Junmin Deng

Sunho Yoon

Mathieu Pasturel

Sungjun Bae

Hanna Khalil
email: khalil.hanna@ensc-rennes.fr

Interactions between nanoscale zerovalent iron (NZVI) and silver nanoparticles alter the NZVI reactivity in aqueous environments

Keywords: nanoscale zerovalent iron, silver nanoparticles, p-nitrophenol, aggregation

The increasing production and broad application of engineered nanoparticles (ENPs) have led to a substantial release of ENPs into natural systems. Despite their different utilization prospects, ENPs such as nanoscale zero-valent iron (NZVI) and silver nanoparticles (AgNPs) have been co-found in aquatic environments. This study comprehensively examined the effect of AgNPs on the reductive ability of NZVI, which was probed through the conversion of p-nitrophenol (p-NP) to p-aminophenol (p-AP).

Electron microscopic and spectroscopic investigations and Derjaguin-Landau-Verwey-Overbeek (DLVO) calculations showed that the aggregation of NZVI with a low dose of AgNPs dramatically decreased the NZVI reactivity. In contrast, enhanced NZVI reactivity was observed when a high dose of AgNPs was used. The collision between self-assembled aggregates of AgNPs and NZVI can lead to the exposure of inner Fe(0), thereby improving the reductive ability of NZVI. The same behavior was observed for other nanosized (e.g., TiO2) and microsized (e.g., quartz sand) particles, which suggests the predominance of physical processes in NZVI peeling. Further investigations showed that the presence of sulfate and humic acid enhanced the p-NP removal in the NZVI-AgNP suspension, in contrast to other groundwater constituents.

Our findings provide a better understanding of potential ENP interactions, particularly in the context of groundwater remediation.

Graphical Abstract

Interactions between nanoscale zero-valent iron (NZVI) and silver nanoparticles (AgNPs) significantly alter the reductive ability of NZVI

Introduction

Over the past decades, due to increasing production and broad applications of engineered nanoparticles (ENPs), such as textile [START_REF] Lorenz | Characterization of silver release from commercially available functional (nano)textiles[END_REF], electronics (Poudel and Kim, 2022;[START_REF] Poudel | Polyaniline-silver-manganese dioxide nanorod ternary composite for asymmetric supercapacitor with remarkable electrochemical performance[END_REF], pharmaceuticals [START_REF] Sridhar | Electrosprayed nanoparticles and electrospun nanofibers based on natural materials: applications in tissue regeneration, drug delivery and pharmaceuticals[END_REF], environmental remediation (Poudel et al. 2022) and cosmetics [START_REF] Auffan | Structural Degradation at the Surface of a TiO2-Based Nanomaterial Used in Cosmetics[END_REF], a significant amount of ENPs are ended up into natural systems [START_REF] Abbas | Transformation pathways and fate of engineered nanoparticles (ENPs) in distinct interactive environmental compartments: A review[END_REF][START_REF] Danail | Hazards and Risks of Engineered Nanoparticles for the Environment and Human Health[END_REF]. In particular, silver nanoparticles (AgNPs) are one of the most commonly used ENPs. Their catalytic ability, high surface area to volume ratio and antibacterial activity provide them with excellent potential for various industrial [START_REF] Poudel | Silver nanoparticles decorated molybdenum sulfide/tungstate oxide nanorods as high performance supercapacitor electrode[END_REF][START_REF] Shen | Preparation of solid silver nanoparticles for inkjet printed flexible electronics with high conductivity[END_REF], food [START_REF] Echegoyen | Nanoparticle release from nano-silver antimicrobial food containers[END_REF] and medical applications [START_REF] Deng | Mechanistic study of the synergistic antibacterial activity of combined silver nanoparticles and common antibiotics[END_REF]. However, their release and accumulation in the environment (water and soil) [START_REF] Mcgillicuddy | Activated charcoal as a capture material for silver nanoparticles in environmental water samples[END_REF][START_REF] Mahdi | Silver nanoparticles in soil: Aqueous extraction combined with single-particle ICP-MS for detection and characterization[END_REF] is of concern owing to their potential adverse effects on human health and aquatic life [START_REF] Asghari | Toxicity of various silver nanoparticles compared to silver ions in Daphnia magna[END_REF][START_REF] Xu | Silver nanoparticles (AgNPs) cause degeneration of cytoskeleton and disrupt synaptic machinery of cultured cortical neurons[END_REF]. In the meanwhile, nanoscale zero-valent iron (NZVI), as one of the most extensively used ENPs, has been widely applied for groundwater remediation and wastewater treatment because of its high reactivity, costeffectiveness, and environmental friendliness [START_REF] Li | Heavy metal removal using nanoscale zero-valent iron (nZVI): Theory and application[END_REF][START_REF] Phenrat | Electromagnetic Induction of Zerovalent Iron (ZVI) Powder and Nanoscale Zerovalent Iron (NZVI) Particles Enhances Dechlorination of Trichloroethylene in Contaminated Groundwater and Soil: Proof of Concept[END_REF][START_REF] Wei | Influence of nanoscale zero-valent iron on geochemical properties of groundwater and vinyl chloride degradation: A field case study[END_REF][START_REF] Xia | Genesis of pure Se(0) nano-and micro-structures in wastewater with nanoscale zero-valent iron (nZVI)[END_REF]. NZVI can be used to treat various contaminants such as persistent organic pollutants [START_REF] Joo | Oxidative Degradation of the Carbothioate Herbicide, Molinate, Using Nanoscale Zero-Valent Iron[END_REF], toxic inorganic contaminants [START_REF] Ryu | Reduction of highly concentrated nitrate using nanoscale zero-valent iron: Effects of aggregation and catalyst on reactivity[END_REF] and even radioactive nuclides [START_REF] Tsarev | The short-term reduction of uranium by nanoscale zero-valent iron (nZVI): role of oxide shell, reduction mechanism and the formation of U(v)-carbonate phases[END_REF].

When AgNPs are discharged into terrestrial or aquatic environments, their interaction with other co-existing ENPs could influence their environmental behavior, fate, and likely their ecotoxicity. In addition, AgNPs can potentially alter the reactivity of other ENPs used for decontamination, such as NZVI. The injection of NZVI slurry is a common method for the in situ remediations of aquifer and groundwater contamination by organic or inorganic compounds [START_REF] Su | Travel distance and transformation of injected emulsified zerovalent iron nanoparticles in the subsurface during two and half years[END_REF][START_REF] Mueller | Application of nanoscale zero valent iron (NZVI) for groundwater remediation in Europe[END_REF].

Despite the possible co-occurrence of both ENPs in aquatic environments [START_REF] Stefaniuk | Review on nano zerovalent iron (nZVI): From synthesis to environmental applications[END_REF][START_REF] Syafiuddin | Silver Nanoparticles in the Water Environment in Malaysia: Inspection, characterization, removal, modeling, and future perspective[END_REF], there is a lack of knowledge on their potential interactions. Therefore, a comprehensive study on the interactions between NZVI and AgNPs is needed considering the contaminant removal methods used in remediation processes.

Herein, we selected p-nitrophenol (p-NP), as a probe compound because the reductive conversion of p-NP to p-aminophenol (p-AP) is well documented, and the corresponding mass balance is easy to establish owing to the low adsorption of nitrophenol and aminophenol by metal oxides [START_REF] Bae | Effect of NaBH4 on properties of nanoscale zero-valent iron and its catalytic activity for reduction of p-nitrophenol[END_REF][START_REF] Lai | Influence of operating temperature on the reduction of high concentration p-nitrophenol (PNP) by zero valent iron (ZVI)[END_REF][START_REF] Vilardi | P-aminophenol catalysed production on supported nano-magnetite particles in fixed-bed reactor: Kinetic modelling and scale-up[END_REF]. In order to investigate the interactions between NZVI and AgNPs and clearly highlight specific effects of AgNPs on the NZVI reductive ability, we applied a wide concentration range and high-enough dosage of AgNPs (0.5 -100 mg/L) in NZVI suspension. Two concentrations of NZVI (25 and 50 mg/L) within the concentration range typically used in lab-experiments were investigated. The phase transformation and morphological change of NZVI and AgNPs were monitored by X-ray diffraction (XRD), X-ray photoelectron spectroscopy (XPS), and transmission electron microscopy (TEM) with energy-dispersive X-ray spectroscopy (EDS). In addition, a theoretical investigation based on Derjaguin-Landau-Verwey-Overbeek (DLVO) calculations were conducted for the low and high doses of AgNPs. Moreover, to get more insight into the relationship between effects of ENPs and their nature, interactions of NZVI with other widely used ENPs (e.g. nanosized TiO2 particles) and inert mineral surface (e.g. quartz sand) were also investigated. Finally, the NZVI-AgNP interactions were evaluated in the presence of naturally occurring anions and cations and natural organic matter, at levels similar to those encountered in real groundwater.

Materials and methods

Chemicals and materials

Sodium borohydride (NaBH4, ≥99.0%), p-NP (≥99%), p-AP (≥98%), ferric chloride hexahydrate (FeCl3•6H2O, ≥99.0%), hydrochloric acid (HCl, 37%), sodium hydroxide (NaOH, ≥99.0%), sodium chloride (NaCl, ≥99.5%), sodium sulfate (Na2SO4, ≥99.0%), calcium chloride (CaCl2, ≥99.9%), magnesium chloride (MgCl2, ≥98.0%) and sodium metasilicate nonahydrate (Na2SiO3•9H2O, ≥98.0%) were purchased from Sigma-Aldrich, France. Acetonitrile (99.99%, Sigma) and acetic acid (99.7%, ACROS) were used as the mobile phase for the high-performance liquid chromatography (HPLC) analysis. Leonardite humic acid (LHA) standard was purchased from the International Humic Substances Society (IHSS). Silver nano-powder (containing polyvinylpyrrolidone (PVP) as a dispersant, 99.5%, <100 nm, and PZC ~3.3) and TiO2 nano-powder (rutile, 99.5%, 50 m 2 /g, <100 nm, and PZC ~5.9) were purchased from Sigma-Aldrich. Fontainebleau quartz sand (FB sand, 0.06 m 2 /g, 150-300 µm, PZC ~2.7) was obtained from VWR International. Based on a previously published method [START_REF] Deng | Adsorption capacity of the corrosion products of nanoscale zerovalent iron for emerging contaminants[END_REF], NZVI was synthesized by reducing 0.11 M FeCl3•6H2O with 0.9 M NaBH4 in an anaerobic chamber.

The glassware was soaked in 5% (v:v) HCl for at least 48 h and rinsed before use.

Unless specifically stated, all solutions were prepared with deoxygenated deionized water (DDIW, 18.2 MΩ•cm), purged with nitrogen (N2, 99.99%) for 4 h, and stored in an anaerobic chamber (JACOMEX).

Experimental and analytical methods

Batch experiments were conducted in 500 mL flask in an anaerobic chamber under dark conditions. An exact amount of NZVI (0.0125 and 0.025 g) was transferred to the flask containing 500 mL of DDIW, and the mixture was stirred at 500 rpm to prepare the samples with initial NZVI concentrations of 25 and 50 mg/L. To monitor the variations in pH and oxidation-reduction potential (ORP) throughout the experiment, we used a portable pH-ORP meter (HI991003, Hanna). After 1 h, AgNPs were introduced into the suspensions to form a mixture of NZVI and AgNPs at different Ag concentrations (0.5, 2, 10, 25, 50, and 100 mg/L). An equilibration time of 1 hour was chosen because pH and ORP values have reached constant values after 1 h of reaction time.

Although this time is sufficient to obtain an equilibrated NZVI suspension under our experimental conditions, it may not be applicable at the field-scale.

After 15 h, pH and ORP reached an equilibrium, and 2 mL of p-NP stock solution (25 mM) was added to the suspension to initiate the reaction with p-NP (0.1 mM).

Subsequently, 1 mL samples were withdrawn from the suspension and then filtered through a 0.2 μm filter (Whatman) for HPLC analysis. The aqueous concentrations of p-NP and p-AP were determined using HPLC (Waters 600 controller) with a photodiode array detector (Waters 996) and a reversed-phase C18 column (250 mm × 4.6 mm i.d., 5 μm). A mobile phase (acetonitrile: water at 50:50 (v:v)) containing 0.1% formic acid was prepared and used at a flow rate of 1 mL/min in isocratic mode. In addition, mass balance of p-NP and p-AP after reaction was checked and presented in Table S1.

The same procedure was used to investigate the effects of TiO2 nanoparticles, FB sand, and groundwater solutes on the reactivity of NZVI and NZVI/AgNPs. The AgNP dissolution experiments were conducted at pH 9 in absence of NZVI, which is similar to the working pH in presence of NZVI, at AgNP dosage of 0.5, 2, 10, 25, 50, and 100 mg/L. Although the dosage of AgNPs seems to be higher than the frequently detected concentration in natural systems, we have assessed the interactions of NZVI with AgNPs over a wide range of Fe/Ag ratios (0.5 -50) to cover some environmentallyrelevant scenarios.

After 3 d, the samples were filtered through a 0.2 μm filter to measure the amount of Ag + ions by inductively coupled plasma mass spectrometry (ICP-MS, Agilent 7700x).

All experiments were performed in triplicates. The data reported here are the average of three replicated experiments, and the error bars represent the relative standard deviation.

Characterization

XRD (D8, BRUKER) was used to identify the resulting products of NZVI after reaction with p-NP. The resulting suspensions were first collected and washed with DDIW three times, and dried in a vacuum freeze drier (-52 °C, 24 h). The dried samples were then transferred to sample holders and treated with 1:1 (v:v) glycerol solution to avoid surface oxidation during the XRD analysis [START_REF] Bae | Influence of Riboflavin on Nanoscale Zero-Valent Iron Reactivity during the Degradation of Carbon Tetrachloride[END_REF]. XPS (Sigma Probe system, Thermo) analysis was performed for NZVI, AgNPs, and NZVI-AgNP samples. The power source (Al Kα X-ray, 1486.7 eV) was 75 W, and the C 1s peak at 285 eV was used as a reference for the correction of surface charging effects. The morphological characterization of the NZVI and NZVI-AgNP mixtures was performed by TEM-EDS, (JEM-2100, JEOL). The average particle size of the NZVI-AgNP suspension with groundwater solutes and the zeta potential of induvial particles (i.e., NZVI and AgNPs) were measured using a Zetasizer (Matec Applied Sciences). The average hydrodynamic diameter of AgNPs and NZVI were determined using timeresolved dynamic light scattering (DLS, ALV-5000, Langen). The interactions between AgNPs and NZVI were investigated under anoxic conditions by monitoring the variations in pH and ORP throughout the experiment, including at i) the addition of NZVI to the reactor, ii) addition of AgNPs, and iii) removal of p-NP (see Supporting Information (SI), Fig. S1). When NZVI was added to water at 50 mg/L (i.e., first 1 h), there was a clear increase in pH from 6.78 to 9.38 and a decrease in ORP from -54 to -843 mV, which were attributed to the reaction between NZVI and water [START_REF] Bae | Reactivity of nanoscale zero-valent iron in unbuffered systems: effect of pH and Fe (II) dissolution[END_REF]. After the addition of AgNPs at 10 mg/L to the NZVI suspension, the pH and ORP values remained constant, and only slight changes were observed in the pH and ORP after 15 h of reaction. The slight decreasing in pH and increasing in ORP value may be attributed to the consumption of Fe 2+ and OH -by the hydrolysis of ferrous under alkaline conditions [START_REF] Deng | Adsorption capacity of the corrosion products of nanoscale zerovalent iron for emerging contaminants[END_REF]. Afterward, little variations in pH or ORP were observed during the p-NP removal. Only a slight decrease in pH and increase in ORP was observed upon the addition of the p-NP solution.

Results and discussion

Effect of AgNPs on the NZVI reactivity

The removal of p-NP by various materials, including NZVI alone, AgNPs alone, and NZVI/AgNPs mixture, was also evaluated (Fig. 1). Since the standard potential of Ag 0 /Ag + and p-NP/p-AP is 0.79 V and 0.76 V respectively, the conversion of p-NP to p-AP by AgNPs is not likely to take place spontaneously (Hernández-Gordillo et al. [START_REF] Liu | Effect of Particle Age (Fe0 Content) and Solution pH On NZVI Reactivity: H2 Evolution and TCE Dechlorination[END_REF]. However, the working pH was 9 in our experiments, which significantly decrease the generation of hydrogen gas [START_REF] Liu | Effect of Particle Age (Fe0 Content) and Solution pH On NZVI Reactivity: H2 Evolution and TCE Dechlorination[END_REF]. In addition, it has been generally accepted that the active H atom seems to be readily captured by reactive catalysts (e.g., Ni) rather than Fe phase, and reduction of p-NP could not be occurred by dissolved hydrogen alone [START_REF] Du | Synthesis of p-aminophenol from pnitrophenol over nano-sized nickel catalysts[END_REF].

Therefore, hydrogen contribution should be of minor importance in our reaction systems.

The effect of varying AgNP concentrations (0.5-100 mg/L) on the p-NP removal by NZVI (25 and 50 mg/L) was then investigated (Fig. S2). Nearly 72% of p-NP was removed by 25 mg/L of NZVI after 8 h, whereas the addition of 0.5 mg/L of AgNPs to the NZVI suspension decreased the removal to 62%. An increase in the AgNP concentration resulted in a gradual increase in the conversion of p-NP to p-AP (i.e., 64% at 2 mg/L, 84% at 10 mg/L, 89% at 25 mg/L, and 93% at 50 mg/L) (Fig. S2(a)).

Similarly, the removal efficiency of p-NP after 2 h of reaction significantly decreased when 0.5 mg/L of AgNPs was added to a 50 mg/L NZVI suspension, and it continuously increased as the AgNP content increased (Fig. S2(b)).

To further demonstrate the variations in removal rate, we plotted the removal kinetics fitted using the pseudo-first-order equation over the early reaction stage (60 min) and the removal rate constants (kobs) against AgNP dose (Fig. 2 and Fig. S2(c) and (d)). Compared to NZVI alone (0.0085 min -1 at 25 mg/L, and 0.037 min -1 at 50 mg/L), the obtained rate constant was nearly 1.5 and 5.8 times lower after the addition of 0.5 mg/L AgNPs. However, the kobs values gradually increased to 0.0131 and 0.042 min -1 as the AgNP dose increased, at 25 and 50 mg/L NZVI, respectively. Therefore, the results showed that the presence of AgNPs at a low dose significantly inhibited the NZVI reactivity, but this inhibitive effect gradually disappeared as the AgNP concentration increased.

To verify whether this two-step phenomenon was specific to AgNPs, additional investigations were conducted with nanosized TiO2 particles (used as active nanoparticles) and microsized quartz sand particles (used as an inert surface) under dark conditions. For TiO2, the inhibition of p-NP removal was observed at a low dosage (5 mg/L), whereas an increase in TiO2 dose enhanced the p-NP reduction (Fig. S3). It is worth noting that TiO2 alone showed negligible removal of p-NP at pH 9 (See Fig. S4), suggesting that possible adsorption of p-NP on TiO2 can be excluded under our experimental conditions. However, the addition of quartz sand showed a one-step behavior, that is, the removal rate constant continuously increased with increasing sand dose (Fig. S5). Collectively, these results suggest that the increase or decrease in p-NP removal rate was not limited to the effect of AgNPs, and it likely resulted from a combination of physical and chemical processes. The kinetic experiments suggested that the NZVI reactivity can be altered by the interactions between AgNPs and NZVI, which might be caused by potential aggregation and/or redox reactions between the two particles. As previously observed for goethite and hematite [START_REF] Wang | Heteroaggregation and dissolution of silver nanoparticles by iron oxide colloids under environmentally relevant conditions[END_REF], AgNPs might bind to the surface of the oxides layer in the core-shell structure of NZVI. Accordingly, the electrostatic attraction between negatively charged AgNPs (PZC ~3.3) and iron oxides has been used to explain the heteroaggregation process and inhibition of AgNP dissolution [START_REF] Wang | Heteroaggregation and dissolution of silver nanoparticles by iron oxide colloids under environmentally relevant conditions[END_REF]. However, these interactions have been investigated at circumneutral pH values

Heteroaggregation and homoaggregation of AgNPs

(5.5 and 7.5), at which the surface charge of goethite (PZC ~ 9) and hematite (PZC ~ 8) are predominately positive. This emphasizes the role of electrostatic attraction for the adsorption of negatively charged AgNPs onto iron oxides. The zeta potential of NZVI and AgNPs at different pHs were measured to gain more insight into electrostatic interactions between the two types of particle (Fig. S6). However, the working pH of the NZVI suspension lay at 9, at which the surface charge of both NPs are negative.

Indeed, at pH above the PZC of minerals, dissociation of surface functional groups in water will give a negatively charged surface [START_REF] Lützenkirchen | A macroscopic water structure based model for describing charging phenomena at inert hydrophobic surfaces in aqueous electrolyte solutions[END_REF]).

To investigate the interparticle interactions in NZVI-AgNP suspensions under the experimental conditions of the present work, TEM-EDS analyses were conducted. The silver nanoparticles sample used in this study is mostly round shape with particle size in water lower than 100 nm. The NZVI showed chain-like aggregates consisting of spherical particles (60-100 nm) (Fig. 3(a)), which are generally attributed to the high surface energy and intrinsic magnetic interaction among NZVI particles [START_REF] Sohn | Fe (0) nanoparticles for nitrate reduction: stability, reactivity, and transformation[END_REF]. Significantly smaller particles were attached to larger particles (80-100 nm) when 10 mg/L of AgNPs was added (Fig. 3(b)). However, the attachment of small particles was not observed at a high AgNP dose (100 mg/L) (Fig. 3 were observed after the reaction with p-NP (Fig. S7), suggesting no significant modification of the interactions between NZVI and AgNPs.

The enlarged TEM images with EDS mapping revealed a clear difference for different AgNP dosages (Fig. S8). In the absence of AgNPs (Fig. S8(a)), the TEM images of NZVI showed a typical shape of NZVI particles covered by a thin Fe oxide shell. The corresponding elemental mapping (Fe and O) further displayed the core-shell structure of NZVI, showing a strong signal intensity of Fe in the core, whereas O was mostly present in the shell layer (2.5-3.5 nm). However, the thickness of the NZVI oxide shell seems to decrease significantly after the addition of 100 mg/L AgNPs, whereas almost no change was observed with the addition of 10 mg/L AgNPs. Based on the TEM-EDS results, we hypothesized that AgNPs at a low dose attach to the NZVI surfaces (i.e., iron oxides), and at a high dose, they tend to homoaggregate. We also hypothesized that introducing a high dosage of AgNP would lead to a decrease in the thickness of the NZVI oxide layer, which should be further investigated. We also analyzed the attraction and repulsion behaviors of AgNPs and NZVI according to the DLVO theory. This theory considers attractive van der Waals (vdW)

and repulsive electrostatic double-layer (EDL) forces between the particles (ΦDLVO = ΦEDL + ΦvdW). The details of DLVO calculations are provided in the SI. In brief, at a low AgNP dose (10 mg/L), heteroaggregation (AgNP-NZVI) is likely to occur because of the lower repulsive energy barriers of AgNP-NZVI with respect to AgNP-AgNP. In addition, the AgNPs used in this study also contained stabilized PVP, so the surface coatings could serve as a bridge between coated AgNPs and uncoated NZVI to overcome the energy barrier [START_REF] Lin | Polymeric Coatings on Silver Nanoparticles Hinder Autoaggregation but Enhance Attachment to Uncoated Surfaces[END_REF]. Therefore, it is likely that PVP-stabilized AgNPs formed heteroaggregates with NZVI. In contrast, a higher AgNP dose (100 mg/L) led to a significantly higher magnitude of calculated maximum repulsive interaction energies (Φmax), which would not be sufficient to overcome non-DLVO energy. According to the calculations, neither heteroaggregation nor homoaggregation (AgNP-AgNP) would occur spontaneously. However, AgNP homoaggregation was still observed using TEM-EDS. This discrepancy between DLVO calculations and TEM images might be attributed to the potential chemical interactions between AgNPs and NZVI or oxide layer coating NZVI.

A recent study reported that AgNP morphology may change from a dispersed spherical shape to larger aggregates owing to the release of Ag + possibly because of (i) the nano bridges created by adsorption and/or redox conversion of Ag + into Ag 0 at the particle surface to connect nearby particles, and (ii) the growth of reformed small AgNPs around old AgNPs [START_REF] Yu | Highly dynamic PVP-coated silver nanoparticles in aquatic environments: chemical and morphology change induced by oxidation of Ag(0) and reduction of Ag(+)[END_REF]. Other reports showed that Ag + released from AgNPs and possible electron transfer from Fe 0 to Ag + could form a galvanic couple, which may accelerate NZVI-water corrosion [START_REF] Zhuang | Dehalogenation of polybrominated diphenyl ethers and polychlorinated biphenyl by bimetallic, impregnated, and nanoscale zerovalent iron[END_REF][START_REF] Zhu | Rapid Degradation of Carbon Tetrachloride by Microscale Ag/Fe Bimetallic Particles[END_REF]. Therefore, interactions of Ag + ions (E0 Ag/Ag + 0.799 V) with NZVI (E0 Fe/Fe 2+ 0.44 V) and AgNPs through adsorption and/or electron transfer processes may alter the particle aggregation behavior and surface reactivity. In the present work, we have measured Ag + ions in AgNP aqueous suspension but in the absence of NZVI, to avoid sorption of released Ag+ ions on iron oxide layer of NZVI surface. As a result, these dissolution tests were performed under oxic (open atmosphere) or mild reducing conditions (anaerobic chamber), since the strong reducing conditions are ensured by the presence of NZVI. Under such conditions, Ag + can be detected, and released concentration was increased with increasing AgNP dose (See Fig. S10). However, examination of the thermodynamic database by means of Eh-pH diagrams of Ag species suggests that ions released through metal corrosion should be very limited, or even excluded, under the experimental conditions of this study (pH 9, -800 mV) [START_REF] Brookins | Eh-pH diagrams of geochemistry[END_REF]. Indeed, no ions can be detected at alkaline pH values and strongly reducing conditions, thereby ruling out galvanic effects between Ag + ions and inner Fe 0 within iron oxide shell. In addition, the same two-step behavior in NP removal kinetics by NZVI was observed when AgNP has been replaced by TiO2 (Fig. S3), further confirming the minor role of ions release and/or galvanic effect in the investigated systems.

Collectively, these results suggest that heteroaggregation of NZVI-AgNPs which is favorable at lower AgNP dose owing to the relatively lower energy barrier may alter the NZVI reactivity. As the AgNP dose increased, AgNP homoaggregation can be expected, thereby preventing potential alteration of the initial NZVI reactivity.

Spectroscopic investigations

To investigate the oxidation states of Ag and Fe in AgNP and NZVI suspensions, the XPS spectra of the Ag(3d) and Fe(2p) bands of AgNP, NZVI, and NZVI-AgNP suspensions were analyzed (Fig. 4). For AgNPs and NZVI-AgNPs, the Ag(3d5/2) and Ag(3d3/2) peaks at 368.6 and 374.6 eV were assigned to Ag(0) (Fig. 4(a)) [START_REF] Xu | The preparation and antibacterial effects of dopa-cotton/AgNPs[END_REF], and the splitting of the 3d doublet of Ag (6.0 eV) confirmed the presence of Ag(0) [START_REF] Chastain | Handbook of X-ray photoelectron spectroscopy[END_REF][START_REF] Xu | The preparation and antibacterial effects of dopa-cotton/AgNPs[END_REF]. In the pristine NZVI, two distinct peaks were observed at 719.9 and 706.9 eV, which corresponded to the Fe(2p1/2) and Fe(2p3/2) of Fe(0) [START_REF] Luo | Synthesis of nanoscale zero-valent iron immobilized in alginate microcapsules for removal of Pb(ii) from aqueous solution[END_REF], respectively (Fig. 4(b)). In the NZVI-AgNP suspension, Fe(0) peaks were enhanced, indicating the presence of more Fe(0) on the NZVI surface after AgNPs were introduced [START_REF] Li | Solvent-free production of nanoscale zero-valent iron (nZVI) with precision milling[END_REF]. The XPS spectra of Fe(2p3/2) (Fig. 4(c)) showed five peaks at 706.9, 708.2-708.5, 710.3-710.7, 711.5-711.6, and 713.1-713.7 eV, which were assigned to the binding energies of Fe(0) (706.5-707.0 eV), Fe(II)-O (708.2-710.9 eV), and Fe(III)-O (711.0-714.0 eV). The proportion of Fe(0) (10.7%) and Fe(II) (68.1%) on the NZVI surfaces significantly decreased to 3.7% and 58.2% after adding 10 mg/L AgNPs, respectively (Fig. 4(d)). In contrast, their percentages dramatically increased to 11.9% and 84.6% in the sample of NZVI with 100 mg/L AgNPs. These results demonstrate that the presence of AgNPs decreased and increased the Fe(0) content on the NZVI surface, which are in agreement with the results of degradation kinetics in Fig. 2(b). This is also consistent with the TEM-EDS results (Fig. S8(c)), which showed a thinner oxide layer on the NZVI surface at a higher AgNP dose. Previous investigations have shown that the collision and friction between particles might occur in suspensions under stirring conditions [START_REF] Xie | Modeling and simulation of the influences of particle-particle interactions on dense solid-liquid suspensions in stirred vessels[END_REF]. In addition, these particleparticle interactions become much stronger when both particle size and solid loading increased [START_REF] Xie | Modeling and simulation of the influences of particle-particle interactions on dense solid-liquid suspensions in stirred vessels[END_REF]. In our systems, higher AgNP doses enabled homoaggregation and an increase in aggregate size (Fig. S11), likely inducing more collision between large-sized AgNP aggregates and NZVI. This phenomenon likely contributed to the peeling of the NZVI surface layer, which further resulted in the exposure of the NZVI inner core. These results suggest that a higher AgNP dose led to the partial removal of the iron oxide shell of NZVI through collisions between AgNP homoaggregates and NZVI. Self-assembled AgNPs with bigger sizes accelerated the peeling of the Fe oxide shell on the NZVI surface, thereby increasing the exposure of inner Fe(0), and consequently, enhancing the p-NP removal. This also explains the enhanced NZVI reactivity under an increasing dose of sand microparticles (Fig. S5).

As sand microparticles are inert and have a low surface area (0.06 m 2 /g), the enhanced NZVI reductive ability may be attributed to collision-driven processes in an aqueous suspension. In the environment, the concentration of AgNPs is normally much lower than those of NZVI used for remediation purposes [START_REF] Su | Travel distance and transformation of injected emulsified zerovalent iron nanoparticles in the subsurface during two and half years[END_REF][START_REF] Mcgillicuddy | Activated charcoal as a capture material for silver nanoparticles in environmental water samples[END_REF], the inhibition effect of AgNPs to NZVI would be more significant in natural systems. To evaluate the inhibition effects of AgNPs on the NZVI reactivity in conditions that emulate natural systems, the p-NP reduction in NZVI (50 mg/L) and AgNP (10 mg/L) suspensions was investigated in the presence of naturally occurring cations, anions, and organic matter. Experiments were conducted with groundwater constituents and humic acids (HAs) at levels similar to those encountered in real groundwater (Fig. 5(a)) [START_REF] Edmunds | Natural groundwater quality[END_REF]. No impact of NaCl was observed. In the presence of most groundwater solutes, NZVI exhibited good performance for p-NP removal, with 100% removal over 2 h, except in the presence of NO3 -(44%) and SiO3 2-(84%). For NO3 -, the inhibitive result could be caused by a competitive reduction of NO3 -by NZVI [START_REF] Ryu | Reduction of highly concentrated nitrate using nanoscale zero-valent iron: Effects of aggregation and catalyst on reactivity[END_REF]. Compared to NZVI-AgNPs in pure water (85%), the removal efficiency of p-NP by the NZVI-AgNPs with NO3 -, SiO3 2-, Ca 2+ , and Mg 2+ decreased to 21%, 67%, 73%, and 83%, respectively. Furthermore, p-NP was completely removed over 2 h in the presence of SO4 2-and HA, which represents a better performance than that obtained in pure water.

Effects of groundwater constituents on nanoparticles interaction and reactivity

To further clarify the aggregation behavior of the particles, we measured the average particle sizes of NZVI and AgNP-NZVI suspensions in the presence of SO4 2-, SiO3 2-, Ca 2+ , Mg 2+ , and HA (Fig. 5(b)). As expected, the particles sizes in the NZVI-AgNP systems were higher than in systems without AgNPs, which was attributed to the heteroaggregation between AgNPs and NZVI particles. Upon the introduction of Ca 2+

or Mg 2+ , the particle size drastically increased from 593 nm to 1804 and 1727 nm in the NZVI system, and from 1313 nm to 2382 and 2186 nm in the NZVI-AgNP system, respectively. This was attributed to the particle surface charge screening (EDL contraction) by Ca 2+ /Mg 2+ [START_REF] Rogozhnikov | Effect of Alkali Metal Cations on the EDL Structure at a Gold Electrode in Alkali Solutions[END_REF]. EDL contraction can decrease the electrostatic or steric repulsive forces of the nanoparticles, thereby increasing aggregation [START_REF] Petosa | Aggregation and deposition of engineered nanomaterials in aquatic environments: role of physicochemical interactions[END_REF] and further decreasing the NZVI reactivity [START_REF] Ryu | Reduction of highly concentrated nitrate using nanoscale zero-valent iron: Effects of aggregation and catalyst on reactivity[END_REF]. Moreover, Ca 2+ and Mg 2+ showed stronger inhibitive effects in NZVI-AgNPs than in NZVI alone. In the presence of Ca 2+ or Mg 2+ , electrostatic repulsive forces among particles (i.e., NZVI-NZVI, NZVI-AgNP, AgNP-AgNP) favored homoaggregation and heteroaggregation, thus altering the NZVI reactivity.

In contrast, a decrease in the average particle size of NZVI (from 593 to 486 nm)

and NZVI-AgNPs (from 1313 to 877 nm) in the presence of HA was attributed to the interactions of HA with the surfaces, which likely enhanced the electrostatic repulsion decreased [START_REF] Dong | Influence of humic acid on the colloidal stability of surfacemodified nano zero-valent iron[END_REF]. This reduced particle aggregation, which enhanced the removal of p-NP. The particle size decreased more for NZVI-AgNPs than for NZVI, which suggests that less AgNPs were attached to NZVI surfaces, thereby providing a lower inhibition effect of AgNPs in the presence of HA.

XRD analysis was also conducted to analyze the passivation byproducts of NZVI-AgNPs after the reaction (Fig. 5(c)). The byproducts in the pure water system included maghemite (γ-Fe2O3)/magnetite (Fe3O4) and silver (Ag(0)), which were attributed to NZVI oxidation during the p-NP reduction [START_REF] Xiong | Enhancing the efficiency of zero valent iron by electrolysis: Performance and reaction mechanism[END_REF]. Similarly, in the presence of HA, maghemite/magnetite and silver (Ag(0)) phases were also detected after p-NP removal. In contrast, green rust (

[Fe 2+ 4Fe 3+ 2(HO -)12] 2+ •[SO 2- 4•2H2O] 2-
) [START_REF] Tahawy | Exceptionally stable green rust, a mixed-valent iron-layered double hydroxide, as an efficient solar photocatalyst for H2 production from ammonia borane[END_REF] and silver (Ag(0)) were detected in the presence of SO4 2-.

Considering the relatively higher reactivity of green rust than that of magnetite [START_REF] Digiacomo | Immobilization of Cr(VI) by sulphate green rust and sulphidized nanoscale zerovalent iron in sand media: batch and column studies[END_REF][START_REF] Usman | Magnetite and Green Rust: Synthesis, Properties, and Environmental Applications of Mixed-Valent Iron Minerals[END_REF], the enhanced p-NP removal in the NZVI-AgNP-SO4 2-system (Fig. 5(a)) was attributed to the generated green rust. For SiO3 2-, a weak peak of silver was identified as the final product, and there was no obvious Fe phase detected in the XRD analysis. The absence of Fe phases might be attributed to the amorphous coating of Si on the NZVI surface as previously reported, and the surface coated Si would further reduce the amount of active surface sites available for contaminants [START_REF] Meng | Combined effects of anions on arsenic removal by iron hydroxides[END_REF]. Despite the lower particle size of NZVI-AgNP-SiO3 2-, the reactivity of both NZVI and NZVI/AgNPs decreased drastically, which was attributed to the strong binding of silicates to the NZVI surface.

Conclusions

Because of the increasing use of engineered nanoparticles (ENPs) in various industries and consumer items, ENPs are ubiquitously found in aquatic environments.

However, little is known about the interactions among the co-existing ENPs in nature.

This work showed for the first time how two typical engineered nanoparticles, nanoscale zero-valent iron (NZVI) and silver nanoparticles (AgNPs), interact and influence each other under environmentally-relevant conditions. First, we have investigated the potential interactions between AgNPs and NZVI particles using p-NP conversion into p-AP as a probe reaction. At a AgNP load of 10 mg/L, AgNPs covered the reactive NZVI surfaces, which decreased p-NP removal. Whereas at a AgNP load of 100 mg/L, the enhanced homoaggregation of AgNPs led to the exposure of the inner Fe(0) surface of NZVI because of particle collisions, which increased the p-NP removal.

The same behavior was observed with nanosized TiO2 particles and microsized quartz sand particles, thereby excluding the role of potential Ag + ions release, electron transfer process, and/or galvanic effect.

Second, the impact of groundwater constitutes has been evaluated, which showed contrasting effects depending on ion types and the binding affinities to Fe-oxides. These results revealed that not only AgNPs but also other nano/microsized particles (e.g. TiO2 and quartz sand) may influence the NZVI reactivity, which will help to develop decision support and prediction tools for NZVI remediation technologies. Electron microscopic and X-ray spectroscopic investigations and DLVO calculations revealed that heteroaggregation of NZVI and AgNPs resulted in dramatic decrease of the NZVI reactivity, while at higher AgNPs loading, collision between self-assembled aggregates of AgNPs and NZVI can improve the reductive ability of NZVI. These findings shed light on an overlooked aspect of interactions between co-existing nanoparticles, which could influence their fate and mobility in environmental settings. Owing to the increasing occurrence of NPs in aqueous environments, this work calls for more investigations of potential interactions between the co-occurring nanomaterials for accurate assessment of environmental associated risks.

Figure 1 .

 1 Figure 1. Concentrations of p-NP during the removal of p-NP by different nanoparticles. Experimental conditions: [p-NP]initial = 0.1 mM; [NZVI] = 50 mg/L; [AgNPs] = 10 mg/L; pH = 9.0±0.2.

  2014). Indeed, no p-NP removal was observed in the AgNP suspension, which indicated negligible adsorption and/or reduction of p-NP by AgNPs. In contrast, the complete removal of p-NP was observed after 90 min of reaction with NZVI, whereas 85% of p-NP was removed by the NZVI-AgNP suspension over 120 min. The mass balance was evaluated, which confirmed the reductive conversion of p-NP to p-AP (not shown) and the negligible adsorption of p-NP and p-AP on the NZVI surface. It is reported that the hydrogen production caused by anaerobic corrosion of NZVI could potentially affect the reduction of target contaminants

Figure 2 .

 2 Figure 2. Variation in pseudo-first-order kinetic rate constants (kobs, min -1 ) of the

  (c)). Instead, we observed the co-existence of chain-like aggregates and nanoparticle assemblages. The EDS mapping of NZVI-AgNP mixtures showed distinct Fe and Ag nanoparticles highlighted in red and green, respectively, inFigs. 3(d) and (e). As illustrated in Fig.3(d), the Ag-phase was enriched near the surface of Fe aggregates at low AgNP concentrations, which suggests the heteroaggregation of NZVI and AgNPs. The spatial correlation of Ag and Fe was not observed at high AgNP concentrations (Fig.3(e)), but these two elements existed in different locations because AgNPs tended to homoaggregate rather than interact with NZVI at higher AgNP doses. It was previously reported that aggregation behavior of NPs would change along with their oxidation[START_REF] Vilardi | Mathematical modelling of simultaneous nitrate and dissolved oxygen reduction by Cu-nZVI using a bi-component shrinking core model[END_REF], Vilardi et al. 2019). However, similar TEM images of AgNPs and NZVI

Figure 3 .

 3 Figure 3. TEM images of (a) NZVI, (b) AgNP-NZVI at 10 mg/L AgNPs, and (c)

Figure 4 .

 4 Figure 4. XPS spectra of (a) Ag(3d) in AgNP and NZVI-AgNP suspensions, (b) Fe(2p)

Figure 5 .

 5 Figure 5. (a) Removal efficiencies of p-NP (2 h reaction) by NZVI and NZVI+AgNPs

Acknowledgments

We would like to acknowledge the support from Korea Environment Industry & Technology Institute (KEITI) through "Subsurface Environmental Management (SEM) Project", funded by Korea Ministry of Environment (MOE) (2020002480006). We are also grateful for the assistance of Isabelle Soutrel (LC/UV) and Vincent Dorcet (THEMIS platform).

Conflicts of interest

There are no conflicts to declare.


Qinzhi Li

Dieter Schild

Mathieu Pasturel

Johannes Lützenkirchen

Khalil Hanna
email: khalil.hanna@ensc-rennes.fr

Alteration of Birnessite Reactivity in Dynamic Anoxic/Oxic Environments

Keywords: manganese oxide, redox, antibiotics, oxidation, removal

Although the oxidative capacity of manganese oxides has been widely investigated, potential changes of the surface reactivity in dynamic anoxic/oxic environments have been often overlooked. In this study, we showed that the reactivity of layer structured manganese oxide (birnessite) was highly sensitive to variable redox conditions within environmentally relevant ranges of pH (4.0 -8.0), ionic strength (0 to 100 mM NaCl) and Mn(II)/MnO2 molar ratio (0 -0.58) using ofloxacine (OFL), a typical antibiotic, as a target contaminant. In oxic conditions, OFL removal was enhanced relative to anoxic environments under alkaline conditions. Surface-catalyzed oxidation of Mn(II) enabled the formation of more reactive Mn(III) sites for OFL oxidation. However, an increase in Mn(II)/MnO2 molar ratio suppressed MnO2 reactivity, probably because of competitive binding between Mn(II) and OFL and/or modification in MnO2 surface charge. Monovalent cations (e.g., Na + ) may compensate the charge deficiency caused by the presence of Mn(III), and affect the aggregation of MnO2 particles, particularly under oxic conditions. An enhancement in the removal efficiency of OFL was then confirmed in the dynamic two-step anoxic/oxic process, which emulates oscillating redox conditions in environmental settings. These findings call for a thorough examination of the reactivity changes at environmental mineral surfaces (e.g., MnO2) in natural systems that may be subjected to alternation between anaerobic and oxygenated conditions.

Introduction

Manganese dioxides (MnO2) are ubiquitous in a wide range of aquatic and terrestrial environments [START_REF] Post | Manganese oxide minerals: crystal structures and economic and environmental significance[END_REF]. They are widely used in various energy and environmental applications, because of their specific nanoscale properties, high surface areas, low cost, strong oxidation and sorption abilities [START_REF] Post | Manganese oxide minerals: crystal structures and economic and environmental significance[END_REF][START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF].

The layer structured manganese oxide (e.g., birnessite) is structurally similar to biogenically precipitated natural manganese oxides, and thus the common natural occurring MnO2 in surficial environments. It basically consists of randomly stacked edge-sharing MnO6 octahedra and has a large excess of negative charges because of the Mn(IV) vacancies and substitution of Mn(II)/(III) for Mn(IV) [START_REF] Elzinga | Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF]. As a result, MnO2 exhibits a high affinity for cation sorption, but also a non-negligible anion adsorption capacity. Another important feature of MnO2 is the ease of switching the oxidation state, which imparts a good catalytic activity [START_REF] Post | Manganese oxide minerals: crystal structures and economic and environmental significance[END_REF][START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF].

In oxic environments, it has been reported that MnO2 surfaces can catalyze the oxidation of Mn(II) by dissolved oxygen to generate Mn(III)/(IV) [START_REF] Yang | Trivalent manganese on vacancies triggers rapid transformation of layered to tunneled manganese oxides (TMOs): Implications for occurrence of TMOs in low-temperature environment[END_REF][START_REF] Zhao | Redox Reactions between Mn(II) and Hexagonal Birnessite Change Its Layer Symmetry[END_REF], while homogeneous oxidation of Mn(II) by oxygen is considered kinetically sluggish [START_REF] Davies | Manganese(II) Oxidation Kinetics on Metal Oxide Surfaces[END_REF][START_REF] Diem | Is dissolved Mn 2+ being oxidized by O2 in absence of Mn-bacteria or surface catalysts?[END_REF][START_REF] Lan | Mechanisms of Mn(II) catalytic oxidation on ferrihydrite surfaces and the formation of manganese (oxyhydr)oxides[END_REF][START_REF] Nico | Rapid photo-oxidation of Mn(II) mediated by humic substances[END_REF][START_REF] Wilson | Surface and complexation effects on the rate of Mn(II) oxidation in natural waters[END_REF]. Surface oxygen ligands bound to inner-sphere adsorbed Mn(II) may accelerate the electron transfer between Mn(II) and O2 [START_REF] Junta | Manganese (II) oxidation at mineral surfaces: A microscopic and spectroscopic study[END_REF][START_REF] Lan | Mechanisms of Mn(II) catalytic oxidation on ferrihydrite surfaces and the formation of manganese (oxyhydr)oxides[END_REF][START_REF] Luther | Manganese(II) Oxidation and Mn(IV) Reduction in the Environment-Two One-Electron Transfer Steps Versus a Single Two-Electron Step[END_REF][START_REF] Namgung | Macroscopic and microscopic behaviors of Mn(II) (ad)sorption to goethite with the effects of dissolved carbonates under anoxic conditions[END_REF], leading to higher Mn(II) removal under oxic conditions compare to anoxic conditions at circumneutral to alkaline pH [START_REF] Elzinga | Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF][START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF][START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF]. The accumulated Mn(III) can incorporate into Mn octahedral layers and induce the transformation of MnO2 [START_REF] Hinkle | Impact of Mn(II)-Manganese Oxide Reactions on Ni and Zn Speciation[END_REF]Wang et al., 2018a).

Although the influence of phase transformation on the reactivity of birnessite has been already investigated [START_REF] Lefkowitz | Impacts of Aqueous Mn(II) on the Sorption of Zn(II) by Hexagonal Birnessite[END_REF]Wang et al., 2018a;[START_REF] Wang | Effects of metal cations on coupled birnessite structural transformation and natural organic matter adsorption and oxidation[END_REF], very little knowledge exists about the impact of oxygen on the removal of organic compounds interacted with MnO2 surfaces. Furthermore, MnO2 may co-exist with dissolved Mn(II) and emerging contaminants in natural systems which may be subjected to alternation between anaerobic and oxic conditions, such as water-sediment interfaces, water table fluctuations, groundwater recharge [START_REF] Jeong | Aerobic oxidation of mackinawite (FeS) and its environmental implication for arsenic mobilization[END_REF][START_REF] Wu | The Fate of p-Nitrophenol in Goethite-Rich and Sulfide-Containing Dynamic Anoxic/Oxic Environments[END_REF][START_REF] Zhao | Arsenic (III) removal by mechanochemically sulfidated microscale zero valent iron under anoxic and oxic conditions[END_REF]. However, little is known about the potential changes in the surface reactivity of MnO2 in environmental settings under alternating redox conditions. This work has examined for the first time how a shift of redox conditions from anoxic into oxic may alter the reactivity of birnessite.

Ofloxacin (OFL) was chosen here as a target contaminant in order to explore the changes in MnO2 reactivity under dynamic anoxic/oxic conditions. OFL is widely used in human and veterinary medicine, and frequently detected in surface water, groundwater, soils and municipal wastewater systems, with concentration ranging from ng/L to µg/L [START_REF] Zhang | Oxidative Transformation of Fluoroquinolone Antibacterial Agents and Structurally Related Amines by Manganese Oxide[END_REF]. Removal kinetics of OFL in the presence of acid birnessite were performed at pH ranging from 4.0 to 8.0, representing the common pH conditions of suboxic environments [START_REF] Elzinga | 54Mn radiotracers demonstrate continuous dissolution and reprecipitation of vernadite (δ-MnO2) during interaction with aqueous Mn (II)[END_REF]Wang et al., 2018a). The oxidation byproducts of OFL were determined using an ultra performance liquid chromatography-tandem mass spectrometry (UPLC-MS/MS). The oxidative capacity of MnO2 was further monitored for variable Mn(II)/MnO2 molar ratios and ionic strengths under environmentally relevant conditions. In addition, X-ray diffraction (XRD), X-ray photoelectron spectroscopy (XPS) and scanning electron microscopy (SEM) were used to identify the changes in Mn valence states and mineral phase transformation. This work calls for more consideration of alteration of the oxidative capacity of birnessite in redox transition zone in groundwater systems.

Materials and Methods

Materials and Chemicals

All chemicals employed in this study were purchased from Sigma-Aldrich of analytical grade and used without further purification. Ultrapure water (specific resistivity, 18.2 MΩ cm -1 , Milli-Q, Millipore) was used for all solutions and suspensions.

Acid birnessite was synthesized by reacting KMnO4 and HCl as described in previous studies [START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF][START_REF] Mckenzie | The synthesis of birnessite, cryptomelane, and some other oxides and hydroxides of manganese[END_REF][START_REF] Pokharel | Water Flow and Dissolved MnII Alter Transformation of Pipemidic Acid by Manganese Oxide[END_REF]. The purity of mineral phase was confirmed with XRD and SEM (Figure S1). The specific surface area (SSA) measured by Brunauer-Emmett-Teller (BET)-N2 adsorption was 85.0 ± 3.0 m 2 g -1 and the isoelectric point (pHIEP) was around 2.0, respectively (Figure S2). Analysis of XPS data showed that the average oxidation sate of the pristine material was 3.87, with 88.3 % Mn(IV), 10.7 % Mn(III) and 1.0 % Mn(II) (Figure S3).

Batch experiments

All batch experiments were carried out in 500 mL polyethylene bottles, containing varying concentrations of acid birnessite. The solutions were spiked with aliquots of 1.0 mM OFL and 0.05 M MnCl2 stock solution to achieve initial concentration of 10 μM OFL and variable target Mn(II) concentrations ranging from 0 to 200 μM. The ionic strength was fixed using NaCl at various concentrations. The pH of the suspensions was adjusted to the desired values over the reaction time using an auto pH titrator with addition of small volumes of HCl (0.1 M) or NaOH (0.1 M) solutions and maintained at pH 4.0, 6.0 and 8.0 throughout the experiments. No pH buffers were used, to prevent any impact on the birnessite reactivity. The stirring speed of the mixture was maintained at 350 rpm. Batch experiments were carried out under ambient air atmosphere or by bubbling oxygen gas (oxic conditions), or by bubbling nitrogen gas (anoxic conditions).

A two-step anoxic/oxic dynamic process was investigated as follows: after a period of anoxic conditions where the bottle was under N2, an oxic environment was subsequently created by removing the cap of the bottle or by bubbling O2. The same kinetic behavior was observed when either the suspension was stirred under open atmosphere (ambient air) or with bubbling O2 (data not shown). In the open system, CO2 from the atmosphere will generate carbonates in solution at high pH values, but this effect on the reactivity can be excluded according to a previous study [START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF].

Aliquots of solution were periodically withdrawn and filtered through a 0.20 μm membrane filter for analysis of OFL concentration, whereas other aliquots without filtration were immediately mixed with 0.1 M ascorbic acid to quench the reaction. At the end of experiments, the remaining suspensions were centrifuged at 4500 rpm for 20 min and settled particles were freeze-dried, and then stored at 4 ℃ for further solid analysis. In another experimental series involving mixing with 1.0 mM sodium pyrophosphate, a volume of about 3.0 mL sample was periodically withdrawn from the mixed solution, filtered and analyzed immediately by UV-Vis spectrophotometry in the range of 200-600 nm wavelength. Batch experiments were conducted in triplicates and the standard deviation was calculated for each experimental series.

Analytical methods

The concentration of OFL in the samples was analyzed via High Performance Liquid Chromatography analysis with UV-Vis detection (HPLC-UV) equipped with a reversed-phase C18 column (250 mm × 4.6 mm i.d., 5 μm) and a UV-Vis detector (Waters 2489) at a wavelength of 287 nm. The mobile phase consisted of acetonitrile/water (15/85 v/v,) containing 0.1% formic acid at a flow rate of 0.6 mL/min. Filtered suspensions were digested with 2% nitric acid and dissolved Mn(II) concentrations were determined by Atomic Absorption Spectroscopy (AAS, AA140, Varian, Shimadzu) with a detection limit of 0.02 μM.

The oxidation byproducts of OFL were identified using an ultra performance liquid chromatography-tandem mass spectrometry (UPLC-MS/MS) system. An electrospray interface was used for the MS measurements in positive ionisation mode and full scan acquisition.

XPS spectra were recorded on an XPS system PHI 5000 VersaProbe II (ULVAC-PHI Inc.) equipped with a scanning microprobe X-ray source (monochromatic Al Kα, hν = 1486.7 eV). An electron flood gun generating low energy electrons (1.1 eV) and low energy argon ions (6 eV) by a floating ion gun were applied for charge compensation at isolating samples (dual beam technique). Survey scans were recorded with an X-ray source power of 31 W, X-ray spot size diameter 200 µm, and pass energy of the analyzer of 187.85 eV. Narrow scans of the elemental lines were recorded at 23.5 eV pass energy, which yields an energy resolution of 0.69 eV FWHM at the Ag 3d5/2 elemental line of pure silver. Calibration of the binding energy scale of the spectrometer was performed using well-established binding energies of elemental lines of pure metals (monochromatic Al Kα: Cu 2p3/2 at 932.62 eV, Au 4f7/2 at 83.96 eV) [START_REF] Seah | XPS: Binding Energy Calibration of Electron Spectrometers 5ÈRe-evaluation of the Reference Energies[END_REF].

The error in binding energies of elemental lines is estimated to ± 0.1 eV for conductors and ± 0.2 eV for isolating samples. High-resolution XPS spectra were calibrated by the carbon deposit C1s binding energy at 284.8 eV and further analyzed by CasaXPS software. After Shirley background subtractions, all peaks were fitted using 70:30 Gaussian: Lorentzian sum function.

A FEI Quanta 650 FEG environmental scanning electron microscope (now Thermo Fisher Scientific Inc.) was applied to image the sample surfaces.

XRD data were collected by using a D8 Advance (Bruker) diffractometer with Cu Kα radiation over the range of 10-80° 2θ at a step size of 0.02° and phases were determined according to the PDF database in Jade 6 software.

Results and Discussion

Effect of oxygen and pH on the initial removal rate of OFL by MnO2

Removal kinetics of OFL were investigated under anoxic and oxic conditions as a function of MnO2 concentration (173-870 µM or 15 -75.6 mg/L) at three pH values (4.0, 6.0, 8.0 (± 0.1)) over 8 hours of reaction time (Figure S4). The entire kinetic curve could not be properly described by simple equations that include classical exponential functions (e.g., first-or second-order model), probably due to the complexity of involved reactions including accumulation and competition of reaction products (e.g. OFL byproducts, Mn(II) ions, etc.) and/or a gradual change of the reactivity of surface sites. For the sake of simplicity and to overcome possible interference caused by the accumulated byproducts, we have determined the rate constant over the first stage of reaction (i.e., 1h). The initial rate constants k (h -1 ) calculated by linear regression of ln The rate constants exhibited the same order, pH 4.0 > pH 6.0 > pH 8.0, regardless of MnO2 concentration. This is consistent with previous reports [START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF][START_REF] Zhang | Oxidative Transformation of Fluoroquinolone Antibacterial Agents and Structurally Related Amines by Manganese Oxide[END_REF][START_REF] Zhang | Kinetic Modeling of Oxidation of Antibacterial Agents by Manganese Oxide[END_REF], which showed that acidic conditions enabled higher quinolones removal in the presence of MnO2. This can be ascribed to more favorable interactions between the less negatively charged surface (pHIEP of MnO2 ~ 2.0) and OFL (pKa 6.08 and 8.25, Figure S5) at acidic pH values. Furthermore, the decrease in redox potential of MnO2 from 0.99 V to 0.76 V when the pH increased from 4.0 to 8.0 would also explain the enhanced reactivity under acidic conditions, as previously reported [START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF].

Mass balance showed that adsorption and oxidation are involved in the removal of OFL in the presence of MnO2, even though the adsorption appears to be relatively weak for our experimental conditions (e.g., less than 1.5 µmol m -2 , at pH 6.0; less than 0.1 µmol m -2 , at pH 8.0). Based on previous investigations for fluoroquinolones [START_REF] Kamagate | Mineralization Enhancement of Pharmaceutical Contaminants by Radical-Based Oxidation Promoted by Oxide-Bound Metal Ions[END_REF][START_REF] Martin | Kinetics and mechanisms of ciprofloxacin oxidation on hematite surfaces[END_REF][START_REF] Zhou | Water flow variability affects adsorption and oxidation of ciprofloxacin onto hematite[END_REF], OFL may form metal-bonded complexes with surface sites and/or directly hydrogen-bonded complexes with surface hydroxo groups involving both carboxylate and keto groups. However, accurate quantification of the adsorption is not straightforward, because binding of redox-active compounds to birnessite is followed by an electron transfer process resulting in the concomitant oxidation of sorbed compound and reduction of surface-bound metal. LC/MS analysis showed that OFL oxidation passes through partial or full N-dealkylation of the piperazine ring, as two predominant products of m/z = 336 (M-26) and m/z = 279 (M-69) have been detected (See Figure S6 for more details). Indeed, the piperazine ring of OFL can be oxidized through two one-electron transfers from the N atom of piperazine ring to Mn(Ⅳ), and then N-dealkylation and/or C-hydroxylation causes ring opening, as typically observed for similar compounds [START_REF] Zhang | Oxidative Transformation of Fluoroquinolone Antibacterial Agents and Structurally Related Amines by Manganese Oxide[END_REF]. Additionally, trace amounts of M+16, M-44 (decarboxylation), M-101 and M+30 were detected, as previously reported for the Mn-based oxidation of OFL [START_REF] Li | Degradation of ofloxacin by a manganeseoxidizing bacterium Pseudomonas sp. F2 and its biogenic manganese oxides[END_REF][START_REF] Zhang | Oxidative Transformation of Fluoroquinolone Antibacterial Agents and Structurally Related Amines by Manganese Oxide[END_REF]. The effect of oxygen on the removal efficiency of OFL was only observed at high MnO2 concentration (870 µM) or at pH 8.0 ± 0.1 for the full range of MnO2 concentration (Figure 1). In both conditions, the apparent pseudo-first-order rate constant k (h -1 ) increased with increasing MnO2 concentration. It is worth noting that a linear dependence of reactivity on MnO2 amount was only observed at pH 8.0, while a logarithmic relationship was observed at pH 4.0 and 6.0 (Figure 1). Contribution of Fenton-like reaction through, for instance hydroxyl radical formation, in OFL removal is excluded, since the addition of ethanol, used as a hydroxyl radical scavenger, has no effect on the OFL removal (Figure S7). Furthermore, we found that the impact of oxygen on birnessite reactivity at pH 8.0 is not specific to OFL, since similar behavior has been observed with another antibiotic compound, pipemidic acid (see Figure S8).

It is generally known that the reaction of MnO2 with redox-active compounds is complex and involves several reaction pathways that occur simultaneously. The redox byproducts generated through reductive dissolution of MnO2 or OFL oxidation, can bind onto the MnO2 surfaces, and alter the initial surface reactivity (i.e., passivation).

Generally, the generated Mn(II) is sorbed onto the MnO2 surface and then oxidized.

Mn(IV) is reduced, and the resulting product is Mn(III), which is commonly called the comproportionation reaction [START_REF] Wang | Effects of metal cations on coupled birnessite structural transformation and natural organic matter adsorption and oxidation[END_REF][START_REF] Zhao | Redox Reactions between Mn(II) and Hexagonal Birnessite Change Its Layer Symmetry[END_REF]. In addition, the adsorbed Mn(II) on vacant sites can be oxidized by O2 at high pH to yield Mn(III).

Because higher MnO2 concentration enables more OFL oxidation (higher k, Figure 1), more reductive dissolution or Mn(II) production is expected. Dissolved Mn(Ⅱ) was below the detection limit (< 0.02 μM for AAS) in all investigated conditions, confirming that generated Mn(Ⅱ) by reductive dissolution of MnO2 was completely removed by binding to MnO2. Competitive binding between OFL and Mn(II) at MnO2 edge sites, and/or adsorption of OFL oxidation byproducts at the MnO2 surface may also explain the passivation of the Mn-oxide surface. All these reactions are pHdependent, as well as the passivation of MnO2 surfaces. This may explain why the relationship between the oxidation rate constant of OFL and MnO2 content is not linear over the whole pH range investigated (Figure 1). Furthermore, higher surface-oxidation rates of Mn(II) at alkaline pH values [START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF][START_REF] Namgung | Cr(OH)3 (s) Oxidation Induced by Surface Catalyzed Mn(II) Oxidation[END_REF] may explain the pronounced impact of oxygen on birnessite reactivity at pH 8.0 (Figure 1).

Collectively, these results suggest that enhanced removal efficiency of OFL under oxic conditions at high MnO2 concentration or high pH is due to the Mn(III) generated through Mn(II) surface-catalyzed oxidation by oxygen. This is, however, in disagreement with a recent work which reported that oxygen inhibited phenol degradation by δ-MnO2 because the generated Mn(III)-phases at the MnO2 surface may block reactive sites and cause particle aggregation, thereby lowering the surface reactivity [START_REF] Hu | Impact of dissolved O2 on phenol oxidation by δ-MnO2[END_REF]. This discrepancy with the present work may be explained by the difference in experimental conditions including target contaminant characteristics and its affinity to MnO2 surfaces, and target compound/MnO2 molar ratio.

On the other hand, previous works have reported that the Mn(III) oxides or Mn(III) complexes can rapidly oxidize a variety of organic compounds [START_REF] Chen | Reduction of carbadox mediated by reaction of Mn(III) with oxalic acid[END_REF][START_REF] Huang | Effect of MnO2 Phase Structure on the Oxidative Reactivity toward Bisphenol A Degradation[END_REF][START_REF] Klewicki | Dissolution of β-MnOOH particles by ligands: pyrophosphate, ethylenediaminetetraacetate, and citrate[END_REF][START_REF] Nico | Mn (III) center availability as a rate controlling factor in the oxidation of phenol and sulfide on δ-MnO2[END_REF]. Mn(III) is a strong oxidant for one electron transfer reaction [START_REF] Hu | Role of dissolved Mn(III) in transformation of organic contaminants: Non-oxidative versus oxidative mechanisms[END_REF][START_REF] Sun | Degradation of Adsorbed Bisphenol A by Soluble Mn(III)[END_REF], compared to two electron transfer reaction involving Mn(IV).

Mn (III) + e -→ Mn (II) E 0 = +1.51 V (1) Mn(III)OOH + 3H + + e -→ Mn (II) + 2H2O E 0 = +1.50 V (2) Mn(IV)O2 + 4H + + 2e -→ Mn (II) + H2O E 0 = +1.29 V (3) In a heterogeneous system, the greater reactivity of Mn(III)-bearing phases has been explained by longer and weaker Mn(III)-O bonds facilitating the electron transfer with sorbing compounds (e.g., phenol) [START_REF] Huang | Effect of MnO2 Phase Structure on the Oxidative Reactivity toward Bisphenol A Degradation[END_REF][START_REF] Ukrainczyk | Oxidation of Phenol in Acidic Aqueous Suspensions of Manganese Oxides[END_REF].

It seems that direct coordination of Mn(III) with target compounds ensures to facilitate electron transfer without change in the Mn spin state, while the outer sphere electron transfer in surface Mn(IV) requires a change in the Mn spin state, in turn making the electron transfer process more difficult [START_REF] Huang | Effect of MnO2 Phase Structure on the Oxidative Reactivity toward Bisphenol A Degradation[END_REF][START_REF] Ukrainczyk | Oxidation of Phenol in Acidic Aqueous Suspensions of Manganese Oxides[END_REF]. Indeed, Mn(III) was reported to be as a high spin d 4 ion with much lower ligand field stabilization energy and faster ligand exchange rates comparing with Mn(IV) [START_REF] Nico | Importance of Mn (III) availability on the rate of Cr(III) oxidation on δ-MnO2[END_REF]. In contrast, other reports showed that Mn(III) sites are less reactive with respect to target contaminants. For example, lower As(III) oxidation have been attributed to the lower affinity of Mn(III) sites for As(III) adsorption, and slower electron transfer rates with adsorbed As(III) (Lafferty et al., 2010a;Lafferty et al., 2010b;[START_REF] Zhu | Quantum Chemical Study of Arsenic (III, V) Adsorption on Mn-Oxides: Implications for Arsenic(III) Oxidation[END_REF]. Moreover, the surface functional groups on Mn(III) sites are more labile than functional groups on Mn(IV) sites [START_REF] Zhu | Quantum Chemical Study of Arsenic (III, V) Adsorption on Mn-Oxides: Implications for Arsenic(III) Oxidation[END_REF], and therefore strong adsorption via ligand exchange processes on Mn(III) sites than on Mn(IV) sites is expected. Finally, different types of Mn sites, e.g., Mn(III, IV) in sheets, Mn(III, IV) at edges, and Mn(III) in interlayers, likely have varying reactivity [START_REF] Manceau | Structural mechanism of Co 2+ oxidation by the phyllomanganate buserite[END_REF][START_REF] Simanova | Time-Resolved Investigation of Cobalt Oxidation by Mn(III)-Rich δ-MnO2 Using Quick X-ray Absorption Spectroscopy[END_REF][START_REF] Yu | Structural factors of biogenic birnessite produced by fungus Paraconiothyrium sp. WL-2 strain affecting sorption of Co 2+[END_REF]. The interlayer Mn(III) was suggested to be the dominant oxidizing site, and the abundance of Mn(III) with respect to total Mn may be key in promoting the reactivity of birnessite [START_REF] Peng | Redox properties of birnessite from a defect perspective[END_REF]. From a kinetic point of view, Mn(III) at the edges seems to contribute at short reaction times, while Mn(III, IV) in the MnO2 sheets is important at longer times [START_REF] Simanova | Time-Resolved Investigation of Cobalt Oxidation by Mn(III)-Rich δ-MnO2 Using Quick X-ray Absorption Spectroscopy[END_REF]Wang et al., 2018b). Despite this conflicting information about the role of structural Mn(III) in promoting MnO2 reactivity, we may conclude that multiple factors may control the contribution of Mn(III) in enhancing the removal capacity of birnessite. This may include the ratio of Mn(III) to Mn(IV), mineral phase transformation along the redox reaction, and properties of target contaminants and its interaction with Mn sites. To gain more insights into the underlying passivation mechanisms, the effects of Mn(II)/MnO2 ratio on the removal kinetics of OFL were investigated under oxic and anoxic conditions in the following section.

Combined effects of Mn(II) and O2 under a wide range of ionic strength

As an attempt to confirm the impact of surface-catalyzed oxidation of Mn(II) by O2 on the reactivity of MnO2, OFL removal kinetics were investigated with different Mn(II)/MnO2 molar ratios at pH 8.0 (Figure 2). As expected, the initial rate constants of OFL sharply decreased with increasing Mn(II) under both investigated conditions.

As previously reported [START_REF] Chen | Reaction of lincosamide antibiotics with manganese oxide in aqueous solution[END_REF][START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF][START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF], dissolved Mn(II) strongly adsorbed on the negatively charged birnessite surfaces, resulting in competition between OFL and Mn(II) binding to the reactive surface sites of MnO2. This Mn(Ⅱ) binding is much more important at alkaline pH values [START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF]. The adsorbed Mn(II) on vacant sites at high pH is oxidized to Mn(III) through surface-catalyzed oxidation of Mn(Ⅱ) enabling more Mn(III) in the vacancies [START_REF] Zhu | Cation Effects on the Layer Structure of Biogenic Mn-Oxides[END_REF]. This explains why the dissolved Mn(II) concentration was found here below the AAS detection limit, even for the highest Mn(II) amount (200μM).

Time-dependent UV-Vis measurements for birnessite suspensions containing 200 μM Mn(II), suggested more generated Mn(III) in the presence of oxygen relative to anoxic conditions (Figure S9). Indeed, the intensity of adsorption bands at 258 nm for the potential Mn(III)-PP complex (λ1 = 480 nm, ε1 = 65 M -1 ; λ2 = 258 nm, ε2 = 6750 M -1 ) [START_REF] Hu | Role of dissolved Mn(III) in transformation of organic contaminants: Non-oxidative versus oxidative mechanisms[END_REF] was found to be higher under oxic conditions.

It should be noted that the greater reactivity of birnessite under oxic conditions relative to anoxic systems is only observed at lower Mn(II)/MnO2 molar ratio (≤ 0.1). At higher ratio, strong competition between Mn(II) and OFL may act as a counterweight to the Mn(III) contribution, thereby reducing the overall reactivity of MnO2.

When removal kinetics was investigated at higher ionic strength (10 or 100 mM of NaCl), a decrease in the initial rate constants of OFL was observed, mainly under oxic conditions and at low Mn(II)/MnO2 molar ratio (Figure 2). Increasing ionic strength may increase the size of MnO2 via oriented aggregations [START_REF] Burrows | Effect of ionic strength on the kinetics of crystal growth by oriented aggregation[END_REF][START_REF] Yang | Metal Adsorption Controls Stability of Layered Manganese Oxides[END_REF], thereby reducing the surface reactivity. This decrease in surface reactivity has also been previously observed during aging of manganese oxides through coalescence of particles [START_REF] Stone | Reduction and Dissolution of Manganese(III) and Manganese(IV) Oxides by by organics. 1. Reaction with hydroquinone[END_REF]. Furthermore, Na + ions may compete with the cationic Mn(II) and/or zwitterionic form of OFL (pKa,1 = 6.08 and pKa,2 = 8.25) for negatively charged adsorptive sites on MnO2 (pHIEP ~ 2.0). Na + can occupy the interlayers of birnessite, potentially competing with Mn(II) and/or influencing the diffusion process of Mn(II) at reactive surface sites [START_REF] Yang | Metal Adsorption Controls Stability of Layered Manganese Oxides[END_REF].

Previous work reported similar competitive binding due to the presence of monovalent or divalent cations, resulting in decrease in the rate and magnitude of organic compound transformation [START_REF] Chen | Reaction of lincosamide antibiotics with manganese oxide in aqueous solution[END_REF]. Under oxic conditions, oxidation of Mn(II) into Mn(III) and subsequent incorporation of Mn(III) into the layers may occur, resulting in birnessite transformation from hexagonal to orthogonal layer symmetry [START_REF] Zhao | Redox Reactions between Mn(II) and Hexagonal Birnessite Change Its Layer Symmetry[END_REF]. In this case, Na + ions could compensate for the charge deficiency caused by the presence of Mn(III), and subsequently act to stabilize Mn(III) in the layers [START_REF] Drits | Birnessite polytype systematics and identification by powder X-ray diffraction[END_REF][START_REF] Zhu | Cation Effects on the Layer Structure of Biogenic Mn-Oxides[END_REF]. This may explain the pronounced effect of increasing NaCl concentration on the MnO2 surface reactivity in the presence of O2. To check whether transition zones at oxic-anoxic interfaces may influence the reactivity of birnessite, removal kinetics experiments were conducted for 24 h under N2, with subsequent ambient air conditions or bubbling O2 for 48h (Figure 3). Under anoxic conditions, a plateau was observed at around 16h, and then continued over 72 h. When suspensions were exposed to ambient air or O2 after a pre-equilibration period of 24h under N2, OFL was further removed up to 65% in suspension, which could be ascribed to the Mn(III) generated through Mn(II) surface-catalyzed oxidation by oxygen. As mentioned above, these newly generated Mn(III) are more reactive for OFL oxidation.

It should be noted that strong oxygenation of solution with bubbling O2 has no greater effect on the birnessite reactivity, as compared to the solution exposed to ambient air. The same two-step anoxic/oxic process was repeated but in the presence of 5 μM Mn(II).

This Mn(II) amount was chosen to cover the typical range of Mn concentration found in surface and groundwaters (~ 4 μM) [START_REF] Namgung | Cr(OH)3 (s) Oxidation Induced by Surface Catalyzed Mn(II) Oxidation[END_REF]. The presence of 5 μM Mn(II) did not significantly affect the removal rate whether under oxic or anoxic conditions. Only a slight inhibition was observed especially for longer reaction times , likely because of competitive effects of Mn(II) and OFL for binding onto reactive sites.

In all conditions, removal kinetic exhibited a plateau under N2, followed by a drop in the aqueous concentration of OFL when the reaction media was exposed to air or O2 (Figure 3). 385

XRD and XPS investigations

To evaluate the role of Mn speciation on reactivity changes and the passivation of the MnO2 surface, XPS analysis of initial and reacted MnO2 samples were performed under defined experimental conditions.

First, the total contents of major elements are reported in Table S1. In the absence of Mn(II), a decrease in Mn content in reacted samples was observed, likely resulting from the presence of various adsorbing ligands at the surface (Table S1) [START_REF] Shaikh | Spectroscopic Investigation of Interfacial Interaction of Manganese Oxide with Triclosan, Aniline, and Phenol[END_REF].

Higher Na content under oxic conditions also proved the compensation effect of Na + as discussed above. Increasing Mn(II) concentrations from 50 μM to 200 μM enhanced Mn percentage, likely because of the adsorption of Mn(II) onto the MnO2 surfaces. K content was below detection limit, likely due to the large adsorption amounts of Mn exchanging with interlayer K + ion.

Second, the high resolution spectra of Mn 2p3/2 was then fitted and the BE (binding energy) positions of Mn(II), Mn(III), Mn(IV) ranged from 640.1 to 641.4 eV, from 642.1 to 642.4 eV and from 642.6 to 643.1 eV respectively, consistent with previously reported values (Figure 4, Table S2) [START_REF] Sun | Antimony oxidation and sorption behavior on birnessites with different properties (δ-MnO2 and triclinic birnessite)[END_REF][START_REF] Tang | Effect of Surface Manganese Valence of Manganese Oxides on the Activity of the Oxygen Reduction Reaction in Alkaline Media[END_REF][START_REF] Tang | Oxalate route for promoting activity of manganese oxide catalysts in total VOCs'oxidation: effect of calcination temperature and preparation method[END_REF]. We have subsequently determined the average oxidation state (AOS) of all samples, and found that the AOS (initial value 3.87) decreased for all reacted samples (Table S2).

As expected, increasing the Mn(II) concentration from 0 μM to 200 μM enhanced the amount of Mn(II) bond to MnO2 surfaces (Figure 4). An increase in the percentage of Mn(III) was also observed in the presence of oxygen relative to anoxic conditions, while that of Mn(IV) decreased. As previously explained, Mn(II) can undergo surfacecatalyzed oxidation by molecular oxygen and/or comproportionation with structural Mn(IV) forming Mn(III) sites. Both processes can enhance the Mn(III) content in MnO2, particularly under oxic conditions (Table S2). The high-resolution spectra of O1s were then evaluated to assess potential transformation of birnessite surfaces (Table S3 and Figure S10). The compositions of oxygen bonding types consisted of a main peak located at around 530.2 eV corresponding to lattice oxygen with metal (Olatt), a shoulder peak near 531.6 eV attributing to the adsorbed oxygen species on MnO2 surface (Oads) in the form of hydroxide oxygen (OH -) and the weakest peak at higher binding energy assigning to the adsorbed oxygen in water molecule (Osurf) [START_REF] Tang | Effect of Surface Manganese Valence of Manganese Oxides on the Activity of the Oxygen Reduction Reaction in Alkaline Media[END_REF][START_REF] Zhang | Enhanced removal of heavy metal ions from aqueous solution using manganese dioxide-loaded biochar: Behavior and mechanism[END_REF].

After reaction, the increased proportions of H2O suggested that the surface became progressively hydroxylated and hydrated. It is worth noting that in the presence of 200 μM Mn(II), high contents in Mn(III) and OH -under anoxic condition have been considered as characteristic of intermediate reaction products Mn(III)-hydroxide, here represented as MnOOH [START_REF] Nesbitt | XPS study of reductive dissolution of 7Åbirnessite by H3AsO3, with constraints on reaction mechanism[END_REF]. Moreover, the content of Olatt decreased to 49 % (anoxic) and 65 % (oxic), which suggests that lattice oxygen may participate in the oxidation of Mn(II) with the formation of oxygen vacancy [START_REF] Cheng | Effects of dissolved oxygen on the start-up of manganese oxides filter for catalytic oxidative removal of manganese from groundwater[END_REF].

Once the Olatt is consumed, dissolved oxygen could compensate [START_REF] Cheng | Effects of dissolved oxygen on the start-up of manganese oxides filter for catalytic oxidative removal of manganese from groundwater[END_REF], which would explain the difference in Olatt content between anoxic and oxic environments.

The produced Mn(III) may incorporate into the layer of birnessite and induce the changes of mineral structure by converting birnessite into lower valence Mn phases [START_REF] Zhao | Redox Reactions between Mn(II) and Hexagonal Birnessite Change Its Layer Symmetry[END_REF]. To check possible phase transformation, XRD analysis of MnO2 samples under various Mn(II)/MnO2 molar ratios (R) and under anoxic and oxic conditions was conducted (Figure 5). Samples reacted with OFL in the presence of dissolved Mn(II) (50 μM, R = 0.14) showed no bulk transformation. However, XRD data of samples with higher R (i.e. 0.29 or 0.58) showed different peaks, both under anoxic and oxic conditions, which could suggest the presence of lower valence Mn oxides, e.g. feitknechite (β-MnOOH) (Wang et al., 2018a;[START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF]. This is consistent with SEM images where needle-shaped particles were observed amid nanoflower-shaped particles of birnessite, especially at higher Mn(II) concentration and under oxic conditions (Figure S11). These needle-shaped crystals can be attributed either to the presence of manganite (γ-MnOOH) as previously reported by e.g. [START_REF] Larsen | Manganite reduction by Shewanella putrefaciens MR-4[END_REF], or more certainly according to XRD data, to feitknechtite as previously reported by [START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF]. [START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF] the * symbol has previously been attributed to hausmannite (Wang et al., 2018a) but in the absence of the main Bragg peaks of this phase, it would rather belong to an hydrated -MnO2 phase [START_REF] Rossouw | Alpha manganese dioxide for lithium batteries: a structural and electrochemical study[END_REF].

Conclusions

As a result of natural or anthropogenic activities, such as groundwater recharge, sediment dredging, water table fluctuations, etc., alternating redox conditions commonly occur in natural systems. Here, we have demonstrated for the first time that the oxygen can alter the redox reactivity of MnO2 under environmentally-relevant conditions of Mn(II) dosage, pH and ionic strength. A decline in the oxidation rate constant with increasing pH was observed, likely because of the decrease in the redox potential of MnO2 at higher pH and modification in surface interactions with the target compound. At high pH, the adsorbed Mn(II) on vacant sites can be oxidized by O2 or Mn(IV) surrounding the vacant sites into Mn(III) (comproportionation reaction).

Therefore, OFL adsorbed at vacant sites can be oxidized by the newly generated Mn(III) sites, which seems to be strong oxidizing species. However, excess dissolved Mn(II) would compete with compounds for binding at reactive sites, thus decreasing the removal capacity of birnessite. The decreased reactivity of MnO2 with higher ionic strength also suggested that the sorption of alkali cation (Na + ) on the surface may control the birnessite reactivity. Furthermore, sequential anoxic-oxic dynamic process may significantly impact the fate of existing contaminants in the presence of birnessite.

Combining anoxic and oxic environments into one process may be relevant in many subsurface environments, for instance, when anoxic aquifers and/or sediments come into contact with oxygen. Further redox oscillation studies are needed to explore the fate and mobility of contaminants and co-existing natural compounds (e.g., dissolved organic matter), in the presence of manganese oxides.

[

  OFL] /[OFL]o versus time, were plotted against MnO2 concentration in Figure 1. The correlation coefficients (R 2 ) of linear regression (ln C /Co = -kt) were more than 0.90 for all experiments.

Figure 1 .

 1 Figure 1. Initial removal rate constants of OFL versus MnO2 concentration at three pH values 4.0, 6.0, 8.0 (± 0.1). 173, 345, 870 μM (15, 30, 75.6 mg/L) MnO2; 10 μM OFL; 10 mM NaCl; under anoxic (N2) and oxic (O2) conditions. Dashed lines are the best correlation obtained between k and MnO2 amount. Correlation coefficients (R 2 ) are indicators of the strength of the linear or non-linear relationship.

Figure 2 .

 2 Figure 2. Initial removal rate constants of OFL as a function of Mn(II)/MnO2 molar ratio under anoxic or oxic conditions. pH 8.0 (± 0.1), 345 μM (30 mg/L) acid birnessite. 0, 5, 50, 100, 200 μM Mn(II); 10 μM OFL; (a) 0 mM NaCl, (b) 10 mM NaCl, (c) 100 mM NaCl. Dashed lines are only as visual guide.

Figure 3 .

 3 Figure 3. Removal kinetics of OFL in dynamic same two-step anoxic/oxic process: (a) 0 μM Mn(II); (b) 5 μM Mn(II). pH 8.0 (± 0.1); 10 μM OFL; 10 mM NaCl; 345 μM acid birnessite. The red vertical line indicates the beginning of oxygenation through exposure to ambient air or with bubbling O2.

Figure 4 .

 4 Figure 4. Fitting of Mn 2p3/2 XPS high resolution spectra of the synthesized pristine acid birnessite and samples treated with Mn(II) and OFL under anoxic and oxic conditions. The R represents the Mn(II)/MnO2 molar ratio. 345 μM acid birnessite; 10 μM OFL; 10 mM NaCl; 0, 50, 200 μM Mn(II); 48 h reaction time; pH 8.0 (±0.1). The numbers inset are the percentage of Mn(II) (red), Mn(III) (green) and Mn(IV) (blue).

Figure 5 .

 5 Figure 5. XRD patterns of the MnO2 samples treated with Mn(II) and OFL under oxic and anoxic conditions. 345 μM acid birnessite; 10 μM OFL; 10 mM NaCl; 0, 50, 100, 200 μM Mn(II); 48 h or 6 d reaction time; pH 8.0 (± 0.1). The R represents the Mn(II)/MnO2 molar ratio. The f symbols indicate XRD peaks visible on the feitknechtite (β-MnOOH) sample in[START_REF] Lefkowitz | Influence of pH on the Reductive Transformation of Birnessite by Aqueous Mn(II)[END_REF] the

Acknowledgments

Q.L thanks the Chinese Scholarship Council of PR China for her Ph.D. grant and Rennes Métropole (France) for a mobility grant for an extended research visit at KIT (Germany). We want to acknowledge Valerie Courousse (AAS) and Kouhail Yasmine (XRD) for providing help in experiments.