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numerals. Arabic numerals were also used when both solid and aqueous species are possible. ≡CeIVO2+ Fe2+ (aq)↔ Ce3+ (aq)(or ≡CeIII2O3) + ≡Fe3+

208

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(1)


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X-ray absorption spectroscopy (XAS) was carried out to determine the Ce K-edge and Fe

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K-edge spectra for the solid CeO2 NPs before and after reaction in the FeCl2 and control systems

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(Figure 2). X-ray absorption near edge structure (XANES) results for Ce indicated no significant

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change in the oxidation state after reaction. In other words, there is no CeIII contribution in both

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initial CeO2 NPs and reacted CeO2 NPs (Figure 2A). This could be due to negligible amounts of

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CeIII on the surface compared to bulk CeIV. Considering that the presence of surface CeIII2O3 was

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not detected, and the aqueous Ce concentration was significantly enhanced in the presence of Fe2+

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(as will discussed shortly after), we suggest that aqueous Ce3+ is the more predominant reaction

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product than solid CeIII2O3. For iron, XAS results (Figure 3B) indicate that all iron on the reacted

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CeO2 NP surface was oxidized to FeIII. Furthermore, the iron spectrum for the reacted CeO2 NPs

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is consistent with that of the ferrihydrite standard (Figure 2B).

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Aqueous Ce levels were monitored for the first 6 hours of reaction in the 250 mg/L CeO2

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system (Figure 1B) because the most vigorous interfacial reaction occurs during the early stage.

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Based on the concentration of dissolved Ce from CeO2 NPs in the supernatant, we also found that

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the dissolved Ce concentration was 15 times higher in the presence of aqueous Fe2+ than in the

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control system in the absence of Fe2+ over the course of the reaction period. Because the solubility

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of Ce3+ is 25 orders of magnitude higher than CeIVO2, and it was reported that under similar

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aqueous conditions to our experimental condition, more than 99% dissolved Ce was Ce3+,36, 49 we

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assumed that the aqueous Ce in our system was primarily Ce3+. Thus, the increased concentration

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of dissolved Ce could result from the reduction of CeIVO2 by Fe2+. Then, the redox reaction is also

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likely to accelerate the precipitation of Fe(III)(hydr)oxides on the surface of CeO2 NPs while the

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dissolved Ce3+ ions are released from the CeIVO2 NP surfaces.

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Reaction Pathways: Adsorption, Redox Reaction, and Surface Precipitation of Fe(III)(Hydr)oxides on the Surfaces of CeO2 NPs and Dissolution of CeO2 NPs

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A series of reactions could contribute to the enhanced colloidal stability of CeO2 in the

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presence of Fe2+ ions. MgCl2 extraction revealed that the adsorbed Fe2+ concentration on the

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surface of CeO2 NPs was 3.75 mg/L, which is equivalent to 0.05% of the total iron (Figure S5,

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Supporting Information). The Fe2+ can be adsorbed on the CeO2 NPs by ion exchange (eq. (2)),

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and such a process is expected to shift the zeta potential of CeO2 NPs to be more positive and

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release H+ into the solution:50

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Fe2+ Sorption: ≡CeOH + Fe2+ ↔ ≡CeOFe+ + H+

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On the other hand, the adsorbed FeII can react further with CeIV on the surface of CeO2 NPs

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rather than existing as ferrous state, due to the large enough difference in the standard redox

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potential for Ce4+/Ce3+ (1.44 V) and Fe3+/Fe2+ (0.77 V).28 As a control experiment, we tested the

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reaction of aqueous Ce4+ using Ce(SO4)2 and Fe2+ ions in solution and observed a significant

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decrease of Fe2+ concentration by the Ferrozine method after 1 d (Table S2, Supporting

245

Information). Furthermore, the XANES experiments of the CeO2 NP‒Fe2+ system also revealed

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oxidized Fe(III) on the surface of CeO2 NPs under anaerobic conditions, providing direct evidence

247

of surface redox reactions (Figure 2).

(2)

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In the experimental systems, the pH of the reaction solution (CeO2 NPs + FeCl2) was measured

249

to be lower than either CeO2 NPs or FeCl2 solutions alone, and decreased during the first few hours

250

of reaction (Figure S6, Supporting Information). The decrease of solution pH can be attributed to

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both the adsorption of Fe2+ through ion exchange (eqn. (2)) and the hydrolysis of Fe3+, the redox

252

reaction product, which releases more H+ into solution than is released by Fe2+ sorption. The Fe2+12



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only system also has slightly decreased pH due to minor Fe2+ hydrolysis (eqn. 3a), but the rate of

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pH decrease is much less than that of the CeO2/Fe2+ system, which undergoes the hydrolysis of

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ferric ions (eqn. 3b).

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Hydrolysis: Fe2+ + H2O  Fe(OH)+(aq) + H+ or

(3a)

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Fe3+ + 3 H2O  Fe(OH)3 (aq) + 3 H+

(3b)

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Once hydrolysis of ferric iron forms the monomer (eqn. 3b), then dimers and polymers can

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form through continuous olation (hydroxo–bridging) and oxolation (oxo–bridging) reactions.44

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Once the polymeric cluster size is larger than the critical nucleus size, stable nuclei form, and

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Fe(III) (hydr)oxide precipitation occurs.45

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To observe the particle formation of Fe(III) (hydr)oxides and their morphological changes

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in Fe2+–CeO2 NP systems, AFM and GISAXS were used on a CeO2-sputtered Si wafer reacted in

264

the same solution conditions. AFM images show significant precipitation of FeIII solid phase on

265

the CeO2-sputtered Si wafer substrate after 6 h (Figure 3). The particle size of precipitated Fe(III)

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(hydr)oxides on the CeO2 substrates increased from 4.8 ± 0.7 nm (based on 20 particle analysis)

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to 10.6 ± 2.8 nm after 6 h, and 25.9 ± 5.4 nm after 1 day (based on 50 particle analysis each). In

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contrast, precipitation of Fe(III) (hydr)oxide on a clean Si substrate was negligible within 1 day.

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GISAXS results also showed Fe(III) (hydr)oxide nucleation on the surface of CeO2-sputtered

270

substrates (Figure 4). Figure 4A revealed that particles with an Rg of 1.72 nm were observed on

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the surface of CeO2 substrates in the presence of 3 mM Fe2+ in 1 h. In contrast, no particles formed

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in the absence of Fe2+due to no scattering increase for the CeO2 only system. SAXS images in

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Figure 4B indicates that no particles formed in solution in the absence of CeO2 (e.g., Fe2+ only

274

system) as well. TEM analysis of CeO2 NPs revealed that nucleated Fe(III) (hydr)oxide covered 13


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the surfaces of CeO2 NPs and appeared to increase the CeO2 NP particle–particle distance (Figure

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S7). Furthermore, electron diffraction analysis of the Fe(III) (hydr)oxide products from the CeO2‒

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Fe2+ system revealed the formation of 6-line ferrihydrite, which is consistent with XAS results. It

278

is important to note that artifacts and phase transformation might be introduced during TEM

279

sample preparation due to the drying process.43 Therefore, the morphology of the in situ Fe(III)

280

(hydr)oxide particles might be different than what is observed using TEM. However, the multiple

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complementary techniques used in this study provide convincing and consistent evidence that the

282

precipitation of iron on the CeO2 surfaces greatly affected the surface properties and colloidal

283

stability of CeO2 NPs. Based on the above discussion, eqn. (1) can be rewritten as follows:

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≡CeIVO2+ Fe2+ (aq)↔ ≡CeIVO2 + Ce3+ (aq) + ≡Fe(III)(hydr)oxide (e.g., ferrihydrite).

285

(4)

Enhanced Stability by Iron (Hydr)oxide Surface Precipitation

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First, we analyzed the change in stability of CeO2 NPs using classical Derjaguin–Landau–

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Verwey–Overbeek (DLVO) interactions, comprised of electrostatic repulsion forces and van der

288

Waals attractions. The electrostatic interaction increases with the absolute value of surface

289

potential (eqns. (S1-3) in Section S4 of the Supporting Information). Nanoparticles with higher

290

zeta potential have higher electrostatic repulsions each other, and therefore, are more stable in

291

suspension. Hence, the increased zeta potential due to the adsorption of Fe2+ and precipitation of

292

Fe(III)(hydr)oxides on the CeO2 NP surfaces can lead to more significant electrostatic repulsions.

293

In addition, the values of the Hamaker constant for iron oxide phases (magnetite, maghemite, and

294

hematite) in the literature range from 1.3 to 4.5 × 10-20 J.51 Although the Hamaker constant for

295

ferrihydrite is not available, the highest reported value for an iron oxide related phase (4.5 × 10-20

296

J) is still smaller than that of CeO2 (5.6–6.0×10-20 J).2, 52 A smaller Hamaker constant signifies that 14



Page 16 of 31

297

there are less attractive forces between Fe(III) (hydr)oxide–Fe(III) (hydr)oxide than that of the

298

CeO2 NP–CeO2 NP. The precipitated Fe(III)(hydr)oxides on the surface of CeO2 will most likely

299

increase the electrostatic repulsion and decrease the van der Waals attractions.

300

Non-DLVO interactions also possibly play a significant role in the interactions of CeO2

301

NPs after being reacted with Fe2+. We hypothesize that the enhanced Ce stability in the presence

302

of Fe2+ may also result from Fe(III) (hydr)oxide precipitation changing the CeO2 NP surface

303

hydrophilicity. Due to the distinctive electronic structure of rare earth atoms, the CeO2 surface has

304

intrinsically less hydrophilic properties.53 In a cerium atom, the outer full octet of electrons in the

305

5s2p6 shell shields the unfilled 4f orbitals. As a result, CeO2 tends to not exchange electrons and

306

form hydrogen bonds with surrounding water molecules, making it less hydrophilic than other

307

metal oxides.53 The precipitated Fe(III) (hydr)oxides could, therefore, make the NP surface more

308

hydrophilic than the CeO2 surface.53 As a result, the coated NPs would be more stable in the

309

aqueous phase. To test the extent of hydrophilicity changes of CeO2 NPs by Fe(III) (hydr)oxide

310

surface coatings, we conducted surface angle measurements. It appeared that the surface of CeO2-

311

sputtered wafers became more hydrophilic after the precipitation of Fe(III) (hydr)oxides within a

312

3 day reaction period. (Figure S8, Supporting Information)

313

ENVIRONMENTAL IMPLICATIONS

314

In this study, we found that redox reactions between CeO2 NPs and Fe2+ lead to the

315

formation of 6-line ferrihydrite on the CeO2 surface, which enhanced the colloidal stability by

316

increasing the zeta potential and hydrophilicity of CeO2 NPs. The findings of this work suggest

317

longer retention periods and farther transport distances of CeO2 NPs in aquatic environments

318

containing Fe2+ ions. This study calls for immediate attention to the significant effects of redox

15


Page 17 of 31


319

reactions and surface precipitation on the fate, transport, and bioavailability of engineered NPs in

320

aquatic environments. These nanoparticles could act as heterogeneous nucleation sites and

321

adsorption sites when released into the environment, incorporating toxic elements and molecules

322

into a “hybrid” engineered/natural nanoparticle composite.54 Redox reactive elements, e.g., Fe2+,

323

Mn2+, and contaminants, such as As, Cr, and U, can be adsorbed on and react with redox reactive

324

engineered NPs such as CeO2 NPs.55, 56 In addition, the positively charged CeO2 can adsorb and

325

aggregate with negatively charged natural colloids or polymers. In the presence of Fe2+, the new

326

hybrid Fe(III) hydroxide coated CeO2 nanoparticles become more positively charged and may take

327

longer to be destabilized by the natural colloids than in the absence of Fe2+ ions. In addition, the

328

interactions of adsorbed natural polymers with Fe2+ (e.g., complexation with natural polymers)

329

can make the redox reaction of CeO2 more complicated.57 The findings of the current study provide

330

an important starting point for investigating the long term influences of Fe2+ on CeO2 fate, transport,

331

and toxicity. These reactions have a significant impact on the transport and transformation of both

332

NPs and contaminants.

333

Consequential changes in the physicochemical properties of CeO2 NPs can affect the

334

toxicity by altering factors such as the surface charge, particle size, production of reactive oxygen

335

species (ROS), dissolution of reactive ions, hydrophilicity, and surface functionality or coatings.

336

For instance, positively charged CeO2 NPs were found to penetrate C. elegans cell membranes

337

more easily than the neutral and negatively charged CeO2 NPs, making them more toxic.8 Because,

338

in general, dissolved Ce3+ is far more toxic than CeO2 NPs,27 the presence of Fe2+can also enhance

339

their toxicity due to increased dissolution. Conversely, the surface of CeO2 NPs may become

340

passivated with time by Fe(III) (hydr)oxide coatings. This would retard further dissolution and

341

could help mitigate CeO2 NP toxicity. CeO2 NP redox processes with Fe2+ and other redox active 16



Page 18 of 31

342

species could also interact with natural organic matter in wastewater treatment plants, potentially

343

affecting their toxicity to microorganisms in wastewater treatment plants.14, 39 In addition to CeO2

344

NPs, many engineered NPs may be subject to surface reactions and precipitation due to ubiquitous

345

redox reactive elements in the environment. These interactions need to be considered

346

comprehensively while evaluating the fate of engineered CeO2 NPs and assessing their risk in

347

aquatic environments.

348

Supporting Information Available

349

Supporting information includes experimental descriptions, UV-Vis spectra, DLS data, zeta

350

potential measurements, TEM images, an ED pattern, and pH monitoring data. This material is

351

available free of charge via the Internet at http://pubs.acs.org.

352

Acknowledgments

353

This work is supported by the National Science Foundation’s Environmental Chemical Science

354

Program (CHE-1214090) and Washington University’s Faculty Startup. JRR was supported by the

355

Environmental Protection Agency STAR Fellowship and CWN was supported by the Mr. and Mrs.

356

Spencer T. Olin Fellowship. We would like to thank Dr. Seonke Seifert of the Advanced Photon

357

Source Sector 12-ID-C at Argonne National Laboratory, supported by the U.S. Department of

358

Energy, Office of Science, Office of Basic Energy Sciences, under Contract No. DE-AC02-

359

06CH11357 and the Institute of Materials Science and Engineering and Nano Research Facility at

360

WUStL for experimental support. XAS work was performed at GeoSoilEnviroCARS (Sector 13),

361

Advanced Photon Source (APS), Argonne National Laboratory. GeoSoilEnviroCARS is supported

362

by the National Science Foundation-Earth Sciences (EAR-1128799) and Department of Energy-

363

GeoSciences (DE-FG02-94ER14466). We thank Dr. Matt Newville and Dr. Tony Lanzirotti for 17


Page 19 of 31


364

their help with XAS experiments. We appreciate the assistance and constructive suggestions from

365

our colleagues in the Environmental NanoChemistry Lab (ENCL) at WUStL.

18



366

TOC Art

367

19


Page 20 of 31

Page 21 of 31


368

LIST OF FIGURES

369

Figure 1.(A) Sedimentation kinetics of CeO2 NPs at 10 mM IS and pH 4.8 in the presence and

370

absence of aqueous Fe2+ ions. The percentages were obtained from suspended nanoparticle

371

concentration normalized by the initial suspended concentration measured by UV-Vis. The error

372

bars represent the standard deviation of CeO2 concentration from triplicate experiments. (B)

373

Dissolved Ce concentrations from CeO2 NPs in the presence and absence of Fe2+ at pH 4.8.

374

Figure 2. XAS spectra for CeO2 NPs reacted in the presence of 3 mM FeCl2 (A: Ce K-edge and

375

B: Fe K-edge). The increase in energy of the K-edge position compared to the FeCl2 standard

376

indicates oxidation of Fe(II) to Fe(III). XAS results showing (A) cerium K-edge and (B) iron K-

377

edge spectra for reacted samples and standards. Ce K-edge results show no detectable Ce(III) on

378

the surface. Fe K-edge results show the most similarity between ferrihydrite and the Fe(III) phase

379

formed on the CeO2 NPs after the 6 h reaction.

380

Figure 3. Representative AFM images for the precipitation of iron oxide particles on (A) CeO2-

381

sputtered Si substrates and (B) pure Si control substrates. Height scale (HS) is 10 nm unless

382

otherwise noted.

383

Figure 4.1D reduced grazing incidence small angle X-ray scattering (GISAXS) data of

384

Fe3+precipitation on the surfaces of CeO2 sputtered wafers ( , Figure 4A) and data from the CeO2

385

NP control system ( , Figure 4A). SAXS raw data of the Fe2+ only system is depicted in Figure

386

4B. Figure 4A revealed that particles were observed in CeO2 + Fe2+ system with Rg = 1.72 nm (at

387

1 hr), and no particles formed in the absence of Fe2+ because there was no scattering increase for

388

the CeO2 only system. Figure 4B (Fe2+ only system) indicates that no particles formed without

389

CeO2 as well. In Figure 4B, no background subtraction was performed due to no increase of 20



390

intensity throughout the reaction and some signal around 0.035 Å-1 results from the beamline setup

391

rather than actual particle contribution.

21


Page 22 of 31

Page 23 of 31


CeO2 conc. (%)

100

3 mM FeCl2 + 1 mM NaCl 10 mM NaCl

80 60 40 20 0

A 0

20

40

60

Time (h)

393

80

100

Dissolved Ce conc. (mg/L)

392

2.0 1.5 1.0 0.5 0.0

B 0

1

2

3

4

5

Time (h)

394 395

Figure 1

22


6


Page 24 of 31

B

Normalized absorbance (a.u.)

A CeO2 standard

CeO2 control 2+ CeO22++Fe Fe(II) CeO

5700

396 397

5750 5800 Energy (eV)

Hematite

Normalized absorbance (a.u.)

Ce(III) Nitrate standard

5850

Goethite Lepidocrocite

Ferrihydrite CeO Fe(II) CeO Fe2+ 2+ 2+

7100

398

Figure 2

23

Fe(II) chloride


7150 7200 Energy (eV)

7250

Page 25 of 31


399

400

Figure 3

401

24



Page 26 of 31

4

1

A

Intensity (counts)

Intensity (counts)

10

3

10

Rg = 1.72 nm

2

10

1

10

0

CeO2 + Fe(II) CeO Fe2+ CeO22 CeO

10 10

-1 6

2

0.01

B 4 2

0.1 4 2

0.01

3 4 56

2

3 4 5

2

0.1

0.01

-1

q_xy (Å )

402 403

Figure 4

25


3

4 5 6

0.1 -1 q_xy (Å )

2

3

4

Page 27 of 31


404

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405

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Enhanced Colloidal Stability of CeO2 Nanoparticles by Ferrous Ions

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