I degradation of carbon tetrachloride in the 𝛽-MnO2 activated H2O2 system through the generation of reactively transient species [12]. Zhu et al. employed the β-MnO2 nanorods to activate PDS for the removal of phenol. Efficient degradation of phenol was achieved in β-MnO2/PDS system through the generation of singlet oxygen ( 1 O2) [13]. Zhou et al. indicated the higher catalytic property of α-MnO2 than δ-MnO2 in PMS activation for 4-nitrophenol degradation because α-MnO2 owns more active sites, larger Brunauer-Emmett-Teller (BET) area, faster electron transfer rate, and better adsorption performance [14]. Moreover, it was reported that the structure of manganese oxides and the contents of Mn(III) on the surface of manganese oxides play a critical role in the oxidative and catalytic reactivity of manganese oxides. For instance, Huang et al. investigated the activation of PMS by MnO2 with different crystal phases (i.e. α-, β-, γ-, and δ-MnO2) [15]. The results demonstrated the important role of crystalline structure and Mn(III) contents on the catalytic reactivity of MnO2. Saputra et al. investigated the effect of Mn oxidation states (such as MnO, Mn2O3, Mn3O4, and MnO2) on the activation of PMS for phenol degradation [16]. The results showed that Mn2O3 has the highest ability on PMS activation among these manganese oxides. He et al. also demonstrated the efficient degradation of 2,4-dichlorophenol in the Mn III OOH activated PMS system [17].

Although the good activity of Mn(III/IV) oxides for radical precursors activation has been demonstrated, the underlying mechanisms of activation and generation of reactive species are not fully understood. Moreover, the previous studies focused on the elimination efficiency of pollutants, and overlooked the influence of oxidants (for example H2O2) on the oxidative activity of Mn(III/IV) oxides. Furthermore, most of the previous literature were conducted in the pure water. Indeed, the degradation efficiency of organic pollutants by manganese oxidesbased AOPs in the real wastewaters need to be evaluated. Therefore, the purpose of this study was to 1) investigate the impacts of H2O2 and oxyanions on the reactivity of acid birnessite. 2) investigate the activation mechanism of various radical precursors by Mn2O3 at near-neutral pH solution. 3) investigate the reactivity of manganite (𝛾-MnOOH) in PMS activation and evaluate the efficiency of manganite/PMS system in the removal of multi-pollutants in the real wastewater.

Introduction

With the increase of urbanization, industrialization, and the number of people on Earth, a large number of chemical products are consumed and discharged, leading to the occurrence and increase of emerging pollutants (EPs) in various water compartments [1]. The concentration of EPs detected in waters ranges from ng/L to µg/L [2]. The presence of EPs in the environment poses a threat to the safety of the ecosystem and the health of human beings [3,4]. However, the conventional wastewater treatment technologies (such as adsorption, precipitation, and filtration) have shown low efficiency on the removal of EPs [5]. EPs can enter the environment repeatedly through the discharging of the sewage treatment plants (STP), leading to the bioaccumulation of EPs in the various organism and water compartments [6]. Advanced oxidation processes (AOPs) have attracted much attention due to their high efficiency in the degradation of recalcitrant pollutants. AOPs are able to mineralize, through the generation of reactive radicals including sulfate and hydroxyl radicals, various organic pollutants [7].

The generation of reactive species in AOPs can be achieved through the activation of radical precursors (such as hydrogen peroxide (H2O2), peroxymonosulfate (PMS, HSO 5 -), and peroxydisulfate (PDS, S 2 O 8 2-)) by various techniques including ultraviolet light (UV), heat, electrochemistry, ultra-sonication, and transition metals [8]. Among these methods, the transition metals-based AOPs have been widely used in the treatment of wastewater due to the advantage of simple implementation. Fenton reaction is one of the typical transition metalsbased AOPs, which uses the Fe 2+ to activate H2O2 for the generation of hydroxyl radicals. The latter is an high oxidant species and is able to degrade organic pollutants [9]. However, the strict reaction conditions of the Fenton reaction (such as low pH) limited its application in practical water treatment [10]. In comparison, manganese (Mn) has shown a good reactivity in transition metal-based AOPs due to its various valence states and wider pH stability [11].

In the natural environment, manganese species mainly exist in the oxidation states of Mn(II), Mn(III), and Mn(IV). The activation of radical precursors by various manganese oxides for the removal of organic pollutants have been widely reported. For example, Watts et al. reported the In this work, different organic pollutants such as bisphenol A (BPA), caffeine (CAF), 17βestradiol (E2), and 17α-ethinylestradiol (EE2) were used as the model pollutants to investigate the activation mechanism of various radical precursors by manganese (III/IV) oxides. The main experimental contents were summarized as follows:

In the first part of our work, the impacts of environmental levels of hydrogen peroxide and oxyanions on the redox activity of MnO2 particles have been investigated. Specifically, the effect of different parameters such as the H2O2 concentration, pH values, Mn(II) contents, and oxyanion types (such as silicate, nitrate, and phosphate) on the reactivity of acid birnessite was evaluated.

In the second part of this manuscript, the degradation of caffeine by Mn2O3 activated peroxymonosulfate or peroxydisulfate system is described. In this study, the effects of solution pH, Mn2O3 dosages, PDS and PMS concentrations on the reactivity of Mn2O3 were reported.

The primary reactive species generated in Mn2O3 activated PDS or PMS system were identified using chemical scavenging experiments and electron paramagnetic resonance (EPR) spectroscopy. The mineralization efficiency of Mn2O3 + PDS or PMS system was monitored by measuring the total organic carbon (TOC) decrease as a function of reaction time. Moreover, the performance of Mn2O3 activated PDS or PMS to degrade CAF in the real conditions was evaluated using sewage treatment plant (STP) water.

In the third part of our work, the removal of three estrogenic compounds (BPA, E2, and EE2) by manganite-activated peroxymonosulfate was investigated. The effect of manganite and PMS dosages on the degradation rate of selected pollutants was studied. The scavenging tests, EPR and X-ray photoelectron spectroscopy (XPS) experiments were conducted to identify the main reactive species generated in the reaction and Mn oxidation state changes. The influence of various inorganic ions on the system of manganite/PMS was also studied and application for the degradation of three pollutants in the STP water was reported.

degradation in aqueous solutions, Applied Catalysis B: Environmental. [START_REF] Shah | Hydroxyl and sulfate radical mediated degradation of ciprofloxacin using nano zerovalent manganese catalyzed S2O8 2[END_REF] Chapter I Bibliography 1.1 Advanced oxidation processes (AOPs)

General description of AOPs

The wide occurrence of emerging pollutants (EPs) in the aquatic environment has raised much concern about water security. It is suspected that the increase of some incidence and diseases (such as heart diseases, diabetes, and cancer) is related to the presence of EPs in the environment [1]. Due to the chemical stability of EPs and the low efficiency of conventional water treatment technologies (such as physicochemical adsorption, coagulation precipitation, and membrane filtration), EPs can be persistent in various waters and bioaccumulate through the biological cycles in different organisms [2]. Therefore, it is necessary to develop an efficient water treatment technology to efficiently remove EPs in waters. In the past decades, advanced oxidation processes (AOPs) have attracted much attention due to their excellent efficiency in the removal of refractory pollutants.

AOPs are defined as the processes involving the generation of highly reactive species which can degrade organic substances [3]. AOPs can remove and mineralize most of the organic contaminants into less toxic products (e.g. CO2, H2O) [4]. The mentioned processes also exhibit excellent performance in the treatment of multiple organic pollutants in simultaneous time. The good oxidation reactivity of AOPs is attributed to the generation of highly reactive oxygen or sulfur species (ROS/RSS), for example, hydroxyl radical (HO • ) and sulfate radical (SO 4 •-).

Hydroxyl radical has a high standard redox potential (E0 = 2.80 V vs NHE) which can mineralize various organic pollutants without selectivity [5]. For sulfate radical, it exhibits high activity on the oxidation of various organic substances relying on its high standard reduction potential (E0 = 2.60 V vs NHE) [5]. The pathways of organic pollutants degradation by HO • and SO 4

•-include hydrogen abstraction, electron transfer, carbon double bonds addition, and radical interaction [6,7].

The way of AOPs activation can be classified into homogeneous and heterogeneous processes based on the uniformity of substances which participating in the reaction. For example, Fenton reaction (Fe 2+ /H2O2), Fenton-like reaction (Fe 3+ /H2O2), H2O2 photolysis (UV/H2O2), ultrasonic irradiation, and ozonation (O3) are all homogeneous processes [8]. In comparison, heterogeneous catalysis usually refers to the radical generation processes which use transition metal oxides (such as Fe3O4, Fe2O3, Co3O4, CuO, MnO2, and Mn2O3) as catalysts [4]. The efficiency of radical formation in homogeneous and heterogeneous processes can be improved by adding different radical precursors such as hydrogen peroxide (H2O2), peroxydisulfate (PDS, S 2 O 8 2-), and peroxymonosulfate (PMS, HSO 5 -). To better understand the activation mechanism of oxidants in AOPs, the properties of H2O2, PDS and PMS were summarized in the next section.

The commonly used oxidants in AOPs

In this section, the primary properties of H2O2, PDS, and PMS are presented.

The chemical and physical properties of H2O2

Hydrogen peroxide is a colorless liquid, less volatile, and more viscous than water [9]. The primary physiochemical properties of H2O2 are listed in Table 1-1. H2O2 has a threedimensional structure which is popularly known as an "open book" structure (Figure 1-1(a)) [10]. H2O2 is highly unstable and can be slowly decomposed in the presence of light. The decomposition of H2O2 into oxygen and water can be accelerated with the occurrence of metals [11]. Therefore, H2O2 is typically stored with a stabilizer in a weakly acidic solution in the plastic bottle. H2O2 exhibits different chemical properties in different pH environments, for example, H2O2 acts as an oxidizing agent (E0 = 1.77 V) in the acidic medium while in the basic environment, H2O2 acts as a reducing agent [12,13]. The application of H2O2 is diverse, for instance, oxidizer, bleaching agent, detergent, disinfectant, antiseptic, and propellant in the rocket [14]. In the study of pollutants removal by AOPs, H2O2 is usually used as the oxidant to promote the generation of radicals. However, it is important to realize that the wide occurrence of H2O2 in water may have impacts on the natural degradation processes of organic pollutants by oxidative minerals due to the reactions between H2O2 with natural metal oxides. 2 PMS refers to Oxone (KHSO5 • 0.5KHSO4 • 0.5K2SO4). 3 The values of water solubility are achieved at 20℃. 4 The density is achieved in 20℃ with 30% (w/w) H2O2 solution.

The chemical and physical properties of PDS and PMS

Peroxydisulfate (PDS) can be associated with three different salts i.e. sodium (Na2S2O8), potassium (K2S2O8), and ammonia ((NH4)2S2O8). The application of K2S2O8 and (NH4)2S2O8

in the in situ remediation has been restricted due to some disadvantages [15]. For example, K2S2O8 has very low solubility, and the application of (NH4)2S2O8 can cause residual ammonia and secondary contamination. In comparison, Na2S2O8 exhibits good performance for in situ chemical oxidation (ISCO) treatment. The main chemical and physical properties of PDS have been listed in Table 1-1. It is worth noting that PDS is composed of a symmetrical O-O bond (Figure 1-1(b)). The bond dissociation energy and bond length of O-O bond was respectively determined to be 92 kJ mol -1 and 1.497 Å [15].

Peroxymonosulfate (PMS) originates from Caro's acid (H2SO5) but is difficult to handle owing to its reactive and explosive characteristics. HSO 5 -has been identified as the dominant species under neutral and weakly acidic conditions due to its high dissociation constants (pKa = 9.3) [16]. To generate a more stable form, a triple salt was synthesized with a commercial name Oxone. The artificial Oxone contains three salts 2KHSO5, KHSO4, K2SO4 or KHSO5, 0.5KHSO4, 0.5K2SO4. The physical and chemical properties of PMS have been listed in Table 1-1. As shown in Figure 1-1(c), PMS possesses unsymmetrical reactive O-O bond whose length (1.46 Å) is shorter than PDS while the bond dissociation energy (377 kJ mol -1 ) is higher than PDS [15]. Theoretically, PMS requires more energy to break the O-O bond, however, in the actual catalysis process, PMS is easier than PDS to be activated. This phenomenon might be related to the unsymmetrical structure of PMS. The standard oxidation-reduction potential for the reduction of PDS anion (Eq. (1-1)) equals 2.01 V, hence it is higher than that of PMS i.e.

1.4 V (Eq. (1-2)) [15].

Transition metals-based AOPs

The application of transition metals-based AOPs (TMs-AOPs) in water or wastewater treatment has gained increasing attention due to their good performance on recalcitrant pollutants degradation. The excellent catalytic reactivity of transition metals is attributed to the presence of multiple oxidation states which is resulted from the special arrangement of electrons in the electric orbital of transition metal atoms [17]. The AOPs based on transition metals activation can be classified into two categories including homogeneous and heterogeneous TMs-AOPs according to the form of transition metals participated in the reaction [18]. The homogeneous TMs-AOPs employ transition metal ions such as Fe 2+ , Fe 3+ , Mn 2+ , and Co 2+ as the catalysts, while in heterogeneous TMs-AOPs, the transition metal oxides (e.g. Fe2O3, Fe3O4, MnO2, Co3O4, and CuO) are usually used as the radical activators. The radical precursors commonly used in TMs-AOPs include H2O2, PDS, and PMS. In this section, the introduction of TMs-AOPs is presented by summarizing the mechanism of homogeneous and heterogeneous TMs-AOPs.

The homogeneous TMs-AOPs

The activation of radical precursors such as H2O2, PDS, and PMS by transition metal ions has been widely reported [19,20]. The transition metals commonly used in AOPs activation include iron (Fe), manganese (Mn), copper (Cu), cerium (Ce), chromium (Cr), cobalt (Co), silver (Ag), nickel (Ni), vanadium (V) and ruthenium (Ru). The representative homogeneous TMs-AOPs is Fenton reaction which employs ferrous ion (Fe 2+ ) as the catalyst to activate H2O2 aiming to generate reactive hydroxyl radicals (Eq. (1-3)) [20]. Since the dual roles of H2O2, the formed Fe 3+ can be reduced to Fe 2+ by H2O2 realizing the cycle of iron ions (Eq. (1-4)) [20].

Fe 2+ + H2O2 → Fe 3+ + HO • + OH ─ (1-3)

Fe 3+ + H2O2 → Fe 2+ + HO 2 • + H + (1-4)
The mechanism of H2O2 activation by other transition metals such as Mn, Co, Cu, and V are similar to the Fenton reaction, thus one representative activation process has been proposed, as shown in Eqs. (1-5) - (1)(2)(3)(4)(5)(6). It is worth noting that although the activation processes are uniform, the valence states of transition metals involved in the H2O2 activation are different, as example, Fe 2+ /Fe 3+ , Ce 3+ /Ce 4+ and Co 2+ /Co 3+ . A list of valence states of transition metals that participated in the activation of oxidants (such as H2O2, PDS, and PMS) is presented in Table 1-2. The transition metals activated H2O2 processes have shown many advantages in the process of nonselective degradation of stubborn pollutants. However, the homogeneous TMs-AOPs still exists some disadvantages in the practical water treatment. For example, in the Fenton reaction, the narrow optimal pH range (2.5-3.5), the difficult recovery of iron catalysts, and the excess iron sludge generation restricts its application in the practical water treatment processes [21]. The application of the Cr 6+ /H2O2 system in wastewater treatment has been blocked by the risk of leaking harmful Cr 6+ [20]. To minimize the influence of transition metal ions and broaden the application scope of solution pH, different alternatives have been developed. For example, the introduction of various energy resources (such as ultraviolet light (UV), solar light, ultrasound, and electricity) and employing metal oxides (such as Fe3O4, Fe2O3, Mn2O3, MnO2) as the catalysts [4]. In addition, the employment of PMS or PDS is also a good substitute for improving the efficiency of pollutants degradation.

M n+ + H2O2 → M (n+1)+ + HO • + OH ─ (1-5) M (n+1)+ + H2O2 → M n+ + HO 2

• + H + (1-6)
Table 1-2. The valence states of transition metals involved in oxidants activation.

Transition metals Main oxidation state Redox potential vs NHE Reference

Fe +2, +3 E 0 (Fe 3+ /Fe 2+ ) = 0.771 V [22] Ce +3, +4 E 0 (Ce 4+ /Ce 3+ ) = 1.72 V [20] Ag 0, +1 E 0 (Ag + /Ag 0 ) = 0.799 V [23] Ru +2, +3 E 0 (Ru 3+ /Ru 2+ ) = 1.29 V [24] Mn +2, +3, +4 E 0 (Mn 3+ /Mn 2+ ) = 1.54 V E 0 (Mn 4+ /Mn 2+ ) = 1.24 V [25,26] Cr +3, +6 E 0 (Cr 6+ /Cr 3+ ) = 1.35 V [27] Cu +1, +2 E 0 (Cu 2+ /Cu + ) = 0.153V [28] Co +2, +3 E 0 (Co 3+ /Co 2+ ) = 1.92 V [29] V +3, +4 E 0 (V 4+ / V 3+ ) = 0.34 V [19] The activation of PDS and PMS by transition metals has provided one new way to degrade the organic pollutants in water. According to the previous literature, the activation mechanism of PDS and PMS by transition metals can be described by the equations (1-7)-(1-10) [7]. Generally, sulfate radical can be generated through the braking of the O-O bonds of persulfate anions activated by transition metals with lower valance while higher valance of them can promote the formation of SO 5

•-and S 2 O 8 •-radicals. The homogeneous activation of PDS and PMS by transition metals has been reported by previous studies [19]. For instance, Anipsitakis and Dyonysiou investigated the activation of PDS and PMS by nine transition metal ions including Ag(I), Ce(III), Co(II), Fe(II), Fe(III), Mn(II), Ni(II), Ru(III), and V(III) [19]. It was found that Co(II) and Ru(III) are the best transition metal catalysts for the activation of PMS. Ag(I) showed the best results toward activating PDS. In comparison, Mn(II) showed low efficiency in PMS activation for 2,4-dichlorophenol (2,4-DCP) degradation, and no 2,4-DCP removal was observed in the system of Mn(II)/PDS in the neutral pH. The results indicated that Mn(II) can activate PMS to generate reactive species while PDS is difficult to be catalyzed by Mn(II) at pH 7. This activation mechanism of PDS and PMS by Mn(II) has been verified in many related studies. For instance, Zhu et al. reported that the degradation of organic pollutants in crystallographic MnO2/PDS system was not related to the leaching of Mn 2+ since Mn 2+ is not reactive for catalytic PDS activation [30]. Zhou et al. demonstrated that the increased degradation efficiency of 4-nitrophenol in the α-MnO2/PMS system was attributed to the presence of Mn 2+ and Mn 3+ which can improve the generation of reactive radicals [31]. The different activation performance of Mn 2+ on PDS and PMS might be attributed to the different chemical structures of oxidants as discussed in section 1.1.2.

In summary, the degradation efficiency of recalcitrant pollutants in homogeneous TMs-AOPs depends on the employed oxidants and transition metals. Fe(II) and Fe(III) are the most efficient transition metals for the activation of H2O2. Co(II) and Ru(III) show good performance on PMS activation while Ag(I) is the best activator for PDS decomposition. The concerned Mn(II)

shows less reactivity in H2O2 and PDS activation but exhibited good ability for PMS activation.

M n+ + HSO 5 -→ M (n+1)+ + SO 4 •-+ OH ─ (1-7) M n+ + S 2 O 8 2-→ M (n+1)+ + SO 4 •-+ SO 4 2- (1-8) M (n+1)+ + HSO 5 -→ M n+ + SO 5 •-+ H + (1-9) M (n+1)+ + S 2 O 8 2-→ M n+ + S 2 O 8 •- (1-10)

The heterogeneous TMs-AOPs

The heterogeneous AOPs based on transition metal oxides have been widely applied in the treatment of recalcitrant pollutants due to their various advantages, for example, high efficiency, low cost, and environmental-friendly [32]. Moreover, the development of heterogeneous TMs-AOPs efficiently solves the drawbacks of the homogeneous TMs-AOPs, for example, the low recovery of metal ions, high cost, restrict pH condition, and secondary treatment for residual metal ions. The high catalytic performance of transition metal oxides on oxidants can be attributed to the presence of more active sites on the surface of heterogeneous catalysts which can activate radical precursors to produce more reactive species. The amounts of active sites on the surface of transition metal oxides were believed to be related to the specific surface area and the crystal structure of catalysts [33]. Various transition metal oxides can activate oxidants (e.g. H2O2, PDS, and PMS) to produce the reactive oxidation species such as SO 4

•-, HO • , and singlet oxygen ( 1 O2). The commonly reported transition metal oxides used in radical precursors activation include iron oxides (such as Fe3O4, 𝛼-Fe2O3, 𝛼-FeOOH and FeS2) [34,35], Copper oxide (such as CuO) [36,37], Cobalt oxide (such as Co3O4) [38], manganese oxides (such as MnO2, Mn2O3, Mn3O4, and 𝛾-MnOOH) [39,40].

Iron-based materials are widely employed for catalysis because they are cost-effective, ecofriendly, and efficient materials [32]. The activation of H2O2 by iron oxides (commonly named as Fenton-like reaction) has been widely used in the treatment of recalcitrant pollutants. For instance, Zhang et al. reported the efficient oxidation of phenolic and aniline compounds by nano-Fe3O4 activated Fenton-like reaction in the neutral solution [41]. Wang et al. demonstrated that α-FeOOH could efficiently degrade methyl orange through activating H2O2 [42]. Although iron is fixed in the catalyst and can effectively produce HO • by activating H2O2 without producing the iron hydroxide precipitation in heterogeneous Fenton-like reactions, there is still an iron leaching problem, especially when the catalyst is reused. Thus, some strategies were developed to improve the iron redox cycle in heterogeneous Fenton system catalyzed by iron materials, for example, introducing semiconductors, the modification with other elements, the application of carbon materials as carriers, the introduction of metal sulfides as co-catalysts, and the direct reduction with reducing substances [4].

As mentioned above, Co 2+ exhibited good performance on PMS activation, thus the application of cobalt oxide such as Co3O4 in PMS activation for organic pollutants degradation has been reported. For example, Chen et al. reported that Co3O4 nano-spheres were most effective in the degradation of Acid Orange 7 and were more stable at the neutral condition with dissolved Co concentration below 0.05 mg L -1 [38]. Saputra et al. also reported that Co3O4 nanoparticles could completely remove phenol in about 20 min in the presence of PMS [43]. However, the undesirable Co leaching during the catalysis of PMS with cobalt oxides restricts its application in practical water treatment [44]. To reduce the Co leaching problem, Co3O4 has been decorated on various supports including inert 1/2/3-D materials [45], metal oxides [46], magnetic particles [47], and molecular sieves [48]. In addition, the activation of PMS by other common transition metal oxides such as CuO, Fe 0 , Fe2O3, Fe3O4, MnO2, and Mn2O3 for organic pollutants degradation are also reported [34,[49][50][51][52].

In comparison with PMS, the redox recycling of the metal-based catalyst in the PDS system cannot be achieved easily due to the nature of PDS which cannot act as a reductant [19]. Among the transition metal catalysts, the zero-valent iron (Fe 0 ) is the most widely used heterogeneous catalyst for PDS activation due to its advantages such as low-cost, versatile, and eco-friendly [53]. The system of Fe 0 /PDS has been reported in the elimination of various organic pollutants such as pharmaceuticals [54], phenolic compounds [55], arsenic [56], pesticides [57], and industrial waste [58]. Furthermore, the magnetic Fe3O4 [59], Co3O4 [60], MnO2 [30], Zn 0 [61],

and Ni 0 [62] have also been investigated as PDS activator. It is worth noting that except for the generation of reactive radicals, the non-radical species such as the catalyst surface activated PDS molecules and singlet oxygen were also identified in the system of transition metal oxidesactivated PDS. For example, Zhang et al. reported the efficient degradation of 2,4dichlorophenol by CuO-activated PDS under mild conditions with the formation of outersphere intermediates between PDS and CuO [63].

In conclusion, the common oxidants (i.e. H2O2, PDS, and PMS) can be efficiently activated by most transition metal oxides to remove the organic pollutants in water. Moreover, both radical and non-radical mechanisms were involved in the activation of oxidants by transition metal oxides. The leaching of metal ions from heterogeneous catalysts has become one concerning problem in the application of heterogeneous TMs-AOPs on the treatment of wastewater. Thus, various substitute strategies were developed to control the release of metal ions and improve the catalytic efficiency of transition metal oxides.

Manganese oxides

Oxidation states

Manganese (Mn) is one of the most abundant elements on the earth. The average concentration of Mn in the Earth's crust is 11 g/kg [64]. Mn exists in different oxidation states from Mn(II)

to Mn(VII). Among them, Mn(II), Mn(III) and Mn(IV) are the dominant Mn valences that are discovered in the natural environment [65]. Generally, Mn(II) ions are soluble in water and can be oxidized to form various manganese oxides (MnOx) on exposure to oxygen and moisture [66]. Mn(IV) oxides as the oxidation products of Mn(II) are always presented in solution with insoluble forms [64]. For Mn(III) species, the existing situation is more complicated than Mn(II) and (IV) species. On one hand, Mn(III) species are not stable in solution and can combine with other parts of Mn(III) species to generate Mn(II) and Mn(IV) [67]. On the other hand, the complex generated between Mn(III) and organic substances (for example, humic compound, pyrophosphate chemicals) can be dissolved and stable in solution [START_REF] Madison | Abundant porewater Mn(III) is a major component of the sedimentary redox system[END_REF].

In nature, MnOx can be obtained through the oxidation of Mn(II) species. The ways of Mn(II) oxidation can be classified into abiotic and biogenic processes [START_REF] Learman | Formation of manganese oxides by bacterially generated superoxide[END_REF]. Firstly, the formation of MnOx based on the abiotic processes was considered as a long-term formation procedure due to its strict reaction conditions such as the high activation energy, the suitable solution pH value, and Mn concentrations [START_REF] Sunda | Photoreduction of manganese oxides in seawater and its geochemical and biological implications[END_REF]. It was reported the half-life of Mn(II) in oxic seawater is 350 days in the absence of catalysts for its oxidation [START_REF] Morgan | Kinetics of reaction between O2 and Mn(II) species in aqueous solutions[END_REF]. In addition, the oxidation of Mn(II) on the surface of minerals was also found to have a contribution to the formation of MnOx [START_REF] Gilkes | Geochemistry and Mineralogy of Manganese in Soils[END_REF].

Although the speed of mineral surface-triggered Mn(II) oxidation is more rapid than the homogeneous Mn(II) oxidation, the rate is still slower than the biogenic Mn(II) oxidation in five orders of magnitude [64]. The biogenic processes are induced by widespread fungi and bacteria with multifunctional enzymes, which are the main pathways to the formation of natural MnOx [START_REF] Miyata | Manganese(IV) Oxide Production by Acremonium sp. Strain KR21-2 and Extracellular Mn(II) Oxidase Activity[END_REF]. In biogenic processes, the microorganisms serve as the primary catalysts for the oxidation of Mn(II). The formed MnOx further assists the oxidation of Mn(II) on its surface.

Moreover, most of the low-crystalline MnOx formed by biogenic processes will be transferred to crystalline minerals during the aging process [START_REF] Manceau | Short-range and long-range order of phyllomanganate nanoparticles determined using high-energy X-ray scattering[END_REF].

Previous literature reported that more than 30 kinds of crystalline and non-crystalline MnOx have been found on the Earth's surface or crust [64]. Some representative MnOx minerals and their chemical formulas are listed in Table 1-3. MnOx is ubiquitous in the natural environments and occurs in a wide range of environmental settings, for example, fine-grained aggregates, veins, marine, and fresh-water nodules and concretions, crusts, dendrites, ocean floors, soils, and sediments [64]. Owing to the variable redox states and unique tunnel/layer structures, MnOx has been applied in many areas, such as lithium-ion batteries, magnetic materials, chemical catalysts, and energy storage [START_REF] Perner | A New MnOx Cathode Material for Rechargeable Lithium Batteries[END_REF][START_REF] Zener | Interaction Between the d Shells in the Transition Metals[END_REF][START_REF] Ko | Electrochemical properties of MnO2/activated carbon nanotube composite as an electrode material for supercapacitor[END_REF]. Moreover, to improve the oxidative and catalytic performance of MnOx, various nanostructures of MnOx have been synthesized through different artificial methods (i.e. hydrothermal, solvothermal, sol-gel, co-precipitation, electrochemical, and thermal annealing techniques) [START_REF] Zhu | Nanostructured manganese oxides: natural/artificial formation and their induced catalysis for wastewater remediation[END_REF]. To better understand the reactivity of MnOx, their structures are introduced in the following section. It was reported that recalcitrant pollutants in wastewater can be efficiently removed by different MnOx due to their good oxidative and catalytic performance. However, the reactivity of MnOx can be substantially affected by its structure. The basic building block for most of the MnOx is the MnO6 octahedral. The basic unit MnO6 can be assembled through sharing edges and/or corners to construct tunnel or layer structure MnOx [64].

For the formation of tunnel MnOx, the MnO6 octahedral units are firstly gathered via sharing edges, which leads to the generation of chains with different lengths (such as, single, double, and triple chains) [START_REF] Zhu | Nanostructured manganese oxides: natural/artificial formation and their induced catalysis for wastewater remediation[END_REF]. Then the various chains can be linked by sharing the corners resulting in the fabrication of tunnel MnOx with distinct shapes including square, rectangular, and crossing shapes (Figure 12)). The scale of tunnel MnOx depends on the length of formed chains. Therefore, the size of tunnel MnOx ranges from 1×1 to 3×3 chains. Moreover, different metal ions or water molecules can be hosted in tunnels leading to the formation of different tunnel MnOx. The classical tunnel-type MnOx found in nature include ramsdellite, pyrolusite, nsutite, romanechite, hollandite, and todorokite [64].

Layer-type MnOx is widespread in the environment. It was reported that layer-type MnOx exhibits excellent oxidative reactivity in organic pollutants removal [START_REF] Lin | Oxidative Removal of Bisphenol A by Manganese Dioxide: Efficacy, Products, and Pathways[END_REF]. To construct layer-type MnOx, the MnO6 octahedral units are first arranged to form sheets through sharing the edges, then the formed sheets are stacked together with the insertion of different metal ions or water molecules in the interlayer (Figure 1-2D) [64]. The performance of layer-type MnOx can be influenced by various factors, such as the number of vacancy sites, the concentration of Mn(III) in layer and interlayer, and the type of insert metal ions. The commonly found layertype MnOx include birnessite and birnessite-like minerals [65]. To better understand the structure of MnOx, some representative tunnel, and layer-type MnOx will be introduced in detail.

Pyrolusite. Pyrolusite ( 𝛽 -MnO2) is the most stable and abundant Mn(IV) mineral compared with other MnO2 such as ramsdellite and nsutite. In pyrolusite, a (1x1) tunnel structure was constructed by basic MnO6 units through the way of sharing the corners between the formed single chains (Figure 1-2A) [64]. The tunnel size of pyrolusite is too small to host other chemical species leading to the composition of pyrolusite has high similarity with pure MnO2. The existing forms of pyrolusite in nature are the low-temperature hydrothermal deposits or the replacements of other MnOx.

Ramsdellite.

The crystal phase of ramsdellite is 𝛾-MnO2. Ramsdellite shows the (1x2)

tunnel structure with a rectangular-shaped cross-section (Figure 1-2B) [64]. Minor amounts of water, Na, and Ca ions are detected in ramsdellite tunnels. Ramsidellite occurs as the lowtemperature hydrothermal deposits and commonly transfers to pyrolusite.

Hollandite.

Hollandite is a group of MnO2 minerals that has the same chemical formula (R0.8-1.5[Mn(IV), Mn(III)]8O16, R = Ba, Pb, K or Na). Hollandite has the crystal phase of 𝛼-MnO2. In hollandite, double chains of edge-sharing MnO6 octahedral are formed and construct a (2x2) tunnel structure with the square cross-section (Figure 1-2C) [64]. The tunnels of hollandite are filled with large uni-or divalent cations and some water molecules. The increase of hollandite charge with the introduction of cations can be balanced by the substitution of Mn(IV). Hollandite is mainly found in the oxidized zones of Mn deposits and ores.

Birnessite. Birnessite and birnessite-like minerals are ubiquitous in the environment and widely used in the elimination of organic pollutants [65]. In birnessite, the basic MnO6 octahedral are assembled to form sheets by sharing the edges. Then the sheets are stacked to construct the layer-type structure (Figure 1-2D) [64]. The interlayers are filled with various cations and water molecules. Various terms such as birnessite, acid birnessite, 𝛿-MnO2, and buserite are used interchangeably in referring to the layered Mn(IV) oxides that contain small contents of Mn(III) in layers and Na + as the interlayer cation [START_REF] Saratovsky | Manganese oxides: parallels between abiotic and biotic structures[END_REF]. Mn(III) oxide and hausmannite. The structure of Mn(III) oxide (α-Mn2O3) has been recognized as the body-centered cubic bixbyite phase (Figure 1-3B) [START_REF] Ghosh | Diversity in the Family of Manganese Oxides at the Nanoscale: From Fundamentals to Applications[END_REF], whereas the hausmannite (Mn3O4) exhibits a normal spinel structure with the formula Mn 2+ (Mn 3+ )2O4 where the Mn 2+ and Mn 3+ ions occupy the tetrahedral and octahedral sites, respectively (Figure 1-3C) [64]. In addition to crystal phases, the morphology and crystal facets also can affect the reactivity of MnOx. It was reported that MnOx can present in different dimensions including onedimensional (1D) tunnel structures, two-dimensional (2D) layered phases, and threedimensional (3D) spinels [START_REF] Zhu | Nanostructured manganese oxides: natural/artificial formation and their induced catalysis for wastewater remediation[END_REF]. The dimensional effects of MnOx not only impact the porous structure and specific surface areas (SSA), but also regulate the surface-atom arrangement, redox capacity of Mn species, and exposed facets [32]. The morphologies of α-, β-, δ-MnO2 are identified as nanowires, nanorods, and layers by high-resolution scanning electron microscopy (SEM) (as shown in Figure 1-4) [START_REF] Li | The Regular/Persistent Free Radicals and Associated Reaction Mechanism for the Degradation of 1,2,4-Trichlorobenzene over Different MnO2 Polymorphs[END_REF]. In addition, the SEM images of synthesized Mn2O3 with different facets are also demonstrated in previous literature. Figure 1-5 shows that α-Mn2O3octahedral (Mn2O3-O), α-Mn2O3-cubic (Mn2O3-C), and α-Mn2O3-truncated octahedral (Mn2O3-T) have the facets of ( 111), (001) and the mixture of ( 111) and (001), separately [START_REF] Yan | Tailoring Facets of α-Mn2O3 Microcrystalline Catalysts for Enhanced Selective Oxidation of Glycerol to Glycolic Acid[END_REF].

Trivalent manganese oxides activating sulfur radical-AOPs

Sulfur radical-based advanced oxidation processes (SR-AOPs) have attracted increasing attention due to their high efficiency in organic pollutants degradation. Mn(III) species have been proved to play an important role in the oxidative and catalytic reactivity of manganese oxides. The commonly used Mn(III) (oxyhydr)oxides in the activation of peroxydisulfate (PDS) and peroxymonosulfate (PMS) include α-Mn2O3, γ-MnOOH, and Mn3O4. In this section, the effect of structure on the reactivity of Mn(III) (oxyhydr)oxides is presented. Then, the mechanisms of PMS and PDS activation by different single or combined Mn(III) (oxyhydr)oxides are discussed in detail. Both radicals (sulfate and hydroxyl radicals) and nonradical (singlet oxygen and surface activated complex) mechanisms are presented. Finally, the effect of solution pH and inorganic ions on the catalytic performance of Mn(III)

(oxyhydr)oxides are discussed.

Effect of structure on the reactivity of Mn(III) (oxyhydr)oxides

The oxidative and catalytic performance of manganese oxides can be affected by various structural factors including crystal phases, morphologies, crystal facets, and structural dimensionalities [START_REF] Wu | Recent Progress of Thermocatalytic and Photo/Thermocatalytic Oxidation for VOCs Purification over Manganese-based Oxide Catalysts[END_REF]. For instance, Huang et al. reported that δ-MnO2 showed higher oxidative activity than α-, β-,γ-, λ-MnO2 on bisphenol A (BPA) oxidation due to the occurrence of more accessible active sites in layered δ-MnO2 than other tunnel structured MnO2 [START_REF] Huang | Effect of MnO2 Phase Structure on the Oxidative Reactivity toward Bisphenol A Degradation[END_REF]. The authors also demonstrated the effects of structured MnO2 on peroxymonosulfate (PMS) activation, and the low reactivity of δ-MnO2 was attributed to its less crystallinity [40].

The influence of structures in the reactivity of common Mn(III) (oxyhydr)oxides was summarized in Table 1- octahedral-and octahedral-structure and investigated the effect of crystal facets on the combustion of soot [START_REF] Cheng | Crystal facetdependent reactivity of α-Mn2O3 microcrystalline catalyst for soot combustion[END_REF]. The results show that the soot combustion efficiency followed the order of α-Mn2O3-cubic>α-Mn2O3-truncated octahedral>α-Mn2O3-octahedral. The enhanced reactivity of α-Mn2O3-cubic was explained by the fact that the exposed (001) surface facets of α-Mn2O3-cubic have higher amounts of low-coordinated surface oxygen sites which are capable of facilitating the oxygen activation and improving the surface redox properties.

In addition to α-Mn2O3, the oxidative and catalytic performances of Mn3O4 and γ-MnOOH were also reported to be affected by their structures. For example, Ji et al. reported that hexagonal nanoplate Mn3O4 exhibited superior catalytic performance on diesel soot combustion compared to the octahedral and nanoparticle Mn3O4, and the finding was explained by the improved amount of surface Mn 4+ species and surface reactive oxygen species due to the increased fraction of exposed [START_REF] Khan | Highly efficient α-Mn2O3@α-MnO2-500 nanocomposite for peroxymonosulfate activation: comprehensive investigation of manganese oxides[END_REF] facets in hexagonal nanoplate Mn3O4 [START_REF] Ji | Promoting diesel soot combustion efficiency by tailoring the shapes and crystal facets of nanoscale Mn3O4[END_REF]. The effect of morphology was also discovered by Liu et al. which demonstrated that the nanoflake Mn3O4 (exposure of (001) facet) has the highest oxygen reduction reactivity in comparison with nanoparticle Mn3O4 and nanorod Mn3O4 (exposure of (101) facet) [START_REF] Liu | Activating Mn3O4 by Morphology Tailoring for Oxygen Reduction Reaction[END_REF]. In addition, He et al.

investigated the activation of PMS by γ-MnOOH with different shapes, and the results showed that the catalytic activity of γ-MnOOH followed the order of nanowires > multi-branches > nanorods [START_REF] He | New insights into MnOOH/peroxymonosulfate system for catalytic oxidation of 2,4-dichlorophenol: Morphology dependence and mechanisms[END_REF]. Different physicochemical parameters such as specific surface area, Lewis sites, zeta-potential, and redox potential were measured to study the reason for the different catalytic performances of γ-MnOOH with distinct morphologies. It was found that the charge density on the surface played a crucial role in the interfacial reactivity between PMS and γ-MnOOH. In summary, the reactivity of Mn(III) (oxyhydr)oxides on radical precursors activation and pollutants oxidation can be deeply affected by their structures. The desirable morphologies and facets (such as cubic structure with (001) facet exposure) can improve the reactivity of Mn(III) (oxyhydr)oxides. [52]. The simultaneous formation of Mn(II) and Mn(IV) and the conversion of Mn ions with different oxidation states explained well the good performance of Mn(III) (oxyhydr)oxides on PMS activation. Except for the above-mentioned processes, the direct generation of HO • by Mn(III) activation of PMS was also reported by some researchers (Eq.

(1-15)) [START_REF] Wang | Insights into the transformations of Mn species for peroxymonosulfate activation by tuning the Mn3O4 shapes[END_REF][START_REF] Shokoohi | Degradation and mineralization of methylene blue dye by peroxymonosulfate/Mn3O4 nanoparticles using central composite design: Kinetic study[END_REF][START_REF] Wang | Facile synthesis of three-dimensional Mn3O4 hierarchical microstructures for efficient catalytic phenol oxidation with peroxymonosulfate[END_REF][START_REF] Zhao | Efficient removal of ciprofloxacin by peroxymonosulfate/ Mn3O4 -MnO2 catalytic oxidation system[END_REF][START_REF] Hu | Peroxymonosulfate activation by Mn3O4/metalorganic framework for degradation of refractory aqueous organic pollutant rhodamine B[END_REF][START_REF] Shokoohi | Comparing the performance of the peroxymonosulfate/Mn3O4 and threedimensional electrochemical processes for methylene blue removal from aqueous solutions: Kinetic studies[END_REF]. In comparison with SO 4

•-, SO 5

•-has been regarded as a low oxidative activity for organic pollutants removal due to its low reduction potential (E0 = 1.10 V vs NHE) [5]. Nevertheless, the transformation from SO 5

•-to SO 4

•-in Mn(III)Ox/PMS system still makes some contribution to the degradation of organic pollutants (Eq. (1-16)) [START_REF] Yermakov | The branching ratio of peroxomonosulfate radicals (SO5 -) self-reaction aqueous solution[END_REF].

In addition, the conversion from SO 4

•-to HO • in water should not be neglected (Eq. (1-17)), especially, when the solution is in the alkaline environment (Eq. (1-18)) [39].

Mn(III) + HSO 5 -→ Mn(IV) + SO 4

•-+ OH - (1-11)
Mn(III) + HSO 5 -→ Mn(II) + SO 5

•-+ H + (1-12)
Mn(II) + HSO 5 -→ Mn(III) + SO 4

•-+ OH - (1-13)
Mn(IV) + HSO 5 -→ Mn(III) + SO 5

•-+ H + (1-14) Mn(III) + HSO 5 -→ Mn(IV) + SO 4 2-+ HO • (1-15) SO 5 •-+ SO 5 •-→ O2 + 2 SO 4 •- (1-16) SO 4 •-+ H2O → SO 4 2-+ HO • + H + (1-17) SO 4 •-+ OH -→ SO 4 2-+ HO • (1-18)
In addition to the active radicals, the generation of non-radical species (such as 1 O2) in the Mn(III) (oxyhydr)oxides-activated PMS system was also reported.

HO 2 • ↔ H + + O 2 •- pKa = 4.88 (1-23) O 2 •-+ HO • → 1 O2 + OH - (1-24)
Currently, the Mn-based oxide composites have attracted increasing attention due to their various advantages such as more oxygen vacancies, higher surface oxygen mobility, and enforced synergistic effects. For instance, Chen et al. prepared the Fe2O3/Mn2O3 composite and studied its activity on PMS activation for tartrazine (TTZ) degradation. The results showed that 97.3% removal of TTZ was achieved in 30 min in the Fe2O3/Mn2O3/PMS system. The efficient degradation of TTZ originated from the generation of active species (e.g. SO 4

•-, HO • ) and the synergistic effect between iron and manganese ions [START_REF] Chen | Degradation of tartrazine by peroxymonosulfate through magnetic Fe2O3/Mn2O3 composites activation[END_REF]. γ-MnOOH coated nylon membrane was synthesized and applied in the activation of PMS towards the removal of 2,4dichlorophenol (2,4-DCP). The deep removal of 2,4-DCP was explained by the synergetic "trap-and-zap" process which improved the stability and catalytic reactivity of γ-MnOOH [START_REF] Zhang | A novel MnOOH coated nylon membrane for efficient removal of 2,4-dichlorophenol through peroxymonosulfate activation[END_REF].

In conclusion, the activation of PMS by Mn(III) (oxyhydr)oxides including pure Mn(III) oxides and Mn(III) containing composites is favorable. The degradation of various pollutants in the Mn(III) (oxyhydr)oxides/PMS system can be achieved through the generation of active radicals and non-radical species. [112]

HO • O 2 •- 1 O2 [95] α-Mn2O3-cubic Phenol [Catalyst] = 0.4 g/L [PMS] = 2 g/L [Phenol] = 25 ppm 100% removal of phenol in 1 h SO 4
HO • O 2 •- 1 O2 [93]

Activation of PDS by Mn(III) (oxyhydr)oxides

Single or combined Mn(III) (oxyhydr)oxides have been employed to activate PDS to remove different organic pollutants, such as phenol, p-chloroaniline (PCA), 2,4-dichlorophenol (2,4-DCP), and organic dyes (Table 123456). The activation pathway of PDS varies with the different types of Mn(III) (oxyhydr)oxides. For example, Shabanloo et al. reported the generation of active SO 4

•-radicals in the nano-Mn3O4/PDS system [START_REF] Shabanloo | Heterogeneous persulfate activation by nano-sized Mn3O4 to degrade furfural from wastewater[END_REF]. Since both Mn(II) and Mn(III) species are identified in the Mn3O4 structure, the formation of SO 4

•-was mainly attributed to the activation of PDS by Mn(II) (Eq. (1-25)). In contrast, the persulfate radical (S 2 O 8

•-) was produced by the reaction of PDS and Mn(III) (Eq. (1-26)). For the system of Mn2O3/PDS, it is believed that the singlet oxygen was the primary active species that was responsible for the degradation of organic pollutants [START_REF] Khan | High performance of the A-Mn2O3 nanocatalyst for persulfate activation: Degradation process of organic contaminants via singlet oxygen[END_REF]. As demonstrated by Khan et ) was believed to be responsible for the removal of phenol. Although the mentioned report explained well the oxidation performance of γ-MnOOH/PMS for phenol removal, the information regarding the mechanism of PDS activation on the surface of γ-MnOOH was not given in detail. Considering this, Xu et al.

conducted a further investigation focusing on the catalytic mechanism of PDS by γ-MnOOH [START_REF] Xu | Activation of persulfate by MnOOH: Degradation of organic compounds by nonradical mechanism[END_REF]. Based on the results of chemical scavenging and ESR experiments, a non-radical mechanism was proposed. Generally, the non-radical mechanism in PDS activation was attributed to three aspects including the generation of 1 O2, the electron transfer process, and the catalyst surface-activated intermediates [START_REF] Zhang | Efficient Peroxydisulfate Activation Process Not Relying on Sulfate Radical Generation for Water Pollutant Degradation[END_REF][START_REF] Wang | Nitrogen-Doped Reduced Graphene Oxide as a Bifunctional Material for Removing Bisphenols: Synergistic Effect between Adsorption and Catalysis[END_REF][START_REF] Lee | Activation of persulfates by carbon nanotubes: Oxidation of organic compounds by nonradical mechanism[END_REF][START_REF] Ahn | Activation of Peroxymonosulfate by Surface-Loaded Noble Metal Nanoparticles for Oxidative Degradation of Organic Compounds[END_REF][START_REF] Cheng | Non-photochemical production of singlet oxygen via activation of persulfate by carbon nanotubes[END_REF]. However, the 1 O2 production and electron transfer process mechanism were excluded according to the results of ESR and linear sweep voltammetry (LSV) experiments. Therefore, the γ-MnOOH surface-activated PDS molecule was verified as the main active species for the degradation of PCA. The corresponding diagram of PDS activation on the γ-MnOOH surface was shown in Figure 123456. The enhanced degradation of MB was attributed to the hampered electron-hole recombination due to the loading of Ag and graphene. Furthermore, the studies regarding the application of modified Mn2O3 in oxidants (such as PMS, H2O2) activation for contaminants removal were also reported [START_REF] Saputra | Egg-shaped core/shell α-Mn2O3@α-MnO2 as heterogeneous catalysts for decomposition of phenolics in aqueous solutions[END_REF][START_REF] Yang | Controllable redox-induced in-situ growth of MnO2 over Mn2O3 for toluene oxidation: Active heterostructure interfaces[END_REF][START_REF] Hazarika | Highly dispersed Mn2O3-Co3O4 nanostructures on carbon matrix as heterogeneous Fentonlike catalyst[END_REF][START_REF] Wang | Role of oxygen vacancies and Mn sites in hierarchical Mn2O3/LaMnO3-δ perovskite composites for aqueous organic pollutants decontamination[END_REF]. For example, Saputra et al. prepared an egg-shaped core/shell α-Mn2O3@α-MnO2 catalyst via a hydrothermal process and investigated the catalytic activity of α-Mn2O3@α-MnO2 in heterogeneous Oxone ® activation for phenol degradation [START_REF] Saputra | Egg-shaped core/shell α-Mn2O3@α-MnO2 as heterogeneous catalysts for decomposition of phenolics in aqueous solutions[END_REF]. The loaded α-MnO2 improved the generation of Mn(III) species through the reaction with PMS. The amount of SO 4

•-and HO • was then increased leading to the enhanced degradation of phenol.

The efficient degradation of organic dye pollutants (such as Rhodamine B (RB) and Congo Red (CR)) by bimetallic Mn2O3-Co3O4/carbon catalyst activated Fenton-like reaction was also reported [START_REF] Hazarika | Highly dispersed Mn2O3-Co3O4 nanostructures on carbon matrix as heterogeneous Fentonlike catalyst[END_REF]. The superior reactivity of Mn2O3-Co3O4/C catalyst in H2O2 activation for RB and CR degradation was attributed to the good synergistic effect between Co3O4 and Mn2O3 as well as the interaction between metal oxides and carbon. However, the investigation regarding the activation of PDS by modified α-Mn2O3 has been less reported. The same effect was also •- [START_REF] Shabanloo | Heterogeneous persulfate activation by nano-sized Mn3O4 to degrade furfural from wastewater[END_REF] Table 1- The Mn(III) (oxyhydr)oxides -mediated activation of PDS/PMS can be affected by solution pH in different ways. For example, influencing the property of charge on the surface of catalysts, changing the ionic forms of PDS/PMS and pollutants molecules, as well as altering the reduction potential of active radicals.

First, the solution pH can affect the interaction between catalyst and PDS/PMS and pollutants through changing the electrostatic effect. The point of zero charges (PZC) of catalyst and the acid dissociation constant (pKa) of radical precursors and contaminants are two important parameters that are used to recognize the charge type on the surface of catalysts and the ionic situation of oxidants and pollutants in solution. For instance, when the solution pH is equal to the PZC value of catalyst, the amounts of positive and negative charges on the surface of catalyst are equal (i.e. the surface charge of catalyst is zero). When the solution pH is higher than the PZC value of catalyst, the surface charges of catalyst are negative. On the contrary, if the solution pH is lower than catalyst PZC value, the surface of catalyst will be positively charged [START_REF] Dong | Oxidation of bisphenol A by persulfate via Fe3O4-α-MnO2 nanoflower-like catalyst: Mechanism and efficiency[END_REF]. The same situation is suitable for the analysis of the ionic form of oxidants and pollutants. The PZC values of commonly used Mn(III) (oxyhydr)oxides and the pKa values of PMS/PDS and some typical pollutants are summarized in Table 1-7. The impacts of solution pH on the interaction between Mn(III) (oxyhydr)oxides and PDS/PMS and pollutants have been reported. For example, Zhao et al. reported that the adsorption and degradation of ciprofloxacin (CIP) by synthesized Mn3O4-MnO2 composite were facilitated at neutral pH solution [START_REF] Zhao | Efficient removal of ciprofloxacin by peroxymonosulfate/ Mn3O4 -MnO2 catalytic oxidation system[END_REF]. The results were explained by the enhanced electrostatic attraction between Mn3O4-MnO2 and CIP.

The PZC value of Mn3O4-MnO2 composite was measured at 2.5, thus, in the solution pH 7, the surface of the catalyst was negatively charged. In comparison, the pKa of CIP was 8.7-10.58 leading to the formation of positively charged CIP ions in the neutral pH solution. Therefore, the electrostatic attraction between the negative catalyst and the positive CIP was improved resulting in a facilitating degradation of CIP. The same phenomenon was also reported in the studies of PDS activation by γ-MnOOH/α-Mn2O3 for pollutants degradation [START_REF] Khan | High performance of the A-Mn2O3 nanocatalyst for persulfate activation: Degradation process of organic contaminants via singlet oxygen[END_REF][START_REF] Xu | Activation of persulfate by MnOOH: Degradation of organic compounds by nonradical mechanism[END_REF].

Second, the transformation of radicals also influenced the reactivity of Mn(III) (oxyhydr)oxides for pollutants degradation. For instance, the reported conversion of SO 4

•-to HO • under the basic solution (as shown in (Eq. (1-18)) can have a significant impact. Since the reduction potential value of HO • under natural pH is lower than that in acidic solution (1.8 vs 2.7V) [START_REF] Buxton | Critical Review of rate constants for reactions of hydrated electrons, hydrogen atoms and hydroxyl radicals (⋅OH/⋅Oin Aqueous Solution[END_REF],

and the lifetime of HO • is shorter than SO 4

•-(20 ns vs 30-40 μs) [START_REF] Wang | Activation of persulfate (PS) and peroxymonosulfate (PMS) and application for the degradation of emerging contaminants[END_REF], thus the transformation from SO 4

•-(E = 2.6 V) to HO • under alkaline solution might lead to a decrease of pollutants degradation. In addition, the leaching of Mn 2+ from Mn(III) (oxyhydr)oxides in an acidic condition also should be taken into consideration for the activation of sulfate compounds (PMS/PDS).

The effect of inorganic anions

Inorganic anions are ubiquitous in various aquatic compartments. It is reported that inorganic anions can suppress the degradation of pollutants in Mn(III) (oxyhydr)oxides activated PMS/PDS systems through competing with pollutants for radicals' reactivity. Thus, to evaluate the applicability of the Mn(III) (oxyhydr)oxides + PMS/PDS system in different water matrices, the influence of inorganic anions on the removal of pollutants has been investigated by many researchers [START_REF] Zhang | A novel MnOOH coated nylon membrane for efficient removal of 2,4-dichlorophenol through peroxymonosulfate activation[END_REF][START_REF] Li | Efficient degradation of bentazone via peroxymonosulfate activation by 1D/2D γ-MnOOH-rGO under simulated sunlight: Performance and mechanism insight[END_REF][START_REF] Khan | High performance of the A-Mn2O3 nanocatalyst for persulfate activation: Degradation process of organic contaminants via singlet oxygen[END_REF][START_REF] Rizal | Photocatalytic decomposition of methylene blue by persulfate-assisted Ag/Mn3O4 and Ag/Mn3O4/graphene composites and the inhibition effect of inorganic ions[END_REF][START_REF] Huang | Catalytic degradation of ciprofloxacin by magnetic CuS/Fe2O3/Mn2O3 nanocomposite activated peroxymonosulfate: Influence factors, degradation pathways and reaction mechanism[END_REF]. In this section, the effect of inorganic anions such as carbonate/bicarbonate ions (CO 3 2-/HCO 3 -), chloride ions (Cl -), and nitrate (NO 3 -)/nitrite ions (NO 2 -) on the reactivity of Mn(III) (oxyhydr)oxides was summarized.

Carbonate (CO 3 2-)/bicarbonate (HCO 3 -) can react with SO 4

•-and HO • to generate less reactive carbonate radical (CO 3

•-) and bicarbonate radical (HCO 3 • ) (Eqs. (1-28)-(1-32)) leading to the inhibited degradation of pollutants [START_REF] Shah | Hydroxyl and sulfate radical mediated degradation of ciprofloxacin using nano zerovalent manganese catalyzed S2O8 2[END_REF]. However, although the redox potential of CO 3

•-is low (1.59V vs NHE), it can still selectively degrade some organic pollutants with a reaction rate of 10 3 -10 9 M -1 s -1 [START_REF] Huie | Electron transfer reaction rates and equilibria of the carbonate and sulfate radical anions[END_REF][START_REF] Chen | Reactivity of the carbonate radical toward aromatic compounds in aqueous solution[END_REF]. In addition, the presence of carbonate and bicarbonate ions can affect the stability of oxidants. For example, PDS can be activated by HCO 3 -to generated peroxymonocarbonate (HCO 4 -) (Eq. (1-33)) [START_REF] Jiang | Bicarbonate-activated persulfate oxidation of acetaminophen[END_REF]. Similarly, PMS can be catalyzed by both

SO 4 •-+ CO 3 2-→ SO 4 2-+ CO 3 •- 6.1 × 10 6 M -1 s -1 (1-28) SO 4 •-+ HCO 3 -→ SO 4 2-+ HCO 3 • 9.1 × 10 6 M -1 s -1 (1-29) HO • + CO 3 2-→ OH -+ CO 3 •- 3.9 × 10 8 M -1 s -1 (1-30) HO • + HCO 3 -→ H2O + HCO 3 • 8.5 × 10 6 M -1 s -1 (1-31) HCO 3 • = H + + CO 3 •- pKa = 9.6 (1-32) S 2 O 8 2-+ HCO 3 -+ 2 OH -→ HCO 4 -+ 2 SO 4 2-+ H2O (1-33) HSO 5 -+ CO 3 2-+ H + → SO 4 •-+ 2 OH -+ CO2 (1-34) HSO 5 -+ HCO 3 -→ SO 4 •-+ 2 OH -+ CO2 (1-35) HSO 5 -+ HCO 3 -→ SO 4 2-+ HCO 4 -+ H + (1-36)
Chloride ion (Cl -) exists widely in various water bodies including surface water, groundwater, and industrial wastewater [START_REF] Fang | Sulfate radical-based degradation of polychlorinated biphenyls: Effects of chloride ion and reaction kinetics[END_REF]. The influence of Cl -on the degradation of organic pollutants by sulfate radical-based AOPs (SR-AOP) was reported by previous studies [START_REF] Qi | Enhanced degradation of organic contaminants in water by peroxydisulfate coupled with bisulfite[END_REF][START_REF] Lei | Heterogeneous Degradation of Organic Pollutants by Persulfate Activated by CuO-Fe3O4: Mechanism, Stability, and Effects of pH and Bicarbonate Ions[END_REF][START_REF] Liang | Influences of carbonate and chloride ions on persulfate oxidation of trichloroethylene at 20 °C[END_REF][START_REF] Zhou | Transformations of chloro and nitro groups during the peroxymonosulfate-based oxidation of 4-chloro-2nitrophenol[END_REF].

Generally, Cl -can react with SO 4 •-by Cl -leads to the decrease of organic pollutants degradation [START_REF] Yang | Comparison of Halide Impacts on the Efficiency of Contaminant Degradation by Sulfate and Hydroxyl Radical-Based Advanced Oxidation Processes (AOPs)[END_REF][START_REF] Grebel | Effect of Halide Ions and Carbonates on Organic Contaminant Degradation by Hydroxyl Radical-Based Advanced Oxidation Processes in Saline Waters[END_REF]. However, Cl • was believed to own higher selectivity on electron-rich compounds than nonselective SO 4

•-, which can offset the negative effect of Cl -on SO 4

•- [START_REF] Yang | Effect of matrix components on UV/H2O2 and UV/S2O8 2advanced oxidation processes for trace organic degradation in reverse osmosis brines from municipal wastewater reuse facilities[END_REF]. Therefore, the conflicting effect of Cl -on organic pollutants in SR-AOP might be attributed to the different reactivity of pollutants with Cl • and Cl 2

•-. In addition, the reactivity of HO • can also be suppressed by Cl -due to the formation of low active radical ClOH •-(Eq.(1-39)) [START_REF] Minakata | Mechanistic Insight into the Reactivity of Chlorine-Derived Radicals in the Aqueous-Phase UV-Chlorine Advanced Oxidation Process: Quantum Mechanical Calculations[END_REF].

SO 4 •-+ Cl -→ SO 4 2-+ Cl • 3.1×10 8 M -1 s -1 (1-37) Cl • + Cl -→ Cl 2 •- 1.0×10 5 M -1 s -1 (1-38) HO • + Cl -→ ClOH •- 4.3×10 9 M -1 s -1 (1-39)
Nitrate (NO 3 -) and nitrite (NO 2 -) can be commonly found in various water matrices [START_REF] Fang | Sulfate radical-based degradation of polychlorinated biphenyls: Effects of chloride ion and reaction kinetics[END_REF]. Both NO 3 -and NO 2 -can react with SO 4

•-to generate low reactive NO 3

• (E0 = 2.3-2.5 V) and NO 2 • radicals (E0 = 1.03 V) (Eqs. (1-40)-(1- 41
)) [7]. The reaction rate constants of SO 4

•-with NO 3 -and NO 2 -are 5 × 10 4 M -1 s -1 and 8.8 × 10 8 M -1 s -1 , respectively [START_REF] Huang | Mn-based catalysts for sulfate radical-based advanced oxidation processes: A review[END_REF]. Thus, NO 2 -, compared with NO 3 -, has higher reactivity with SO 4

•-and increased its consumption. In addition and in a similar way, NO 2 -was also reported as the sink of HO • radicals (Eq. (1-42)) [START_REF] Wang | Effect of inorganic anions on the performance of advanced oxidation processes for degradation of organic contaminants[END_REF].

SO 4 •-+ NO 3 -→ SO 4 2-+ NO 3 • 5 × 10 4 M -1 s -1 (1-40) SO 4 •-+ NO 2 -→ SO 4 2-+ NO 2 • 8.8 ×10 8 M -1 s -1 (1-41) HO • + NO 2 -→ OH -+ NO 2 • (1-42)
Given the conducted summary, some further perspectives are proposed. First, the activation mechanism of PDS by α-Mn2O3 needs to be further investigated. Although the previous study already identified the generation of 1 O2 in α-Mn2O3/PDS system using the ESR and quenching experiments, the detailed catalytic process of PDS on the surface of Mn2O3 was not reported.

Second, the application of α-Mn2O3 or γ-MnOOH based composites in PDS activation can be investigated to understand the synergistic performance of α-Mn2O3 or γ-MnOOH with other loaded materials (such as active carbon, graphene, and transition metals) and improve the reactivity of α-Mn2O3/γ-MnOOH in PDS activation.

Manganese dioxide activated AOPs

The activation of radical precursors (e.g. H2O2, PDS, and PMS) by manganese dioxides (MnO2)

for recalcitrant pollutants degradation has been widely reported. MnO2 possesses different structures (such as 1D, 2D, 3D-MnO2), crystal phases (such as 𝛼-, 𝛽-, 𝛾-, 𝛿-MnO2), morphologies (such as rod, wire, tube, and flower) and crystal facets (such as 100, 110 and 310) [START_REF] Zhu | Nanostructured manganese oxides: natural/artificial formation and their induced catalysis for wastewater remediation[END_REF]. The reactivity of MnO2 is affected by a combination of several structural factors, including surface area, crystallinity, crystal facets, and the fraction of Mn 3+ /Mn 4+ [START_REF] Huang | Mn-based catalysts for sulfate radical-based advanced oxidation processes: A review[END_REF]. Therefore, in this section, the activation mechanism of different oxidants (i.e. H2O2, PDS, and PMS) by various MnO2 will be discussed. Moreover, the limitation and prospective of MnO2-activated AOPs will be presented.

Activation of H2O2 by MnO2

Manganese oxides with various surface and crystal morphologies exhibit high activities for H2O2 decomposition to degrade recalcitrant organic contaminants under mild conditions.

Moreover, due to the wide existence of manganese oxides in the natural environment, the understanding of H2O2 decomposition by manganese oxides toward organic pollutants degradation is of crucial significance.

The mechanism of H2O2 activation by MnO2 has been investigated by many researchers. Watts et al. studied the primary reactive species generated in the system of pyrolusite (𝛽-MnO2)/ H2O2 using carbon tetrachloride (CT) as the chemical probe [START_REF] Watts | Destruction of a Carbon Tetrachloride Dense Nonaqueous Phase Liquid by Modified Fenton's Reagent[END_REF]. The high degradation of CT in the system of 𝛽-MnO2/H2O2 suggested that a transient species other than hydroxyl radical was generated in the reaction, because CT was unreactive with hydroxyl radical. The transient species generated in 𝛽-MnO2/H2O2 system was further identified as the reductants including superoxide (O 2

•-) and hydroperoxide anion (HO 2 -) at a near neutral pH [START_REF] Watts | Oxidative and Reductive Pathways in Manganese-Catalyzed Fenton's Reactions[END_REF]. The same activation mechanism of H2O2 by 𝛽-MnO2 was also reported by Do et al., and the application of the 𝛽-MnO2/H2O2 system for the oxidation of toluene sorbed on activated carbon (AC) was also investigated [START_REF] Do | Hydrogen peroxide decomposition on manganese oxide (pyrolusite): Kinetics, intermediates, and mechanism[END_REF][START_REF] Do | The oxidation of toluene sorbed on activated carbon in the presence of H2O2 and manganese oxide[END_REF]. He et al. employed the 𝛽-MnO2/H2O2 system to efficiently clean the fouled ultrafiltration membranes through simultaneously generating reactive free radicals and oxygen [START_REF] He | Novel H2O2-MnO2 system for efficient physico-chemical cleaning of fouled ultrafiltration membranes by simultaneous generation of reactive free radicals and oxygen[END_REF]. In addition to 𝛽-MnO2, 𝛾-MnO2 also exhibited the reactivity on H2O2 As discussed before, the reactivity of manganese oxides on oxidants activation was influenced by their crystal structures. Kim et al. investigated the degradation of methylene blue (MB) dye by various manganese oxide nanorods (𝛼-, 𝛽-, 𝛾-, 𝛿-MnO2) activated H2O2 [START_REF] Kim | Manganese oxide nanorods as a robust Fenton-like catalyst at neutral pH: Crystal phase-dependent behavior[END_REF]. The results showed that the catalytic activity of MnO2 nanorods followed the order of 𝛾-MnO2 > 𝛽-MnO2 > 𝛼-MnO2 > 𝛿-MnO2. The high reactivity of 𝛾-MnO2 was assigned to its largest surface area and pore volume which facilitating more efficient mass transport and reactants diffusion. However, the catalytic reactivity of MnO2 is not simply function of the surface area. The crystallinity of catalyst was considered as the key factor of affecting the reactivity of MnO2. For example, the crystalline 𝛽-MnO2 with lowest surface area (9.2 ± 0.8 m 2 g -1 ) showed the greater activity than amorphous 𝛿-MnO2 with the surface area of 47.8 ± 2.9 m 2 g -1 . Therefore, the reactivity of MnO2 in the activation of H2O2 is highly dependent on their crystal structures. In addition, Sabri et al.

investigated the effect of structure and disorder on the oxidative and catalytic reactivity of manganese (III/IV) oxides against three substrates including methylene blue, hydrogen peroxide and water [START_REF] Sabri | Oxidant or Catalyst for Oxidation? A Study of How Structure and Disorder Change the Selectivity for Direct versus Catalytic Oxidation Mediated by Manganese(III, IV) Oxides[END_REF]. The results showed that disorder destabilizes the materials thermodynamically, making them stronger chemical oxidants but not necessarily better catalysts. The schematic diagram of the reaction of H2O2 and MnO2 was presented in Figure 1 Natural H2O2 can be formed in oceanic and atmospheric systems as well as the subsurface environment with the assistance of light, natural organic matter, and biological-mediated processes [START_REF] Begg | Plutonium desorption from mineral surfaces at environmental concentrations of hydrogen peroxide[END_REF][START_REF] Cooper | Hydrogen peroxide concentration in a northern lake: photochemical formation and diel variability[END_REF][START_REF] Garg | Photochemical production of superoxide and hydrogen peroxide from natural organic matter[END_REF]. The concentration of H2O2 in aquatic environments ranges from nM to µM [START_REF] Garg | Photochemical production of superoxide and hydrogen peroxide from natural organic matter[END_REF][START_REF] Yuan | Production of Hydrogen Peroxide in Groundwater at Rifle[END_REF][START_REF] Roe | Measurement of dark, particlegenerated superoxide and hydrogen peroxide production and decay in the subtropical and temperate North Pacific Ocean, Deep Sea Research Part I[END_REF][170]. Natural manganese oxides such as birnessite have highly catalytic and oxidative reactivity in the degradation of organic pollutants. Despite the catalyzed decomposition of hydrogen peroxide by MnO2 has been well investigated [START_REF] Watts | Oxidative and Reductive Pathways in Manganese-Catalyzed 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 H2O2 -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 H2O2 on the reactivity of birnessite at environmental levels of H2O2.

Thus, the influence of H2O2 on the reactivity of birnessite under environmentally relevant conditions was investigated in this thesis using bisphenol A as the model pollutant.

Activation of PMS and PDS by MnO2

In the past two decades, the sulfate radical-based advanced oxidation processes (SR-AOPs)

have drawn increasing attention, and manganese oxides have been proven to be effective catalysts for activating PMS and PDS to degrade various recalcitrant organic pollutants. This •-, HO • , and 1 O2. On the contrary, the formed reactive complexes between amorphous MnO2 and PMS was believed to be responsible for BPA degradation [START_REF] Wang | Oxidation of bisphenol A by nonradical activation of peroxymonosulfate in the presence of amorphous manganese dioxide[END_REF].

section
2 MnO2 + HSO 5 -→ Mn2O3 + SO 5 •-+ OH - (1-43) Mn2O3 + HSO 5 -→ 2 MnO2 + SO 4 •-+ H + (1-44)
The catalytic and oxidative reactivity of MnO2 can be influenced by several factors, such as the crystalline structure, morphology, and dimension. For example, Saputra et al. reported that 𝛼-MnO2 exhibited higher activity than β-and 𝛾-MnO2 in PMS activation for phenol degradation [START_REF] Saputra | Different Crystallographic One-dimensional MnO2 Nanomaterials and Their Superior Performance in Catalytic Phenol Degradation[END_REF]. The differences in polymorphs were attributed to the distinct tunnel sizes, oxygen liability, and MnO6 edges. A similar conclusion was observed by Huang et al. that showed the reactivity of MnO2 in PMS activation followed a decreased order: 𝛼-MnO2 > 𝛾- [40]. To reveal the different reactivity of MnO2 in PMS activation, some physiochemical properties of MnO2 were investigated. The results showed that the catalytic reactivity of MnO2 was linearly correlated with the Mn(III) content, Mn average oxidation state (AOS) and conductivity. MnO2 morphologies also affected the exposed crystal facets and specific surface areas (SSA), which further impacted the PMS activation activity. degradation [START_REF] Deng | Degradation of ciprofloxacin using α-MnO2 activated peroxymonosulfate process: Effect of water constituents, degradation intermediates and toxicity evaluation[END_REF]. The result showed that the catalytic activity of 𝛼-MnO2 followed an order of α-MnO2 nanoflowers > α-MnO2 nanorods > α-MnO2 nanoparticles > commercial MnO2. The variances were explained by the difference in surface area and crystallinity. Based on the above studies, one short conclusion can be provided that the catalytic activity of MnO2 in PMS activation relies on various physicochemical properties including the morphology of nanostructures, crystal structure, exposed crystal facets, surface area, Mn(III) content, and oxygen species.

MnO2 > β-MnO2 > 𝛿-MnO2
The activation of PDS by MnO2 for organic pollutants degradation was reported by a few papers. β-MnO2-activated PDS, and the primary reactive oxygen species was identified as singlet oxygen [30]. The generation of 1 O2 was originated from the direct oxidation or recombination of superoxide ions (Eqs. (1-47)-(1-50)). However, in comparison to PMS, the investigation of PDS activation by MnO2 was less reported. To better understand the catalytic activity of various MnO2 in PDS activation, more experiments should be conducted.

≡Mn(IV) + S 2 O 8 2-→ ≡Mn(III) + S 2 O 8 •- (1-45) ≡Mn(III) + S 2 O 8 2-→ ≡Mn(IV) + SO 4 •-+ SO 4 2- (1-46) 2 [≡Mn(IV)-OH] III + HS 2 O 8 -→ 2 [≡Mn(IV)-O-O-SO3] II + 3 H + (1-47) 2 [≡Mn(IV)-O-O-SO3] II + 4H2O +S 2 O 8 2-→2[≡Mn(III)-OH] II + SO 4 2-+ 2O 2 •-+ 8 H + (1-48) [≡Mn(IV)-O-O-SO3] II + O 2 •-+ OH -→ [≡Mn(III)-OH] II + SO 4 2-+ 1 O2 (1-49) 2 O 2 •-+ 2 H2O → 1 O2 + H2O2 + 2 OH - (1-50)
In summary, the common oxidants (i.e. H2O2, PDS, and PMS) can be activated by various MnO2 to generate reactive species for the removal of organic pollutants. The oxidative and catalytic activity of MnO2 can be influenced by various factors, such as the crystalline structure, morphology, and exposed crystal facets. With the increase of H2O2 concentration in the natural environment, the influence of H2O2 on the reactivity of MnO2 for organic pollutant removal should be considered. Moreover, more investigation regarding the activation of PDS by different structural MnO2 should be conducted to improve the understanding of the PDS activation mechanism.

Emerging pollutants

Emerging pollutants (EPs) have been defined as a group of natural or synthetic chemicals occurring in water bodies throughout the globe, which are potential to cause environmental problems and pose a threat to the safety of the ecosystem and the health of human beings [START_REF] Riva | Risk assessment of a mixture of emerging contaminants in surface water in a highly urbanized area in Italy[END_REF].

The EPs reported in previous literature include different types of organic pollutants, for example, disinfection by-products, pesticides, pharmaceuticals and personal care products (PPCPs), endocrine-disrupting compounds (EDCs), industrial chemicals, artificial sweeteners and food additives, sunscreens, and flame retardants [START_REF] Geissen | Emerging pollutants in the environment: A challenge for water resource management[END_REF]. With the progress of urbanization, industrialization, and number of people, a large number of EPs-containing products were consumed and discharged. The concentration of EPs detected in waters ranges from ng/L to µg/L [START_REF] Li | A review on removing pharmaceutical contaminants from wastewater by constructed wetlands: Design, performance and mechanism[END_REF]. However, the conventional treatment technologies in wastewater treatment plants (WWTPs) cannot completely degrade these recalcitrant pollutants leading to the occurrence and increase of EPs in waters [START_REF] Sauvé | A review of what is an emerging contaminant[END_REF]. Among the EPs, the PPCPs and EDCs have attracted increasing attention due to their close connection with the daily life of humans.

PPCPs represent a diverse array of chemical materials which are used for therapeutic purposes in humans and animals and used in personal hygiene and for make-up [2]. Common PPCPs include caffeine (CAF), carbamazepine, diclofenac, ibuprofen, iopamidol, musk, clofibric acid, triclosan, and phthalates [2]. PPCPs can be released from various anthropogenic activities such as industry, agriculture, farm, and domestic life [START_REF] Korekar | Occurrence, fate, persistence and remediation of caffeine: a review[END_REF]. The ubiquitous presence and ecological risk of PPCPs have attracted growing attention from environmental research. The widespread of PPCPs in the environment can be attributed to the overuse of PPCPs, the low removal efficiency of WWTPs, and the stable structures of PPCPs [START_REF] Deo | Pharmaceuticals in the Surface Water of the USA: A Review[END_REF]. Although the amount of PPCPs detected in water is in trace range, the detrimental effect of PPCPs has been reported through the toxic tests in fishes and plants [START_REF] Melvin | Diurnal activity patterns as a sensitive behavioural outcome in fish: effect of short-term exposure to treated sewage and a sub-lethal PPCP mixture[END_REF][START_REF] Sun | Pharmaceutical and personal care products-induced stress symptoms and detoxification mechanisms in cucumber plants[END_REF].

Endocrine-disrupting compounds (EDCs) are defined by the Environmental Protection Agency (EPA, U.S.) as a series of chemical substances which can interfere with the normal hormone function in bodies of aquatic organisms and human beings [START_REF] Kavlock | Research needs for the risk assessment of health and environmental effects of endocrine disruptors: a report of the U.S. EPA-sponsored workshop[END_REF]. Trace concentrations of EDCs have been detected in various waters including wastewater, sediments, drinking water, groundwater, and surface water [START_REF] Fernandez | An assessment of estrogenic organic contaminants in Canadian wastewaters[END_REF][START_REF] Liu | Removal mechanisms for endocrine disrupting compounds (EDCs) in wastewater treatment -physical means, biodegradation, and chemical advanced oxidation: A review[END_REF]. The persistence of EDCs in water poses a high risk to the security of the environment, especially to aquatic organisms and humans [START_REF] Campbell | Biologically directed environmental monitoring, fate, and transport of estrogenic endocrine disrupting compounds in water: A review[END_REF]. It was reported that the increase of some incidence and diseases (such as heart diseases, diabetes, metabolic disorders, and cancer) was related to the presence of EDCs in the environment [START_REF] Korekar | Occurrence, fate, persistence and remediation of caffeine: a review[END_REF].

The EDCs can be divided into natural substances, synthetic substances, and heavy metals [START_REF] Liu | Removal mechanisms for endocrine disrupting compounds (EDCs) in wastewater treatment -physical means, biodegradation, and chemical advanced oxidation: A review[END_REF].

The sources of EDCs are attributed to various ways including (1) food contact materials and diet; (2) cosmetics and personal care products; (3) consumer goods and cleaning products; (4) pesticides; (5) drugs [START_REF] Kabir | A review on endocrine disruptors and their possible impacts on human health[END_REF].

Source and harm of BPA, E2 and EE2

BPA, E2, and EE2 as the representative EDCs were largely adopted as the model contaminants for the investigation of catalysts' reactivity. The physicochemical properties of these three EDCs are listed in Table 1-8. BPA has been wildly used as an industrial additive and raw material to fabricate polymer materials such as polycarbonates, epoxy resins, and polysulfone resins [START_REF] Yang | Assessment of bisphenol A alternatives in paper products from the Chinese market and their dermal exposure in the general population[END_REF].

All of these polymer materials are further used to produce industrial products including beverage containers, drink bottles, baby bottles, food cans, food containers, dental sealants, thermal paper, flame retardants, building materials, electronic products, and medical equipments [START_REF] Staples | A review of the environmental fate, effects, and exposures of bisphenol A[END_REF]. The produced BPA-containing products greatly improve the life quality of humans. However, with the rise of anthropogenic activities, the consumption and discharging of BPA products was increased leading to the widespread distribution of BPA in the natural environment. The concentrations of BPA detected in surface waters, groundwaters, and drinking waters were listed in Table 1-9. It was reported that BPA may cause different diseases, such as human breast cancer, hormonal disorder, metabolic imbalance, diabetes, and abnormal reproductive and developmental behavior [START_REF] Rochester | Bisphenol A and human health: A review of the literature[END_REF]. The exposure route of humans to EDC is summarized in Figure 123456789.

E2 and EE2 are natural and synthetic estrogens separately with strong estrogenic activity even at very low concentrations [START_REF] Tang | Occurrence and removal of 17αethynylestradiol (EE2) in municipal wastewater treatment plants: Current status and challenges[END_REF]. Natural E2 plays an important role in the development of female secondary sex characteristics and reproduction. Synthetic EE2 and E2 are widely used to produce oral contraceptives [START_REF] Aris | Occurrence of 17α-ethynylestradiol (EE2) in the environment and effect on exposed biota: a review[END_REF]. The occurrence of E2 and EE2 has been identified in various waters including wastewater, surface water, underground water, and drinking water (Table 123456789) [START_REF] Lei | Estrogens in municipal wastewater and receiving waters in the Beijing-Tianjin-Hebei region, China: Occurrence and risk assessment of mixtures[END_REF]. The presence of E2 and EE2 in water poses a high risk to the health of humans due to their strongly adverse effects. It was reported that the sexuality of fish can be changed on the exposure of E2, even at a very low concentration (0.5 ng/L) [START_REF] Hu | Products of Aqueous Chlorination of 17β-Estradiol and Their Estrogenic Activities[END_REF]. EE2 was also reported to have detrimental effects on fishes, for example, hormonal imbalance, intersex, increased egg mortality, and infertility [START_REF] Chen | High estrogen concentrations in receiving river discharge from a concentrated livestock feedlot[END_REF]. To avoid the potential harm of EE2 to humans, EE2 has been listed in the latest Japanese Drinking Water Quality Standard [START_REF] Yuan | Fast trace determination of nine odorant and estrogenic chloro-and bromo-phenolic compounds in real water samples through automated solid-phase extraction coupled with liquid chromatography tandem mass spectrometry[END_REF]. Different exposure routes of humans to EDCs [START_REF] Kabir | A review on endocrine disruptors and their possible impacts on human health[END_REF].

Source and harm of CAF

Caffeine (CAF) is a substance that belongs to the methylxanthines class of drugs [START_REF] Rigueto | Alternative techniques for caffeine removal from wastewater: An overview of opportunities and challenges[END_REF]. The physicochemical properties of CAF are listed in Table 1-10. CAF can be found in various beverages (e.g. coffee, cocoa, tea, chocolate, soft drinks, and energy drinks) and prescription or over-the-counter drugs (e.g. cold medicines, analgesics, stimulants, and illegal drugs) [START_REF] Rigueto | Alternative techniques for caffeine removal from wastewater: An overview of opportunities and challenges[END_REF]. It was evaluated that the average caffeine consumption is around 70 mg per person in the world [START_REF] Rigueto | Alternative techniques for caffeine removal from wastewater: An overview of opportunities and challenges[END_REF]. Over 90% of adults are regular consumers of CAF, with a mean daily intake of 227 mg [START_REF] Nehlig | Interindividual Differences in Caffeine Metabolism and Factors Driving Caffeine Consumption[END_REF]. However, it is worth noting that approximately 5% of ingested caffeine cannot be digested and absorbed by human bodies and eventually excreted in urine [START_REF] Rigueto | Alternative techniques for caffeine removal from wastewater: An overview of opportunities and challenges[END_REF]. Given the above, with the increase in the consumption of caffeinated products, a large of CAF has been discharged into the waters. The concentration of CAF detected in different water matrices ranges from ng/L to µg/L (Table1-11). The adverse effects of CAF on aquatic organisms and human beings have been widely reported, such as depression, cardiovascular disorders, and hyperactivity [START_REF] Buerge | Caffeine, an Anthropogenic Marker for Wastewater Contamination of Surface Waters[END_REF]. An appropriate CAF intake for adults recommended by the European Food Safety Authority (EFSA) is 400 mg per day [START_REF]Outcome of a public consultation on the draft Scientific Opinion of the EFSA Panel on Dietetic Products, Nutrition and Allergies (NDA) on the essential composition of infant and follow-on formulae[END_REF].

Waters

CAF concentration Reference

Rivers and lakes 130-370 ng L -1 (USA) [START_REF] Siegener | Caffeine in Boston Harbor seawater[END_REF] median 530 ng L -1 , up to 880 ng L -1 (Germany) [START_REF] Ternes | Determination of neutral pharmaceuticals in wastewater and rivers by liquid chromatography-electrospray tandem mass spectrometry[END_REF] up to 171 ng L -1 (Netherlands) [START_REF] Hendriks | Monitoring response of XAD-concentrated water in the rhine delta: A major part of the toxic compounds remains unidentified[END_REF] 160 ng L -1 (Greece) [START_REF] Patsias | Development of an automated on-line solidphase extraction-high-performance liquid chromatographic method for the analysis of aniline, phenol, caffeine and various selected substituted aniline and phenol compounds in aqueous matrices[END_REF] Groundwater 10-80 ng L -1 , median 30 ng L -1 (USA) [START_REF] Burkhardt | Determination of Submicrogram-per-Liter Concentrations of Caffeine in Surface Water and Groundwater Samples by Solid-Phase Extraction and Liquid Chromatography[END_REF] <40 ng L -1 (wells deeper than 10 m, USA) [START_REF] Seiler | Caffeine and Pharmaceuticals as Indicators of Waste Water Contamination in Wells[END_REF] Seawater 2-16 ng L -1 (North Sea) [START_REF] Weigel | Drugs and personal care products as ubiquitous pollutants: occurrence and distribution of clofibric acid, caffeine and DEET in the North Sea[END_REF] 5-71 ng L -1 (Massachusetts Bay) [START_REF] Siegener | Caffeine in Boston Harbor seawater[END_REF] 140-1600 ng L -1 (Boston Harbor) [START_REF] Siegener | Caffeine in Boston Harbor seawater[END_REF] WWTPs, influent water 20 μg L -1 (USA) [START_REF] Siegener | Caffeine in Boston Harbor seawater[END_REF] 147 μg L -1 (Germany) [START_REF] Ternes | Determination of neutral pharmaceuticals in wastewater and rivers by liquid chromatography-electrospray tandem mass spectrometry[END_REF] WWTPs, effluent water 6.7 μg L -1 (USA) [START_REF] Siegener | Caffeine in Boston Harbor seawater[END_REF] 0.19 ± 0.09 μg L -1 , up to 1.9 μg L -1 (Germany) [START_REF] Ternes | Determination of neutral pharmaceuticals in wastewater and rivers by liquid chromatography-electrospray tandem mass spectrometry[END_REF] 2 μg L -1 (Sweden) [START_REF] Paxéus | Organic pollutants in the effluents of large wastewater treatment plants in Sweden[END_REF] 1.

Treatment technologies for EDCs and CAF

The occurrence of EDCs and CAF in water has attracted considerable research attention owing to the potential safety risks they posed to the ecosystem and human beings. MnO2/PMS [40] CuFe2O4/PMS [START_REF] Dong | Monodispersed CuFe2O4 nanoparticles anchored on natural kaolinite as highly efficient peroxymonosulfate catalyst for bisphenol A degradation[END_REF] Electro-Fenton [START_REF] Burgos-Castillo | Application of electrochemical advanced oxidation to bisphenol A degradation in water. Effect of sulfate and chloride ions[END_REF] Fe(III)-EDDS*/H2O2 [START_REF] Huang | Assessment of the Fe(III)-EDDS Complex in Fenton-Like Processes: From the Radical Formation to the Degradation of Bisphenol A[END_REF] E2/EE2 Fenton [START_REF] Li | Removal of steroid estrogens from waste activated sludge using Fenton oxidation: Influencing factors and degradation intermediates[END_REF] Ozonation [START_REF] Esplugas | Ozonation and advanced oxidation technologies to remove endocrine disrupting chemicals (EDCs) and pharmaceuticals and personal care products (PPCPs) in water effluents[END_REF] UV/H2O2 [START_REF] Ma | Simultaneous Degradation of Estrone, 17β-Estradiol and 17α-Ethinyl Estradiol in an Aqueous UV/H2O2 System[END_REF] UV/PDS [START_REF] Gabet | Hydrogen peroxide and persulfate activation using UVA-UVB radiation: Degradation of estrogenic compounds and application in sewage treatment plant waters[END_REF] Photo-Fenton [START_REF] Feng | Degradation of estrone in aqueous solution by photo-Fenton system[END_REF] Bio-electro-Fenton [START_REF] Xu | Bio-electro-Fenton system for enhanced estrogens degradation[END_REF] CAF O3/UV [START_REF] Souza | Degradation of Caffeine by Advanced Oxidative Processes: O3 and O3/UV[END_REF] UV/H2O2 [236] UVC/PDS [START_REF] Starling | Intensification of UV-C treatment to remove emerging contaminants by UV-C/H2O2 and UV-C/S2O8 2-: Susceptibility to photolysis and investigation of acute toxicity[END_REF] Solar/TiO2 [START_REF] Ghosh | Solar photocatalytic degradation of caffeine with titanium dioxide and zinc oxide nanoparticles[END_REF] Photo-Fenton [START_REF] Miralles-Cuevas | Comparison of UV/H2O2, UV/S2O8 2-, solar/Fe(II)/H2O2 and solar/Fe(II)/S2O8 2at pilot plant scale for the elimination of micro-contaminants in natural water: An economic assessment[END_REF] Solar/Fe/H2O2/EDDS [START_REF] Miralles-Cuevas | EDDS as complexing agent for enhancing solar advanced oxidation processes in natural water: Effect of iron species and different oxidants[END_REF] *EDDS: Ethylene diamine-N,N'-disuccinic acid

Chapter II

Materials and Methods

Chemicals and instruments

Chemicals

Preparation of materials and solutions

Synthesis of acid birnessite

Acid birnessite (AB) was prepared following the procedures of McKenzie [1]. 166 mL of concentrated HCl was added dropwise to 2.5 L of 0.4 M KMnO4. The suspension was stirred vigorously and kept at 90°C during the HCl addition. The reaction was maintained at 90 °C for a further 10 min. Then, the obtained slurry was stirred vigorously at room temperature for 15 h.

The precipitate was collected by centrifugation, washed with ultrapure water repeatedly until the conductivity was close to 0.6 μS cm -1 . The suspensions were stored in polypropylene containers at 4 °C for further use.

Synthesis of Mn(III) rich birnessite

Mn(III)-rich birnessite was synthesized according to previous published methods [2,3]. Briefly, about 2 g of acid birnessite was incubated with 80 mL of 10 mM sodium 4-(2-hydroxyethyl) piperazine-1-ethanesulfonate (HEPES) and 10 mM NaCl solutions at pH 6.5 ± 0.05 at a temperature of 25 ℃ for 48 h. The precipitate was collected by centrifugation and washed with ultrapure water.

Synthesis of manganite

Manganite (γ-MnOOH) particles were synthesized as described by Yu and co-workers [4].

Manganite was synthesized by adding 300 mL of 0.2 M NH3 to a solution of 20.4 mL 30%

H2O2 and 1000 mL 0.06 M MnSO4. The mixture was heated and kept at 95 °C for 6 h under constant stirring. Thereafter, the hot suspension was filtered through a G4 glass filter, washed with equal amounts (with respect to the initial solution) of hot deionized water, collected on watch glass and dried at room temperature under rough vacuum (10 -3 Torr) in presence of P2O5 for 48 h. The manganite crystals dispersed well in suspensions, although they formed aggregates when dried.

Sewage treatment plant water collection

The sewage treatment plant water (STPW) was collected from the urban treatment plant named "3 rivières" in Clermont-Ferrand, France in December 2019. Prior to use, the STP water was filtered using a syringe filter (CHROMAFIL® Xtra RC-45/25) and then use for experiments.

The main physico-chemical parameters relevant to this study are reported in

Reaction process

Acid birnessite reduction by H2O2

BPA degradation was tested in this system. BPA kinetic experiments were carried out in a 150 mL serum bottle in aerated solution at 293 K. Reactivity change assessment was investigated as following: 100 mL aqueous suspension of acid birnessite corresponding to an initial concentration of 345 μM (30 mg L -1 ) was first prepared and then an appropriate amounts of stock solutions of H2O2, BPA, and/or Mn(II) were added to start the reaction. The pH was adjusted to the desired value (6.5, unless otherwise indicated) using a pH meter by adding 0.1 M NaOH and/or HCl. The suspension was stirred with a speed of 350 rpm. At predetermined time interval, 1 mL of aqueous sample was withdrawn and filtered through 0.22 μm PTFE filter and immediately mixed with 200 μL of methanol to quench the reaction, and then analyzed.

The same batch experiments were carried out in the presence of silicate or phosphate to evaluate the combined effects of anions and H2O2 on BPA oxidation.

PDS and PMS activation by Mn2O3

The chemical reactions were performed in a borosilicate brown bottle in which 100 mL of solution was kept at room temperature (298 ± 2 K) under stirring (700 rpm) to ensure homogeneity. 25 µM of CAF was added to the 100 mL reaction solution containing a certain concentration of Mn2O3, then a predetermined amount of oxidants (i.e. H2O2, PS, PMS, or HSO3 -) was added to start the reaction. The solution pH was fixed by adding 0.1 M NaOH and/or HClO4. 2 mL of samples were withdrawn at the predetermined interval times and were filtered (0.20 µm PTFE filter) to remove Mn2O3 particles. The reaction was stopped by the addition of 100 µL of methanol (HPLC grade, Sigma-Aldrich) and then the samples were analyzed by HPLC.

PMS activation by γ-MnOOH

PMS was added to the solution containing selected ECDs and γ-MnOOH just before reaction starting. All reactions were performed in a brown glass bottle (125 mL) at room temperature (293 ± 2 K) and under stirring to ensure the homogeneity. Unless otherwise stated, experiments were performed at pH 6.5 ± 0.2. For HPLC analysis, 2 mL of solution were withdrawn at fixed time intervals, filtered using (0.20 µM PTFE filter) and reaction was immediately quenched with addition of 20% methanol (HPLC grade, Sigma-Aldrich). Sample solution were kept at 277 K and analysis were performed within 4 hours. Furthermore, to identify the generated radicals, EtOH, TBA and FFA were added to the solution before PMS addition and a kinetic approach based on the second-order rate constants and pseudo-first order decay determination was used to assess the radical species involvement.

Analysis methods

Determination of target pollutants concentration

The concentration of target pollutants (i.e. BPA, CAF, E2 and EE2) were identified using the and (E) for BPA using fluorescence detector.

H2O2 quantification

H2O2 concentration was determined using spectrofluorimetric detection [5]. The analysis is based on the reaction between H2O2 and 4-hydroxyphenylacetic acid which form the a stable

Determination of PDS and PMS concentration

The concentrations of PDS and PMS can be measured using a modified iodometric titration method [7]. The reaction mechanisms are as follow: PDS or PMS reacts with potassium iodide (KI) in the presence of sodium bicarbonate (NaHCO3) to generate iodine (I2). Then, I2 further reacts with iodide (I -) leading to the formation of triiodide (I 3 -).

Identification of caffeine byproducts

The identification of CAF transformation intermediates was obtained by UPLC-MS system coupled with an ESI ion source (MSD VL). The column was a Phenomenex Kinetex C18 column (100 mm × 2.1 mm, particle and pore sizes 2.6 μm × 100 Å) using an initial gradient of 5% acetonitrile and 95% water acidified with 1% formic acid, followed by a linear gradient to 95% acetonitrile within 15 min and kept constant during 10 min. The flow rate was 0.2 mL min -1 and the injection volume was 10 μL.

Identification of hydroxyl radicals

Terephthalic acid (TA) was reported to have the ability to trap hydroxyl radicals (HO • ) leading to the formation of 2-hydroxyterephthalic acid (TAOH) (Figure 2345). [8]. TAOH is a florescent product with an excitation wavelength at 320 nm and an emission wavelength at 425 nm, while TA cannot give any florescence. Therefore, the HO • generated in reactions can be quantified through detecting the product of HO • samples with TA using a fluorescence spectrophotometer.

Identification of reactive oxygen species

The common reactive oxygen species (such as sulfate radical, hydroxyl radical and singlet oxygen) can be identified using the electron paramagnetic resonance spectroscopy (EPR) with

Determination of Mn 2+ concentration

The concentration of dissolved Mn 2+ in solution was measured using an Atomic absorption spectroscopy (AAS). The parameters were settled as follows: the signal was set at AA (atomic absorption) with corrected background signal; the slit was set at 1.8 mm (width) / 2.3 mm (height); the detection wavelength was set at 279.48 nm; The lamp was installed for analysing Mn element with a courant of 30 mA; oxygen and acetylene flows were 10 and 2.5 L/min; integration time was 3s. Mn 2+ standard solution for AAS (TraceCERT®, 1000 mg L -1 Mn in 2% nitric acid) was used for making calibration curve of Mn 2+ (Figure 23456). The detection rages of Mn 2+ in AAS was 0.2 to 5 mg L -1 . In the study, all the Mn 2+ containing samples were first filtered through a 0.2 µm PTFE filter, then were acidified with the addition of 2% nitric acid (65%).

Determination of Mn average oxidation state

A. Determination of Mn AOS using oxalic acid-permanganate back-titration method

The oxalic acid-permanganate back-titration method was employed to measure the average oxidation state (AOS) of acid birnessite [10]. The process was conducted as follow: 0.1 g of the acid birnessite were dissolved in 2.5 mL of 0.5 M H2C2O4 and 5 mL of 1 M H2SO4 to reduce all highly charged manganese (such as Mn 3+ and Mn 4+ ) to Mn 2+ . The excess C2O4 2- AAS was used to determine the total Mn content (mMn). After that, according to the titration result and total Mn content, the AOS of acid birnessite was calculated (Eq. (2-3)).

2 MnO 4

-+ 5 C 2 O 4 2-+ 16 H + → 2 Mn 2+ + 10 CO 2 + 8 H 2 O (2-1) nMn = 2 × (C H2C2O4 × V H2C2O4 -2.5 C KMnO4 × V KMnO4 ) × 10 -3 (2-2)
X-ray powder diffraction can provide the detailed information about the crystallographic structure, chemical composition, and physical properties of materials. In the study, the solid was characterized using XRD spectra obtained with the Bruker AXS D8 Advance diffractometer (θ-2θ Bragg-Brentano geometry). The monochromatized CuKα1 (1.54Å) was used to irradiate over the range of 10º-100º 2θ at a step size of 0.02º. Phases were determined according to PDF database in Jade 6 software.

C. Microscopy-based techniques

Scanning electron microscopy (SEM) can provide a high-quality image of the material surface.

Thus, the morphology of materials (such as rod, flower, and wire) can be identified using the SEM. The High-resolution transmission electron microscopy (HRTEM) provides more information about the internal structure of materials. The presented SEM and TEM images of this thesis were obtained using 1) SEM (JEOL JSM-7100F) and 2) HRTEM (JOEL 2100 LaB6).

Determination of total organic carbon

The total organic carbon (TOC) is an important parameter to estimate the mineralization efficiency of organic pollutants. The TOC value can be determined as the difference between the total carbon (TC) and the inorganic carbon (IC) values. In the treatment process, the decreasing of TOC value indicates that the organic pollutant has been converted into carbon dioxide and carbonates ions. In this study, the TOC was determined using a TOC-L CPH CN200, SHIMADZU instrument. All the samples were filtered using the 0.45 𝜇m PTFE filter. The detection range of TC was 0-15 mgC L -1 . For the IC, the detection range was 0-4 mgC L -1 . The calibration curve of TC and IC with 50 𝜇L injection volume were shown in Figure 234567. investigated the degradation of toluene sorbed on activated carbon in the system of β-MnO2/H2O2. However, the previous studies focused on the elimination efficiency of pollutants, and overlooked the influence of H2O2 on the oxidative activity of MnO2. H2O2 has an ambivalent redox activity and has been commonly found in aquatic environments at concentration ranging from nM to µM. Birnessite-type MnO2 play a key role in the global carbon cycle as well as pollutant dynamics. Therefore, it is significant to investigate the impact of environmental levels of H2O2 and oxyanions on the oxidative reactivity of acid birnessite.

Introduction

Layered structure manganese oxides (e.g. birnessite) are ubiquitous in a wide range of aquatic and terrestrial environments [1]. 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 [2,3]. 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 [4][5][6][7][8][9]. Acid birnessite has been widely investigated because it is structurally similar to biogenic "natural" manganese oxides [10][11][12][13]. The birnessite reactivity is mainly affected by solution pH, surface properties and composition of MnO2 and structural characteristics of redox-active contaminants [4][5][6]12].

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 (H2O2) has an ambivalent redox activity and is commonly found in aquatic environments at concentration ranging from nM to µM [14][15][16][17][18][19][20][21][22][23][24][25][26].

In environmental systems, photodependent reactions mediated by natural organic matter and biological-mediated processes dominate H2O2 production [26]. Sunlight-induced photochemical reactions mediated by organic ligands and/or biochemical metabolic processes can result in amounts of H2O2 up to 20 µM [14][15][16][17][18]. Input from rainwater can reach up to 40 μM [20,21], while hydrogen peroxide production has been often observed in irradiated seawater [19][20][21]. Light-independent generation of H2O2 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 [22,23]. For example, natural hydroquinones or hydroquinone moieties ubiquitously present in reduced organic matter can donate electrons to O2 to generate H2O2 [24,25]. H2O2 production through microbial mechanism has also been shown to enhance iodide oxidation and organo-iodine formation in soils and sediments [27]. All these studies suggested that H2O2 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 H2O2 production [28].

Despite the catalyzed decomposition of hydrogen peroxide by MnO2 has been well investigated [29][30][31], knowledge is very limited on the effect of H2O2 on the reactivity of birnessite at environmental levels of H2O2. In engineering applications, metal oxides such as MnO2 are generally investigated for catalytic decomposition of H2O2 and then formation of reactive transient species using very high concentrations of hydrogen peroxide (H2O2) (≥ 0.5 mM) [29][30][31]. While these remediation studies aimed to eliminate pollution in contaminated systems, they overlooked the influence of H2O2 on the electron transfer heterogeneous reaction and then the oxidative activity of birnessite. In addition, little is known about the influence of H2O2 on the birnessite composition and interactions of redox products (e.g. Mn(II)) with MnO2 surfaces.

In this study, we examine how H2O2 affects the reactivity of birnessite under environmentally relevant conditions. To monitor H2O2-induced changes in birnessite reactivity, we used Bisphenol A (BPA) as a model compound because it has a strong reactivity with the birnessite (δ-MnO2), with a well-documented underlying redox mechanism [32][33][34][35]. 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 H2O2 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 (H2O2/MnO2 molar ratio, dissolved Mn(II) amount and pH value). To check whether the impact of H2O2 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 H2O2 concentrations as low as 10 µM or with H2O2/MnO2 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.

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 with a speed of 350 rpm. 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 or phosphate to evaluate the combined effects of anions and H2O2 on BPA oxidation.

Potential generation of hydroxyl radical was monitored through the fluorescence emission spectrum of 2-hydroxyterephthalic acid (λex 320 nm, λem 425 nm) using a spectrofluorometer (Shimadzu RF-5301PC) (more information is given in the 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.

H2O2 concentration was determined using a spectrofluorometric detection as described in our previous work [39]. Briefly, samples containing H2O2 were mixed with 4-hydroxyphenylacetic 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 H2O2 [40].

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 (HNO3, 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 lied at 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 41 while phosphate and nitrate concentrations by ion chromatography [42]. 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 N2:H2 98:2). Prior for anoxic experiments, all solutions were sparged with nitrogen to remove oxygen. These preliminary tests showed that there is no effect of oxic conditions on the BPA removal in presence of H2O2 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, and given in the caption of each figure.

Results and Discussion

Effect of H2O2 on the oxidative ability of birnessite

BPA removal was monitored in the presence of acid birnessite (AB) at different H2O2 concentrations at pH 6.5 (Figure 1). In the absence of H2O2, almost complete removal of BPA was observed after 24 h of reaction time. Mass balance showed that the oxidation reaction was mainly involved in the removal of BPA in presence of MnO2, while adsorption was very low under our experimental conditions (i.e. less than 5% of initial BPA), which is consistent with previous investigations [6,32,33]. Indeed, BPA containing two hydroxyphenyl functionalities is known to weakly interact with mineral surfaces, resulting in lower adsorption affinity to metal-oxides [32]. As previously reported [6,33,43], 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 O2 + BPA + 4 H + → Mn II + BPAox + 2 H2O

Electron exchange of BPA with MnO2 forms a radical, followed by a series of reactions including radical coupling, fragmentation, substitution and elimination to form multiple byproducts [6,32,33]. 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 [44]. In the presence of hydrogen peroxide, the oxidative removal of BPA gradually decreased with increasing in H2O2 concentration, and completely suppressed at higher dose of H2O2 (200 µM equivalent to a H2O2/MnO2 molar ratio ~ 0.6). In these experiments, the H2O2 decomposition is relatively fast whatever the investigated pH, since complete decomposition of H2O2 was achieved after 15 min of reaction, even for the highest H2O2 dose (See Figure S2 andS3).

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 (1 mM) or t-butanol (1 mM) as hydroxylradical scavengers confirmed that the BPA removal is a no radical-based oxidation process regardless of the presence or absence of H2O2.

The suppressive effect of H2O2 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 (see example in Figure S4). The initial rate constants for all investigated H2O2 amounts exhibited the same order, pH 4 > pH 6.5 > pH 8 (Figure 2). In the absence of H2O2, 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 MnO2 [33,34,45]. We attribute this to variability in two pH-dependent factors: 1) speciation of BPA that affects binding to MnO2 surfaces and then oxidation (Figure S5), and 2) redox-potential of MnO2 that decreases when the pH increased from 4 to 8 [46]. At relatively low pH, favorable electrostatic interactions between neutral BPA molecules and the negatively charged surface of MnO2 (PZC of MnO2 is ~ 2.3-2.9) may exist. On the other hand, reductive conversion of MnO2 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 H2O2, the removal rate constants of BPA sharply decreased with increasing in H2O2 amount whatever the investigated pH value (Figure 2). As for BPA, it is previously reported that hydrogen peroxide can also reduce Mn(IV) to Mn(III) and then Mn(II) as follows [47,48]:

Mn IV O2 + H2O2 + 2 H + → Mn II + O2 + 2 H2O (2)
To check this possibility under our experimental conditions, the reductive dissolution of MnO2 in presence of various amounts of H2O2 was monitored at three pH values (Figure 3). In both Eq. 1 and Eq. 2, protons are directly involved in the oxidation of BPA as well as the reductive dissolution of MnO2. In the absence of H2O2, dissolved Mn(II) was only detected at pH 4, while it was below the detection limit at pH 6.5 or 8. Increasing H2O2 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 H2O2 decomposition over time, where the complete decomposition of H2O2 observed after approximately 15 min of reaction time coincides with the peak observed for dissolved Mn(II) generation (Figure S3). This behavior also suggests that the production of Mn(II) via H2O2-induced reduction of MnO2 is faster than the Mn(II) adsorption onto birnessite. Once the added H2O2 is fully decomposed, the generated dissolved Mn(II) will in turn sorb onto MnO2 surfaces. Because higher pH implies more Mn(Ⅱ) binding [49][50][51], greater amounts of Mn(II) were observed at low pH values (Figure 3). When a further addition of H2O2 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. This suggests that the reductive dissolution of MnO2 by H2O2 is a fast process, as compared to the reaction between MnO2 and BPA.

To check whether Eq.1 and Eq.2 may occur simultaneously during the heterogeneous reactions with MnO2, we compare the standard potential of involved reactions at the working pH range.

MnO2 has standard (reduction) potential varying from 0.99 to 0.76 V when the pH rises from 4 to 8 for MnO2 + 4 H + + 2 e -→ Mn II + 2 H2O. This value is higher than the standard oxidation potential of O2/H2O2 (O2 + 2 H + + 2e -→H2O2 0.68 V vs. NHE) or of BPA which varies between 0.75 V and 0.52 V in the pH range 4-8 [52][53][54]. 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 MnO2 is considered in excess, the reaction rate depends on the concentration of reactants (i.e. H2O2 or BPA). When the H2O2 amount increases, the reaction 2 should become much faster than the reaction 1. This was experimentally shown in Figures 1 and3 and Figure S2, where both reactions Eq. 1 and Eq. 2 occurred simultaneously until complete inhibition of Eq. 1 at the highest H2O2 amount (200 µM). Though H2O2 was full decomposed after 15 min of reaction time (Figure S2), the suppressive effect on the birnessite reactivity was still observed over 24 h of reaction time (Figure 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 MnO2 and H2O2 were allowed to react first (i.e. pre-equilibration step of 30 min) until total decomposition of H2O2 before addition of BPA (Figure S6). Therefore, the H2O2-mediated reduction of MnO2 (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 MnO2 surfaces and then be oxidized into higher valence Mn:

Mn II + MnO2 → Mn II /Mn III -MnO2

This oxidation reaction may be made possible by the residual H2O2 and/or generated O2 (through Eq. 2), which is pH-dependent. The contribution of O2 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 MnO2 particles and then fast sedimentation was observed upon addition of H2O2. 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 MnO2 surfaces, as previously reported [55]. Furthermore, disproportionation /comproportionation reaction (Eq. 4) may occur within the MnO2, 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 [12,32,35,51,56]:

Mn II + Mn IV ↔ 2 Mn III (4)
Since this reaction is pH-dependent, we have measured the average oxidation state (AOS) of birnessite over the whole investigated pH range (4)(5)(6)(7)(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 MnO2 upon addition of 200 µM of H2O2 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 [57]. The effect of H2O2 on the structure of birnessite will be discussed in the following section.

Collectively, these results suggest that the H2O2 induced reduction of MnO2 is an effective process, and the fast generation of Mn(II) is key for suppressed oxidative capacity of MnO2 towards BPA. To confirm the impact of Mn(II) on the reactivity of MnO2, BPA removal kinetics were investigated in Mn(Ⅱ)-amended birnessite suspensions in the following section.

Effects of Mn(Ⅱ) on the removal capacity of MnO2

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 (Figure 4). 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 MnO2 surface [58,59]. Similarly, the Mn(Ⅱ) adsorption should lead to a decrease in the negative surface charge, thereby altering the binding capacity of MnO2. When the Mn(II) amount further increases, we speculate a competition between BPA and Mn(Ⅱ) to bind at reactive sites of MnO2 during the first kinetic phase, as recently observed for quinolones [43].

As for H2O2, 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 (Figure 4). In addition to the comproportionation reaction that is pH-dependent, the decrease in the oxidative ability of MnO2 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 MnO2 surfaces suggests competition of BPA and Mn(Ⅱ) for surface sites, which is pH dependent [33,45]. Adsorption tests showed that 100 µM of Mn(Ⅱ)

were completely removed at both pH 6.5 and 8 (i.e. dissolved Mn(II) was below the detection limit), while 70 % of the added Mn(II) was removed 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 (Figure S7). 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 [49,50]. 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 H2O2. As expected, Mn(Ⅱ) binding to MnO2 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 MnO2 surfaces [49,50].

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) [49,50]. 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)/MnO2 molar ratio = 0.3) showed no notable transformation of MnO2 over the investigated pH range (Figure 5). However, titration experiments showed that addition of 100 µM of dissolved Mn(Ⅱ) dropped down the AOS of MnO2 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)/MnO2 ratios (0.09 and 0.13), while secondary phases such as MnOOH and Mn3O4

can be generated at high Mn(II)/MnO2 ratios (0.5 and 1) [60]. 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)/MnO2 molar ratio = 0.3) or H2O2 (H2O2/MnO2 molar ratio = 0.6), suggesting changes in the structure or ordering of the birnessite mineral sheets (Figure 5). We also observed these broad bands in the XRD pattern of synthetic Mn(III)-rich birnessite sample (see Figure 5). This is consistent with previous studies, which suggested that Mn III formation can be made in edge-sharing complexes on MnO2 edge sites or around vacancy sites in octahedral layers of MnO2 [5]. 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 [49].

Synergistic effects of anions and H2O2 on the oxidative activity of MnO2

To check whether the suppression of reactivity at environmental relevant concentrations of H2O2 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 (Figure 6). At 0 µM of H2O2, 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 MnO2 surfaces under the experimental conditions of this study (pH 6.5, 100 µM of anion, 345 µM of MnO2). However, significant decrease in kinetic rate constants was observed in presence of silicate and phosphate when H2O2 was added Chapter III 123 in the reaction medium (H2O2/MnO2 molar ratio between 0.03 and 0.3), while no impact on the reactivity was observed in presence of nitrate (Figure 6). These results can be explained if the Mn(II) generated through H2O2-induced reduction of MnO2 renders the MnO2 surfaces more able to bind anionic ligands, thereby altering the surface reactivity towards BPA. It was previously reported that unlike nitrate, silicate or phosphate may interact with MnO2 through hydrogen bonding or formation of outer-sphere complexes with surface hydroxyl groups of MnO2 [59][60][61][62]. On the other hand, previous studies reported that the presence of dissolved silicate may decrease the oxidation rate constants of MnO2 toward chlorinated compounds [63,64]. They have attributed the decrease in MnO2 reactivity to the surface-bound silicate but none of them have provided adsorption data of silicate onto MnO2. In the present work, we showed that silicate or phosphate has a very low affinity to the negatively charged MnO2 surfaces under our experimental conditions, but the presence of 100 µM of Mn(II) significantly enhanced their adsorption amounts onto MnO2 (Figure S8). This increased adsorption might be due to the changes in both the surface charge of MnO2 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 [65]. However, no precipitation is possible under our experimental conditions due to the low degree of saturation, resulting from the very low aqueous concentration of ligand and complete removal of dissolved Mn(II) by adsorption (log Ks Mn(Ⅱ)/phosphate = -27.07) [66]. Furthermore, Mn(Ⅱ) binding should lead to a decrease in the negative surface charge, thereby enhancing ability to bind anions through electrostatic interactions [55,58,59]. Therefore, the MnO2-bound Mn(II) system could adsorb more effectively anions such as phosphate or silicate through cation bridging, as recently reported for humic acid [55]. 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 MnO2 surfaces.

It should be noted that the decomposition of H2O2 (200 µM) was found to be similar in presence or absence of silicate or phosphate (100 µM) (Figure S9). 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 [29][30][31]. Despite silicate adsorption on MnO2 surfaces was found negligible, there was an inhibition effect on the H2O2 decomposition rate over the investigated silicate concentration range (0 to 1.5 mM) [30,67]. Although these works have used different experimental conditions (e.g. much higher concentration of ligands or H2O2), and different MnO2 types, the present findings call for indepth consideration of the combined/synergistic effects that the cations and anions co-presents in the reaction medium may have on the MnO2 reactivity. Overall, these findings suggest that Mn(II) generation through H2O2-mediated reduction may alter both adsorption and redox transformation of environmental compounds.

Conclusions

Photochemical and dark production of H2O2 make it ubiquitous not only in oceanic and atmospheric systems but also in the subsurface environment. Here, we have notably demonstrated that environmental levels of hydrogen peroxide can induce reductive dissolution of MnO2 into Mn(II), thereby affecting the surface reactivity of MnO2. As the H2O2 decomposition is a fast process, the generated dissolved Mn(II) binds to MnO2 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

Motivation

The key role of trivalent manganese species in promoting sulfate radical-based advanced oxidation processes (SR-AOPs) has been reported. For example, Huang et Although the excellent performance of Mn(III) in radical precursors activation has been well highlighted in the previous literature, the activation mechanism of PDS and PMS by Mn(III) oxides (such as Mn2O3 and manganite) is not clear. Moreover, some important influence factors for the reactivity of Mn2O3 such as solution pH were not investigated. The comparison of PMS activation mechanism by different Mn(III) oxides was not reported. Furthermore, most of former investigations were conducted in the pure water, the efficiency of Mn(III) oxides + PDS/PMS system in the practical wastewater was not evaluated.

Considering the above-mentioned questions, the degradation of organic pollutants in water by Mn(III) oxides activated PDS and PMS were investigated. In the Chapter IV, the results of my study are divided into two parts. In the part 1, the caffeine degradation using PDS and PMS in the presence of Mn2O3 was investigated. In the part 2, the removal of estrogenic compounds in water by manganite-activated PMS was studied. Advanced oxidation processes (AOPs) constitute a promising technology for the remediation of environmental compartments [1]. AOPs are efficient in the removal of harmful and persistent organic pollutants in water due to the generation of strongly reactive species [2]. Hydroxyl radical (HO • ) belongs to the reactive oxygen species (ROS) and has a high oxidation potential (E 0 = 2.73 V vs NHE) ) [3]. Similarly, the reactive sulfur species (RSS) such as sulfate radical (SO4 •-) (E 0 = 2.6 V vs NHE) have also caught the attention of the scientific community due to their efficient oxidation abilities [4]. Generation of ROS and RSS is usually performed through the activation of radical precursors by the catalytic processes, such as direct photolysis [5],

homogeneous metal ions activation [6], and heterogeneous metal oxides catalysis [7]. The metal oxides-based heterogeneous AOPs represent one of the most investigated strategies for organic pollutants removal in the laboratory and pilot scales [8]. Among various metal oxides, iron oxides are by far the most studied metal oxides for reactive radicals generation in water [9].

However, other metal oxides such as manganese oxides (MnOx) are gaining stronger interests due to their versatile oxidation states (II, III, IV, and VII) [10].

Among the different MnOx, manganese dioxide (MnO2) has attracted increasing attention due to its good oxidative and catalytic performance for the elimination of pollutants [11]. For ( 1 O2) [13]. Moreover, the influence factors for the good catalytic and oxidative performance of various MnO2 were investigated by Huang et al., and the results showed that the efficient oxidation of pollutants by MnO2 with or without oxidants was attributed to the occurrence of Mn(III) on the surface of MnO2 [14,15]. Therefore, it is believed that the presence of Mn(III) species in MnOx can improve the activation efficiency of PS/PMS. Given the above, the direct activation of PS/PMS by Mn(III) oxides was investigated by previous studies. For example, Xu et al. reported the efficient elimination of p-chloroaniline in the γ-MnOOH/PS system through the formation of active intermediates between γ-MnOOH and PS [16]. Wang et al. indicated the good reactivity of octahedral Mn3O4 on PMS activation for ciprofloxacin degradation due to the presence of high contents of Mn(III), Mn(IV) and hydroxyl groups (-OH) [17]. Moreover, the degradation of organic pollutants by PMS using synthesized Mn2O3 as the catalyst was also reported. For instance, Saputra et al. indicated that Mn2O3 compared with MnO, Mn3O4, MnO2

showed the highest reactivity in PMS activation for phenol degradation due to the formation of both SO4 •-and SO5 •-radicals [18]. The effect of Mn2O3 structures on PMS activation was also investigated by Saputra et al., and the results showed cubic-Mn2O3 has higher PMS catalysis efficiency due to its high surface area and surface atoms arrangement [19]. These experiments demonstrated well the effect of Mn oxidation states and catalyst structures on the activation of PMS. However, the catalytic mechanism of Mn2O3 on PMS is not clear, and some important influence factors for the reactivity of Mn2O3 such as solution pH were not investigated.

Furthermore, Khan et al. investigated the activation of PS using synthesized Mn2O3 and MnO2

for the degradation of phenol in solution [20]. The results showed that phenol was efficiently removed in A-Mn2O3/PS system due to the generation of singlet oxygen ( 1 O2). Although the authors adopted different experimental methods to identify the generated reactive species and investigated various influence parameters for the reactivity of A-Mn2O3/PS system, it should be noted that the experiments were conducted in the acidic environment (pH~3.2). Moreover, since different reactive species were generated in γ -MnOOH/PS system under different solution pH [21], thus it is reasonable to deduce that the type of generated reactive species in the Mn2O3/PS system can be influenced by solution pH. However, the authors only studied the generated reactive species in the acidic solution. Therefore, it is meaningful to investigate the mechanism of PS/PMS activation by Mn2O3 in the near-neutral solution and evaluate the suitability of the Mn2O3 activated PS/PMS system in real polluted waters. Caffeine (CAF) was used as a model of organic contaminants since its adverse effects on aquatic organisms and human beings [22]. Moreover, CAF as the emerging contaminant has been widely detected in different water matrices at concentrations ranging from ng/L to µg/L [23].

In this work, we investigated and compared the ability of Mn2O3 to activate PS and PMS under different experimental conditions, i.e. pH, Mn2O3 dosage, and radical precursor concentrations.

Reactive radical species and non-radical pathways were identified and their contribution to the pollutant removal was quantified using chemical scavenging experiments, electron paramagnetic resonance spectroscopy (EPR) and X-ray photoelectron spectroscopy (XPS).

Moreover, the CAF degradation pathway, main degradation products and mineralization efficiency were reported. Finally, the ability of both systems was compared in real sewage treatment plant water, where 80 % of CAF degradation was obtained by the proposed method after 20 hours.

Materials and Methods

Chemicals

All chemicals were reagent-grade (see Text S1 in Supplementary Materials) and used as received. Ultrapure water (Millipore, resistivity 18.2 MΩ cm) was used in the preparation of all solutions.

Catalyst characterization

The purity of the used Mn2O3 sample (Sigma-Aldrich) was confirmed by powder X-ray diffraction (XRD) (Figure S1). The average oxidation state (AOS) of Mn after reaction was obtained using X-ray photoelectron spectroscopy (XPS). The analysis conditions of XRD and XPS were presented in Text S2.

Experimental procedures

The chemical reactions were performed in a borosilicate brown bottle in which 100 mL of solution was kept at room temperature (298 ± 2 K) under stirring (700 rpm) to ensure homogeneity. 25 µM of CAF was added to the 100 mL reaction solution containing a certain concentration of Mn2O3, then a predetermined amount of oxidants (i.e. H2O2, PS, PMS, and HSO3 -) was added to start the reaction. The solution pH was fixed by adding 0.1 M sodium hydroxide and perchloric acid. 2 mL of samples were withdrawn at the predetermined interval times and were filtered (0.20 µm PTFE filter) to remove Mn2O3 particles. The reaction was stopped by the addition of 100 µL of methanol (HPLC grade, Sigma-Aldrich) and then the samples were analyzed by HPLC.

To determine the involvement of reactive radical species leading to the degradation of CAF, a kinetic competition study was performed using different radical scavengers such as tert-butanol (TBA) and 2-Propanol (2-Pr) to quench hydroxyl radical (HO • ) and sulfate radical (SO4 •-) [24,25]. Furfuryl alcohol (FFA) was used to trap singlet oxygen ( 1 O2) [26]. The contribution of main radical species to the degradation of CAF was estimated considering the second-order rate constants between the selected radicals and quenchers or CAF (see Table S1). The quenching efficiency was determined as C1k1/(C1k1 + C2k2). Where C1 was the initial concentration of organic scavengers and k1 was the second-order rate constant with selected oxidant species (i.e. HO • , SO4 •-, or 1 O2), C2 and k2 were the concentration of CAF and reactivity constant with the oxidant species.

Sewage treatment plant water (STPW) was collected in December 2019 from the urban treatment plant named "3 rivières" in Clermont-Ferrand and used after 0.45 µm filtration using PTFE filter. The main physicochemical characteristics are reported in Table S2.

Chemical analysis and kinetics

CAF concentration was determined by a high-performance liquid chromatography system (HPLC, Alliance 2695 Waters) equipped with a photodiode array detector (DAD) and C18 column (Agilent Eclipse XDB, 250 × 4.6 mm, 5 μm of particle size). The mobile phase was a mixture of acetonitrile (ACN) and milli-q water acidified with 0.1 % phosphoric acid in gradient eluent mode (VACN/ Vwater from 20/80 to 70/30 in 6 min). The flow rate was set at 0.8 mL min - 1 and the detection wavelength of CAF was fixed at 272 nm. The pseudo-first-order rate constant of CAF (k'CAF) was determined from the slope of ln (C/C0) vs time plot, where C0 and C are the concentration of CAF at time 0 and t respectively. The error bars associated to the degradation constant data represent 3σ, derived from the scattering of the experimental data around the fit curves (intra-series variability).

To investigate the mineralization of CAF, total organic carbon (TOC) analysis (TOC-L CPH PS and PMS concentrations were measured using a modified spectrophotometric method [27].

Briefly, 100 µL of PS or PMS sample was first transferred into 5 mL of the mixture of potassium iodide (KI, Sigma-Aldrich) and sodium bicarbonate (NaHCO3, Sigma-Aldrich) to generate iodine (I2). Then, I2 further reacted with iodide (I -) leading to the formation of complex (I 3 -).

Finally, I 3 -was detected using a Varian Cary 3 UV-visible spectrophotometer (Abs at 351 nm) and quantified using a calibration curve.

The identification of radicals was performed by electron paramagnetic resonance spectroscopy (EPR). The X-band cw-EPR spectra (modulation frequency of 100 kHz) were monitored at 293 K with the EMXplus spectrometer (Bruker) equipped with the High Sensitivity Probe-head (Bruker) in the small quartz flat cell (Wilmad-LabGlass, WG 808-Q) using 5,5-dimethyl-1- The identification of CAF transformation intermediates was obtained by an LC-MS system coupled with an ESI ion source (MSD VL). The column was a Phenomenex Kinetex C18 column (2.6 μm × 100 Å; 100 × 2.1 mm) using an initial gradient of 5% ACN and 95% water acidified with 1% formic acid (Sigma-Aldrich), followed by a linear gradient to 95% ACN within 15 min and kept constant during 10 min. The flow rate was 0.2 mL min -1 and the injection volume of 10 μL.

Results and discussion

Reactivity of Mn2O3 with PS and PMS

The reactivity of Mn2O3 with different radical precursors (i.e. H2O2, PS, PMS, and HSO3 -) was investigated. As shown in Figure 1, the CAF degradation was observed only in the Mn2O3 activated PS and PMS (5 mM) systems. Up to 40% of removal extent was observed after 3 h of reaction time in the unbuffered solutions. Although the solution pH dropped to 5.5 ± 0.2 and 2.7 ± 0.2 with the addition of PS and PMS, the purpose of checking the reactivity of Mn2O3 with PS and PMS was achieved (Table S3). Moreover, blank and adsorption tests showed that CAF removal by adsorption on Mn2O3 or direct oxidation by radical precursors was negligible whatever the investigated pH (Figure S2 andS3). S3), and T = 298 K.

The formation of radical species and their relative contribution to the CAF degradation were determined by chemical scavenging and EPR experiments. Before starting the scavenging tests, suitable concentrations of scavengers were determined through the calculation of scavenging efficiency with reactive oxygen species (ROS). As shown in Table S4, in the presence of 10 mM of TBA, 10 mM of 2-Pr and 1 mM of FFA, a scavenging efficiency of 97.4, 93.2 and 99.4 % of the corresponding ROS was achieved. Thus, the proposed scavenger concentration was enough to inhibit almost the whole radical reactions. In addition, to verify if the degradation of CAF was originated from the heterogeneous reaction between PS/PMS and Mn2O3, the leaching of Mn 2+ in Mn2O3 activated PS/PMS reactions and the homogeneous activation of PS/PMS by Mn 2+ were studied. As shown in Figure S4. A negligible Mn 2+ concentration (~ 1 µM) was determined when 0.4 g L -1 of Mn2O3 was suspended in solution. Moreover, the degradation efficiency of CAF in the Mn 2+ activated PS/PMS system was far lower than Mn2O3, even the employed concentration of Mn 2+ was 10-1000 times higher than leached Mn 2+ . Thus, the removal of CAF was attributed to the activation of PS/PMS by Mn2O3.

Reactivity of Mn2O3 with PS

In the system of Mn2O3/PS, only a negligible effect of TBA and 2-Pr was observed indicating that HO • and SO4 •-were not the main radical species involving in the CAF degradation (Figure 2A). However, in the presence of 1 mM of FFA, the CAF degradation constant (k'CAF) dropped from 2.4 × 10 -3 min -1 to 2.6 × 10 -5 min -1 , corresponding to 99% of inhibition after 8 h (Figure S5). The inhibition of CAF degradation in the presence of FFA can be due to the quenching of 1 O2. To identify whether 1 O2 was generated in the Mn2O3/PS system, the EPR experiments were conducted using TEMP as the spin trapping agent and no signal corresponding to the formation of TEMPO was observed (Figure S6), thus indicating the inhibition of CAF degradation by FFA was not related to the scavenging of 1 O2. Other reasons can explain the inhibition of CAF by FFA, such as the scavenging of HO • , the increased PS consumption, and the occupied reactive sites on the Mn2O3 surface [16,26,29]. The scavenging of HO • was excluded since CAF degradation was not inhibited in the presence of TBA (Figure 2A) while CAF removal was completely suppressed after the simultaneous addition of 2-Pr and FFA (Figure S7). Moreover, the PS decomposition due to the reactivity with FFA was negligible (Figure S8) compared to the reactivity in the presence of PMS [29]. Therefore, the inhibition of CAF degradation in presence of FFA was attributed to the adsorption of FFA on the surface of Mn2O3 and competing with PS ions for the active sites. The results are in agreement with previous studies which reported FFA can affect the surface properties of catalysts [16,30].

Therefore, a non-radical mechanism of PS activation is proposed and it generally includes the generation of 1 O2 (as mentioned just above, this hypothesis is excluded), the surface-activated PS species, or the electron transfer process [31]. It is worth noting that the electron transfer mechanism in PS activation usually happens when carbon-based materials were used as catalysts [32]. Therefore, the removal of CAF in the Mn2O3/PS system was attributed to the formation of Mn2O3 surface-activated PS. The result is in agreement with previous studies which indicated no 1 O2 was generated but the active intermediates formed between PS and catalyst were responsible for the degradation of pollutants [16,21].

To better understand the role of PS, the concentration of PS in different systems was monitored (Figure 2B). The depletion of PS by CAF alone was negligible, whereas a small amount of PS was consumed by Mn2O3 in the absence of CAF. This result indicates that PS was absorbed on the active sites present on the surface of Mn2O3. CAF as an electron donor leads to the increased consumption of activated PS since in the Mn2O3/PS system the depletion of PS was enhanced.

The non-radical mechanism of PS activation was further verified through the measurement of the Mn oxidation state in the Mn2O3/PS system. The average oxidation state (AOS) of samples was calculated from the binding energy difference (∆E) of Mn 3s doublet splitting as following: AOS = 8.956 -1.126 × ∆E [33]. Figure S9 shows that the AOS of Mn2O3 did not change after reaction with PS (AOS ∼3.00), indicating the activation of PS by Mn2O3 was a non-radical catalytic process.

Reactivity of Mn2O3 with PMS

In the PMS-containing system, the presence of TBA and 2-Pr decreased the k'CAF from 6.5 × 10 -3 min -1 to 4.0 × 10 -3 min -1 and 7.0 × 10 -4 min -1 , respectively (Figure 3A and Figure S10)

indicating the predominant role of SO4 •-compared to HO • . PMS can react with both ≡Mn III (Eq.

1) and ≡Mn IV (Eq. 2) leading to the formation of SO4 •-and SO5 •- [34]. Moreover, SO5 •-selfreaction in water can lead to the formation of SO4 •-(Eq. 3) [35]. Formation of HO • occurs through SO4 •-reaction with water and hydroxide, although the latter reaction is only relevant in alkaline solution (Eqs. 4 and 5) [36]. The generation of SO4 •-and HO • was further identified by EPR experiments. As shown in Figure 3B, the EPR spectrum represents a superposition of two individual signals (according to the results of the best simulation fit), including the nine-line signal of 5,5-dimethyl-2-oxo-pyrroline-1-oxyl ( • DMPO-X) with the characteristic spin-Hamiltonian parameters (aN = 0.732 mT, aH  (2H) = 0.406 mT; g = 2.0070  0.0001), which is often produced by over-oxidation of DMPO in the presence of reactive oxygen species (e.g. HO • or 1 O2) [37]. The second signal is characterized by the hyperfine coupling constants (aN = 1.373 mT, aH  = 1.019 mT, aH  = 0.135 mT, aH  = 0.073 mT), which correlate well with

• DMPO-SO4 -spin-adduct with limited stability in aqueous media [38].

≡Mn III + HSO5 -→ ≡Mn IV + SO4 •-+ OH - (1)
≡Mn IV + HSO5 -→ ≡Mn III + SO5 Since the addition of FFA exhibited a strong inhibition effect (~ 85%) on CAF degradation (Figure S11A), thus, to verify whether the generated • DMPO-X was derived from the reaction with 1 O2, the EPR experiment was further performed. As shown in Figure S11B, in the presence of 25 mM of TEMP, a typical TEMPO signal (three-line signal with 1:1:1 intensity ratio) was observed in the presence of PMS alone which can be attributed to the self-decomposition of PMS [39]. In Mn2O3/PMS system, the TEMPO signal was approximately equal to the TEMPO intensity from PMS alone, indicating that no further 1 O2 was generated in the Mn2O3/PMS system. Moreover, the addition of FFA in the Mn2O3/PMS system also showed a negligible effect on the intensity of the TEMPO signal. Thus, the formation of • DMPO-X in the system of Mn2O3/PMS cannot be contributed to 1 O2. On the other side, the generated HO • might be responsible for the formation of • DMPO-X, consistent with the observations of Yang et al. [39].

•-+ H + (2) SO5 •-+ SO5 •-→ O2 + 2 SO4 •- (3) SO4 •-+ OH -→ SO4 2-+ HO • (4) SO4 •-+ H2O → SO4 2-+ HO • + H + (5)
The decomposition of PMS was further investigated. Figure S11C shows that PMS can be decomposed by FFA alone, and the depletion efficiency of PMS is higher than Mn2O3, suggesting the inhibition of CAF degradation in presence of FFA might be also attributed to the increased consumption of PMS.

Besides, the AOS of Mn in the Mn2O3/PMS system was measured, as shown in Figure S9C.

The AOS of Mn2O3 was increased to 3.10 after reacting with PMS, suggesting the generation of Mn (IV). It is known that Mn (IV) can oxidize organic pollutants compounds through a heterogeneous electron transfer process [14], thus the non-radical degradation of CAF in the Mn2O3/PMS system might be related to the formation of Mn (IV).

In summary, different activation mechanisms of PS and PMS by the same Mn2O3 were exhibited. In the Mn2O3/PS system, the catalyst surface-activated PS molecules were the primary reactive species that are responsible for the degradation of pollutants. This can be attributed to the symmetric structure of PS which facilities the contained O-O bonds binding to Mn2O3 surface through the substitution with surface hydroxyl groups (-OH) [13,16,40]. In comparison, the activation of PMS by Mn2O3 was mainly processed with the generation of radicals. It can be explained by the fact that the unsymmetric structure of PMS can be easier activated by catalysts, the causes are in agreement with the results of previous research [41][42][43].

Effect of pH

The CAF removal by Mn2O3-mediated activation of PS or PMS was also investigated over a wide range of initial pH (3)(4)(5)(6)(7)(8) and degradation pseudo-first order rates were plotted versus pH (Figure 4). In the presence of PS, k'CAF decreased from 4.7 × 10 -3 to 2.6 × 10 -4 min -1 from pH 3 to 8. It should be noted that the neutral form of CAF predominates over the investigated pH range (pKa = 10.4). The greater efficiency observed at low pH might result from the favorable electrostatic interactions between positively charged Mn2O3 (PZC ~5.9) and PS [44]. The system Mn2O3/PMS exhibited an improved k'CAF with increasing the solution pH from 2.2 × 10 -3 min -1 to 6.4 × 10 -3 min -1 , followed by a decrease to 2.4 × 10 -3 min -1 (Figure 4). The low efficiency at low pH may come from the impact of H + concentration on the reactivity of SO4 •-, resulting in the low CAF degradation [45]. Since the oxidation potential of HO • is lower than SO4 •-under alkaline environments, the transformation from SO4 •-to HO • may explain the low efficiency at high pH values. To further explore this transformation process, the scavenging tests with the addition of TBA and 2-Pr were conducted under various initial solution pH (Figure S12). The contribution of HO • and SO4 •-in the degradation of CAF was also assessed using Eq. S1 (Text S3). As shown in Figure S12D, the radical involvement percentage (RI) of HO • after 6 h of CAF degradation first decreased from 10% to 7% and then increased to 34% as the initial solution pH increased from 3 to 8. In the meanwhile, the RI percentage of SO4

Effects of Mn2O3 and oxidant dosages

The effect of Mn2O3 concentration on CAF degradation in presence of PS under initial solution pH 5 was then investigated (Figure 5A and Figure S13A). An increase of Mn2O3 dosage from 0.05 to 1.0 g L -1 led to an increase of k'CAF from 5.4 × 10 -4 min -1 to 4.6 × 10 -3 min -1 . Similarly, the removal of CAF by Mn2O3 activated PMS was increased with the increase in Mn2O3 concentration (Figure S13B). This effect can be attributed to the increase of the active surface area of Mn2O3, leading to more active sites for radical precursor activation. It is worth noting that complete removal of CAF was obtained for 1.0 g L -1 Mn2O3 and 5 mM PMS after 2 h, while 8 h reaction time is required for the Mn2O3/PS system. The degradation of CAF (k'CAF) by Mn2O3 (1.0 g L -1 ) with PMS (5 mM) was 3.9 × 10 -2 min -1 , which results in an 8-fold increase compared to that in Mn2O3/PS system (4.6 × 10 -3 min -1 ). Figure 5B shows the increase of the k'CAF from 1.6 × 10 -4 min -1 to 4.7 × 10 -3 min -1 with the increase of PS concentration from 1 to 10 mM. The effects of PMS concentration on CAF degradation in the presence of Mn2O3 (0.4 g L -1 ) were similar, where an increase of k'CAF from 3.5 × 10 -4 min -1 to 8.5 × 10 -3 min -1 with increasing PMS concentration from 1 to 10 mM. These enhanced degradations with the increase of radical precursor concentrations (≤ 10 mM) can be attributed to the promotion of the number of reactive radicals as discussed in section 3.1.

CAF degradation pathway and mineralization

The generated from the decarboxylation of P1 through the attack of radicals and followed by the addition of hydroxyl group was identified.

In pathway 2, the C=C bond of product P1 was attacked by radical species, opening the sixmember ring and leading to the formation of two carbonyl groups (-C=O) on the carbon bonding with nitrogen [48,49]. Then the newly formed carbonyl groups might be further attacked by S5). P2 and P4 were identified only in the presence of PS (see table S5).

The mineralization efficiency of Mn2O3 + PS and PMS systems was monitored by measuring TOC over reaction time. As shown in Figure 7, 34% and 48% of TOC degradation were reached in Mn2O3-activated PS and PMS system in 70 h, while complete removal of CAF was achieved before 8 h of reaction time (Figure S14). However, the oxidation process remains active for a long time, over 70 h, suggesting the long-term efficiency of the investigated process.

Degradation efficiency in sewage treatment plant waters (STPW)

To assess the performance and ability of Mn2O3 + PS and PMS to degrade CAF in the real water body, the degradation was evaluated in STPW (Total Organic Carbon (TOC) = 39.7 mg C L -1 , Inorganic Carbon (IC) = 4.5 mg C L -1 , pH = 8.1) and compared to milli-q water system using 0.4 and 1.0 g L -1 of Mn2O3 (Figure S15). The degradation efficiency of CAF (DE) was determined from the percentage of CAF degraded after 20 hours of reaction and inhibition in STPW was observed in all experiments (Figure 8). Generally, this effect can be attributed to the radical scavenging ability of natural organic matter and inorganic carbon (i.e. hydrogen carbonate ions) of STPW that are recognized to be the main inhibitors of radicals in AOPs [15].

However, in Mn2O3/PS system, the main reactive species are the Mn2O3 surface-activated PS.

The inhibition of CAF degradation in STPW was mainly attributed to the increased ionic strength [50].

The inhibition effect (IE) of STPW was determined as IE = (1 -𝐷𝐸 𝑆𝑇𝑃𝑊 𝐷𝐸 𝑚𝑖𝑙𝑙𝑖-𝑞 ⁄ ) × 100, in which 𝐷𝐸 𝑆𝑇𝑃𝑊 and 𝐷𝐸 𝑚𝑖𝑙𝑙𝑖-𝑞 were the CAF degradation efficiencies in STPW and milliq water after 20 hours. It is interesting to note that, in the presence of PS, about 95 % of inhibition was reported using two different Mn2O3 concentrations. The results verified again the non-radical mechanism in the Mn2O3/PS system and the formation of Mn2O3 surfaceactivated PS. However, a decrease of inhibition effect with the increase of Mn2O3 concentration was observed using PMS. The CAF degradation efficiency drops from 88.6 to 33.6 % (IE of 62%) when 0.4 g L -1 of Mn2O3 was used, while only 20% of inhibition in STPW was determined in the presence of 1.0 g L -1 . Increasing in the production of reactive species is expected when a higher amount of solid (higher exposed surface area) is used. This finding highlights the crucial role of Mn in promoting oxidation reactions and suggests that manganese-mediated advanced oxidation processes (Mn-AOPs) could be a promising process for wastewater treatment.

Conclusion

Mn2O3 exhibited higher activation capability for PS and PMS compared with H2O2 and HSO3 -.

CAF degradation in the Mn2O3-activated PS and PMS was highly influenced by the initial solution pH. PS was efficiently activated by Mn2O3 under acidic pH, while under neutral or alkaline conditions, environmentally relevant pH values, higher degradation was observed using PMS. Chemical quenching and EPR experiments demonstrated that HO • and SO4 •-were the main radicals generated in the system with PMS. In addition, a non-radical mechanism could contribute to the degradation of CAF in the Mn2O3/PS system. The degradation efficiency was tested in real STPW and, despite the inhibition of CAF degradation was observed due to the water constituents (mainly inorganic and organic carbon), the central role of Mn2O3 activation for highly reactive species formation was demonstrated in particular when PMS was used (about 80 % of CAF degradation was obtained in STPW after 20 hours). Comparison of obtained results in pure and real STPW indicates that a lower inhibition effect (20 %) was obtained using PMS compared to the PS system. Given the ability of Mn2O3 to activate PS and PMS, the manganese-based oxidation processes (Mn-AOPs) reactions can be considered as a promising treatment technology for the removal of organic pollutants in water. Text S3. To evaluate the contribution of each radical to the CAF degradation, the radical involvement (RI) was determined as follows: RI (%) = P scavenger -P no scavenger 1-P no scavenger × 100 S1

Recherches en

Where P scavenger and P no scavenger are the percentage of CAF degradation with and without scavengers.

Table S1 Second-order rate constants (𝑘") of scavengers and CAF with different radical species.

Chemicals

𝒌" 𝑯𝑶 Table S3 The initial (pH0h) and final (pH3h) pH values of different processes.

Reaction systems pH0h pH3h

Mn2O3 + CAF Endocrine-disrupting chemicals (EDCs) affect the normal functioning of endogenous hormones [1]. EDCs consist of various types of chemicals such as natural and synthetic hormones, plasticizers, pesticides, and flame retardants [2]. With the development of industry and the pursuit of high-quality life of humans, the EDCs-containing products have been consumed in a large amount. Over the past 10 years, EDCs as a major concern of water quality have received extensive attention [3][4][5]. The regularly detected amount of EDCs in surface and drinking waters are up to ng/L while higher concentrations in municipal wastewater, an important source of anthropogenic pollution release into the environment, were also reported [3,6].

To decrease the concentration of EDCs in wastewater, different processes based on thermal and photochemical degradation were reported [7][8][9]. Advanced oxidation processes (AOPs) are innovative technologies used to improve the degradation of the recalcitrant pollutants based on the formation of highly oxidative species [7,8]. Reactive oxygen and sulfur species (ROS and RSS) with high oxidation potentials such as hydroxyl (HO • ) and sulfate (SO4 •-) radicals are generated through photochemical or electron transfer reaction [10]. Such radicals are capable of inducing hazardous pollutants degradation till complete mineralization in water. Among the different AOPs systems, transition metal-based activation of radical precursors represents a low-energy cost and efficient process [11,12]. Iron (Fe) is one of the most used metal in homogeneous and heterogeneous applications and the redox cycle between Fe(II) and Fe(III) provide a very efficient system for activation of radical precursors and abatement of organic pollutant concentration in waters [13]. However, organic complexes such as aminopolycarboxylic are regularly used to improve the iron stability toward a large range of pH values [14,15].

Recently, other metals such as manganese (Mn) has been used for different water treatment applications. Mn is a promising metal due to its multiple oxidation states (+II, +III, +IV, +VII)

and high oxidative and catalytic reactivity. Most of published reports have focused on manganese dioxide-based AOPs (MnO2-AOPs), to remove pollutants, such as phenol and 2,4dichlorophenol, using MnO2 activated peroxymonosulfate (PMS) or persulfate (PS) in water [16,17]. Huang et al. (2019) investigated the effects of properties of MnO2 such as crystal structure, morphology, and surface Mn oxidation states on PMS activation. They showed that the catalytic activity of MnO2 is positively corrected with the surface Mn(III) content [18].

Other studies have used Mn(III) oxides as a mean to activate oxidants (e.g. PMS and PS) and then remove organic contaminants, such as α-Mn2O3 [19] and MnOOH [20]. He et al.

demonstrated that nanowires MnOOH have higher activity on PMS activation for 2,4dichlorophenol oxidation as compared with multi-branches and nanorods MnOOH [21].

Moreover, Li et al. showed that sulfate and hydroxyl radicals can be generated from manganiteactivated PS under pH 11 [22]. Although the good ability of Mn III -oxides for PS or PMS activation has been investigated the underlying mechanisms of activation and generation of reactive species remain elusive. In addition, most of available data on the MnOOH/PMS activation process has been obtained in ultrapure water and very little is known about the application of this process in real wastewater. Therefore, the ability of γ-MnOOH for PMS activation under real-world conditions for wastewater treatment merits investigation.

In the present study, we investigated the degradation of three EDCs (Bisphenol A (BPA), 17βestradiol (E2) and 17α-ethinylestradiol (EE2)) using γ-MnOOH-activated PMS process. We have combined radical scavenging approaches, EPR and XPS analyses, to determine the different activation and removal mechanisms (i.e. radical and non-radical pathways). Moreover, the influence of key parameters such as the initial concentration of γ-MnOOH and PMS, and the impact of water inorganic constituents were also evaluated. Finally, the efficiency of the degradation system was evaluated in real sewage treatment plant water, collected from the "3 rivières" urban treatment plant in Clermont-Ferrand, France.

Material and Methods

Chemicals and reagents

Bisphenol A (BPA, 99%), 17β-estradiol (E2, 98%), 17α-ethynylestradiol (EE2, 98%), potassium peroxymonosulfate (PMS, purchased in the form of KHSO5 • 0.5KHSO4 • 0.5K2SO4), sodium nitrite (NaNO2, 97%), sodium bicarbonate (NaHCO3, 99.7%), and sodium chloride (NaCl, 99%) were purchased from Sigma-Aldrich Company, France. Sodium nitrate (NaNO3, 99%) was bought from Fluka. Manganese (II) chloride (MnCl2) was purchased from Riedel-de Haen. HPLC grade acetonitrile solvent (ACN) was purchased from CARLO ERBA Reagents.

Sodium hydroxide (NaOH, 99.9%) and perchloric acid (HClO4, 72%) which were used to adjust solution pH were purchased from Sigma-Aldrich. Ethanol (96%, VWR Chemicals), t-Butyl alcohol (TBA), and furfuryl alcohol (FFA) from Sigma-Aldrich were employed as radical and singlet oxygen scavengers. Ultrapure water (Millipore, resistivity 18.2 MΩ cm) was used in the preparation of all solutions. All chemicals were used as received without further purification.

γ-MnOOH synthesis and characterization

Manganite (γ-MnOOH) particles were synthesized as described by Yu and co-workers [23].

The nature of γ-MnOOH was confirmed by X-ray diffraction (XRD) and Transmission Electron Microscopy-Selective Area Electron Diffraction (TEM-SAED) using an accelerating voltage of 200 kV (TEM Jeol JEM 2100 HR) (Figure S1). TEM images of the as-prepared manganite display wire-like morphology with diameters of 20-40 nm, and the SAED pattern confirmed the nature of γ-MnOOH. For XPS analysis, 0.5 g/L of solid were put in contact with 5 mM of PMS with and without BPA (200 µM). After 1 h of reaction, the solution was centrifuged and water was withdrawn and the remaining water completely removed under the N2 stream. Then, the solid particles were collected and analyzed in a UHV chamber (pressure 10 -7 Pa) facilitated with an XPS system (hemispherical electron energy analyzer OMICRON EA125, Germany).

Data analysis, curve fitting and quantification of the Mn XPS spectra were performed using CasaXPS software. The binding energy (BE) scale was referenced to the adventitious C1s peak at 284.8 eV. High-resolution XPS spectra were processed using Shirley background subtraction.

After background subtraction, the spectra were smoothed and fitted by using of Gaussian-Lorentzian peak shape (70:30). For each sample, fitting parameters (FWHM, the full width at half maximum; H, the peak height; E, the peak position center) were adjusted to minimize the value of residual standard deviation (STD).

Real water matrix

The sewage treatment plant water (STPW) was collected from the urban treatment plant named "3 rivières" in Clermont-Ferrand, France in December 2019. Prior to use, the STP water was filtered using a syringe filter (CHROMAFIL ® Xtra RC-45/25) and then use for experiments.

The main physico-chemical parameters relevant to this study are reported in Table S1.

Experimental procedure

PMS was added to the solution containing selected EDCs and 𝛾-MnOOH just before the reaction starts. All reactions were performed in a brown glass bottle (125 mL) at room temperature (293 ± 2 K) and under stirring to ensure homogeneity. Unless otherwise stated, experiments were performed at pH 6.5 ± 0.2. For HPLC, 2 mL of solution were withdrawn at fixed interval times, filtered using 0.20 µM PTFE filter and the reaction was immediately quenched with the addition of 20% methanol (HPLC grade, Sigma-Aldrich). Samples were kept at 277 K and analysis was performed within 4 hours. Furthermore, to identify the generated radicals, EtOH, TBA and FFA were added to the solution before PMS addition and a kinetic approach based on the second-order rate constants and pseudo-first order decay determination was used to assess the radical species involvement.

Analytical methods

BPA, E2 and EE2 quantification during reactions was performed by high performance liquid chromatography (HPLC) equipped with a diode array detector (Waters 2475, USA). The injection volume was 50 µL and a flow rate of 1 mL min -1 with a mobile phase consisting of a mixture of water and ACN (50/50, v/v) were used. An Agilent Eclipse XDB C18 of 250 mm × 4.6 mm × 5 µm column was adapted as the analytic column. Under these conditions, the detection wavelength and retention time were: for BPA (277 nm) at 5.4 min, for E2 (280 nm) at 6.2 min, for EE2 (280 nm) at 7.3 min (Figure S2). To investigate the mineralization, the total organic carbon (TOC) analysis (TOC-L CPH CN200, Shimadzu) was performed.

The pseudo-first order decay of selected pollutant (𝑘′ 𝑝 ) was determined from the slope of ln(C/C0) vs reaction time as following: 𝑙𝑛 𝐶 𝐶 0 = 𝑘′ 𝑝 × 𝑡. Where C0 and C are the initial and remaining concentrations in solution and t is the reaction time.

PMS was quantified using the iodometric titration in which generated I3 -was quantified at 348 nm using a Cary 300 Varian Spectrophotometer [24]. Calibration was performed using different PMS concentrations and data are presented in the supplementary Materials (Figure S3). Mn in 2% nitric acid, Sigma-Aldrich) was used for making the calibration curve of Mn 2+ . The detection wavelength was set at 279.48 nm.

The identification of reactive oxygen species was performed on the electron paramagnetic resonance spectroscopy (EPR, Burker) using 5,5-dimethyl-1-pyrroline N-oxide (DMPO) as the spin trapping agents. Typically, 375 µL of sample was withdrawn and immediately mixed with 125 µL of DMPO (100 mM), then the mixture was transferred into an EPR quartz capillary tube (0.5 mm, 100 mm) for measurement. The EPR spectrometer settings were as follows:

modulation frequency, ~100 kHz; microwave frequency, ~9.87 GHz; microwave power, ~6.87 mW; center field, ~352.7 mT; sweep width, 20 mT; sweep time, 30 s; sample g-factors, 2.0; modulation amplitude, 1.0 G.

Results and discussion

Effect of γ-MnOOH and PMS dosage

Preliminary experiments showed that no significant degradation was observed with PMS or γ-MnOOH, while efficient removal was reached when both γ-MnOOH and PMS were used simultaneously (Figure 1). In this system, BPA was completely removed after 2 h with an initial pseudo-first order degradation rate (k'BPA) of 0.058 ± 0.001 min -1 showing that reactive species (such as SO4 •-and HO • ) may be generated through the reaction between γ-MnOOH and PMS in solution. Moreover, the adsorption of BPA on γ-MnOOH was negligible (< 3 %) after 3 h of reaction. When various concentrations of γ-MnOOH were used in the presence of 0.5 mM PMS (Figure 2A), a different trend can be observed. Increasing the γ-MnOOH concentration up to 0.1 g L -1 , faster degradation of BPA was determined and the corresponding kinetic rate constant of BPA degradation (𝑘′ 𝐵𝑃𝐴 ) increased up to 0.163 ± 0.001 min -1 . At higher γ-MnOOH concentrations, the BPA degradation decreased and almost no degradation was observed using 0.5 g L -1 of γ-MnOOH. Similar trend was observed when different PMS concentrations were used in the presence of 0.1 g L -1 of γ-MnOOH (Figure 2B). From 0 to 2 mM, 𝑘′ 𝐵𝑃𝐴 increased almost linearly to 0.719 ± 0.007 min -1 , while higher PMS concentrations inhibited the degradation.

This behavior suggests that Mn(III)-oxide and PMS would scavenge the generated reactive species when they are used at high concentrations. This effect can be attributed to the quenching of excessive γ-MnOOH and PMS on the active species [25,26]. Thus, an optimal concentration of γ-MnOOH (0.1 g/L) and PMS (2 mM) was adopted for the followed study.

Activation pathway and reactive species formation

In order to verify the generation of reactive species resulted from the activation of γ-MnOOH, the leaching of dissolved Mn 2+ in γ-MnOOH/PMS system was measured, and the homogenous activation of PMS by Mn 2+ was evaluated. The total amount of leached and adsorbed Mn 2+ was determined at around 34 µM after 30 min of reaction. However, a high amount of dissolved Mn 2+ (100 µM) can only remove about 20% of target compound in 1 h reaction (Figure S4).

Therefore, the removal of BPA was mainly attributed to the heterogeneous activation of PMS by γ-MnOOH.

To identify the reactive species generated in the γ-MnOOH/PMS system, chemical scavenging tests and EPR experiments were conducted. Firstly, the generation of reactive oxidant species was investigated using different chemical quenchers. TBA was used as the HO • scavenger due to its higher selectivity toward HO • (𝑘 𝑇𝐵𝐴, 𝐻𝑂 • = 6.0 × 10 8 M -1 s -1 ) compared to SO4 •- (𝑘 𝑇𝐵𝐴, 𝑆𝑂 4 •-= 8.4 × 10 5 M -1 s -1 ) [27]. Theoretically, the addition of 10 mM of TBA can suppress 96% of BPA degradation considering the reaction rate constants of TBA, PMS and BPA with HO • (𝑘 𝑃𝑀𝑆, 𝐻𝑂 • = 1  10 7 M -1 s -1 and 𝑘 𝐵𝑃𝐴, 𝐻𝑂 • = 8.6  10 9 M -1 s -1 ) (Table S2) [28,29].

However, the 𝑘′ 𝐵𝑃𝐴 was not inhibited when 10 mM of TBA were added to the solution, indicating that no HO • was generated in the system (Figure 3). Comparatively, using 10 mM of EtOH can efficiently quench both HO • and SO4 •-radicals (𝑘 𝐸𝑡𝑂𝐻, 𝐻𝑂 • = 1.9 × 10 9 M -1 s -1 and 𝑘 𝐸𝑡𝑂𝐻, 𝑆𝑂 4 •-= 4.3 × 10 7 M -1 s -1 ) [30,31]. The 𝑘′ 𝐵𝑃𝐴 decreases to 40 % of the initial value (0.719 min -1 ) showing that SO4 •-are generated and contributed to the BPA degradation. Further increasing EtOH concentration up to 50 mM, the BPA degradation decreased to 39 %. The generation of reactive radicals were also confirmed by the EPR experiments using DMPO as the complexing agents. Figure S5 shows one clear seven-peaks signal was formed in the γ-MnOOH/PMS system. This multiplet peaks signal is recognized as the 5,5-dimethyl-2-oxopyrroline-1-oxyl (DMPOX) which can be generated by the overoxidizing of DMPO by reactive species such as HSO5 ─ , HO • and SO4 •-radicals. Since no HO • was generated in the reaction,

and HSO5

─ has low reactivity to oxidize DMPO [32], the formation of DMPOX can be reasonably attributed to the reactivity of DMPO with SO4 •-radicals.

Furthermore, under the adopted experimental conditions we can estimate that, considering only SO4 •-involved during the degradation system, quenching of 78.5 and 94.8% of BPA degradation should be expected with the addition of 10 and 50 mM of EtOH (Table S2). However, this difference has not been observed in our experiments, which suggest the occurrence of other reactive species un-quenchable by the organic scavengers. Thus, one non-radical mechanism was assumed to be responsible for the degradation of the rest part of BPA. He et al. reported the generation of singlet oxygen ( 1 O2) in the system of γ-MnOOH/PMS [21]. To verify if 1 O2 was generated, the degradation of BPA was followed in the presence of 1 mM FFA as 1 O2 quencher (𝑘 𝐹𝐹𝐴,1 𝑂 2 = 1.2 × 10 8 M -1 s -1 ) in addition to 50 mM of EtOH [33]. Since FFA has high reactivity with HO • (𝑘 𝐹𝐹𝐴, 𝐻𝑂 • = 1.5 × 10 10 M -1 s -1 ) [34], the simultaneous addition of FFA and EtOH was planned to consume all the possible HO • radicals using EtOH, and then verify the inhibition of BPA degradation caused by FFA addition was only attributed to the generation of 1 O2. As reported in Figure 3, almost no effect was observed with the addition of FFA indicating that 1 O2 was not generated. To further shed light on the non-radical mechanism in the γ-MnOOH/PMS system, more attention had been paid to the oxidation states of Mn species changing during the reaction. Since the direct oxidation of the organic compounds by Mn(IV) through electron transfer process has been widely reported [18,35], we can suppose that higher valence state Mn species (e.g. Mn IV ) can be formed, and thus contribute to the degradation of BPA. In order to get more insights on the redox process involving Mn species, XPS analysis of γ-MnOOH surfaces was performed during different reaction steps. The chemical oxidation state of Mn in the external layers was determined using the Mn 3s multiplet splitting method and the position and shape of the Mn 2p3/2 spectra (Figure 4) [36,37]. The peak position and relevant parameters from curve fitting of Mn 3s and Mn 2p3/2 high resolution spectra are listed in Table 1. The average oxidation state (AOS) of samples was calculated from the % of each Mn species from the fitting of Mn 2p3/2 spectra and from the binding energy difference of Mn 3s doublet splitting as following: AOS = 8.956 -1.126 × ∆E [25,26]. Using the multiplets fitting of the Mn 2p3/2 peaks, the percent contents of Mn(IV), Mn(III) and Mn(II) in the γ-MnOOH sample were 7%, 85% and 8%, respectively and AOS was 2.99. After reaction with PMS, the relative contents of Mn(IV), Mn(III) and Mn(II) changed and Mn(IV) percentage increased to 24.5 consistently with an increase of AOS up to 3.10. In the sample γ-MnOOH + PMS + BPA, Mn(II) percentage increased to 37% while Mn(IV) decreased to 9% with a decrease of AOS down to 2.72. AOS values were consistent with those determined from the fitting of Mn 3s multiplet splitting data (Table 1). The increase of Mn(IV) in the presence of PMS followed by a decrease when BPA was present in solution, suggests that freshly generated Mn(IV) may contribute to the degradation of BPA. Such a process, called metal-based oxidation, was previously indicated as relevant surface-mediated oxidation of organic pollutants in the presence of Mn(IV) oxides [38].

Huang et al. demonstrated the efficient oxidation of BPA by MnO2 with different structures such as α-, β-, γ-, δ-MnO2 with and without PMS [18,35]. Wang et al. reported the removal of BPA was due to a complex between amorphous MnO2 and PMS [39]. Collectively, quenching experiments and XPS analysis suggest that the first electron transfer reaction occurring between ≡Mn(III) and PMS on the surface of oxide led to the formation of ≡Mn(IV) and SO4 •-(R1). SO4 •-can oxidize BPA (R4) (Table S2) while Mn(IV) can further react with PMS leading to the formation of ≡Mn(III) and SO5 •-(R2) or directly with BPA (R3). The involvement of SO5 •-during the BPA degradation (R5) is reasonably considered less relevant due to the typical second-order rate constants between SO5 •-and organic molecules (10 4 -10 6 M -1 s -1 ) [40].

≡ 𝑀𝑛(𝐼𝐼𝐼) + 𝐻𝑆𝑂 5 -→ ≡ 𝑀𝑛(𝐼𝑉) + 𝐻𝑂 -+ 𝑆𝑂 4

•-

(R1)

≡ 𝑀𝑛(𝐼𝑉) + 𝐻𝑆𝑂 5 -→ ≡ 𝑀𝑛(𝐼𝐼𝐼) + 𝑆𝑂

Degradation and mineralization of BPA, E2 and EE2

Degradation and mineralization efficiencies of γ-MnOOH/PMS system were tested on BPA, E2 and EE2 mixing solution at pH 6.5. In Figure 5, the degradation of three EDCs (5 µM each)

and PMS (2 mM) consumption were followed during the reaction time in the presence of γ-MnOOH (0.1 g L -1 ). Despite no degradation in the presence of γ-MnOOH alone and less than 10 % after 2 hours in the presence of PMS alone (Figure S6), almost complete degradation of three pollutants can be observed after 30 min. Degradation of BPA, E2 and EE2 follows a monoexponential decay and after 30 min of reaction, ~ 100 µM of PMS were consumed. We can then estimate an efficiency of degradation system considering the stoichiometric efficiency value (η) that is calculated as the ratio between degraded EDCs vs consumed PMS (η =

∆[𝐸𝐷𝐶𝑠] ∆[𝑃𝑀𝑆]
). In the presence of 0.1 g L -1 γ-MnOOH and 2 mM of PMS a value of η = 15% was determined at pH 6.5.

Effect of inorganic ions

Inorganic anions are commonly detected in natural waters and they could affect the radicalmediated degradation efficiency of organic pollutants. In the γ-MnOOH / PMS system, the effect of chloride (Cl -), bicarbonate (HCO3 -), nitrite (NO2 -) and nitrate (NO3 -) anions was investigated and BPA, E2, and EE2 degradations were followed using 0.1 g L -1 of γ-MnOOH and 2 mM of PMS.

As illustrated in Figure 7A, increasing the concentration of Cl -from 0 to 5 mM, the BPA, E2, and EE2 degradation rates slightly increased indicating a positive effect of Cl -on the oxidation system. The results are in agreement with previous studies which indicated that improvement of organic pollutants degradation efficiency in the presence of Cl -was attributed to the generation of reactive chlorine species such as Cl • and HOCl (R7 to R9) [42]. In fact, despite a different reactivity constant between Cl • and Cl 2 •-with organic compounds (𝑘 𝐶𝑙 • = 10 9 -10 10 M -1 s -1 , 𝑘 𝐶𝑙 2 •-= 10 7 -10 8 M -1 s -1 ), their lower reactivity was partially counterbalanced by direct activation occurring in the presence of PMS leading to the formation of oxidant species such as HOCl as observed at high Cl -concentration in absence of γ-MnOOH (R9) (Figure S7) [43]. Bicarbonate ions (HCO3 -) are generally considered as interfering species during oxidative processes based on radical species generation due to their reactivity with radicals and generation of less oxidative species such as CO 3

•- [30,31]. However, a negligible impact on the EDCs degradation was observed with the addition of HCO3 -(Figure 7B). Such effect can be explained considering that the inhibition effect on the sulfate radical reactivity (R10-R11) is counterbalanced by the direct activation of HCO3 -by PMS leading to the formation of HCO4 - as reported in Figure S7 (R12) [42].

Efficiency in the real water matrix

To evaluate the efficiency of γ-MnOOH/PMS system in realistic conditions, the treatment of three EDCs in filtered sewage treatment plant water (STPW) was conducted. The initial TOC was determined to be 10.8 mgC L -1 (in which 2.8 mgC L -1 come from the addition of 3 EDCs).

As expected, the removal rates and efficiencies of EDCs and TOC in STP water were slower in comparison with results in milli-Q water (Figure 8). This effect can be attributed to the organic

Conclusion

In this work, the degradation of three estrogens: BPA, E2 and EE2 was performed using manganite (γ-MnOOH) in the presence of peroxymonosulfate ions (PMS). The rapid degradation in water was attributed to the generation of sulfate radical and the direct oxidation mediated by Mn(IV). In fact, about 60% of estrogens degradation was attributed to the generation of sulfate radical, while ~ 40 % to the direct oxidation mediated by Mn(IV). The effect of inorganic ions at environmentally relevant concentrations showed only limited inhibition effect on EDCs degradation in ultrapure water and real sewage treatment plant water, indicating that the γ-MnOOH/PMS system can be applied for real wastewater treatment. The Mn(III)/(IV) cycle represents the main reaction step leading to the formation of sulfate radical but also mediated Mn(IV) species have high oxidation ability toward the removal of selected pollutants. These results have strong implications for the development of Mn-mediated oxidation processes in wastewater remediation technologies.

Table S2. Radical reactivity (%) of each chemical species determined from the second order rate constants from literature [1][2][3] and initial concentrations.

Radical species

Concentrations Second order rate constants ( M -1 s -1 ) systems and associated impacts on the redox-active minerals should be considered in contaminant fate and transport assessment.

Calculated

In the second part of this dissertation, the removal of caffeine (CAF) in water using PDS or PMS in the presence of Mn2O3 was investigated. Mn2O3 exhibited higher activation capability for PDS and PMS compared with H2O2 and HSO3 -. Moreover, the degradation of CAF in the Mn2O3-activated PDS or PMS system was highly influenced by the initial solution pH. PDS was efficiently activated by Mn2O3 under acidic pH, while under neutral or alkaline conditions, environmentally relevant pH values, higher degradation was observed using PMS. Chemical quenching and EPR experiments demonstrated that HO • and SO4 •-were the main radicals generated in the system with PMS. In addition, a non-radical mechanism could contribute to the degradation of CAF in the Mn2O3/PDS system. The degradation efficiency was tested in the real STP water, and despite the inhibition of CAF degradation was observed due to the water constituents (mainly inorganic and organic carbon), the central role of Mn2O3 activation for highly reactive species formation was demonstrated in particular when PMS was used (about 80 % of CAF degradation was obtained in STP water after 20 hours). Comparison of obtained results in pure and real STP waters indicates that a lower inhibition effect (20 %) was obtained using PMS compared to the PDS system. Given the ability of Mn2O3 to activate PDS and PMS, the manganese-based oxidation processes (Mn-AOPs) reactions can be considered as a promising treatment technology for the removal of organic pollutants in water.

In the third part of this dissertation, the degradation of three estrogens: BPA, E2 and EE2 was performed using manganite (γ-MnOOH) in the presence of peroxymonosulfate ions (PMS).

The rapid degradation in water was attributed to the generation of sulfate radical and the direct oxidation mediated by Mn(IV). In fact, about 60% of estrogens degradation was attributed to the generation of sulfate radical, while ~ 40 % to the direct oxidation mediated by Mn(IV). The effect of inorganic ions at environmentally relevant concentrations showed only limited inhibition effect on EDCs degradation in ultrapure water and real sewage treatment plant water, indicating that the γ-MnOOH/PMS system can be applied for real wastewater treatment. The Mn(III)/(IV) cycle represents the main reaction step leading to the formation of sulfate radical but also mediated Mn(IV) species which have high oxidation ability toward the removal of selected pollutants. These results have strong implications for the development of Mn-mediated oxidation processes in wastewater remediation technologies.

Given the fruitful results, some future outlooks are proposed.

(1) Introducing the ultraviolet light (UV) into the homogeneous Mn-mediated AOPs. For example, Mn(II) can activate radical precursors to generate radicals and high-valent Mn species such as Mn(V). The presence of UV light can assist the recycle of Mn from Mn(V) to Mn(II), which resulting in the improvement of degradation efficiency of organic pollutants.

(2) Synthesizing the Mn containing-composites. The synthesis of bimetallic Fe-Mn oxides has been reported. However, the activation mechanism of radical precursors by Fe-Mn oxides was not clear. Moreover, in addition to the synthesizing of Fe-Mn oxides, there are many other transition metals (such as Co, Cu) that can be used to produce composites with Mn. Furthermore, the synthesis of Mn containing photocatalytic materials also can be one good option for the degradation of organic pollutants in water.

(3) Introducing chelators into homogeneous and heterogeneous Mn-mediated AOPs. Mn(III)

plays an important role in the oxidative and catalytic reactivity of MnO2. However, the stability of Mn(III) in solution is weak. The presence of chelators can complex Mn(III) species leading to the increase of pollutant degradation efficiency.

(4) The application of Mn-mediated AOPs is not restrict to the removal of organic pollutants in wastewater. The system can be used for the removal of recalcitrant contaminants in groundwater and soil as well. Moreover, the combining of Mn-mediated AOPs with other water treatment technologies (such as membrane technology) will provide new solutions for the elimination of pollutants in sewage and remediation of groundwater and soil.

Figure 1 - 1 .

 11 Figure 1-1. The chemical structures of H2O2 (a), PDS (b), and PMS (c).

Figure 1 - 2 .

 12 Figure 1-2. The crystal structures of (A) pyrolusite, (B) ramsdellite, (C) hollandite, (D) birnessite, stacked layers of MnO6 are blue and the yellow balls represent the disordered H2O or Na + [64].

Figure 1 - 3 .

 13 Figure 1-3. The crystal structures of (A) γ-MnOOH, (B) α-Mn2O3 and (C) Mn3O4. The red, blue and white balls represent oxygen, manganese and hydrogen atoms, respectively. The black dashed lines represent the single unit cell. The crystalline parameters of Mn(III) (oxyhydr)oxides were taken from the crystallography open database (COD), and the COD ID of α-Mn2O3, γ-MnOOH, and Mn3O4 are 2105791, 1011012, and 1514121, separately.

Figure 1 - 4 .

 14 Figure 1-4. The SEM images of the α-, β-, and δ-MnO2 [82].

Figure 1 - 5 .

 15 Figure 1-5. The SEM images of the (a) Mn2O3-O, (b) Mn2O3-T, and (c) Mn2O3-C [83].

4 .

 4 For instance, Saputra et al. investigated the effect of morphology on the oxidation of phenol by Mn2O3 activated PMS. The results showed that cubic-Mn2O3 has the highest reactivity on PMS activation in comparison with octahedral-and truncated octahedral-Mn2O3 and it was due to the high surface area and distinct surface atoms arrangement of cubic-Mn2O3 [86]. Similarly, Cheng et al. successfully prepared three α-Mn2O3 in cubic-, truncated

Figure 1 - 6 .

 16 Figure 1-6. The picture of PDS activation on the surface of γ-MnOOH. The red, blue, and white balls represent oxygen, manganese, and hydrogen atoms, respectively [116].

  observed for the γ-MnOOH-based composites. This might be attributed to the distinct activation way of PDS by α-Mn2O3 or γ-MnOOH compared with Mn3O4.In summary, Mn3O4 can activate PDS to generate SO 4•-through radical mechanisms, while the activation of PDS by α-Mn2O3 and γ-MnOOH is processed in a non-radical mechanism with the generation of 1 O2 and catalyst surface-activated PDS substances. For the activation of PDS by Mn(III) (oxyhydr)oxides composites, the Mn3O4-based composites have shown good catalytic performance in PDS activation for pollutants degradation. In comparison, the activation of PDS by modified α-Mn2O3 or γ-MnOOH catalysts needs to be further investigated.

Figure 1 - 7 .

 17 Figure 1-7. The diagram of PDS activation by Mn3O4/C composite towards 2,4-D degradation [122].

  of furfural removal in 60 min SO4

CO 3 2 -and HCO 3 -

 23 to form active radicals and HCO 4 -(Eqs. (1-34)-). Furthermore, the solution pH can be changed in the presence of carbonate/bicarbonate ions, which can affect the reactivity of Mn(III) (oxyhydr)oxides in PMS/PDS activation as discussed in section 1.3.4.1.

  activation. For instance, Mitchell et al. reported the efficient degradation of perfluorooctanoic acid (PFOA) in the 𝛾-MnO2 activated H2O2 system at neutral pH, and the dominate reactive species were identified as O 2 •-and HO 2 -[162]. Furman et al. reported that 𝛾-MnO2-catalyzed decomposition of H2O2 promoted the degradation of the superoxide probe hexachloroethane (HCA), and the electron spin resonance spectroscopy (ESR) confirmed that superoxide was the dominant reactive species generated in 𝛾-MnO2/H2O2 system.

- 8 .

 8 In summary of these studies, it can get one clear conclusion that the crystal structure of MnO2 plays a vital role in the activation of H2O2. Specifically, the crystal MnO2 has the good ability in catalyzing H2O2, while the disordered MnO2 is a good oxidant in H2O2 oxidation. This conclusion was also verified by other studies. For instance, Huang et al. reported that in comparison to 𝛼-, β-, 𝛾-, λ-MnO2, 𝛿-MnO2 showed the highest efficiency in bisphenol A oxidation[START_REF] Huang | Effect of MnO2 Phase Structure on the Oxidative Reactivity toward Bisphenol A Degradation[END_REF].

Figure 1 - 8 .

 18 Figure 1-8. The schematic diagram of the reaction of H2O2 with MnO2. (a) disproportionation of H2O2 in which MnO2 acts as a catalyst and (b) oxidation H2O2 in which MnO2 acts as a sacrificial oxidant that is concomitantly reduced to Mn 2+ .

  For example, Zhao et al. investigated the degradation of 2,4-dichlorophenol in an aqueous solution by α-MnO2 nanowire-activated PDS. The results of quenching tests and ESR experiments showed that the SO 4 •-and HO • were responsible for the removal of 2,4dichlorophenol in the system of α-MnO2/PDS. The activation mechanism is as shown in Eqs. (1-45)-(1-46) and Eqs. (1-17)-(1-18). Zhu et al. reported the efficient degradation of phenol by

Figure 1 - 9 .

 19 Figure 1-9. Different exposure routes of humans to EDCs [191].

× 100 ,

 100 High-Performance Liquid Chromatography system (HPLC, Alliance 2695 Waters) equipped with Diode Array Detector (DAD) and multi 𝜆 Fluorescence Detector (FD). An Agilent Eclipse XDB C18 column (250 × 4.6 mm, 5 μm particle size) was adapted to analyze CAF, E2, EE2 and BPA(using FD detector). Another analytic column Macherey Nagel Nucleodur 100-5 C18 column (150 × 4.6 mm, 5 μm particle size) was used to analyze BPA (using DAD detector).The temperatures of column and sample were set at 30 and 15 °C, respectively. The injection volume was set at 50 µL. Isocratic and gradient elution modes with different mobile phase compositions were used for the chromatographic separation. The parameters of flow rate and mobile phase compositions are presented in Table2-4. For the measurement of E2, EE2 and BPA using the fluorescence detector, the excitation and emission wavelengths were set at 280 and 305 nm, respectively. The calibration curves of target compounds with different concentration were presented in Figure2-1. According to the achieved calibration curves, the concentration of remaining pollutants in different reaction time can be obtained, and the corresponding degradation efficiency (DE) can be calculated using the equation of 𝐶 0 -𝐶 𝑡 𝐶 0 where C 0 and C t represent the initial and the residual concentration of the pollutants at time t, respectively.

Figure 2 - 1 .

 21 Figure 2-1. The calibration curves of target compounds. (A) BPA; (B) CAF; (C) E2; (D) EE2;

4 -

 4 hydroxyphenyl acetic acid dimer in the presence of peroxidase (Figure 2-2) [6]. 4hydroxyphenylacetic acid dimer is highly fluorescence compound with a maximum of excitation wavelength at 320 nm and emission at 420 nm. The calibration curve of H2O2 was obtained using different concentrations of H2O2 (Figure 2-3).

Figure 2 - 2 .

 22 Figure 2-2. Enzyme peroxidase catalyzing reduction of H2O2.

Figure 2 - 3 .

 23 Figure 2-3. Calibration curve of H2O2 quantifications.

I 3 -

 3 can be quantified using a UV-visible spectrophotometer. Specifically, 100 µL of sample containing PDS or PMS were transferred into 5 mL of a solution containing KI (0.6 M) and HCO 3 -(0.06 M). The absorption of solution at 348 nm was measured after 10 min of reaction and PDS/PMS concentration (CPDS/PMS, (mM)) was determined as 𝐶 𝑃𝐷𝑆/𝑃𝑀𝑆 = 𝑉 𝑡𝑜𝑡 𝑉 𝑃𝐷𝑆/𝑃𝑀𝑆 × 𝐴𝑏𝑠 348 𝑛𝑚 𝑆 . Where VPDS/PMS and Vtot are the volume of sample (100 µL) and total volume (5.1 mL), S was the slope of the linear correlation between the PDS/PMS concentration and value of absorption at 348 nm. In this study, SPDS and SPMS were 0.56 and 0.50, respectively. Calibration curves of PDS and PMS were performed using different PDS and PMS concentrations as shown in Figure 2-4.

Figure 2 - 4 .

 24 Figure 2-4. Calibration curve of PDS and PMS concentrations.

Figure 2 - 5 .

 25 Figure 2-5. Reaction between TA and hydroxyl radicals.

5, 5 -

 5 dimethyl-1-pyrroline N-oxide (DMPO) or 2,2,6,6-tetramethyl-4-piperidinol (TEMP) as the spin trapping agents. DMPO is commonly used as the spin traps for sulfate and hydroxyl radicals. The typical hyperfine splitting constants of DMPO-HO • are aN = aH = 14.9 G. For DMPO-SO 4 •-, the hyperfine splitting constants are aN = 13.2 G, aH = 9.6 G, aH = 1.48 G, and aH = 0.78 G [9]. TEMP is a commonly used spin trap for singlet oxygen. The characteristic peaks of TEMP-1 O2 are the triplet signals (1:1:1) [9]. The procedure operated in the measurement were as follow, 375 µL of sample was withdrawn and immediately mixed with 125 µL of DMPO (100 mM) or TEMP (100 mM), then the mixture was transferred into an EPR quartz capillary tube (0.5 mm, 100 mm) for measurement. The EPR spectrometer was set as following: modulation frequency, ~100 kHz; microwave frequency, ~9.87 GHz; microwave power, ~6.87 mW; center field, ~352.7 mT; sweep width, 20 mT; sweep time, 30 s; sample g-factors, 2.0; modulation amplitude, 1.0 G.

Figure 2 - 6 .

 26 Figure 2-6. Calibration curve of Mn 2+ concentrations.

  was then back titrated with standardized 0.02 M KMnO4 solution at 75-85 C until the color of solution turned into reddish and lasted for 30 seconds without fading, to obtain the oxidation number of Mn (nMn) (see Eqs. (2-1) and (2-2)). Then, 0.1 g of the acid birnessite sample was dissolved in 20 mL of 0.25 M hydroxylamine hydrochloride and was diluted to 2 L with ultrapure water.

Figure 2 - 7 .

 27 Figure 2-7. Calibration curves of total carbon and inorganic carbon.

Figure 1 .

 1 Figure 1. Effect of H2O2 concentration on the BPA removal at pH 6.5 at room temperature. Experimental conditions: [AB] = 345 µM, [BPA] = 25 µM. H2O2/MnO2 ratio = 0 -0.6. The relative experimental error lied at 5 % for BPA.

Figure 2 .

 2 Figure 2. Removal rate constants (h -1 ) of BPA as a function of H2O2 dose at three pH values (4, 6.5 and 8). Experimental conditions: [AB] = 345 µM, [BPA] = 25 µM, [H2O2] = 0 -200 µM. H2O2/MnO2 ratio = 0 -0.6, room temperature. The relative experimental error lied at 5 %.

Figure 3 .

 3 Figure 3. Mn(II) formation under different H2O2 concentrations and pH 4.0, 6.5 and 8.0. Experimental conditions: [AB] = 345 µM and [BPA] = 25 µM. H2O2/MnO2 ratio = 0 -0.6, room temperature. The relative experimental error lied at 7 % for Mn(II).

Figure 4 .

 4 Figure 4. Removal rate constants (h -1 ) of BPA as a function of Mn(II) concentration at three pH values (4, 6.5 and 8). Experimental conditions: [AB] = 345 µM, [BPA] = 25 µM, [Mn(II)] = 0 -100 µM, room temperature. The relative experimental error lied at 5 %.

Figure 5 .

 5 Figure 5. XRD patterns of birnessite before and after reaction in the presence of H2O2 or Mn(Ⅱ) (two pH values, 24 h of reaction time), and of Mn(III) rich Birnessite. The black dashed lines indicate the peaks of acid birnessite, while the red dashed lines indicate broad bands appeared in reacted samples. Experimental conditions: [AB] = 345 µM and [BPA] = 25 µM. H2O2/MnO2 ratio = 0.6. Mn(II)/MnO2 ratio = 0.3. The solid was characterized using X-ray powder diffraction (XRD) with a Bruker AXS D8 Advance diffractometer (θ-2θ Bragg-Brentano geometry) using monochromated Cu Kα1 (1.54Å) radiation.

Figure 6 .

 6 Figure 6. Removal rate constants (h -1 ) of BPA as a function of H2O2 dose in absence or presence of silicate (Si), nitrate (N) or phosphate (P): Experimental conditions: pH 6.5, [AB] = 345 µM, [BPA] = 25 µM; [Na2SiO3] = 100 μM; [NaNO3] = 100 μM; [NaH2PO4] = 100 μM. H2O2/MnO2 ratio = 0 -0.3, room temperature. The relative experimental error lied at 5 %.

  to those encountered in natural waters further decreased the reactivity of MnO2 in presence of H2O2. 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 H2O2 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 H2O2 and co-exiting anions on the global carbon cycle. Furthermore, the widespread presence of H2O2 in surface and ground water systems and associated impacts on the redox-active minerals should be considered in contaminant fate and transport assessment. µM) in the 100mL ultrapure water. At different time interval, 2 mL of sample was withdrawn and detected by a fluorescence spectrophotometer (Varian Cary Eclipse, USA). The limit of detection for hydroxyl radicals using fluorescence method was in the low nano-molar range.

Figure S1 .

 S1 Figure S1. SEM (a) and TEM (b) images of the synthesized acid birnessite.

Figure S2 .

 S2 Figure S2. H2O2 decomposition versus time with two H2O2 concentrations and linear regression of ln [H2O2] / [ H2O2]o versus time. Experimental conditions: [AB] = 345 µM, [H2O2] = 100 and 200 µM, pH = 6.5. H2O2/MnO2 ratio = 0.3 and 0.6.

Figure S3 .

 S3 Figure S3. H2O2 decomposition vs time at pH 4, 6.5 and 8, and linear regression of ln [H2O2] /[ H2O2]o versus time. Experimental conditions: [H2O2] = 200 µM; [AB] = 345 µM; [BPA] = 25 µM. H2O2/MnO2 ratio = 0.6.

Figure S4 .Figure S5 . 1 = 9 .

 S4S519 Figure S4. Linear regression of ln [BPA] /[BPA]o versus time. Experimental conditions: pH 6.5, [AB] = 345 µM, [BPA] = 25 µM. H2O2/MnO2 ratio 0 and 0.03.

Figure S6 .

 S6 Figure S6. Effect of pre-equilibration H2O2/MnO2 on the BPA removal. Experimental conditions: [AB] = 345 µM, [BPA] = 25 µM.

Figure S7 .

 S7 Figure S7. Mn(II)-birnessite sorption isotherms at three pH values. Experimental conditions: [AB] = 345 μM; T= 20 ± 1°C; reaction time = 48 h; aerobic conditions.

4 Figure S8 .Figure S9 .

 4S8S9 Figure S8. (a)Adsorption of Si or P at pH 6.5 in presence or absence of dissolved Mn(II). Chemical speciation of (b) silicates Si (Na2SiO3) and (c) phosphates P (NaH2PO4) versus pH. pKas of silicate (pKa,1 = 9.82 and pKa,2 = 13.45) and phosphate (pKa,1 = 2.12, pKa,2 = 7.21 and pKa,3 = 12.31) at infinite dilution were obtained from conditional pKa values and the Davies equation. Experimental conditions: [AB] = 345 μM; [Mn(Ⅱ)] = 100 μM; [Na2SiO3] = 100 μM; [NaH2PO4] = 100 μM; reaction time = 24 h.

  instance, Zhou et al. indicated the efficient degradation of 4-nitrophenol by α-and δ-MnO2 activated PMS through the generation of SO4 •-and HO • [12]. Zhu et al. demonstrated the rapid removal of bisphenol A (BPA) in the β-MnO2/PS system due to the formation of singlet oxygen

  pyrroline N-oxide (DMPO, Sigma-Aldrich) and 2,2,6,6-tetramethyl-4-piperidinol (TEMP, Sigma-Aldrich) as the spin trapping agents. Typically, 375 µL of sample was withdrawn and immediately mixed with 125 µL of DMPO (100 mM) or TEMP (100 mM), then the mixture was transferred into an EPR quartz capillary tube (0.5 mm, 100 mm) for measurement. The EPR spectrometer settings were as follows: microwave frequency, ~ 9.87 GHz; microwave power, ~6.87 mW; center field, ~352.7 mT; sweep width, 20 mT; sweep time, 30 s; sample gfactors, 2.0; modulation amplitude, 1.0 G. The experimental EPR spectra were analyzed by the WinEPR software (Bruker) and the calculations of spin-Hamiltonian parameters and relative concentrations of individual DMPO-adducts were performed with the EasySpin toolbox working on MatLab  platform [28].

Figure 1 .

 1 Figure 1. CAF (25 µM) degradation in the presence of 0.4 g L -1 of Mn2O3 and different oxidants at 5 mM, natural pH (see TableS3), and T = 298 K.

Figure 2 .

 2 Figure 2. Determination of ROS species for CAF oxidation in Mn2O3/PS. A: CAF degradation by Mn2O3 with PS in the presence of different radical scavengers ([TBA]0 = [2-Pr]0 = 10 mM, [FFA]0 = 1 mM). B: PS decomposition in different processes. [Mn2O3]0 = 0.4 g L -1 , [CAF]0 =

Figure 3 .

 3 Figure 3. Determination of reactive radical species for CAF degradation in Mn2O3/PMS system. A: Quenching tests of CAF degradation in the process of Mn2O3 + PMS with various radical scavengers ([TBA]0 = [2-Pr]0 = 10 mM). B: The normalized experimental (black line) and simulated (red line) EPR spectra measured in the suspension of PMS/Mn2O3/DMPO/water ([DMPO]0 = 25 mM). [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PMS]0 = 5 mM, the pH of solution was 5.0 and T = 298 K.

  first increased from 55% to 63% and then decreased to 50%. The results fully and directly demonstrated the transformation between HO • and SO4 •-with the change of solution pH. In summary, neutral or basic conditions are suitable for the activation of PMS by Mn2O3, while PS activation by Mn2O3 was enhanced in acidic conditions only.

Figure 4 .

 4 Figure 4. Degradation constant of CAF (k'CAF) in the systems Mn2O3 + PS and Mn2O3 + PMS under different pH values. Experimental conditions are [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PS]0 = [PMS]0 = 5 mM, T = 298 K.

Figure 5 .

 5 Figure 5. Degradation of CAF (k'CAF) in the system Mn2O3 with PS and PMS. A: Effect of Mn2O3 concentration in the presence of 5 mM PS and PMS. B: Effect of PS and PMS concentrations using 0.4 g L -1 of Mn2O3. Initial pH = 5.0, T = 298 K.

  main transformation intermediates of CAF in Mn2O3-activated PS and PMS systems were identified by LC-MS analysis. The experimental conditions were: [Mn2O3]0 = 0.4 g L -1 , [CAF] = 25 µM and [PS] = [PMS]0 = 5 mM, initial solution pH = 5.0 and analysis was performed after 2 hours of reaction. The molecular formulas, structures, retention time, and positive/negative molecular ions (m/z) of intermediates are summarized in Table S5, and the possible degradation pathways of CAF are shown in Figure 6. CAF was identified at [M+H] + = 195.08 m/z. The first transformation step was the hydrogen abstraction (C8 position) and then water molecule addition generating an intermediate with a keto-enol equilibrium. Further oxidation steps lead to the formation of product P1 at [M-H] -= 228.19 m/z. Similar transformation processes in CAF oxidation were reported in the previous study [46,47]. Subsequently (pathway 1), the intermediate P2 ([M+H] + = 202.15 m/z, [M+H-H2O] + = 184.13 m/z, and [M-H] -= 200.13 m/z)

  radicals and the C-N broke leading to the loss of nitrogen derivative (-N-CH3-COOH) and formation of product P3 ( [M+H] + = 188.13 m/z, [M+H-H2O] + = 170.12 m/z and [M-H] -= 186.11 m/z) [48]. The intermediate P4 at [M+H] + = 174.11 m/z and [M-H] -= 172.09 m/z were further identified, which was originated from P3 through the demethylation process. It is worth noting that not all the intermediates (P1-P4) were observed in the Mn2O3/PMS system and this can be explained by the different types of active radicals generated in Mn2O3 activated PS and PMS systems (see Table

Figure 6 .

 6 Figure 6. Proposed degradation pathways of CAF in Mn2O3-activated PS and PMS systems.

Figure 7 .

 7 Figure 7. TOC decrease with time in Mn2O3-activated PS and PMS systems. Experimental conditions: [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PS]0 = [PMS]0 = 5 mM, initial solution pH = 5.0, T = 25 ± 2 °C.

Figure 8 .

 8 Figure 8. The degradation efficiency of CAF after 20 hours of reaction in milli-q and STP waters (STPW) using 0.4 and 1.0 g L -1 of Mn2O3. PS and PMS concentrations were 5 mM and the pH of solutions was 8.0 ± 0.2.

  Environnement" through the CPER "Environment" founded by the Region Auvergne, the French government and the FEDER (European Community), PRC program CNRS/NSFC n°270437 and CAP 20-25 I-site project. The authors gratefully acknowledge financial support from China Scholarship Council provided to Daqing Jia (No, 201806920034) to study at the University Clermont Auvergne in Clermont-Ferrand, France. Vlasta Brezová is acknowledged for useful discussion on EPR spectra interpretation.

  b [CAF]0 = 25 µM; c [TBA]0 = d [2-Pr]0 = 10 mM; e [FFA]0 = 1 mM;

Figure S1 .

 S1 Figure S1. Powder XRD pattern of Mn2O3 used in the present study.

Figure S2 .

 S2 Figure S2. The changes in the relative concentration of CAF were monitored in the presence of 5 mM of H2O2, HSO3 -, PS, and PMS with the natural pH of 6.8, 4.0, 5.6, 2.7, respectively.

Figure S3 .

 S3 Figure S3. The changes in the relative concentration of CAF were monitored in the presence of Mn2O3 (0.4 g L -1 ) under different initial pH values.

Figure S4 .

 S4 Figure S4. The leaching of Mn 2+ in Mn2O3 (0.4 g L -1 ) activated PS/PMS (5 mM) reactions (A) and the homogeneous activation of PS/PMS by Mn 2+ (B). The pH of the solution was 5.0 and T = 25 ± 2 °C.

Figure S5 .

 S5 Figure S5. Kinetic values of CAF degradation by Mn2O3 with PS in the presence of different scavengers. Experimental conditions: [TBA]0 = [2-Pr]0 = 10 mM, [FFA]0 = 1 mM, [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PS]0 = 5 mM, The pH of solution was 5.0 and T = 25 ± 2 °C.

Figure S6 .

 S6 Figure S6. ERP spectra of different systems in the presence of TEMP. Experimental conditions: [Mn2O3]0 = 0.4 g L -1 , [PS]0 = 5 mM, [TEMP]0 = 25 mM.

Figure S7 .

 S7 Figure S7. CAF degradation in the process of Mn2O3/PS with the addition of 2-Pr and FFA simultaneously. Experimental conditions: [2-Pr]0 = 10 mM, [FFA]0 = 1 mM, [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PS]0 = 5 mM, The pH of solution was 5.0 and T = 25 ± 2 °C.

Figure S8 .

 S8 Figure S8. PS (5 mM) decomposition in the presence of 1 mM of FFA at solution pH 5.0.

Figure S9 .

 S9 Figure S9. Fitting of XPS Mn 3s spectra of samples. (A) Mn2O3; (B) Mn2O3 with PS; (C) Mn2O3 with PMS. Experimental conditions: [Mn2O3]0 = 0.4 g L -1 , [PS]0 = [PMS]0 = 5 mM, pH0 = 5.

Figure S10 .

 S10 Figure S10. Kinetic values of CAF degradation by Mn2O3 with PMS in the presence of various scavengers. Experimental conditions: [TBA]0 = [2-Pr]0 = 10 mM, [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PMS]0 = 5 mM, The pH of solution was 5.0 and T = 25 ± 2 °C.

Figure S11 .

 S11 Figure S11. (A) the effect of FFA (1mM) on CAF degradation in Mn2O3/PMS system. (B) EPR signals of different systems in the presence of TEMP (25 mM). (C) PMS decomposition in the process of Mn2O3/PMS. Experimental conditions: [Mn2O3]0 = 0.4 g L -1 , [CAF]0 = 25 µM, [PMS]0 = 5 mM, The pH of solution was 5.0 and T = 25 ± 2 °C.

Figure S12 .

 S12 Figure S12. The changes of the RI percentages of HO • and SO4 •-radicals in Mn2O3 (0.4 g L -1 ) /PMS (5 mM) process under different initial pH values.

Figure S13 .

 S13 Figure S13. Effect of Mn2O3 loading on CAF degradation in the presence of 5 mM PS (A) and PMS (B). Experimental conditions: [Mn2O3] = 0.05 -1.0 g L -1 , [CAF]0 = 25 µM, [PS]0 = [PMS]0 = 5 mM, initial pH = 5.0, T = 25 ± 2 °C.

Figure S14 .

 S14 Figure S14. Effect of radical precursor concentrations on CAF degradation by Mn2O3 with PS

Figure S15 .

 S15 Figure S15. Degradation of CAF in milli-q (circles) and STP (triangles) waters using 5 mM of

5 .

 5 Briefly, 100 µL of sample was withdrawn at different reaction times and added to 5 mL of a solution containing KI (0.6 M) and HCO3 -(0.06 M). The Abs of solution at 348 nm was measured after 10 min of reaction and PMS concentration (CPMS, (mM)) was determined as 𝐶 𝑃𝑀𝑆 = Where VPMS and Vtot are the volume of sample (100 µL) and total volume (5.1 mL), 0.5 was the slope of the linear correlation between the PMS concentration and value of Abs at 348 nm.The Atomic absorption spectroscopy (AAS, PerkinElmer) was employed to determine the concentration of dissolved Mn 2+ in the aqueous solution. The procedure involved the filtration of the samples through a 0.2 μm PTFE filter and the addition of 2% nitric acid (HNO3, 65%, Sigma-Aldrich) before detection. Mn 2+ standard solution for AAS (TraceCERT ® , 1000 mg L -1

Figure 1 .

 1 Figure 1. Degradation of BPA (25 µM) in the presence of γ-MnOOH (0.3 g L -1 ) and PMS (0.5 mM) at pH 6.5.

Figure 2 .

 2 Figure 2. Pseudo-first order degradation rate of BPA using different γ-MnOOH concentrations in the presence of 0.5 mM of PMS (A) and in the presence of γ-MnOOH (0.1 g L -1 ) and different PMS concentrations (B). Error bars are determined from the σ-level uncertainty of the data fit used to obtain k'BPA.

Figure 3 .

 3 Figure 3. BPA degradation (k'BPA) in the system γ-MnOOH 0.1 g L -1 + PMS 2 mM in the absence and presence of different scavengers at pH 6.5. Error bars are determined from the σlevel uncertainty of the data fit used to obtain k'BPA.

Figure 4 .

 4 Figure 4. XPS Mn 2p and Mn 3s spectra of γ-MnOOH (A and B), γ-MnOOH + PMS (C and D) and γ-MnOOH + PMS + BPA (E and F) Initial conditions are [BPA] = 5 µM, [γ-MnOOH] = 0.1 g L -1 and [PMS] = 2 mM at pH 6.5.

Figure 5 .Figure 6 .

 56 Figure 5. Degradation of BPA, E2 and EE2 in the system γ-MnOOH/PMS. The solid lines show the fit of data using a mono-exponential decay function. Blue points show the PMS consumption during the reaction. Initial conditions are [BPA] = [E2] = [EE2] = 5 µM, [γ-MnOOH] = 0.1 g L -1 and [PMS] = 2 mM at pH 6.5.

  nitrogen anions (NO2 -and NO3 -) on BPA, E2 and EE2 degradation were investigated, individually. As shown in FigureS8, no effect of both NO2 -and NO3 -on the degradation of the three EDCs was observed.

Figure 7 .

 7 Figure 7. Effect of chloride ions (Cl -) (A) and bicarbonate ions (HCO3 -) (B) on the degradation of BPA, E2, and EE2 using γ-MnOOH/PMS system. Experimental conditions are: [γ-MnOOH] = 0.1 g L -1 , [PMS] = 2 mM, [BPA] = [E2] = [EE2] = 5µM, pH = 6.5 (for Cl -) and pH = 8.0 (for HCO3 -). Error bars are determined from the σ-level uncertainty of the data fit used to obtain k' of estrogens.

  matter and inorganic anions present in STP water, which can compete with degradation EDCs during the γ-MnOOH/PMS process. It is interesting to notice that EDCs initial degradation in STP water is lower during the first 15 min and rapidly increases to almost complete removal after 35-40 min. 18 % of TOC removal was also observed after 30 min of reaction and TOC removal continues after EDCs degradation with a rate of 0.067 ± 0.007 mgC min -1 .

Figure 8 .

 8 Figure 8. BPA, E2 and EE2 degradation by γ-MnOOH/PMS system in STPW and TOC removal. Experimental conditions are: [γ-MnOOH] = 0.1 g L -1 , [PMS] = 2 mM, [BPA] = [E2] = [EE2] = 5µM, and pH = 6.5. Error bars on TOC analysis was 0.2 mgC L -1 .

1 EtOH ( 10 Figure S1 .

 110S1 Figure S1. XRD patterns of γ-MnOOH. γ-MnOOH indexed with the manganite structure (* indicates an XRD contribution from the stored salt NaNO3 of manganite suspension). For clarity, only the main Bragg peaks are indexed. TEM image of the as-synthesized manganite (MnOOH) nanowires. SAED pattern. Structure of γ-MnOOH confirmed by the electronic diffraction diagrams. Interreticular distances: 3,472 Å (111); 2,466 Å (002) ; 2,244 Å (121).

Figure S2 .

 S2 Figure S2. HPLC-fluorescence chromatography of 5 𝜇M BPA, E2 and EE2.

Figure S3 .

 S3 Figure S3. Calibration curve of PMS quantification using iodometric titration.

Figure S6 .

 S6 Figure S6. BPA, E2, and EE2 (5 µM) degradation in the presence of 0.1 g/L γ-MnOOH and 2 mM PMS at pH 6.5.

Figure S7 .

 S7 Figure S7. BPA, E2, and EE2 degradation in the presence of PMS without (solid line) and with (dashed lines) and Cl -or HCO3. Experimental conditions are: [PMS] = 2mM, [BPA] = [E2] = [EE2] = 5 µM, [NaCl] = 2 mM, [NaHCO3] = 5 mM and pH = 6.5.

Figure S8 .

 S8 Figure S8. Effect of NO2 -(A) and NO3 -(B) on the degradation of BPA, E2, and EE2 using γ-MnOOH/PMS system. Experimental conditions are: [γ-MnOOH] = 0.1 g L -1 , [PMS] = 2 mM, [BPA] = [E2] = [EE2] = 5µM, pH = 6.5. Error bars are determined from the σ-level uncertainty of the data fit used to obtain k' of estrogens.

  -143 (2013) 729-735. https://doi.org/10.1016/j.apcatb.2013.06.004. [17] D. He, Y. Li, C. Lyu, L. Song, W. Feng, S. Zhang, New insights into

	MnOOH/peroxymonosulfate system for catalytic oxidation of 2,4-dichlorophenol: Morphology
	dependence and mechanisms, Chemosphere. 255 (2020) 126961. https://doi.org/10.1016/
	j.chemosphere.2020.126961.

Table 1 - 1 .

 11 The physical and chemical properties of H2O2, PDS, and PMS[14,15].

	Priorities	H2O2	PDS 1	PMS 2
	CAS number	7722-84-1	7775-27-1	70693-62-8
	Chemical formula	H2O2	NaS2O8	KHSO5 • 0.5KHSO4 • 0.5K2SO4
	Molar weight (g mol -1 )	34	238.1	307.28
	Water Solubility 3 (g L -1 ) miscible	730	298
	Density (g cm -3 )	1.11 4	2.59	1.15
	pKa	11.75	-3.5	9.3

1

PDS refers to sodium peroxydisulfate.

Table 1 -

 1

	Mineral name	Chemical formula	Mn valence
	Manganosite	MnO	+II
	Pyrochroite	Mn(OH)2	+II
	Rhodocrosite	MnCO3	+II
	Hausmannite	Mn3O4	+II, +III
	Bixbyite	Mn2O3	+III
	Groutite	𝛼-MnOOH	+III
	Feitknechite	𝛽-MnOOH	+III
	Manganite	𝛾-MnOOH	+III
	Birnessite	(Na, K, Ca)4Mn7O14 • 2.8 H2O	+II, +III, +IV
	Vernadite	MnO2 • nH2O
	Lithophorite	LiAl2(Mn 2 4+ Mn 3+ )O6(OH)6	+III, +IV
	Manjiroite	Nax(Mn 4+ , Mn 3+ )8O16	+III, +IV
	Chalcophanite	ZnMn3O7 • 3 H2O	+II, +IV

3. Representative MnOx minerals

[64]

.

Table 1 -

 1 3 (continued)

	Mineral name	Chemical formula	Mn valence
	Todorokite	(Ca,Na,K)X(Mn 4+ , Mn 3+ )6O12 • 3.5 H2O	+III, +IV
	Romanechite	Ba0.66(Mn 4+ , Mn 3+ )5O10 • 1.34 H2O	+III, +IV
	Coronadite	Pbx(Mn 4+ , Mn 3+ )8O16	+III, +IV
	Cryptomelane	Kx(Mn 4+ , Mn 3+ )8O16	+III, +IV
	Hollandite	Bax(Mn 4+ , Mn 3+ )8O16	+III, +IV
	Nsutite	Mn(O,OH)2	+II, +III, +IV
	Ramsdellite	𝛾-MnO2	+IV
	Pyrolusite	𝛽-MnO2	+IV
	1.2.2 Structures

Table 1 -4. The

 1 effect of structures on the reactivity of Mn(III) (oxyhydr)oxides.

	Catalysts	Structure	Initial conditions	Reactivity	Mechanism	Ref.

Table 1 -4 (continued)

 1

	Catalysts	Structure	Initial conditions	Reactivity	Mechanism	Ref.
				100 mg of catalysts
				10 mg of soot	96.3, 89.7 and 85.2% of soot combustion
		cubic
				5% v/v of O2	efficiencies were observed with the	the exposed (001) facet
	α-Mn2O3	octahedral				[87]
				0.25% v/v of NO	catalysis of α-Mn2O3-cubic, -truncated	of cubic Mn2O3
		truncated octahedral
				N2 as balance gas	octahedral, -octahedral
				9990 h -1 of GSHV
				[Catalyst] = 0.3 g/L
					98%, 88%, and 55% removal of 2,4-DCP
		nanowires	[PMS] = 12 mM		higher zeta-potential
					was achieved in γ-MnOOH nanowires,
	γ-MnOOH	multi-branches	[2,4-DCP*] =100		value of nanowires γ-	[90]
					multi-branches, and nanorods activated
		nanorods	mg/L		MnOOH
					PMS systems, separately
				Initial pH = 7
				[Catalyst] = 0.2 g/L
		nano-cubic			lager surface Mn(IV)
				[PMS] = 0.65 mM	100% CIP removal in 80 min by Mn3O4
	Mn3O4	nano-plate			contents of Mn3O4	[92]
				[CIP*] = 10 mg/L	nano-octahedral
		nano-octahedral			nano-octahedral
				pH = 7.7
	* CIP: ciprofloxacin.	*2,4-DCP: 2,4-dichlorophenol

.3.2 Activation of PMS by Mn(III) (oxyhydr)oxides

	The Mn(III) (oxyhydr)oxides/PMS system has been applied for the removal of various
	contaminants, such as phenol, bisphenol A, 2,4-dichlorophenol, ciprofloxacin, and organic dyes
	[92-97]. Different studies involving PMS activation by Mn(III) (oxyhydr)oxides are gathered
	in Table 1-5. According to the literature, the efficient degradation of organic pollutants is
	generally attributed to the generation of active species such as SO 4 •-, HO • , and 1 O2. The
	activation mechanisms of PMS by Mn(III) (oxyhydr)oxides are proposed as shown in Eqs. (1-
	11)-(1-14)

  For example, He et al.

	demonstrated the contribution of 1 O2 for the degradation of 2,4-dichlorophenol in the γ-
	MnOOH/PMS system. The generation of 1 O2 was attributed to two pathways including the
	decomposition of PMS and the reaction of O 2 •-with HO • (Eqs. (1-19)-(1-24)) [90,102,103].
	Chen et al. synthesized one new Mn3O4 nanodots-g-C3N4 nanosheet (Mn3O4/CNNS) and
	investigated its performance on PMS activation for 4-chlorophenol (4-CP) degradation [104].
	The chemical scavenging tests and electron spin resonance (ESR) experiments confirmed the
	contribution of 1 O2 for the removal of 4-CP.
	HSO 5 -+ SO 5 2-→ HSO 4 -+ SO 4 2-+ 1 O2	(1-19)
	HSO 5 -+ H2O → HSO 4 -+ H2O2	(1-20)
	H2O2 ↔ H + + HO 2 -	pKa = 11.6	(1-21)
	H2O2 + HO • → H2O + HO 2 •	(1-22)

Table 1 -5. Summary

 1 of PMS activation by Mn(III) (oxyhydr)oxides.

	Catalyst	Pollutant	Initial conditions	Reactivity	Active species Ref.
	Mn2O3	Phenol	[Catalyst] = 0.4 g/L	100% removal of phenol in 60 min	SO 4 •-	[52]
			[PMS] = 2 g/L
			[Phenol] = 25 mg/L
	Mn3O4	Phenol	[Catalyst] = 0.4 g/L	100% removal of phenol in 20 min	SO 4 •-	[43]
			[PMS] = 2 g/L
			[Phenol] = 25 mg/L
	Mn3O4 nanoparticle	Methylene blue	[Catalyst] = 0.12 g/L	86.71% removal of MB in 20 min	SO 4 •-	[94]
		(MB)	[PMS] = 0.94 g/L
			[MB] = 62 mg/L
			pH = 4
	Mn3O4 nano-octahedral Ciprofloxacin	[Catalyst] = 0.2 g/L	100% removal of CIP in 80 min	SO 4 •-	[92]
		(CIP)	[PMS] = 0.65 mM		HO •
			[CIP] = 10 mg/L
			pH = 7.7

Table 1 -5 (continued) Catalyst Pollutant Initial conditions Reactivity Active species Ref.

 1

	yolk-shell Mn3O4	Bisphenol A	[Catalyst] = 0.1 g/L	87.7% of removal of BPA in 60 min	SO 4 •-	[97]
		(BPA)	[PMS] = 0.3 g/L		HO •
			[BPA] = 10 mg/L
			pH = 5.3
	3D hierarchical Mn3O4 Phenol	[Catalyst] = 0.2 g/L	100% removal of phenol in 60 min	SO 4 •-	[96]
			[PMS] = 6.5 mM		HO •
			[Phenol] = 20 ppm
			pH = 6.8
	dumbbell-like Mn2O3	Rhodamine B	[Catalyst] = 0.25 g/L	100% of removal of RhB in 30 min	SO 4 •-
		(RhB)	[PMS] = 0.75 g/L
			[RhB] = 10 mg/L

Table 1 -5 (continued) Catalyst Pollutant Initial conditions Reactivity Active species Ref.

 1

	γ-MnOOH nanowire 2,4-dichlorophenol	[Catalyst] = 0.3 g/L	98% removal of 2,4-DCP in 6 h	SO 4 •-	[90]
		(2,4-DCP)	[PMS] = 12 mM		HO •
			[2,4-DCP] = 100 mg/L		O 2 •-
			pH = 7		1 O2
	Ce-Mn2O3	2,4-dichlorophenol	[Catalyst] = 0.2 g/L	100% removal of 2,4-DCP in 90	SO 4 •-	[106]
		(2,4-DCP)	[PMS] = 1.0 g/L	min	HO •
			[2,4-DCP] = 50 mg/L		1 O2
			pH = 7
	Mn3O4-GO	Orange II	[Catalyst] = 50 mg/L	100% removal of Orange II in	SO 4 •-	[107]
			[PMS] = 1.5 g/L	120 min under neutral pH
			[Orange II] = 30 mg/L
	MnOOH@nylon	2,4-dichlorophenol	[Catalyst]=0.76 mg/cm 2	97.9% removal of 2,4-DCP in 2 h	SO 4 •-
		(2,4-DCP)	[PMS] = 138 mg/L
			[2,4-DCP] = 25 mg/L
			pH = 6.0-6.4

Table 1 -5 (continued) Catalyst Pollutant Initial conditions Reactivity Active species Ref.

 1

	γ-MnOOH-rGO Bentazone	[Catalyst] = 0.075 g/L	96.1% removal of Bentazone in 90 min	HO •	[108]
		[PMS] = 0.615 g/L		1 O2
		[Bentazone] = 10 mg/L
		pH = 7, sunlight

Table 1 -5 (continued) Catalyst Pollutant Initial conditions Reactivity Active species Ref.

 1

	Mn3O4/MOF	Rhodamine B	[Catalyst] = 0.4 g/L	98% removal of RhB in 60 min	SO 4 •-	[99]
		(RhB)	[PMS] = 0.3 g/L		HO •
			[RhB] = 10 mg/L
			pH = 5.18
	Fe3O4/Mn3O4/GO	Methylene Blue	[Catalyst] = 100 mg/L	98.8% removal of MB in 30 min	SO 4 •-
		(MB)	[PMS] = 0.3 g/L		HO •
			[MB] = 50 mg/L
			pH = 7

Table 1 - 6 .

 16 Summary of PDS activation by Mn(III) (oxyhydr)oxides.

	Catalyst	Pollutant	Initial conditions	Reactivity	Active species	Ref.

Table 1 -

 1

	Catalysts	PZC	Reference
	α-Mn2O3	4.7	[114,132]
	γ-MnOOH	3.4	[116]
	Mn3O4	5.6-7.34	[98,113,133]
	Oxidants	pKa	Reference
	PMS	9.4	[134]
	PDS	-3.5	[135]
	Pollutants	pKa	Reference
	Phenol	9.98	[136]
	Bisphenol A	9.6-10.2	[137]
	2,4-dichlorophenol	9.4	[109]
	Ciprofloxacin	8.70-10.58	[98,138]
	p-Chloroaniline	4.2	[116,139]
	4-Chlorophenol	9.29	[140]
	Orange II	11.4	[128]

7.

The PZC values of Mn(III) (oxyhydr)oxides and pKa values of PMS/PDS and pollutants.

  •-to generate Cl • which can react with another Cl -to form Cl 2 •-(Eqs. (1-37)-(1-38)) [149]. Both Cl • and Cl 2 •-have low reduction potentials (E0 = 2.4 & 2.0 V) in comparison with SO 4 •-, thus the consumption of SO

  4

  focuses on the discussion of the activation mechanism of PDS and PMS by various MnO2. Both radical and nonradical mechanisms are presented, and the influence factors of the catalytic and oxidative reactivity of MnO2 are also discussed.

Manganese dioxides have been used to activate PMS to remove various contaminants, such as bisphenol A, phenol, ibuprofen, and organic dyes

[40,[START_REF] Liu | Degradation of Rhodamine B by the α-MnO2/Peroxymonosulfate System[END_REF][START_REF] Saputra | α-MnO2 activation of peroxymonosulfate for catalytic phenol degradation in aqueous solutions[END_REF]

. It is generally considered that sulfate and hydroxyl radicals were generated in the system of MnO2/PMS, as shown in Eqs. (1-43)-(1-44) and Eqs. (1-17)-(1-18) [39]. The presence of SO 4 •-and HO • was not verified in the early study of the activation process of PMS by MnO2 [174,175]. However, with the development of technology, the identification of reactive radicals (such as SO 4 •-and HO • ) and non-radical species (such as singlet oxygen and reactive surface complex) were achieved in various MnO2/PMS systems [176,177]. For instance, Huang et al. reported the generation of SO 4 •-, HO • , and 1 O2 in the activation of PMS by 𝛼-, β-, 𝛾-, 𝛿-MnO2 under acidic pH [40]. Wang et al. demonstrated that the degradation of BPA by amorphous MnO2-activated PMS was not attributed to the generation of SO

  4

  Wang et al. compared the catalytic reactivity of 𝛼-MnO2 with different shapes (i.e. nanowire, nanorod and nanotube) in PMS activation[START_REF] Wang | New insights into heterogeneous generation and evolution processes of sulfate radicals for phenol degradation over one-dimensional α-MnO2 nanostructures[END_REF]. The results showed that the 𝛼-MnO2 in the

shape of nanowire exhibited a better catalytic performance in PMS activation than 𝛼-MnO2 in nanorod and nanotube shape, which can be ascribed to the higher surface areas and exposed high-energy facet (001) of 𝛼-MnO2 (nanowire). The nanostructures of MnO2 in different dimensions also exerted a considerable effect on the catalytic activity. For instance, Deng et al.

synthesized different dimensional 𝛼-MnO2 including 0D nanoparticle, 1D nanorod, and 3D nanoflower and investigated the reactivity of 𝛼-MnO2 in PMS activation for ciprofloxacin (CIP)

Table 1 - 8 .

 18 The physicochemical properties of BPA, E2, and EE2[START_REF] Liu | Removal mechanisms for endocrine disrupting compounds (EDCs) in wastewater treatment -physical means, biodegradation, and chemical advanced oxidation: A review[END_REF].

	Priorities	Bisphenol A (BPA) 17β-estradiol (E2)	17𝛼-Ethynylestradiol (EE2)
	CAS number	80-05-7	50-28-2	57-63-6
	Chemical formula C15H16O2	C18H24O2	C20H24O2
	Chemical structure
	Molar weight	228.9 g mol -1	272.4 g mol -1	296.4 g mol -1
	Water Solubility 1	120 mg L -1	3.6 mg L -1	11 mg L -1
	Log Kow 2	3.32	4.01	3.67
	pKa	9.6	10.71	10.4

1

The values of water solubility are achieved at 25℃.

2

Log Kow: Octanol-water partition coefficient.

Table 1 - 9 .

 19 The occurrence of EDCs in various water systems.

	Target EDCs		Concentration (ng L -1 )
		Surface water	Groundwater	Drinking water
	BPA	< 8.8-1000 [201]	n.d-7 [202]	0.5-2.0 [203]
		0.5-14 [203]	n.d-930 [204]	5.0 [205]
		65-295 [205]
	E2	< 0.3-5.5 [206]	n.d-45 [207]	0.2-2.1 [203]
		0.6-1.0 [208]	6-66 [209]	2.6 [210]
		< 0.8-1.0 [201]	13-80 [211]	n.d-2.6 [212]
	EE2	< 0.1-4.3 [206]	n.d-0.94 [204]	0.15-0.5 [203]
		< 0.3-0.4 [201]
		0.1-5.1 [203]

Table 1 - 10 .

 110 The physicochemical properties of CAF[START_REF] Rigueto | Alternative techniques for caffeine removal from wastewater: An overview of opportunities and challenges[END_REF].

	Priorities	Caffeine (CAF)
	CAS number	58-08-2
	Chemical formula	C8H10N4O2
	Chemical structure
	Density	1.23 g cm -3
	Molar weight	194.19 g mol -1
	Water Solubility	20 g L -1
	Log Kow	-0.04 to 0.01
	Melting point	235 to 238℃
	pKa	10.4

Table 1 -

 1

11. CAF concentration in different water matrices.

Table 1 -

 1 To improve the removal efficiency of EDCs and CAF in waters, various water treatment technologies have been developed as shown in Table 1-12. The efficient degradation of EDCs and CAF by these technologies demonstrated the potential of AOPs in the treatment of wastewater. Moreover, EDCs and CAF are proved to be good probes for testing the efficiency of various wastewater treatment technologies. 12 Technologies for EDCs and CAF degradation.

	Pollutants	Technologies	Reference
	BPA	UV/TiO2	[225]

Table 2 - 1 .

 21 Chemical reagents.

	Reagent name	Chemical formula	CAS number	Grade	Company
	Hydrogen peroxide	H2O2	7722-84-1	30 wt.% in H2O Sigma-Aldrich*
	Sodium persulfate	Na2S2O8	7775-27-1	98%	Sigma-Aldrich
	Oxone	KHSO₅•0.5KHSO₄•0.5K₂SO₄ 70693-62-8	-	Sigma-Aldrich
	Sodium bisulfite	NaHSO3	7631-90-5	ACS reagent	Sigma-Aldrich
	Manganese(III) oxide Mn2O3	1317-34-6	99%	Sigma-Aldrich
	Manganese chloride	MnCl2	7773-01-5	99%	Riedel-de Haen
	Perchloric acid	HClO4	7601-90-3	72%	Sigma-Aldrich
	Sodium hydroxide	NaOH	1310-73-2	97%	Sigma-Aldrich
	Hydrochloric acid	HCl	7647-01-0	37%	Sigma-Aldrich
	2-propanol	C3H8O	67-63-0	99%	Sigma-Aldrich
	tert-butanol	C4H10O	75-65-0	99%	Acros organics
	Ethanol	C2H6O	64-17-5	96%	VWRchemicals
	Furfuryl alcohol	C5H6O2	98-00-0	98%	Sigma-Aldrich
	Sodium chloride	NaCl	7647-14-5	99%	Sigma-Aldrich
	Sodium bicarbonate	NaHCO3	144-55-8	99.7%	Sigma-Aldrich
	Sodium nitrite	NaNO2	7632-00-0	97%	Sigma-Aldrich
	Sodium nitrate	NaNO3	7631-99-4	99%	Fluka
	Sodium phosphate	NaH2PO4	7558-80-7	99%	Sigma-Aldrich
	Sodium metasilicate	Na2SiO3	6834-92-0	-	Sigma-Aldrich
	Phosphoric acid	H3PO4	7664-38-2	85%	Sigma-Aldrich

*Sigma-Aldrich refers to the Sigma-Aldrich company located in France.

Table 2 -

 2 1 (continued)

	Reagent name	Chemical formula	CAS number	Grade	Company
	Acetonitrile	C₂H₃N	75-05-8	HPLC grade	Carlo Erba Reagents
	Potassium permanganate	KMnO4	7722-64-7	99%	Sigma-Aldrich
	HEPES	C8H18N2O4S	7365-45-9	99%	Sigma-Aldrich
	Oxalic acid	H2C2O4	144-62-7	-	Sigma-Aldrich
	Sulfuric acid	H2SO4	7664-93-9	98%	Sigma-Aldrich
	Hydroxylamine hydrochloride	NH2OH•HCl	5470-11-1	98%	Sigma-Aldrich
	Terephthalic acid	C8H6O4	100-21-0	98%	Sigma-Aldrich
	Potassium iodide	KI	7681-11-0	99%	Sigma-Aldrich
	2-hydroxy-terephthalic acid	C8H6O5	636-94-2	97%	Sigma-Aldrich
	4-hydroxyphenylacetic acid	C8H8O3	156-38-7	98%	Sigma-Aldrich
	Peroxidase	-	9003-99-0	-	Sigma-Aldrich
	Nitric acid	HNO3	7697-37-2	65%	Sigma-Aldrich
	Caffeine (CAF)	C8H10N4O2	58-08-2	-	Sigma-Aldrich
	Bisphenol A (BPA)	C15H16O2	80-05-7	99%	Sigma-Aldrich
	17β-estradiol (E2)	C18H24O2	50-28-2	-	Sigma-Aldrich
	17α-ethinylestradiol (EE2)	C20H24O2	57-63-6	98%	Sigma-Aldrich
	DMPO 1	C6H11NO	3317-61-1	-	Sigma-Aldrich
	TEMP 2	C9H19NO	2403-88-5	98%	Sigma-Aldrich

Table 2 - 2 .

 22 Equipment and instruments.

	Instruments

Table 2 -

 2 3.

Table 2 - 3 .

 23 Characteristics of the sewage treatment plant water (STPW).

	Parameters	Values
	pH	8.5
	TOC*	8.2 mgC L -1
	IC*	30.35 mgC L -1
	Cl -	0.77 mM
	NO3 -	0.41 mM
	SO4 2-	0.29 mM
	PO4 3-	0.09 mM

*TOC: Total organic carbon; *IC: inorganic carbon.

Table 2 - 4 .

 24 Parameters used for chromatographic separation of compounds on HPLC.

	Chemicals Detector Flow rate	1 DW	Elution mode & Mobile phase
	BPA	DAD	0.8 mL min -1 277 nm Gradient elution mode
					2 MeOH	Water
					0 min	70%	30%
					7 min	90%	10%
					9 min	90%	10%
					9.5 min 70%	30%
	CAF	DAD	0.8 mL min -1 272 nm Gradient elution mode
					3 ACN	Water
					(0.1% H3PO4)
					0 min	20%	80%
					6 min	70%	30%
					8 min	70%	30%
					8.5 min 20%	80%
	4 E2/EE2	FD	1.0 mL min -1 280 nm isocratic ACN	Water (50%)
					(50%)
	BPA	FD	1.0 mL min -1 277 nm isocratic ACN	Water (50%)
					(50%)

1 DW: detection wavelength. 2 MeOH: methanol. 3 ACN: acetonitrile.

4

E2/EE2: The retention time of E2 and EE2 are 6.2 and 7.3 min, respectively.

  al. indicated the efficient degradation of bisphenol A by crystalline MnO2 in the presence of PMS was positively correlated to the amount of Mn(III). The investigation of PDS and PMS activation by Mn(III) oxides was also reported. For instance, Saputra et al. demonstrated that Mn2O3 compared with MnO, Mn3O4, and MnO2 showed the highest reactivity in PMS activation for phenol degradation. He et al. suggested that nanowires MnOOH have higher activity on PMS activation for 2,4-dichlorophenol oxidation as compared with multi-branches and nanorods MnOOH.

Table S2

 S2 Physico-chemical properties of the sewage treatment plant water (STPW)

	Parameter	Values
	pH	8.5
	TOC*	0.68 mM
	IC*	2.53 mM
	Cl -	0.77 mM
	NO3 -	0.41 mM
	SO4 2-	0.29 mM
	PO4 3-	0.09 mM

*TOC: Total organic carbon; *IC: inorganic carbon.

  Table S4 Scavenging ratios between CAF and scavengers with ROS a .

				% of radical
	ROS			C1k1 / (C1k1 + C2k2)	reacting with
				scavenger
	HO •	[TBA] c × 𝑘 HO • ,TBA /( [TBA] c × 𝑘 HO • ,TBA + [CAF] b × 𝑘 HO • ,CAF )	97.4
	SO4 •-	[2-Pr] d × 𝑘 𝑆𝑂 4 •-,2-Pr / ([2-Pr] d × 𝑘 𝑆𝑂 4 •-,2-Pr +[CAF] b × 𝑘 𝑆𝑂 4 •-,CAF )	93.2
	1 O2	[FFA] e × 𝑘 1𝑂 2 ,FFA / ([FFA] e × 𝑘 1𝑂 2 ,FFA + [CAF] b × 𝑘 1𝑂 2 ,CAF )	99.4
	a ROS: Reactive Oxygen Species; 1 for scavengers and 2 for CAF;
			7.1	7.0
	Mn2O3 + H2O2 + CAF	7.8	7.5
	Mn2O3 + HSO3 -+ CAF	3.5	2.2
	Mn2O3 + PS + CAF	5.7	5.3
	Mn2O3 + PMS + CAF	2.7	2.6

Table S5

 S5 Degradation intermediates of CAF in Mn2O3-activated PS and PMS systems.

Table 1 .

 1 Parameters from curve fitting of Mn 3s and Mn 2p3/2 XPS spectra and average oxidation state.

					γ-	γ-
		Sample		γ-MnOOH
					MnOOH+PMS	MnOOH+PMS+BPA
		Mn(II)	BE (eV)	639.60	639.64	640.54
			%	8	14	37
		Mn(III)	BE (eV)	640.41	641.00	640.88
	Mn 2p		%	85	61	54
		Mn(IV)	BE (eV)	643.10	642.85	643.18
			%	7	25	9
			AOS	2.99	3.10	2.72
		Multiplet Splitting
				5.29	5.18	5.61
	Mn 3s	(eV)
			AOS	3.00	3.12	2.64

DMPO: 5,5-dimethyl-1-pyrroline N-oxide;

TEMP: 2,2,6,6-tetramethyl-4-piperidinol;

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.

Acknowledgments

This work was supported by the Region Council of Auvergne Rhône Alpes, the "Féderation des

Acknowledgements

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

AOS = 2 + 54.94 ×

n Mn (mol) m Mn (g) (2)(3)

B. Determination of Mn AOS using X-ray photoelectron spectroscopy method

The determination of Mn AOS using x-ray photoelectron spectroscopy (XPS) can be achieved through analyzing the Mn 3s multiplet splitting spectra. The equation used for the calculation of Mn AOS is shown in Eq. (2)(3)(4) [11]. The ∆𝐸 𝑠 can be obtained from the energy difference between the main peak and its satellite in Mn 3s. AOS = 8.95 -1.13 ∆𝐸 𝑠 (2)(3)(4) In addition, the Mn 2p spectrum can be used to obtain the abundance of Mn(II), Mn(III), and Mn(IV) on the surface of manganese oxides. The binding energy of Mn(II), Mn(III), and Mn(IV)

in Mn 2p spectrum are 639, 640, and 643eV, separately. It is worth noting the binding energy of Mn ions may shift in a small scale due to the different reactions.

Data analysis, curve fitting and quantification of the Mn XPS spectra were performed using the CasaXPS and xpsPEAK4 software. The binding energy (BE) scale was referenced to the adventitious C1s peak at 284.8 eV. High-resolution XPS spectra were processed using Shirley back-ground subtraction. After background subtraction, the spectra were smoothed and fitted by using of Gaussian-Lorentzian peak shape. For each sample, fitting parameters (FWHM, the full width at half maximum; H, the peak height; E, the peak position center) were adjusted to minimize the value of residual standard deviation (STD).

Characterization of manganese oxides

A. Brunauer-Emmett-Teller (BET) surface measurement

The surface area of manganese oxide was measured using the BET method. Specifically, it allows the determination of the overall specific external and internal surface area of nano-or porous solids by measuring the amount of physically absorbed gas according to the BET method.

A micrometrics ASAP 2010 instrument was employed to measure the adsorption of nitrogen.

Prior to adsorption, the electrode materials were out-gassed under vacuum at room temperature.

B. X-ray diffraction (XRD) analysis

Chapter III

Supplementary Materials

Acid birnessite synthesis. Acid birnessite (AB) was prepared following the procedures of

McKenzie [1]. 166 mL of concentrated HCl was added dropwise to 2.5 L of 0.4 M KMnO4.

The suspension was stirred vigorously and kept at 90°C during the HCl addition. The reaction was continued at 90 °C for a further 10 min. Then, the obtained slurry was stirred vigorously at room temperature for 15 h. The precipitate was collected by centrifugation, washed with ultrapure water repeatedly until the conductivity was close to 0.6 μS cm -1 . The suspensions were stored in polypropylene containers at 4 °C for further use.

Preparation of Mn(III)-rich birnessite. Mn(III)-rich MnO2 was synthesized according to previous published methods [2,3]. Briefly, about 2 g of acid birnessite was incubated with 80 mL of 10 mM HEPES and 10 mM NaCl solutions at pH 6.5 ±0.05 at a temperature of 25 ℃ for 48h. The precipitate was collected by centrifugation and washed with ultrapure water.

Oxalic acid-permanganate back-titration method. 0.1 g of the acid birnessite was dissolved in 2.5 mL of 0.5 M H2C2O4 and 5 mL of 1 M H2SO4 to reduce all highly charged manganese ions to Mn(Ⅱ). After all samples to be dissolved, the excess C2O4 Hydroxyl radical detection method. Potential generation of hydroxyl radical was monitored by using a fluorescence method using terephthalic acid (TA) as a probe [4]. The detection mechanism is that nonfluorscence TA reacts with hydroxyl radical to generate fluorescent 2hydroxyterephthalic acid which can be measured at 425 nm under the excitation of 320 nm.

Specifically, in this work, 345 µM of acid birnessite was mixed with TA (1 mM) and H2O2 ( 200Chapter IV 2θ at a step size of 0.02º. The average oxidation state (AOS) of Mn after the reaction was obtained using XPS (hemispherical electron energy analyzer OMICRON EA125, Germany).

The data analysis, curve fitting, and quantification of the Mn XPS spectra were performed using XPSPEAK 4.1 software. The C 1s signal at 284.8 eV was taken as a reference for binding energy calibration.

Supplementary Material

Table S1. Characteristics of the sewage treatment plant water (STPW).

Parameters

Chapter V General conclusion and outlook

During this PhD, we have investigated the ability of various manganese oxides used to degrade the organic pollutants in water. The activation mechanism of radical precursors (such as H2O2, PDS, and PMS) by Mn(III/IV) oxides was investigated. Different influence factors (such as solution pH, catalysts and oxidants dosages, presence of inorganic anions) on the oxidative and catalytic activity of Mn(III/IV) oxides were studied. Moreover, the application of PDS or PMS activated by Mn(III) oxides system in the sewage treatment plant (STP) water was also evaluated.

In the first part of this dissertation, the impacts of environmental levels of hydrogen peroxide and oxyanions on the redox activity of MnO2 particles were investigated. We have notably demonstrated that environmental levels of hydrogen peroxide can induce reductive dissolution of MnO2 into Mn(II), thereby affecting the surface reactivity of MnO2. As the H2O2 decomposition is a fast process, the generated dissolved Mn(II) binds to MnO2 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 MnO2 in presence of H2O2. 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 H2O2 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 H2O2 and co-exiting anions on the global carbon cycle. Furthermore, the widespread presence of H2O2 in surface and ground water