Tung Huu Dang

Angela Vasanelli
email: angela.vasanelli@ens.fr

Yanko Todorov

Carlo Sirtori

Yoann Prado

Audrey Chu

Charlie Gréboval

Adrien Khalili

Herve Cruguel

Christophe Delerue

Gregory Vincent

Emmanuel Lhuillier

Bias Tunable Spectral Response of Nanocrystal Array in a Plasmonic Cavity

Keywords: plasmon, hopping transport, nanocrystal, bias-induced spectral shift, infrared

Nanocrystals (NCs) have gained considerable attention for their broadly tunable absorption from the UV to the THz range. Nevertheless, their optical features suffer from a lack of tunability once integrated into optoelectronic devices. Here, we show that tunable absorption with bias can be obtained by coupling a HgTe NC array with a plasmonic resonator. Up to 15 meV blueshift can be achieved from a 3-µm absorbing wavelength structure under a 3 V bias voltage when the NC exciton is coupled with a mode of the resonator. We demonstrate that the blueshift arises from the interplay between hopping transport and inhomogeneous absorption due to the presence of the photonic structure. The observed tunable spectral response is qualitatively reproduced in simulation by introducing a bias-dependent diffusion length in the charge transport. This work expands the realm of existing NC-based devices and paves the way towards light modulators.

Introduction

While nanocrystals (NCs) are claimed for being essential building blocks for optoelectronic devices, the actual number of devices based on these colloidal nanoparticles remains mostly limited to solar cells, [START_REF] Carey | Colloidal Quantum Dot Solar Cells[END_REF][START_REF] Nozik | Quantum Dot Solar Cells[END_REF][START_REF] Ellingson | Highly Efficient Multiple Exciton Generation in Colloidal PbSe and PbS Quantum Dots[END_REF] light-emitting diodes, [START_REF] Wood | Colloidal Quantum Dot Light-Emitting Devices[END_REF][START_REF] Qu | Electroluminescence from HgTe Nanocrystals and Its Use for Active Imaging[END_REF] and, more recently, infrared focal plane arrays. [START_REF] Chu | HgTe Nanocrystals for SWIR Detection and Their Integration up to the Focal Plane Array[END_REF][START_REF] Rauch | Near-Infrared Imaging with Quantum-Dot-Sensitized Organic Photodiodes[END_REF][START_REF] Buurma | MWIR Imaging with Low Cost Colloidal Quantum Dot Films[END_REF] Thus, many concepts of devices remain unexplored in the case of colloidal NCs. One of the most striking successes of NCs comes from their broad spectral tunability. By combining quantum confinement to the extensive library of materials that can be grown under the colloidal form, a significant part of the electromagnetic spectrum -from UV to THz [START_REF] Goubet | Terahertz HgTe Nanocrystals: Beyond Confinement[END_REF] -can be addressed. However, once integrated into a device, their spectral response is set by the choice of the synthesis parameters. This limitation can be overcome by engineering the light-matter coupling with light resonators and cavities. [START_REF] Tang | Towards Infrared Electronic Eyes: Flexible Colloidal Quantum Dot Photovoltaic Detectors Enhanced by Resonant Cavity[END_REF][START_REF] Tang | Thermal Imaging with Plasmon Resonance Enhanced HgTe Colloidal Quantum Dot Photovoltaic Devices[END_REF] This approach has been used to design spectral filters, [START_REF] Tang | Acquisition of Hyperspectral Data with Colloidal Quantum Dots[END_REF] as well as to broaden the absorption [START_REF] Chu | Near Unity Absorption in Nanocrystal Based Short Wave Infrared Photodetectors Using Guided Mode Resonators[END_REF] of NC thin films. [START_REF] Rastogi | Pushing Absorption of Perovskite Nanocrystals into the Infrared[END_REF][START_REF] Gréboval | Ferroelectric Gating of Narrow Band-Gap Nanocrystal Arrays with Enhanced Light-Matter Coupling[END_REF] While this strategy offers post-synthesis tunability, this remains a non-agile method, where device geometrical factors now determine the spectral response. Nevertheless, reconfigurable responses are still of utmost interest not only for light sensing and emission [START_REF] Sergeev | Tailoring Spontaneous Infrared Emission of HgTe Quantum Dots with Laser-Printed Plasmonic Arrays[END_REF] but also for devices such as light modulators. Several strategies to achieve a reconfigurable spectral response have been explored: exploiting the Stark effect, [START_REF] Miller | Band-Edge Electroabsorption in Quantum Well Structures: The Quantum-Confined Stark Effect[END_REF] with a redshift of the band edge under an electric field application, [START_REF] Empedocles | Quantum-Confined Stark Effect in Single CdSe Nanocrystallite Quantum Dots[END_REF][START_REF] Scott | Time-Resolved Stark Spectroscopy in CdSe Nanoplatelets: Exciton Binding Energy, Polarizability, and Field-Dependent Radiative Rates[END_REF][START_REF] Klem | Quantum Dot Electroabsorption Modulation across the Extended Communications Band 1200-1700 nm[END_REF][START_REF] Walters | The Quantum-Confined Stark Effect in Layered Hybrid Perovskites Mediated by Orientational Polarizability of Confined Dipoles[END_REF] or electrochromism, by charging the NC film. [START_REF] Wang | Electrochromic Nanocrystal Quantum Dots[END_REF] The field required to induce a tunable response is extremely large, typically above 100 kV.cm -1 . This strong electric field is generally not compatible with electrical transport due to electrical breakdowns. As a result, only the optical features can be tuned by these approaches. More recently, reconfigurable spectral responses have been obtained from phase change materials [START_REF] Shi | Reconfigurable and Narrowband Mid-Infrared Filter Based on Surface Lattice Resonance in Phase-Change Ge2Sb2Te5[END_REF][START_REF] Pogrebnyakov | Reconfigurable Near-IR Metasurface Based on Ge2Sb2Te5 Phase-Change Material[END_REF] and MEMS actuators, [START_REF] Stark | MEMS Tunable Mid-Infrared Plasmonic Spectrometer[END_REF] but these strategies are not straightforward to be applied to NC-based films. Here, we design a spectrally tunable NC-based device by carefully coupling the electromagnetic field and the charge collection resulting from hopping within a NC assembly. To build such a device, we combine a HgTe NC [START_REF] Gréboval | Mercury Chalcogenide Quantum Dots : Material Perspective for Device Integration[END_REF] thin-film absorbing in the extended short-wave infrared (e-SWIR) with a plasmonic resonator. The latter induces a spatially inhomogeneous spectral response within the device. We show that, by applying a small electric field (<40 kV.cm -1 ), the spectral response of the photodetector is shifted up to 15 meV.

Results and discussion

We design a device based on HgTe NCs. The latter are grown using the Prado et al. procedure. [START_REF] Prado | Seeded Growth of HgTe Nanocrystals for Shape Control and Their Use in Narrow Infrared Electroluminescence[END_REF] Briefly, HgBr2 and TOP:Te are mixed at room temperature and react to form seeds that later grow when heated in a hot oleylamine solution. TEM observations (Figure 1b and S1) show that the obtained particles have a slightly faceted spherical shape, with a typical diameter close to 10 nm. The resulting absorption spectrum (Figure 1b) shows an exciton resonance at 4200 cm -1 (≈2.4 µm) and a narrow doublet at 2900 cm -1 resulting from the C-H bond resonance of the capping ligands. The plasmonic cavity is fabricated on a sapphire substrate to ensure good transparency in the 1-3 µm range, see figure S2-5 for design and S7-9 for fabrication. On top of the substrate, a gold grating is fabricated by e-beam lithography, as sketched in Figure 1a. The period p of the grating is chosen so that a surface plasmon polariton resonance overlaps with the HgTe NC exciton one when it is excited in the vicinity of the first diffraction order of the grating [START_REF] Jouy | Coupling of a Surface Plasmon with Localized Subwavelength Microcavity Modes[END_REF] . The refractive index of HgTe NCs film being 2.35, [START_REF] Rastogi | Complex Optical Index of HgTe Nanocrystal Infrared Thin Films and Its Use for Short Wave Infrared Photodiode Design[END_REF] p is estimated to be around 1350 nm. [START_REF] Jouy | Coupling of a Surface Plasmon with Localized Subwavelength Microcavity Modes[END_REF] This grating is also used as interdigitated electrodes: bias is applied between two adjacent electrodes to generate photocurrent. The NC film is deposited on the top of the grating by spin-coating an ink containing HgTe particles capped with HgCl2 and short thiols. Atomic force microscopy (Figure S9) reveals that NC deposition on the grating is quasi-conformal. Finally, a top gold layer is added to generate the spoof surface plasmon whose wavelength is selected by the bottom grating period. A thin PMMA (20 nm) layer is added between the NC film and the top electrode to be used as a gate and controls the NC film carrier density.

c. Experimental reflectivity spectra measured at room temperature and at 15° incidence angle for a NC film deposited on a gold layer (blue circles), and for the device depicted in part a with s = 500 nm and p = 1350 nm. The top panel presents the experimental (green circles) and simulated (red line) reflectivity spectrum with TM polarized radiation, the bottom one presents the experimental spectrum in TE polarization. The inset also shows the electric field map associated with the resonance at 3200 cm - 1 (spoof surface plasmon resonance) and 2160 cm -1 (patch cavity mode) in TM polarization, and in correspondence to the broad feature observed in TE.

Figure 1c presents the optical properties of our device by comparing the reflectivity of a nanocrystal film deposited on a gold layer (blue symbols) to the reflectivity of the device in TM (top panel) and TE (bottom panel) polarizations, measured at the incident angle of 15°. At 3800 cm -1 , we observe the HgTe exciton. The peak is shifted compared to the value measured in solution due to interparticle coupling generated by the ligand exchange and change of dielectric environment. The coupling with the first diffraction order of the grating enhances the contrast of the exciton resonance in both polarizations. We have verified that this enhancement corresponds to an increased absorptivity induced by the cavity (see Figure S3a). The optical properties of the device are dependent on the light polarization. In the case of the TM polarization, a second feature appears at 3300 cm -1 in the tail of the absorption of the NCs. This resonance is associated with a surface plasmon polariton mode supported by the gold, nanocrystal, sapphire multilayer, excited in the vicinity of the first diffraction order of the grating. As discussed later, the inhomogeneous spatial distribution of the modes will play a central role in achieving bias tunability of the photoresponse. Note that the plasmon mode has a strong dispersion (see Figure S2c), and at normal incidence, it is shifted towards the excitonic peak. At 2200 cm -1 , corresponding to λ=2.neffective.s with neffective the cavity effective index and s the patch size, a weak and broad dip appears. This feature is associated with a resonance of the patch cavity formed between the top and bottom gold electrodes [START_REF] Todorov | Optical Properties of Metal-Dielectric-Metal Microcavities in the THz Frequency Range[END_REF] (right inset of Figure 1c, top panel, and Figure S5).

b. Responsivity (under blackbody illumination at 980 °C) and detectivity (at 1 kHz) as a function of operating temperature for VDS=0.5 V. c. Current as a function of time while the light (λ=1.55 µm) is turned on and off. d. Currents (dark and under illumination) as a function of gate bias, for VDS=0.2 V, at 150 K. The ratio of those two currents is plotted on the right axis as a function of the gate bias.

The designed device presents photoconductive properties, see Figure 2a and S10-15 for characterization. The introduction of the resonator clearly affects the spectral response of the device. A more pronounced photocurrent peak is observed in the plasmonic resonator device (with p=1.5 µm) with respect to a reference device in which the NC film is connected to conventional interdigitated electrodes (i.e., no top gold layer), separated by a period of 20 µm, see figure S3b. In terms of detection performance, the device responsivity under broadband illumination (blackbody at 980 °C) reaches 0.25 A.W -1 at 200 K (Figure 2b), corresponding to an external quantum efficiency EQE≈20%. Below this temperature, the responsivity drops due to the thermal activation of the mobility. [START_REF] Gréboval | Ferroelectric Gating of Narrow Band-Gap Nanocrystal Arrays with Enhanced Light-Matter Coupling[END_REF] Simultaneously the carrier density is also thermally activated, thus the dark current can be drastically reduced thanks to cooling (see figure S15). Consequently, the prevailing 1/f noise current spectral density drops with temperature (Figure S13), and the specific detectivity increases from 10 7 Jones at 240 K to 8x10 [START_REF] Goubet | Terahertz HgTe Nanocrystals: Beyond Confinement[END_REF] Jones at 80 K, see Figure 2b. The time response of the device, as presented in Figure 2c, is around 5 µs, which is quite typical for devices based on HgTe NCs. [START_REF] Gréboval | Mercury Chalcogenide Quantum Dots : Material Perspective for Device Integration[END_REF] The top gate (PMMA) can be used to control the device carrier density. Figure 2d reports the dark current (black line) and the current under illumination (red line) as a function of the gate bias. The dark conduction increases as a larger positive gate bias voltage is applied, indicating the n-type nature of the HgTe NC channel. This result is consistent with the measurement conducted with an electrolyte gated field-effect transistor configuration, see S16. By applying a negative gate bias, the majority carriers are removed from the channel, and the film becomes more intrinsic, leading to an increase of the photocurrent to dark current ratio up to a factor of 20 (green symbols). The most striking feature of this device is its bias tunable photoresponse, see Figure 3a. As a bias is applied between the two channel electrodes, a blueshift of the photoresponse is observed. This blue shift is present in both TM and TE polarizations (see Figure S17 and S18) and is fully reversible when bias is suppressed. Note that the blue shift is observed for both bias polarities (see Figure S19). This blue shift is somehow counterintuitive: usually, applying an electric field on quantumconfined semiconductors leads to a redshift as the wavefunction starts to leak in the barriers. Using tight-binding simulations, we have checked that a redshift should indeed be observed in HgTe NCs under electric field, see Figure S6. However, the calculated magnitude of this redshift is very weak. Tight-binding simulations predict a 6 cm -1 (<1 meV) shift under 100 kV.cm -1 applied electric field, while we observe an inverted shift reaching more than 15 meV under 40 kV.cm -1 (3 V over 0.85 µm electrode spacing) at 50 K, see Figure 3b. We can conclude that the observed shift is not induced by an effect of the electric field on the electronic states of the NCs. HgTe NCs present an inverted temperature dependence of the bandgap, leading to a blueshift of the bandgap upon heating. To quantify this effect, we performed a series of photocurrent spectra at various temperatures and found a 120 µeV.K -1 value for dEg/dT, see Figure S21. As a bias is applied, the current flowing within the device may generate heat due to Joule effect. However, the 15 meV observed shift would correspond to an increase of 120 K in temperature, which is improbable to be induced by the small Joule dissipated power (RI 2 ≈ a few nW under 1 V bias << several mW cooling power of the cryocooler at 80 K) in our device at low temperatures. Moreover, this heating should be more critical when the device resistance is reduced (i.e., at higher temperatures), while in our device, the opposite trend is observed (i.e., more noticeable shift as temperature decreases), see Figure 3b. We then use the gate to see its impact on the spectrum, see Figure 3c. In this case, no shift is observed while the magnitude of the gate-induced electric field (1 V over 20 nm of PMMA) corresponds to 500 kV.cm -1 , which is one order of magnitude larger than the drain-source electric field applied to generate the shift. As a consequence, the observed shift cannot be induced by a variation of the carrier density. In order to understand the origin of the blueshift of the photocurrent spectrum, it should be pointed out that no shift is observed for a HgTe film on conventional interdigitated (no cavity, no plasmon) electrodes under bias voltage application, see Figure S14. Furthermore, even devices with smaller electrode spacing, with ≈50 nm separation between the electrodes, and consequently with even larger applied electric fields, do not present such a bias dependence of the photocurrent spectrum [START_REF] Chu | Infrared Photoconduction at the Diffusion Length Limit in HgTe Nanocrystal Arrays[END_REF] . Therefore, it is suggested that the special device architecture, generating a specific light-matter coupling, is at the origin of this spectral response tunability. The simulated absorption maps, obtained as described in section S3 of the SI, reveal this coupling for TM (Figure 4a) and TE polarizations (Figure S4a). In both cases, the absorption maps appear to be strongly inhomogeneous. In particular, the NC absorption is concentrated between, rather above, the electrodes of the grating. The areas between the electrodes (region 1 of Figure 4a) present a maximum of the NC absorption strongly blue-shifted compared to those above the electrodes (region 2 of Figure 4a). Away from the cavity, when field enhancement is weak, the spectrum is driven by the NC film only. In the cavity, the field magnitude is enhanced and the plasmon resonance is spectrally overlapping with the exciton. As seen from Figure 1c and Figure S3 b, the plasmon has been chosen to be slightly detuned compared to the exciton, thus enhancing the red part of the NC spectrum. As a result, we have coupled the field enhancement with the redshifted spectrum, see Figure 4b. Due to the polycrystalline nature of the film, the transport occurs through a hopping process, leading to short diffusion lengths of the order of a few tens of nm. [START_REF] Chu | Infrared Photoconduction at the Diffusion Length Limit in HgTe Nanocrystal Arrays[END_REF] In a field-assisted hopping model, [START_REF] Nenashev | Effect of Electric Field on Diffusion in Disordered Materials. II. Two-and Three-Dimensional Hopping Transport[END_REF][START_REF] Nenashev | Effect of Electric Field on Diffusion in Disordered Materials. I. One-Dimensional Hopping Transport[END_REF][START_REF] Richert | Diffusion and Drift of Charge Carriers in a Random Potential: Deviation from Einsteins Law[END_REF][START_REF] Arkhipovt | A Model of Weak-Field Quasi-Equilibrium Hopping Transport in Disordered Materials[END_REF] the diffusion coefficient can be written D(F,T)= D0(T)+A(F,T).F 2 (1), where D0(T) is the diffusion coefficient in the zero-field limit, F is the magnitude of the electric field, and A(F,T) is a coupling factor. This diffusion coefficient directly relates the diffusion length with the carrier lifetime 𝜏 through 𝐿 𝐷 (𝐹, 𝑇) = √𝐷(𝐹, 𝑇) ⋅ 𝜏. We see from equation (1) that the application of a bias tends to increase the diffusion length. Thus, carriers that are photogenerated further away from the electrodes can be collected. As a result, the weight from the absorption occurring in region 1 becomes more important, leading to the observed blueshift of the photoresponse. As an illustration, we have plotted a map of the product between the position-dependent absorption spectrum Abs(𝑟, 𝜆) and exp(-r 2 /2LD 2 ), with r the distance to the electrodes. The latter term represents the collection efficiency in a diffusive model. The Abs(𝑟, 𝜆) x exp(-r 2 /2LD 2 ) map is plotted for two values of the diffusion length, and we see an apparent increase of the signal coming from region 1, see Figure 4c and d for TM polarization and figure S4 c-d for TE polarization. Finally, we also have plotted the spectral response of the device, taken as the surface integral of the product Abs(𝑟, 𝜆) x exp(-r 2 /2LD 2 ). A blueshift of the resonance is observed for both polarizations (Figure 4e and S4e) as the diffusion length is increased, which corresponds to an increase of the applied bias voltage. More quantitative estimation of the diffusion length from the spectral shift requires considering the vectorial nature of F, D0 (T), and A(F,T), which is beyond the scope of this paper. Eventually, our model also explains the temperature dependence of the shift. At high temperatures, the term D0(T) prevails in (1), [START_REF] Nenashev | Effect of Electric Field on Diffusion in Disordered Materials. II. Two-and Three-Dimensional Hopping Transport[END_REF][START_REF] Nenashev | Effect of Electric Field on Diffusion in Disordered Materials. I. One-Dimensional Hopping Transport[END_REF] and only a weak shift is experimentally observed, see Figure 3b. At lower temperatures, the term D0(T) is reduced, and the field-assisted component A(F,T).F 2 then prevails, which explains why a more significant shift is observed at low temperatures.

CONCLUSION

To summarize, we have designed a plasmonic resonator coupled to the short-wave infrared absorption of HgTe NCs. Here, we use a plasmonic resonance to introduce an inhomogeneous electromagnetic field within the NC array. The diffusive transport, often seen as a limitation of NCbased devices, is used to collect the charges over different device areas. Under low field, photogenerated charges are only collected from the region close to the electrodes, where the absorption spectrum is redshifted with respect to that away from the electrodes. In this sense, we fully exploit the interplay between charge transport and optical properties, going beyond the usual approach where cavity resonances are spectrally matched with material absorption [START_REF] Palaferri | Room-Temperature 9 µm-Wavelength Photo-Detectors and GHz-Frequency Heterodyne Receivers[END_REF] . In our case, we engineer the photoconduction by coupling on-demand electromagnetic maps and charge collection through hopping. From this strategy, we obtain a bias tunable photodetector, where the photocurrent peak can be tuned over 15 meV for extended short-wave infrared absorption.

Figure 1 :

 1 Figure 1: HgTe NCs coupled to plasmonic resonator. a. Schematic of the plasmonic resonator coupled to a HgTe NC film. b. Room temperature absorption spectrum of the HgTe NCs. The background is a transmission electron microscopy image of those HgTe NCs. c. Experimental reflectivity spectra measured at room temperature and at 15° incidence angle for a NC film deposited on a gold layer (blue circles), and for the device depicted in part a with s = 500 nm and p = 1350 nm. The top panel presents the experimental (green circles) and simulated (red line) reflectivity spectrum with TM polarized radiation, the bottom one presents the experimental spectrum in TE polarization. The inset also shows the electric field map associated with the resonance at 3200 cm - 1 (spoof surface plasmon resonance) and 2160 cm -1 (patch cavity mode) in TM polarization, and in correspondence to the broad feature observed in TE.

Figure 2 :

 2 Figure 2: Detection performance. a. Photocurrent spectra under unpolarized illumination at 80 K for the film in the plasmonic cavity. b. Responsivity (under blackbody illumination at 980 °C) and detectivity (at 1 kHz) as a function of operating temperature for VDS=0.5 V. c. Current as a function of time while the light (λ=1.55 µm) is turned on and off. d. Currents (dark and under illumination) as a function of gate bias, for VDS=0.2 V, at 150 K. The ratio of those two currents is plotted on the right axis as a function of the gate bias.

Figure 3 :

 3 Figure 3: Bias-induced spectral shift. a. Photocurrent spectra for non-polarized light under various applied bias voltages. b. Energy of the first photocurrent peak as a function of the applied bias for various operating temperatures. c. Photocurrent spectra for non-polarized light for two different gate bias voltages.

Figure 4 :,

 4 Figure 4: Origin of the bias-induced shift for TM polarization. a. Simulated absorption map for the TM polarization. For the device depicted in Figure 1a, we define two zones: region 1 is between electrodes, while region 2 is above electrodes. b. Simulated absorption spectra in region 1 and 2 defined on part a. c. (resp. d.) Product of the absorbance by exp(-r 2 /2LD 2 ), with r the distance to the electrode while LD is set equal to 20 nm (resp. 100 nm). e. Normalized photoresponse (calculated by ∬ 𝐴𝑏𝑠(𝑟, 𝜆). 𝑒𝑥𝑝 (-𝑟 2 2𝐿 𝐷 2 ) 𝑑𝑆

ACKNOWLEDGEMENTS

The project is supported by ERC starting grants blackQD (grant n° 756225). We acknowledge the use of clean-room facilities at the "Centrale de Proximité Paris-Centre". This work has been supported by Region Ile-de-France in the framework of DIM Nano-K (grant dopQD). This work was supported by French state funds managed by the ANR within the Investissements d'Avenir programme under reference ANR-11-IDEX-0004-02 and, more specifically, within the framework of the Cluster of Excellence MATISSE and by grants IPER-Nano2 (ANR-18CE30-0023-01), Copin (ANR-19-CE24-0022), Frontal (ANR-19-CE09-0017), Graskop (ANR-19-CE09-0026), NITQuantum (ANR-20-ASTR-0008-01), Bright and MixDferro. AC thanks Agence Innovation Defense for PhD funding.

Supporting Information

HgTe NCs synthesis and their characterization, electromagnetic simulation model, tight-binding simulations, device fabrication process, electrical characterization and the description of associated setup, and finally additional data relative to TE polarization.

COMPETING INTEREST

The authors declare no competing financial interests.


Daqing Jia

Qinzhi Li

Tao Luo

Olivier Monfort

Gilles Mailhot

Marcello Brigante

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

Impacts of environmental levels of hydrogen peroxide and oxyanions on the redox activity of MnO 2 particles

Keywords: birnessite, redox, hydrogen peroxide, oxyanion

Despite the widespread presence of hydrogen peroxide in surface water and groundwater systems, little is known about the impact of environmental levels of H 2 O 2 on the redox activity of minerals. Here we demonstrate that environmental concentrations of H 2 O 2 can alter the reactivity of birnessite-type manganese oxide, an earth-abundant functional material, and decrease its oxidative activity in natural systems across a wide range of pH values (4-8). The H 2 O 2 -induced reductive dissolution generates Mn(II) that will re-bind to MnO 2 surfaces, thereby affecting the surface charge of MnO 2 . Bisphenol A (BPA), as a ubiquitous environmental contaminant with endocrine disruption potential, was chosen to be the target compound, while the competition of BPA and Mn(II) to interact with reactive surface sites may cause suppression of the oxidative ability of MnO 2 . This suppressive effect becomes more effective in presence of oxyanions such as phosphate or silicate at concentrations comparable to those encountered in natural waters. Unlike nitrate, adsorption of phosphate or silicate onto birnessite increased in presence of Mn(II) added or generated through H 2 O 2 -induced reduction of MnO 2 . This suggests that naturally occurring anions and H 2 O 2 may have synergetic effects on the reactivity of birnessite-type manganese oxide at a range of environmentally relevant H 2 O 2 amounts. As layered structure manganese oxides play a key role in the global carbon cycle as well as pollutant dynamics, the impact of environmental levels of hydrogen peroxide (H 2 O 2 /MnO 2 molar ratio ≤ 0.3) should be considered in environmental fate and transport models.

Introduction

Layered structure manganese oxides (e.g. birnessite) 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] The presence of exchangeable hydrated cations within the interlayers combined with a high content of vacancy sites and variable oxidation states of manganese make them powerful sorbents and oxidants. [START_REF] Wang | Diffusion-and pH-dependent reactivity of layertype MnO 2 : Reactions at particle edges versus vacancy sites[END_REF][START_REF] Birkner | Thermodynamics of manganese oxides: Sodium, potassium, and calcium birnessite and cryptomelane[END_REF] As a result, birnessite-type manganese oxides play a significant role in controlling the cycles of several key nutrients as well as the fate and mobility of inorganic and organic contaminants. [START_REF] Wick | Thallium sorption onto manganese oxides[END_REF][START_REF] Lafferty | Arsenite oxidation by a poorly crystalline manganese-oxide. 2. Results from X-ray absorption spectroscopy and X-ray diffraction[END_REF][START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF][START_REF] Pokharel | Water flow and dissolved Mn(II) alter transformation of pipemidic acid by manganese oxide[END_REF][START_REF] Huang | Redox reactions of iron and manganese oxides in complex systems[END_REF][START_REF] Ricko | Reductive transformations of dichloroacetamide safeners: effects of agrochemical co-formulants and iron oxide manganese oxide binary-mineral systems[END_REF] Acid birnessite has been widely investigated because it is structurally similar to biogenic "natural" manganese oxides. [START_REF] Villalobos | Structural model for the biogenic Mn oxide produced by Pseudomonas putida[END_REF][START_REF] Webb | Structural characterization of biogenic Mn oxides produced in seawater by the marine bacillus sp. strain SG-1[END_REF][START_REF] Zhu | Cation effects on the layer structure of biogenic Mn-oxides[END_REF][START_REF] Villalobos | Characterization of the manganese oxide produced by Pseudomonas putida strain MnB1[END_REF] The birnessite reactivity is mainly affected by solution pH, surface properties and composition of MnO 2 and structural characteristics of contaminants (i.e., BPA). [START_REF] Wick | Thallium sorption onto manganese oxides[END_REF][START_REF] Lafferty | Arsenite oxidation by a poorly crystalline manganese-oxide. 2. Results from X-ray absorption spectroscopy and X-ray diffraction[END_REF][START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF][START_REF] Zhu | Cation effects on the layer structure of biogenic Mn-oxides[END_REF] It can also be influenced by the presence of naturally occurring compounds, e.g. cations, anions, natural organic matter and other redox-active compounds. Among the latter, hydrogen peroxide (H 2 O 2 ) has an ambivalent redox activity and is commonly found in aquatic environments at concentration ranging from nM to µM. [START_REF] Cooper | Hydrogen peroxide concentration in a northern lake: Photochemical formation and diel variability[END_REF][START_REF] Zika | Spatial and temporal variations of hydrogen peroxide in gulf of Mexico waters[END_REF][START_REF] Gonzalez-Flecha | Homeostatic regulation of intracellular hydrogen peroxide concentration in aerobically growing, Escherichia Coli[END_REF][START_REF] Guillén | Production of hydrogen peroxide by aryl-alcohol oxidase from the ligninolytic fungus Pleurotus Eryngii[END_REF][START_REF] Garg | Photochemical Production of Superoxide and Hydrogen Peroxide from Natural Organic Matter[END_REF][START_REF] Moffett | An investigation of hydrogen peroxide chemistry in surface waters of vineyard sound with H218O2 and 18O2[END_REF][START_REF] Roe | Measurement of dark, particle-generated superoxide and hydrogen peroxide production and decay in the subtropical and temperate North Pacific Ocean[END_REF][START_REF] Cooper | A chemiluminescence method for the analysis of H2O2 in natural waters[END_REF][START_REF] Yuan | Production of hydrogen peroxide in groundwater at rifle, Colorado[END_REF][START_REF] Liao | Extensive dark production of hydroxyl radicals from oxygenation of polluted river sediments[END_REF][24][START_REF] Page | Dark Formation of hydroxyl radical in arctic soil and surface waters[END_REF][START_REF] Begg | Plutonium desorption from mineral surfaces at environmental concentrations of hydrogen peroxide[END_REF] In environmental systems, photodependent reactions mediated by natural organic matter and biological-mediated processes dominate H 2 O 2 production. [START_REF] Begg | Plutonium desorption from mineral surfaces at environmental concentrations of hydrogen peroxide[END_REF] Sunlight-induced photochemical reactions mediated by organic ligands and/or biochemical metabolic processes can result in amounts of H 2 O 2 up to 20 µM. [START_REF] Cooper | Hydrogen peroxide concentration in a northern lake: Photochemical formation and diel variability[END_REF][START_REF] Zika | Spatial and temporal variations of hydrogen peroxide in gulf of Mexico waters[END_REF][START_REF] Gonzalez-Flecha | Homeostatic regulation of intracellular hydrogen peroxide concentration in aerobically growing, Escherichia Coli[END_REF][START_REF] Guillén | Production of hydrogen peroxide by aryl-alcohol oxidase from the ligninolytic fungus Pleurotus Eryngii[END_REF][START_REF] Garg | Photochemical Production of Superoxide and Hydrogen Peroxide from Natural Organic Matter[END_REF] Input from rainwater can reach up to 40 μM [START_REF] Roe | Measurement of dark, particle-generated superoxide and hydrogen peroxide production and decay in the subtropical and temperate North Pacific Ocean[END_REF][START_REF] Cooper | A chemiluminescence method for the analysis of H2O2 in natural waters[END_REF] , while hydrogen peroxide production has been often observed in irradiated seawater. [START_REF] Moffett | An investigation of hydrogen peroxide chemistry in surface waters of vineyard sound with H218O2 and 18O2[END_REF][START_REF] Roe | Measurement of dark, particle-generated superoxide and hydrogen peroxide production and decay in the subtropical and temperate North Pacific Ocean[END_REF][START_REF] Cooper | A chemiluminescence method for the analysis of H2O2 in natural waters[END_REF] Light-independent generation of H 2 O 2 has been recently reported in river sediments and in groundwater of an alluvial aquifer, which is likely to occur in transitional redox environments where reduced elements react with oxygen. [START_REF] Yuan | Production of hydrogen peroxide in groundwater at rifle, Colorado[END_REF][START_REF] Liao | Extensive dark production of hydroxyl radicals from oxygenation of polluted river sediments[END_REF] For example, natural hydroquinones or hydroquinone moieties ubiquitously present in reduced organic matter can donate electrons to O 2 to generate H 2 O 2 . 24,25 H 2 O 2 production through microbial mechanism has also been shown to enhance iodide oxidation and organo-iodine formation in soils and sediments. [START_REF] Li | Bacterial production of organic acids enhances H2O2-dependent iodide oxidation[END_REF] All these studies suggested that H 2 O 2 can be formed not only in oceanic and atmospheric systems but also in the subsurface environment. Finally, higher concentrations (up to 1 mM) have been reported in nuclear waste repository scenarios when ionizing radiation of groundwater adjacent to spent nuclear fuel result in H 2 O 2 production. [START_REF] Amme | Effects of Fe(II) and hydrogen peroxide interaction upon dissolving UO2 under geologic repository conditions[END_REF] Despite the catalyzed decomposition of hydrogen peroxide by MnO 2 has been well investigated [START_REF] Watts | Oxidative and reductive pathways in manganesecatalyzed Fenton's reactions[END_REF][START_REF] Pham | Inhibitory effect of dissolved silica on H2O2 decomposition by iron(III) and manganese(IV) oxides: Implications for H₂ O₂ -based in situ chemical oxidation[END_REF][START_REF] Kamagate | Mineralization enhancement of pharmaceutical contaminants by radical-based oxidation promoted by oxide-bound metal ions[END_REF] , knowledge is very limited on the effect of H 2 O 2 on the reactivity of birnessite at environmental levels of H 2 O 2 . In engineering applications, metal oxides such as MnO 2 are generally investigated for catalytic decomposition of H 2 O 2 and then formation of reactive transient species using very high concentrations of hydrogen peroxide (H 2 O 2 ) (≥ 0.5 mM). [START_REF] Watts | Oxidative and reductive pathways in manganesecatalyzed Fenton's reactions[END_REF][START_REF] Pham | Inhibitory effect of dissolved silica on H2O2 decomposition by iron(III) and manganese(IV) oxides: Implications for H₂ O₂ -based in situ chemical oxidation[END_REF][START_REF] Kamagate | Mineralization enhancement of pharmaceutical contaminants by radical-based oxidation promoted by oxide-bound metal ions[END_REF] While these remediation studies aimed to eliminate pollution in contaminated systems, they overlooked the influence of H 2 O 2 on the electron transfer heterogeneous reaction and then the oxidative activity of birnessite. In addition, little is known about the influence of H 2 O 2 on the birnessite composition and interactions of redox products (e.g.

Mn(II)) with MnO 2 surfaces.

In this study, we examine how H 2 O 2 affects the reactivity of birnessite under environmentally relevant conditions. To monitor H 2 O 2 -induced changes in birnessite reactivity, we used Bisphenol A (BPA) as a model compound because it has a strong reactivity with the birnessite (δ-MnO 2 ), with a well-documented underlying redox mechanism. [START_REF] Balgooyen | Structural transformation of MnO2 during the oxidation of bisphenol A[END_REF][START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF][START_REF] Huang | Effect of MnO2 phase structure on the oxidative reactivity toward bisphenol A degradation[END_REF][START_REF] Balgooyen | Impact of bisphenol A influent concentration and reaction time on MnO2 transformation in a stirred flow reactor[END_REF] Acid birnessite, a well-established laboratory synthesized analog of layered birnessite mineral, was chosen as a representative manganese dioxide mineral. The removal kinetics of BPA have been assessed in presence of variable amounts of H 2 O 2 under a wide range of pH (4, 6.5 and 8) under aerobic and anaerobic conditions. Alteration of birnessite structure and composition was monitored using XRD and titration experiments under different conditions (H 2 O 2 /MnO 2 molar ratio, dissolved Mn(II) amount and pH value). To check whether the impact of H 2 O 2 can persist in natural waters, changes in birnessite reactivity were investigated in presence of commonly found anionic ligands such as phosphate, nitrate and silicate. We notably demonstrated that at H 2 O 2 concentrations as low as 10 µM or with H 2 O 2 /MnO 2 molar ratio as low as 0.03, the reactivity of acid birnessite can be affected, which could alter biogeochemical cycles as well as pollutant dynamics. and MnCl 2 were purchased from Sigma-Aldrich and were all AR grade. All chemicals were used as received without further purification. All solutions were prepared in ultrapure water obtained from a water purification system (Millipore, resistivity 18.2 MΩ cm).

Materials and Methods

Synthesis and characterization of acid birnessite

Acid birnessite was prepared following the procedure of McKenzie [START_REF] Mckenzie | The synthesis of birnessite, cryptomelane, and some other oxides and hydroxides of manganese[END_REF] , and Mn(III)-rich MnO 2 was synthesized according to previous published methods. [START_REF] Sun | Effects of Fe (II) on Cd (II) immobilization by Mn (III)-rich δ-MnO2[END_REF][START_REF] Simanova | Probing the sorption reactivity of the edge surfaces in birnessite nanoparticles using nickel (II)[END_REF] More details are provided in the Supporting Information (SI). The solid was characterized using X-ray powder diffraction (XRD) using the Bruker AXS D8 Advance diffractometer (θ-2θ Bragg-Brentano geometry) using monochromatized CuKα1 radiation over the range of 10º-100º 2θ at a step size of 0.02º. X-ray diffraction (XRD) confirmed that the only product of the synthesis was poorly-crystalline hexagonal birnessite. The Brunauer-Emmett-Teller (BET) specific surface area of the synthetic birnessite measured by multipoint N 2 adsorption was 60 ± 2.5 (SD) m 2 g -1 . The Average Oxidation State (AOS= 3.98 (±0.02)) of synthetic birnessite and (AOS= 3.65 (±0.02)) of Mn(III)-rich MnO 2 were measured using the oxalic acid-permanganate back-titration method (more details are provided in SI). Scanning Electron Microscope (SEM; JEOL JSM-7100F) and High-resolution Transmission Electron Microscope (HRTEM; JEOL 2100 LaB 6 ) images showed a nanoflower-shaped birnessite consisting of nanoflakes aggregations (see Figure S1).

Kinetics experiments and analyses

Reactivity changes assessment were investigated as following: 100 mL aqueous suspension of acid birnessite (AB) corresponding to an initial concentration of 345 μM was first prepared and then an appropriate amount of stock solutions of H 2 O 2 , BPA, and/or Mn(II) were added to start the reaction at room temperature. Since the involved reactions can consume protons (see below), the pH was adjusted to the desired value and then kept constant throughout the reaction using a pH meter (Cyberscan 510, Thermo Scientific) by adding 0.1 M NaOH or HCl. The suspension was stirred at an agitation speed of 350 rpm during the experiments. 1 mL of aqueous sample was withdrawn at different time interval, filtered through 0.22 μm PTFE filter, and then analyzed. The same batch experiments were carried out in the presence of silicate and phosphate to test the effect of anions and H 2 O 2 on BPA oxidation.

Potential generation of hydroxyl radical was monitored through the fluorescence emission spectrum of 2-hydroxyterephthalic acid (λ ex 320 nm, λ em 425 nm) using a spectrofluorometer (Shimadzu RF-5301PC) (more information at SI).

The BPA concentration was determined by a high-performance liquid chromatography system (HPLC, Waters Alliance) equipped with a diode array detector (DAD).

Chromatographic separation was performed using a Nucleodur 100-5 C18 column (150 × 4.6 mm, 5 μm of particle size, Macherey Nagel). The detection wavelength of BPA was set at 277 nm and the column temperature was kept at 30 °C. Methanol and water were mixed as the mobile phase under a gradient eluent mode, and the percentage of methanol changed with time was as follows: from 0 to 7 min was increased from 70 to 90% and kept constant up to 9 min. The flow rate was set at 0.8 mL min -1 . The byproducts of BPA oxidation were identified using Ultraperformance 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.

H 2 O 2 concentration was determined using a spectrofluorometric detection as described in our previous work. [START_REF] Wu | Fe (III)-EDDS Complex in Fenton and photo-Fenton processes: From the radical formation to the degradation of a target compound[END_REF] Briefly, samples containing H 2 O 2 were mixed with 4hydroxyphenylacetic acid to form the stable 4-hydroxyphenyl acetic acid dimer in the presence of peroxidase (POD, Sigma-Aldrich). The 4-hydroxyphenylacetic acid dimer was quantified using a Cary Eclipse fluorescence spectrophotometer. The excitation wavelength was set at 320 nm and emission wavelength maximum was determined at 420 nm.

Concentration was calculated using calibration curve obtained using different concentrations of H 2 O 2 .

Dissolved Mn(II) concentration in aqueous solution was determined by Atomic

Absorption Spectroscopy (AAS, PerkinElmer). All the samples were first filtered using a PTFE 0.22 μm filter and then mixed with 2% nitric acid (HNO 3 , 65%, Sigma-Aldrich) prior for analyses. Dissolved Mn(II) standard solution for AAS (TraceCERT®, 1000 mg L -1 Mn(II) in 2% nitric acid, Sigma-Aldrich) was used for the calibration curve. The detection wavelength was set at 279 nm and the AAS detection limit is below 0.02 μM. Mn(Ⅱ) removal was calculated as the difference between the initial and final Mn(Ⅱ) solution concentrations.

The impact of anions on reactivity changes was investigated as mentioned above, but with addition of the chosen anion with BPA or Mn(II) to the AB suspension. Silicate concentration was determined by the molybdenum-blue colorimetric method [START_REF] Mullin | The colorimetric determination of silicate with special reference to sea and natural waters[END_REF] while phosphate and nitrate concentrations by ion chromatography. [START_REF] Ruiz-Calero | Ion chromatographic separations of phosphorus species: a review[END_REF] XRD was performed to determine the possible transformation of solids upon Mn(Ⅱ) sorption.

Since Mn(II) oxidation by molecular oxygen may take place, particularly at alkaline pH values, we compared the removal kinetics under aerobic (open atmosphere) vs anaerobic conditions (glove box N 2 :H 2 98:2). Ultrapure water under anoxic environments were sparged with nitrogen to remove oxygen prior to use. These preliminary tests showed that there is no effect of oxic conditions on the BPA removal in presence of H 2 O 2 or Mn(II). Great attention has also been paid to the pH adjustment in all experiments, since the involved reactions are sensitive to pH value (see below). All experiments were conducted in triplicates and the standard deviation was calculated for all experimental series.

Results and Discussion

Effect of H 2 O 2 on the oxidative ability of birnessite

BPA removal was monitored in the presence of acid birnessite (AB) at different H 2 O 2 concentrations at pH 6.5 (Fig. 1). In the absence of H 2 O 2 , almost complete removal of BPA was observed after 24h of reaction time. Mass balance showed that the oxidation reaction was mainly involved in the removal of BPA in presence of MnO 2 , while adsorption was very low under our experimental conditions (i.e. less than 5% of initial BPA), which is consistent with previous investigations. [START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF][START_REF] Balgooyen | Structural transformation of MnO2 during the oxidation of bisphenol A[END_REF][START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF] Indeed, BPA containing two hydroxyphenyl functionalities is known to weakly interact with mineral surfaces, resulting in lower adsorption affinity to metal-oxides. [START_REF] Balgooyen | Structural transformation of MnO2 during the oxidation of bisphenol A[END_REF] As previously reported, [START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF][START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF][START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF] binding to birnessite is followed by an electron transfer process resulting in the concomitant oxidation of sorbed compound and reduction of surface-bound Mn(IV) to yield Mn(III) that can be further reduced to give Mn(II).

Mn IV O 2 + BPA + 4 H + → Mn II + BPA ox + 2 H 2 O (Eq.1)
Electron exchange of BPA with MnO 2 form a radical, followed by a series of reactions including radical coupling, fragmentation, substitution and elimination to form multiple byproducts. [START_REF] Remucal | A critical review of the reactivity of manganese oxides with organic contaminants[END_REF][START_REF] Balgooyen | Structural transformation of MnO2 during the oxidation of bisphenol A[END_REF][START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF] LC/MS analysis identified one predominant species of mass-to-charge ratio of m/z = 135 as the most dominant byproduct of BPA over the first reaction time (90 min), which likely corresponds to 4-hydroxycumyl alcohol according to previous investigations. [START_REF] Im | Identification of 4-hydroxycumyl alcohol as the Major MnO2-mediated bisphenol A transformation product and evaluation of its environmental fate[END_REF] In the presence of hydrogen peroxide, the oxidative removal of BPA gradually decreased amounts exhibited the same order, pH 4 > pH 6.5 > pH 8 (Fig. 2). In the absence of H 2 O 2 , k was decreased almost 20-fold when pH increased from 4 to 8. This suggests that acidic conditions favored BPA removal, a trend also observed for other organic compounds reacting with MnO 2 . [START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF][START_REF] Huang | Effect of MnO2 phase structure on the oxidative reactivity toward bisphenol A degradation[END_REF][START_REF] Zhang | Kinetic modeling of oxidation of antibacterial agents by manganese oxide[END_REF] We attribute this to variability in two pH-dependent factors: 1) speciation of BPA that affects binding to MnO 2 surfaces and then oxidation (Fig. S4), and 2) redox-potential of MnO 2 that decreases when the pH increased from 4 to 8. 45 At relatively low pH, favorable electrostatic interactions between neutral BPA molecules and the negatively charged surface of MnO 2 (PZC of MnO 2 is ~ 2.3-2.9) may exist. On the other hand, reductive conversion of MnO 2 into Mn(Ⅱ) is dependent on the amount of protons, which would result in increase in BPA removal rate when the pH decreases (Eq.1). In the presence of H 2 O 2 , the removal rate constants of BPA sharply decreased with increasing in H 2 O 2 amount whatever the investigated pH value (Fig. 2).

with
As for BPA, it is previously reported that hydrogen peroxide can also reduce Mn(IV) to Mn(III) and then Mn(II) as follows 46,47 :

Mn IV O 2 + H 2 O 2 + 2 H + → Mn II + O 2 + 2 H 2 O (Eq.2)
To check this possibility under our experimental conditions, the reductive dissolution of MnO 2 in presence of various amounts of H 2 O 2 was monitored at three pH values (Fig. 3). In both Eq.1 and Eq.2, protons are directly involved in the oxidation of BPA as well as the reductive dissolution of MnO 2 . In the absence of H 2 O 2 , dissolved Mn(II) was only detected at pH 4, while it was below the detection limit at pH 6.5 or 8. Increasing H 2 O 2 concentration from 50 µM to 200 µM enhanced the amount of dissolved Mn(II), at the three investigated pH values. At each pH value, the amount of dissolved Mn(II) increased first, reached a maximum and then decreased. This two-step behavior is consistent with the H 2 O 2 decomposition over time, where the complete decomposition of H 2 O 2 observed after approximately 15 min of reaction time coincides with the peak observed for dissolved Mn(II) generation (Fig. S3).

Once the added H 2 O 2 is fully decomposed, the generated dissolved Mn(II) will in turn sorb onto MnO 2 surfaces. Because higher pH implies more Mn(Ⅱ) binding [START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Elzinga | Reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Wang | Structural Transformation of Birnessite by Fulvic Acid under Anoxic Conditions[END_REF] , greater amounts of Mn(II) were observed at low pH values (Fig. 3). When a further addition of H S3, where both reactions Eq.1 and Eq.2 occurred simultaneously until complete inhibition of Eq.1 at the highest H 2 O 2 amount (200 µM). Though H 2 O 2 was full decomposed after 15 min of reaction time (Fig. S2), the suppressive effect on the birnessite reactivity was still observed over 24 h of reaction time (Fig. 1). This was confirmed by carrying out sequentially the reactions (Eq.2) and (Eq.1). Indeed, suppression of BPA removal capacity was also observed when MnO 2 and H 2 O 2 were allowed to react first (i.e. preequilibration step of 30 min) until total decomposition of H 2 O 2 before addition of BPA (not shown). Therefore, the H 2 O 2 -mediated reduction of MnO 2 (Eq.2) considerably affects the surface reactivity of birnessite and its ability to remove BPA. This reductive dissolution generates Mn(II) ions which will in turn bind to MnO 2 surfaces and then be oxidized into higher valence Mn:

Mn II + MnO 2 → Mn II /Mn III -MnO 2 (Eq.3)

This oxidation reaction may be made possible by the residual H 2 O 2 and/or generated O 2 (through Eq.2), which is pH-dependent. The contribution of O 2 from ambient air is excluded since reactivity assessment tests investigated under aerobic vs anaerobic conditions showed similar behavior in term of BPA removal.

It is worth noting that strong aggregation of MnO 2 particles and then fast sedimentation was observed upon addition of H 2 O 2 . This phenomenon can be ascribed to charge neutralization (i.e. surface charge switch from negative to positive) upon Mn(II) ions binding to negatively charged MnO 2 surfaces, as previously reported. [START_REF] Chenga | Mn2+ effect on manganese oxides (MnOx) nanoparticles aggregation in solution: Chemical adsorption and cation bridging[END_REF] Furthermore, disproportionation/comproportionation reaction (Eq.4) may occur within the MnO 2 , i.e. Mn II exchanges electrons with Mn(IV) to form two Mn(III) centers and, conversely, two Mn(III) centers can disproportionate to form Mn(II) and Mn(IV) centers [START_REF] Zhu | Cation effects on the layer structure of biogenic Mn-oxides[END_REF][START_REF] Balgooyen | Structural transformation of MnO2 during the oxidation of bisphenol A[END_REF][START_REF] Balgooyen | Impact of bisphenol A influent concentration and reaction time on MnO2 transformation in a stirred flow reactor[END_REF][START_REF] Wang | Structural Transformation of Birnessite by Fulvic Acid under Anoxic Conditions[END_REF][START_REF] Zhao | Redox Reactions between Mn (II) and Hexagonal Birnessite Change its Layer Symmetry[END_REF] :

Mn II + Mn IV ↔ 2Mn III (Eq.4)

Since this reaction is pH-dependent, we have measured the average oxidation state (AOS) of birnessite over the whole investigated pH range (4-8). The AOS of samples reacted with BPA (25 µM) alone only slightly decreased from 3.98 to 3.90 (±0.02). However, partial reduction of MnO 2 upon addition of 200 µM of H 2 O 2 dropped down the AOS to 3.74 (±0.04) at pH 4 and 6.5, and 3.61 (±0.04) at pH 8. This decrease suggests that percentages of Mn(III)

or Mn(Ⅱ) or both are relatively higher at the end of reaction with respect to the initial sample.

The formation of Mn(III) was further confirmed by the detection of Mn(III)-pyrophosphate complex at 480 nm, Mn(III) being stabilized through ligand-binding complexes. [START_REF] Webb | Evidence for the presence of Mn(III) intermediates in the bacterial oxidation of Mn(II)[END_REF] The effect of H 2 O 2 on the structure of birnessite will be discussed in the following section.

Collectively, these results suggest that the H 2 O 2 induced reduction of MnO 2 is an effective process, and the fast generation of Mn(II) is key for suppressed oxidative capacity of MnO 2 towards BPA. To confirm the impact of Mn(II) on the reactivity of MnO 2 , BPA removal kinetics were investigated in Mn(Ⅱ)-amended birnessite suspensions in the following section.

Effects of Mn(Ⅱ) on the removal capacity of MnO 2

The initial rate constants of BPA removal sharply decreased with increasing in dissolved Mn(II) amount at pH 4 and 6.5, while at pH 8 it increased first between 0 and 10 µM but later decreased with increasing in initial Mn(Ⅱ) amount (Fig. 5). It is previously reported that the adsorption of cations such as Ca(Ⅱ) or Mg(Ⅱ) may change the surface charge from negative to positive by exchange of H + on the MnO 2 surface. [START_REF] Takamatsu | The role of Mn(II)-rich hydrous manganese oxide in the accumulation of arsenic in lake sediments[END_REF][START_REF] Yao | Adsorption of phosphate on manganese dioxide in seawater[END_REF] Similarly, the Mn(Ⅱ) adsorption should lead to a decrease in the negative surface charge, thereby altering the binding capacity of MnO 2 . When the Mn(II) amount further increases, we speculate a competition between BPA and Mn(Ⅱ) to bind at reactive sites of MnO 2 during the first kinetic phase, as recently observed for quinolones. [START_REF] Li | Coupled effects of Mn(II), pH and anionic ligands on the reactivity of nanostructured birnessite[END_REF] As for H 2 O 2 , dissolved Mn(Ⅱ) displayed suppressive effects on the removal rate of BPA and the rate constants over the whole Mn(Ⅱ) concentration range (10 -100µM) exhibited the same order, pH 4 > pH 6.5 > pH 8 (Fig. 5). In addition to the comproportionation reaction that is pH-dependent, the decrease in the oxidative ability of MnO 2 in the presence of dissolved Mn(Ⅱ) may result from competition of compound and Mn(Ⅱ) to interact with reactive surface sites. The strong binding of Mn(Ⅱ) to MnO 2 surfaces suggests competition of BPA and Mn(Ⅱ) for surface sites, which is pH dependent. [START_REF] Lin | Oxidative removal of bisphenol A by manganese dioxide: efficacy, products, and pathways[END_REF][START_REF] Zhang | Kinetic modeling of oxidation of antibacterial agents by manganese oxide[END_REF] Dissolved Mn(II) measurements showed that complete removal of dissolved Mn(Ⅱ) was observed at both pH 6.5 and 8 when 100 µM Mn(Ⅱ) was used (not shown), while a concentration of ~30 µM was still measured at pH 4.

To check the affinity of birnessite for Mn(II) binding, sorption isotherms were determined at three pH values (4, 6.5 and 8) under aerobic conditions (Fig. S5). It should be noted that Mn(II) removal under oxic conditions is expected to be higher than under anoxic conditions, especially at alkaline pH values, which has been attributed to surface-catalyzed oxidation of Mn(II) by molecular oxygen. [START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Elzinga | Reductive transformation of birnessite by aqueous Mn(II)[END_REF] Here, only aerobic conditions were tested, since the aim of these sorption isotherms is to assess the Mn(II) affinity under oxidizing conditions imposed by the presence of H 2 O 2 . As expected, Mn(Ⅱ) binding to MnO 2 surfaces increased with pH increasing. Sorption isotherms showed different shape depending on pH value across the Mn(Ⅱ) concentration range. At pH 4 a typical L-shape with a plateau was observed, while the removal amount continuously increased with Mn(Ⅱ) concentration at higher pH values.

This high binding of Mn(II) at high pH value could cause Mn(III) enrichment in MnO 2 surfaces. [START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Elzinga | Reductive transformation of birnessite by aqueous Mn(II)[END_REF] Some studies have shown that at high Mn(II)/Mn(IV) ratio and pH > 7.5, dissolved Mn(II) can interact with hexagonal birnessite and then transfer electron to lattice Mn(IV) producing Mn(III). [START_REF] Lefkowitz | Influence of pH on the reductive transformation of birnessite by aqueous Mn(II)[END_REF][START_REF] Elzinga | Reductive transformation of birnessite by aqueous Mn(II)[END_REF] The buildup of Mn(III) will induce changes in mineral structure and composition by converting birnessite into lower-valence Mn phases. XRD analysis conducted on samples reacted with 100 µM of dissolved Mn(II) (Mn(II)/MnO 2 molar ratio = 0.3) showed no notable transformation of MnO 2 over the investigated pH range (Fig. 4). However, titration experiments showed that addition of 100 µM of dissolved Mn(Ⅱ) dropped down the AOS of MnO 2 from 3.98 to 3.60 (±0.05) at pH 4 and 6.5, and 3.50 (±0.05) at pH 8, suggesting more Mn(III) or Mn(Ⅱ) at the end of reaction with respect to the initial sample. A recent work showed that birnessite transformation into triclinic birnessite and/or 4 × 4 tunneled Mn oxide may occur at low Mn(II)/MnO 2 ratios (0.09 and 0.13), while secondary phases such as MnOOH and Mn 3 O 4 can be generated at high Mn(II)/MnO 2 ratios (0.5 and 1). [START_REF] Yang | Inhibition of Oxyanions on Redox-driven Transformation of Layered Manganese Oxides[END_REF] In our XRD patterns, broad bands (hump) have appeared between 12° and 22° and between 25° and 35° in the samples reacted with Mn(II) (Mn(II)/MnO 2 molar ratio = 0.3) or H 2 O 2 (H 2 O 2 /MnO 2 molar ratio = 0.6), suggesting changes in the structure or ordering of the birnessite mineral sheets (Fig. 4). We also observed these broad bands in the XRD pattern of synthetic Mn(III)-rich birnessite sample (see Fig. 4). This is consistent with previous studies, which suggested that Mn III formation can be made in edge-sharing complexes on MnO 2 edge sites or around vacancy sites in octahedral layers of MnO 2 . [START_REF] Lafferty | Arsenite oxidation by a poorly crystalline manganese-oxide. 2. Results from X-ray absorption spectroscopy and X-ray diffraction[END_REF] In addition, a small peak near 20° only present in the patterns of samples reacted at pH 8 would suggest the presence of feitknechtite (β-MnOOH), as previously reported. 48 [START_REF] Yao | Adsorption of phosphate on manganese dioxide in seawater[END_REF][START_REF] Yang | Inhibition of Oxyanions on Redox-driven Transformation of Layered Manganese Oxides[END_REF][START_REF] Mustafa | pH effect on phosphate sorption by crystalline MnO2[END_REF][START_REF] Balistrieri | Adsorption of selenium by amorphous iron oxyhydroxide and manganese dioxide[END_REF] On the other hand, previous studies reported that the presence of dissolved silicate may decrease the oxidation rate constants of MnO 2 toward chlorinated compounds. [START_REF] Taujale | Impact of interactions between metal oxides to oxidative reactivity of manganese dioxide[END_REF][START_REF] Yu | Dissolved Silicate Enhances the Oxidation of Chlorophenols by Permanganate: Important Role of Silicate-Stabilized MnO2 Colloids[END_REF] They have attributed the decrease in MnO 2 reactivity to the surface-bound silicate but none of them have provided adsorption data of silicate onto MnO 2 . In the present work, we showed that silicate or phosphate has a very low affinity to the negatively charged MnO 2 surfaces under our experimental conditions, but the presence of 100 µM of Mn(II) significantly enhanced their adsorption amounts onto MnO 2 (Fig. S6). This increased adsorption might be due to the changes in both the surface charge of MnO 2 and the solution speciation of anionic ligands in the presence of divalent cations such as Mn(II). Indeed, the ligand speciation may be modified, since different coordination modes of aqueous complexes of phosphate with Mn(Ⅱ) ions have been previously reported. [START_REF] Sharma | Magnetic property studies of manganese-phosphate complexes[END_REF] However, no precipitation is possible under our experimental conditions due to the low degree of saturation, i.e. low aqueous concentration of free Mn(II) or ligand (log K s Mn(Ⅱ)/phosphate = -27.07). [START_REF] Friedl | Solid phases in the cycling of manganese in eutrophic lakes: New insights from EXAFS spectroscopy[END_REF] Furthermore, Mn(Ⅱ) binding should lead to a decrease in the negative surface charge, thereby enhancing ability to bind anions through electrostatic interactions. [START_REF] Chenga | Mn2+ effect on manganese oxides (MnOx) nanoparticles aggregation in solution: Chemical adsorption and cation bridging[END_REF][START_REF] Takamatsu | The role of Mn(II)-rich hydrous manganese oxide in the accumulation of arsenic in lake sediments[END_REF][START_REF] Yao | Adsorption of phosphate on manganese dioxide in seawater[END_REF] Therefore, the MnO 2bound Mn(II) system could adsorb more effectively anions such as phosphate or silicate through cation bridging, as recently reported for humic acid. [START_REF] Chenga | Mn2+ effect on manganese oxides (MnOx) nanoparticles aggregation in solution: Chemical adsorption and cation bridging[END_REF] This surface-Mn(Ⅱ)-ligand ternary complex may act as a barrier to electron transfer between BPA and Mn(IV) sites, thereby altering reactivity of MnO 2 surfaces.

It should be noted that silicate or phosphate (100 µM) did not significantly affect the decomposition rate of H 2 O 2 (200 µM), since total decomposition was achieved after approximately 15 min as previously observed in Figure S2. Previous works have investigated the catalytic activity of manganese oxides for hydrogen peroxide decomposition and generation of reactive oxygen species in the context of environmental remediation studies. [START_REF] Watts | Oxidative and reductive pathways in manganesecatalyzed Fenton's reactions[END_REF][START_REF] Pham | Inhibitory effect of dissolved silica on H2O2 decomposition by iron(III) and manganese(IV) oxides: Implications for H₂ O₂ -based in situ chemical oxidation[END_REF][START_REF] Kamagate | Mineralization enhancement of pharmaceutical contaminants by radical-based oxidation promoted by oxide-bound metal ions[END_REF] Despite silicate adsorption on MnO 2 surfaces was found negligible, the effect of silicate concentration (ranging from 0 to 1.5 mM) on the H 2 O 2 decomposition rate has been found to display relatively inhibition. [START_REF] Pham | Inhibitory effect of dissolved silica on H2O2 decomposition by iron(III) and manganese(IV) oxides: Implications for H₂ O₂ -based in situ chemical oxidation[END_REF] Although these works have used different experimental conditions (e.g. much higher concentration of ligands or H 2 O 2 ), and different MnO 2 types, the present findings call for in-depth consideration of the combined/synergistic effects that the cations and anions co-presents in the reaction medium may have on the MnO 2 reactivity.

Overall, these findings suggest that Mn(II) generation through H 2 O 2 -mediated reduction may alter both adsorption and redox transformation of environmental compounds.

Conclusions

Photochemical and dark production of H 2 O 2 make it ubiquitous not only in oceanic and atmospheric systems but also in the subsurface environment.

Figure 1 .

 1 Figure 1. Effect of H 2 O 2 concentration on the BPA removal at pH 6.5 at room temperature.

Figure 2 .

 2 Figure 2. Removal rate constants (h -1 ) of BPA as a function of H 2 O 2 dose at three pH values

Figure 3 .Figure 4 .

 34 Figure 3. Mn(II) formation under different H 2 O 2 concentrations and pH 4.0, 6.5 and 8.0.

Figure 5 .

 5 Figure 5. Removal rate constants (h -1 ) of BPA as a function of Mn(II) concentration at three

Figure 6 .H 2 O 2 /

 622 Figure 6. Removal rate constants (h -1 ) of BPA as a function of H 2 O 2 dose in absence or

2.1. Chemicals KMnO

  4 , HCl, NaCl, NaOH, HEPES, H 2 C 2 O 4 , H 2 SO 4 ,hydroxylamine hydrochloride, H 2 O 2 , BPA, HNO 3 , terephthalic acid, 2-hydroxyterephthalic acid, Na 2 SiO 3 , NaNO 3 , NaH 2 PO 4

  increasing in H 2 O 2 concentration, and completely suppressed at higher dose of H 2 O 2 (200 µM equivalent to a H 2 O 2 /MnO 2 molar ratio ~ 0.6). In these experiments, the H 2 O 2 decomposition is relatively fast whatever the investigated pH, since complete decomposition of H 2 O 2 was achieved after 15 min of reaction, even for the highest H 2 O 2 dose (See Fig. S2

and S3

). Monitoring of fluorescence emission spectra of suspensions using a terephthalate as a chemical probe indicated that no hydroxyl radical was generated under the experimental conditions of this study. Additional experiments using 1-propanol and t-butanol as hydroxylradical scavengers confirmed that the BPA removal is a no radical-based oxidation process regardless of the presence or absence of H 2 O 2 .

The suppressive effect of H 2 O 2 on the removal rate of BPA was then confirmed at two other pH values (4 and 8). The removal kinetics 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 in the investigated system. Instead, we calculated the initial rate constant (k , h -1 ) by linear regression of ln [BPA] /[BPA] o versus time over the first stage of reaction (i.e., 60 min). The initial rate constants for all investigated H 2 O 2

  This suggests that the reductive dissolution of MnO 2 by H 2 O 2 is a fast process, as compared to the reaction between MnO 2 and BPA.

	To check whether Eq.1 and Eq.2 may occur simultaneously during the heterogeneous
	reactions with MnO 2 , we compare the standard potential of involved reactions at the working
	pH range. MnO 2 has standard (reduction) potential varying from 0.99 to 0.76 V when the pH
	rises from 4 to 8 for MnO 2 + 4 H + + 2 e -→ Mn II + 2 H 2 O. This value is higher than the
	standard oxidation potential of O 2 /H 2 O 2 (O 2 + 2 H + + 2e -→H 2 O 2 0.68 V vs. NHE) or of BPA
	which varies between 0.75 V and 0.52 V in the pH range 4-8.

2 O 2 is made after total decomposition of the initial amount (e.g. after 15 min at pH 4), the amount of generated Mn(II) increased again and reached almost twice the first measured amount within approximately 15 min (not shown).

[START_REF] Yin | Amperometric biosensor based on tyrosinase immobilized onto multiwalled carbon nanotubes-cobalt phthalocyanine-silk fibroin film and its application to determine bisphenol A[END_REF][START_REF] Yin | Electrochemical oxidation behavior of bisphenol A at surfactant/layered double hydroxide modified glassy carbon electrode and its determination[END_REF][START_REF] Karabiberoğlu | Sensitive voltammetric determination of bisphenol A based on a glassy carbon electrode modified with copper oxide-zinc oxide decorated on graphene oxide[END_REF]

Since the latter redox couples have comparable oxidation potential, reactions (Eq.1) and (Eq.2) could simultaneously take place. Assuming that these reactions obey pseudo-first order kinetic equation because MnO 2 is considered in excess, the reaction rate depends on the concentration of reactants (i.e. H 2 O 2 or BPA). When the H 2 O 2 amount increases, the reaction 2 should become much faster than the reaction 1. This was experimentally shown in Figures

1 and 3

and Figure

3.3. Synergistic effects of anions and H 2 O 2 on the oxidative activity of MnO 2

  To check whether the suppression of reactivity at environmental relevant concentrations of H 2 O 2 can persist in presence of naturally occurring compounds, the initial rate constants k (h -1 ) of BPA removal were determined at pH 6.5 in presence of silicate, nitrate and phosphate, commonly found in natural systems (Fig.6). At 0 µM of H 2 O 2 , the presence of 100 µM of anions did not significantly influence the removal rate of BPA, probably because of weaker interactions of these anionic ligands with MnO 2 surfaces under the experimental conditions of this study (pH 6.5, 100 µM of anion, 345 µM of MnO 2 ). However, significant decrease in kinetic rate constants was observed in presence of silicate and phosphate when H 2 O 2 was added in the reaction medium (H 2 O 2 /MnO 2 molar ratio between 0.03 and 0.3), while no impact on the reactivity was observed in presence of nitrate (Fig.6). These results can be explained if the H 2 O 2 -mediated reduction of MnO 2 into Mn(II) make the presence of silicate or phosphate competitive towards BPA interactions with MnO 2 surfaces. This was confirmed by repeating removal kinetics of BPA in presence of 100 µM silicate or phosphate, but with the addition of Mn(II) instead of H 2 O 2 (data not shown). This suggests that presence of Mn(II) added or generated through H 2 O 2 -mediated reduction renders the MnO 2 surfaces more able to bind anionic ligands, thereby altering surface reactivity.

	It is previously reported that unlike nitrate, silicate or phosphate may interact with
	MnO 2 through hydrogen bonding or formation of outer-sphere complexes with surface
	hydroxyl groups of MnO 2 .

  Here, we have notably demonstrated that environmental levels of hydrogen peroxide can induce reductive dissolution of MnO 2 into Mn(II), thereby affecting the surface reactivity of MnO 2 . As the H 2 O 2 decomposition is a fast process, the generated dissolved Mn(II) binds to MnO 2 surfaces, altering further interactions with co-existing organic compounds. This may result from competition of organic compound and Mn(II) to interact with reactive surface sites and/or aqueous complexation with Mn(II). The presence of silicate or phosphate at concentrations comparable to those encountered in natural waters further decreased the reactivity of MnO 2 in presence of H 2 O 2 . Birnessite-bound Mn(II) adsorbed more effectively anionic ligands such as phosphate or silicate and thus reducing interactions with BPA at a range of environmentally relevant pH values. These findings suggest that naturally occurring anions and H 2 O 2 may have synergetic effects on the reactivity of nanostructured birnessite-type manganese oxide.As manganese oxides can break down high molecular weight humic substances into lower molecular weight organic molecules and/or stabilize dissolved organic carbon, these findings call for in-depth consideration of the impacts of environmental levels of H 2 O 2 and co-exiting anions on the global carbon cycle. Furthermore, the widespread presence of H 2 O 2 in surface

Acknowledgments

This work was supported by the Institut Universitaire de France, the Region Council of Auvergne Rhône-Alpes and the CNRS. We gratefully acknowledge the Chinese Scholarship Council of PR China for providing financial support for Daqing Jia and Qinzhi Li.

Conflicts of interest

There are no conflicts to declare.

Supplementary Material

Additional information regarding synthesis and characterization methods, SEM/TEM images, additional data on removal kinetics and H 2 O 2 decomposition, anions speciation and adsorption onto MnO 2 .