Ascorbic acid promoted magnetite Fenton degradation of alachlor: Mechanistic insights and kinetic modeling

Journal Pre-proof Ascorbic Acid Promoted Magnetite Fenton Degradation of Alachlor: Mechanistic Insights and Kinetic Modeling Hongwei Sun, Guihong Xie, Di He, Lizhi Zhang

PII:

S0926-3373(19)31129-4

DOI:

https://doi.org/10.1016/j.apcatb.2019.118383

Reference:

APCATB 118383

To appear in:

Applied Catalysis B: Environmental

Received Date:

27 September 2019

Revised Date:

23 October 2019

Accepted Date:

2 November 2019

Please cite this article as: Sun H, Xie G, He D, Zhang L, Ascorbic Acid Promoted Magnetite Fenton Degradation of Alachlor: Mechanistic Insights and Kinetic Modeling, Applied Catalysis B: Environmental (2019), doi: https://doi.org/10.1016/j.apcatb.2019.118383

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Ascorbic Acid Promoted Magnetite Fenton Degradation of Alachlor: Mechanistic Insights and Kinetic Modeling Hongwei

Suna,b,

Guihong

Xiea,

Hec,*

Di

[email protected],

Lizhi

Zhanga,*

[email protected]

a

Key Laboratory of Pesticide & Chemical Biology of Ministry of Education, Institute of

Environmental & Applied Chemistry, College of Chemistry, Central China Normal University, Wuhan

b

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430079, People’s Republic of China Department of Chemical Engineering and Division of Environmental Science and Engineering,

Pohang University of Science and Technology (POSTECH), Pohang 37673, Korea

Guangdong Key Laboratory of Environmental Catalysis and Health Risk Control, Institute of

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c

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Environmental and Ecological Engineering, Guangdong University of Technology, Guangzhou,

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510006, China

E-mail:. (L. Zhang)

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E-mail:. (D. He)

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* To whom correspondence should be addressed.

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Graphical Abstract

Highlights 

Ascorbic acid greatly enhanced Fenton performance of Fe3O4/H2O2 system.



A kinetic model was developed to quantitatively describe Fe3O4/AA/H2O2 system.



Individual contributions of surface and homogeneous Fenton reaction were simulated.



Surface Fenton reaction mainly contributed to the overall alachlor degradation by more than

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62.6%.

ABSTRACT

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In this study we constructed a heterogeneous Fenton system with ascorbic acid (AA), magnetite (Fe3O4)

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and H2O2 for the alachlor degradation, aiming to clarify the heterogeneous Fenton mechanism. The addition of AA could significantly accelerate the Fenton reaction by promoting the surface

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Fe(III)/Fe(II) redox cycle (iron cycle) of Fe3O4. A kinetic model was successfully developed to quantitatively describe the complicated reactions in the Fe3O4/AA/H2O2 system. We thus employed

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this model to identify the individual contributions of surface and homogeneous Fenton reactions to the

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overall alachlor degradation in the Fe3O4/AA/H2O2 system, and found the surface Fenton reaction was mainly responsible for the alachlor degradation with more than 62.6% of contribution. This work

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offers a new strategy to improve the heterogeneous Fenton activity via promoting surface Fenton reaction, and sheds light on the possibility to quantitatively describe and predict the heterogeneous Fenton processes with first principle kinetic models.

Keywords: Ascorbic acid; Surface Fenton; Magnetite; Iron redox cycle; Kinetics

1. Introduction Fenton reaction is an efficient advanced oxidation process (AOP) for the remediation of organic pollutants [1]. However, the wide application of traditional homogeneous Fenton system is still restricted by drawbacks like narrow working pH (2.0-3.5) and mass iron sludge [2]. One solution to overcome these shortcomings is to develop heterogeneous Fenton systems [3, 4]. Iron bearing minerals such as hematite (α-Fe2O3) and goethite (α-FeOOH) are abundant in earth crust, inexpensive and environmentally benign, thus often serve as heterogeneous Fenton catalysts [5, 6]. In comparison with

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hematite and goethite, magnetite (Fe3O4) might be more attractive because of its intrinsic Fe(II) sites which act as active centers for H2O2 decomposition, and better electron transfer capacity resulting

for practical application and need further improvement.

Fe(II) + H2O2

→ →

Fe(II) + •HO2 + H+ Fe(III) + •OH + OH-

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Fe(III) + H2O2

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from its inverse spinel structure [6]. Nonetheless, the Fenton efficiency of Fe3O4 is still unsatisfactory

(1) (2)

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Similar with classic homogeneous Fenton process, the efficiency of heterogeneous Fenton systems is also limited by the insufficient Fe(III)/Fe(II) cycle [7], because the rate constant of Fe(III) reduction

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to Fe(II) by H2O2 (Eq 1) is approximately 4 orders of magnitude lower than that of Fe(II) oxidation

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by H2O2 (Eq 2), resulting in rapid depletion of Fe(II) during the reaction [6, 8]. To enhance the efficiency of heterogeneous Fenton, researchers employed chelating or reducing reagents to promote

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the Fe(III)/Fe(II) redox cycle (iron cycle). For example, our group found that hydroxylamine could improve the surface Fenton degradation rate of alachlor by 46 times in goethite/H2O2 system without any release of iron ions into solution, via accelerating the iron cycle on the goethite surface [7]. In contrast, other agents like succinate, citrate, oxalate and ethylene-diaminetetraacetic acid (EDTA) increased the heterogeneous Fenton performance along with significant release of iron ions [9], which thus involved both surface and homogeneous Fenton processes. However, the individual contributions

of surface and homogeneous Fenton reactions are not clear yet, going against a better understanding of heterogeneous Fenton mechanism. Normally, the contribution of homogeneous Fenton process was assessed by conducting control experiments with dissolved Fe2+ or Fe3+ of the same concentrations as those released by heterogeneous Fenton catalysts. Obviously, this approach could only offer an approximated assessment of real cases, since the concentrations of iron ions, hydrogen peroxide, and chelating/reducing reagents changed dynamically during the heterogeneous Fenton processes. So it is necessary to establish quantitative

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description of the kinetics in the heterogeneous Fenton system for more precise assessment. Moreover, the development of a first principle kinetic model, which is composed of the essential elementary reactions to simulate the heterogeneous Fenton system, will enable more detailed understanding about

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the thermodynamics and kinetics, such as the rate constants of surface and homogeneous Fenton

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reactions, the rate-limiting steps, and the equilibrium constants of surface complex. In addition, such a model is able to predict the consumption of reactants, the production of reactive species and the

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degradation of pollutants [10], and thus could benefit the optimization of experimental parameters in the heterogeneous Fenton systems. Unfortunately, the mechanism or kinetics of the heterogeneous

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Fenton reactions were often experimentally and qualitatively investigated [7, 11, 12], with very limited

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quantitative comprehension of the detailed mechanistic information of heterogeneous Fenton systems. Different from chelating/reducing reagents such as EDTA or hydroxylamine which are toxic and

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may cause secondary pollution, ascorbic acid (AA), also known as vitamin C, is an eco-friendly reducing and chelating agent widely found in plants, animals and human, and also widely used as dietary supplement or medicine [11]. Our group previously demonstrated that the addition of AA significantly accelerated the iron cycle in hematite and Fe@Fe2O3 nanowires heterogeneous Fenton systems [11, 13], and AA could also induce reductive dissolution of iron oxides [14]. Considering that magnetite possesses more surface Fe(II) than hematite, and is easily available and more inexpensive

than Fe@Fe2O3 nanowires, herein we constructed the Fe3O4/AA/H2O2 system to unveil the contributions of surface and homogeneous Fenton to the removal of alachlor, a recalcitrant herbicide [15, 16]. The mechanisms involved with the consumption of AA and H2O2, the generation of reactive species, the iron cycle, and the release of iron ions were investigated in detail. On the basis of experimental data, we developed a kinetic model to quantitatively describe the Fe3O4/AA/H2O2 system. The stability of magnetite, the degradation pathways of alachlor and AA, as well as their eco-safety

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were also studied carefully.

2. Materials and methods 2.1. Chemicals and Materials

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Magnetite (Fe3O4) was purchased from Shanghai Aladdin Bio-Chem Technology Co., Ltd. H2O2

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(30 wt %), L-ascorbic acid (AA), sodium acetate, benzoic acid, 4-hydroxyl benzoic acid, ethyl alcohol, sodium hydroxide, tert-butyl alcohol, iso-propanol, 1,10-phenanthroline, and 2,2’-bipyridine were of

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analytical grade and bought from Sinopharm Chemical Reagent Co. Ltd., China. Dichloromethane, methanol and acetonitrile were of HPLC grade and supplied by Merck KGaA. Hydroxylamine

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hydrochloride (99.9%), formic acid, acetic acid, and alachlor were purchased from Sigma-Aldrich. All

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the solutions were prepared in deionized water throughout the experiments. 2.2. Materials Characterization

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The powder X-ray diffraction (XRD, D8 Advance, Bruker) analysis was performed with a D/Max-

IIIA X-ray diffractometer, using a Cu Kα source (λ = 0.15418 nm). The morphology of the as-received commercial magnetite sample was examined by scanning electron microscopy (SEM, 6700-F, JEOL). The nitrogen adsorption-desorption isotherm of the Fe3O4 was obtained using a Micromeritics Tristar 3000 instrument at 77 K, and the specific surface area (SSA) was calculated based on the Brunauer−Emmett−Teller (BET) model.

2.3. Alachlor Degradation Experiments Batch trials of alachlor degradation were conducted in 100 mL conical flasks fixed on an orbital shaker at 250 rpm and ambient temperature. Briefly, the reaction systems consisted of 50 mL of alachlor solution (20 mg/L), 0.05 g of Fe3O4, 0.5 mL of 50 mmol/L AA solution, and the reaction was triggered by adding 0.5 ml of 100 mmol/L H2O2 stock solution. At predetermined intervals, 900 μL of the reaction solution was withdrawn and mixed with 100 μL of ethanol to stop the reaction, and the mixture was filtered through 0.22 μm syringe filter membranes to remove the suspended solids before

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high performance liquid chromatography (HPLC) analysis. The initial pH value of the Fe3O4/AA/H2O2 system was 4.0 without adjustment. The catalyst after reaction was collected by magnetic separation, carefully washed with deionized water and ethanol, and finally dried in a vacuum oven overnight to

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test the reusability of Fe3O4.

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2.4. Analytical Methods

The concentration of alachlor was determined with the method described in our previous reports

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and the details described in the supplementary material (SM Text S1) [13]. H2O2 concentration was measured using the iodide colorimetric method [17], while AA concentrations were determined using

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a modified 2, 2’-bipyridine method with a UV-vis spectrometer [18]. Notably, H2O2 and AA showed

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mutual interference of the colorimetric processes in the mixed Fe3O4/H2O2/AA system, but fortunately, the interference was dose-dependent and thus the exact concentrations of H2O2 and AA can be

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obtained through normalization and calibration, the details of which were provided in the SM Text S2, Table S1 and Table S2. Electron spin resonance (ESR) spectra were recorded with a JES FA 200 Xband spectrometer (JEOL, Japan). Accumulative •OH was probed by the well-defined reaction between benzoic acid (BA) and •OH to produce 4-hydroxybenzoic acid (4-HBA), which could be easily quantified by the HPLC method [19]. The elemental chemical states on magnetite surface before and after the Fenton reaction were checked by the X-ray photoelectron spectroscopy (XPS, Thermal

scientific, ESCALAB 250Xi) as described in the SM Text S3. Concentrations of Fe2+ and Fe3+ were detected by a 1, 10-phenanthroline colorimetric method [20]. For the measurement of dissolved Fe3+, hydroxylamine hydrochloride was used for the reduction of Fe3+ into Fe2+ for subsequent 1, 10phenanthroline measurement (SM Text S4). Surface Fe(II) and Fe(III) concentrations were measured by modifying the method described by He et al. [21]. Briefly, 10 mL of the reaction suspension was sampled at fixed intervals and saturated with argon gas for 5 min, then the magnetite was collected by magnetic separation and washed with 5 mL of oxygen-free deionized water to remove the possible

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residues of AA or H2O2 from the surface of magnetite. After the addition of 5 mL 1, 10-phenanthroline reagent, which consisted of 1.5 mL water, 2 mL 2 g/L 1, 10-phenanthroline and 1.5 mL acetate buffer, degassed with argon for 30 min before use, the sample tube was shaken on an orbital shaker for 30

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min. Finally, 3 mL of the suspension was filtered and measured for the absorbance at 510 nm (Abs510nm)

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by the UV−vis spectrophotometer. Subsequently, 100 μL of 100 g/L hydroxylamine hydrochloride solution was added into 3 mL of the filtered suspension to reduce Fe3+ into Fe2+, and Abs510nm was

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measured again after another 30 min. The procedures described above were conducted under argon gas protection with great care to avoid the oxidation of Fe(II). The possible degradation intermediates

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of alachlor were probed by gas chromatography−mass spectrometry (GC–MS, Trace 1300, ISQ,

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Thermo) with a DB–5 column (30 m × 0.25 mm × 0.25 μm). Meanwhile, the AA degradation byproducts were determined using liquid chromatography–mass spectrometry (LC–MS, TSQ

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Quantum MAX, Thermo). The sample pretreatment procedures and instrumental parameters for GC– MS and LC–MS analysis were described in the Text S5 of SM [22]. 2.5. Quantification of Fe surface site concentration The total surface iron concentrations were determined by measuring the density of surface hydroxyls since the surface iron would react with water molecules and become coordinated with hydroxyl groups in aqueous solution (Eq 3). In this study, the stoichiometric ratio between surface Fe

site and hydroxyl group was fixed at 1:1, assuming that the surface hydroxyl groups were singly coordinated with Fe. Notably, doubly or triply coordinated hydroxyl groups also exist, but it is still a challenge to experimentally quantify their proportions and thus was not considered here [23]. This approximation would contribute to the uncertainty of kinetic modeling. The surface hydroxyl groups get protonated in acidic pH (Eq 4) and deprotonated under alkaline condition (Eq 5), therefore could be quantified by an acid-base potentiometric titration method [24]. (3)

≡Fe(II,III)OH + H+ → ≡Fe(II,III)OH2+

(4)

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≡Fe(II,III) + H2O → ≡Fe(II,III)OH + H+

≡Fe(II,III)OH + OH- → ≡Fe(II,III)O-

(5)

1 g/L of Fe3O4 was dispersed in 100 mL 0.01 M NaCl and stirred for 3 h at ambient condition to

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reach equilibrium. The pH of the suspension was first adjusted to 3 with 0.1 M HCl, and the volume

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of 0.1 M HCl needed was recorded as Va. Then the suspension was titrated by using 0.1 M NaOH in dropwise manner until pH = 10.5, with the stepwise pH values and volumes of base (Vb) recorded.

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100 mL of 0.01 M NaCl without Fe3O4 was tested as control, following the same procedure. Nitrogen gas was constantly bubbled to eliminate the possible interference of dissolved CO2 gas. The total

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concentration of protons (Ht) introduced into the titration system was calculated by Eq 6: (6)

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Ht = (Ca × Va - CbVb)/(V0 + Va + Vb)

where Ca, Cb and V0 refer to the concentration of acid, base and the initial volume of Fe3O4 suspension,

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respectively. The pH was plotted as a function of Ht to indicate the titration curves of Fe3O4 and the control sample. The Gran functions (Eqs 7 and 8) of the titration curves were plotted versus Vb, then the linear regression analysis on both acidic and alkaline sides were performed, with the intercepts at x-axis indicating the equivalence points in the acidic (Vea) and alkaline (Veb) titration sides [24, 25]. The surface hydroxyl group concentration (Hs) can be obtained according to Eq 9. Gran = (V0 + Va + Vb) × 10-pH (acidic side)

(7)

Gran = (V0 + Va + Vb) × 10-pH-13.8 (alkaline side)

(8)

Hs = [(Veb -Vea)sample × Cb - (Veb -Vea)blank × Cb]/V0

(9)

2.6. Kinetic Modeling Kinetic modeling of the Fe3O4/AA/H2O2 Fenton system was undertaken using the KinTeck Explorer 8.0 software [26], with the rate constants of the elementary reactions within the kinetic model either obtained from the literature or by fitting the experimental data. 2.7. Ecotoxicity assessment

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The ecotoxicity of the detected intermediates toward typical species at three different trophic levels (fish, daphnia, and green algae) was calculated using ECOSAR program version 1.11 developed by USEPA. For acute toxicity, the endpoints were EC50 (the concentration for 50% growth inhibition in

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96 h) for green algae, and LC50 (the concentration inducing 50% death of the organism, 96 h for fish

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and 48 h for daphnia). For chronic toxicity, the endpoint chronic value (ChV) is defined as the geometric mean of the no observed effect concentration (NOEC) and the lowest observed effect

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3. Results and discussion

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concentration (LOEC).

3.1. Ascorbic acid enhanced alachlor removal in Fe3O4/H2O2 system

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The commercially available magnetite (Fe3O4) was used as heterogeneous Fenton catalyst in this

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study. The XRD results of the as-received magnetite sample coincided well with the standard XRD pattern of Fe3O4 (JCPDS card No. 88-315), showing the chemical purity of the magnetite sample (Figure S1a). The SEM images revealed that the magnetite particles were in the shapes of either octahedrons or spheres, with diameters in the range of 100-200 nm (Figure S1b). According to the N2 adsorption and desorption isotherms, the specific surface area of as-received magnetite was calculated to be 5.5 m2/g by using the Brunauer-Emmett-Teller (BET) model (Figure S2).

The heterogeneous Fenton activity of the magnetite sample was evaluated with the alachlor removal. The alachlor removal efficiency of the Fe3O4/H2O2 system was less than 10% within 1 h, whereas the introduction of ascorbic acid greatly enhanced the removal ratio of alachlor to 69.5% within only 20 min (Fig. 1a). As either the Fe3O4/AA or the H2O2/AA systems could hardly degrade alachlor, AA might enhance the degradation of alachlor in the Fe3O4/AA/H2O2 system by accelerating the Fenton reaction. As a weak acid, AA could lower the pH of the system, which might induce iron dissolution and benefit the Fenton reaction. Lower than the initial pH (6.1) of the Fe3O4/H2O2 system, the initial

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pH of the Fe3O4/AA/H2O2 system was 4 ± 0.1 without any adjustment with this pH constant throughout the reaction (Figure S3). To investigate the role of lower pH on Fenton activity, we used H2SO4 instead of AA to obtain Fe3O4/H2O2/H2SO4 system with initial pH of 4.0. However, the degradation of alachlor

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in the Fe3O4/H2O2/H2SO4 systems did not increase in comparison with the pristine Fe3O4/H2O2 system

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(Figure S4), suggesting that the decreased pH caused by AA did not contribute to the better alachlor removal performance of the Fe3O4/AA/H2O2 system.

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In the Fe3O4/AA/H2O2 system, the degradation of alachlor significantly decelerated after 20 min. To check if the depletion of reactants was responsible for this deceleration, we conducted experiments

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by replenishing different reactants at 60 min. The addition of Fe3O4 or H2O2 did not cause any further

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removal of alachlor, whereas the supplement of AA triggered further removal of alachlor (Figure S5). Therefore, the depletion of AA might be responsible for the deceleration of alachlor degradation after

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20 min. To confirm this assumption, we further monitored the concentration profile of AA during the reaction, and found the complete consumption of AA within 20 min in the Fe3O4/AA/H2O2 system (Fig. 1b). Therefore, AA played the indispensable role in the enhanced Fenton degradation of alachlor. It was reported that AA could reductively dissolve iron minerals like hematite and magnetite, resulting in its oxidation [14]. However, the concentration of AA in the Fe3O4/AA/H2O2 system decreased much more quickly than that in the Fe3O4/AA system, and slightly declined in the AA/H2O2 system. The fast

depletion of AA in the Fe3O4/AA/H2O2 system could be attributed to the AA oxidation by •OH formed via Fenton reactions (kAA-OH = 4.1 × 109 – 1.3 × 1010 M-1 s-1) instead of direct oxidation by Fe3O4 or H2O2 [27]. 3.2. AA promoted H2O2 decomposition and •OH production To further confirm the accelerated heterogeneous Fenton reaction by AA, we monitored the decomposition of H2O2 and the production of •OH. In the Fe3O4/H2O2 system, less than 2% of the H2O2 was consumed within 60 min, whereas the introduction of AA greatly increased the

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decomposition percentage of H2O2 to ca. 26% within 20 min, and the decomposition of H2O2 slowed down thereafter due to the exhaustion of AA (Fig. 2a). The decomposition of H2O2 was fitted by pseudo-first order kinetics (Fig. 2b), and the apparent rate constant (k) of H2O2 decomposition

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increased by 174 times from 9.56×10-5 min-1 in Fe3O4/H2O2 system to 1.66×10-2 min-1 in

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Fe3O4/AA/H2O2 system.

Consequently, the production of •OH was also enhanced by AA. Electron paramagnetic resonance

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(EPR) spectra were first employed to confirm the yield of •OH by using 5,5-dimethyl-L-pyrroline-Noxide (DMPO) as a spin trapper. As depicted by Fig. 3a, typical signals of DMPO-•OH spin adduct,

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quartet peaks with relative intensity ratios of 1:2:2:1, were distinguished from the ESR spectra in both

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Fe3O4/H2O2 and Fe3O4/H2O2/AA systems [28]. Obviously, the addition of AA significantly promoted the generation of •OH. We then measured the accumulative concentrations of •OH in both systems by

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monitoring the production of p-hydroxyl benzoic acid from the reaction between •OH and benzoic acid. In the Fe3O4/H2O2 system, the production of •OH was very limited, with approximate 10.5 μmol/L accumulated in 60 min. In the presence of AA, the accumulation of •OH was greatly accelerated especially within the initial 20 min, with the peak value of 241.6 μmol/L obtained in the subsequent 60 min (Fig. 3b). 3.3. AA promoted the surface iron cycle of Fe3O4

We then investigated the mechanism of AA promoting heterogeneous Fenton activity in the Fe3O4/AA/H2O2 system, regarding the surface and homogeneous Fenton processes, respectively. In a typical surface Fenton system, the generation of •OH originated from the reaction between active sites on the surface of catalysts and H2O2, thus the generation rate of •OH (V•OH) could be defined as: V•OH =

d[•OH] dt

= kactive site, H2 O2 [active site][H2 O2 ]

(10)

where kactive site, H2O2 refers to the rate constants between the surficial active sites and H2O2, while [active site] and [H2O2] represent their concentrations. Thus, the rise of V•OH could be attributed to the

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promotion of kactive site, H2O2, and/or the increased concentration of active sites which referred to faction of Fe(II) species on the surface of the magnetite in the Fe3O4/AA/H2O2 system. AA was able to induce

Fe(II), described by the following steps [14, 29, 30]:

≡Fe(II)

→

→

≡Fe(II) + AAox

≡Fe(III) + Fe2+

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≡Fe(III)AA

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≡Fe(III) + AA → ≡Fe(III)AA

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reductive dissolution of various iron oxides like hematite, goethite and magnetite, to produce surface

(11) (12) (13)

First, ascorbic acid formed surface complexes (≡Fe(III)AA) with the iron atoms on the surface of

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iron oxide (≡Fe(III)), followed by the electron transfer process from AA ligands to iron atoms,

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generating surface ferrous species (≡Fe(II)) and oxidized AA molecules (AAox). The ≡Fe(II) will subsequently release dissolved ferrous ion (Fe2+) into the bulk solution phase slowly. Therefore, AA

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could increase the density of ≡Fe(II), the active center for the decomposition of H2O2 in surface Fenton reaction. To confirm this mechanism, we monitored the fractions of ≡Fe(II) to total surface iron species (sum of ≡Fe(II) and ≡Fe(III)) on Fe3O4 by a 1, 10-phenanthroline dissolution method under anoxic condition. In pristine Fe3O4, the initial ≡Fe(II) fraction was determined to be ca. 10%, lower than the theoretical ferrous fraction (33%) in Fe3O4, possibly due to the oxidation of the surface Fe(II) by air during storage (Fig. 4a). In the Fe3O4/H2O2 system, the ≡Fe(II) fraction gradually increased to ca. 30%

after 25 min and then decreased, indicating that the ≡Fe(III) could be slowly reduced by H2O2 to ≡Fe(II) (Eq 14), which subsequently underwent re-oxidation during the Fenton process (Eq 15). In the Fe3O4/AA system, the ≡Fe(II) fraction increased drastically to nearly 100% within only 5 min, and maintained at this high level thereafter. As for the Fe3O4/AA/H2O2 system, the ≡Fe(II) proportion showed similar sharp rise to 97% within the initial 5 min, and decreased gradually after 15 min, coinciding with the depletion of AA. These results suggested that AA was able to promote the rapid reduction of ≡Fe(III) to produce more ≡Fe(II) that acted as the active center for H2O2 decomposition,

≡Fe(III) + H2O2 ≡Fe(II) + H2O2

→ →

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improving surface Fenton efficiency of Fe3O4. ≡Fe(II) + •HO2 + H+

(14)

≡Fe(III) + •OH + OH-

(15)

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X-ray photoelectron spectroscopy (XPS) analysis of the reacted Fe3O4 samples in different systems

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also supports the proposed mechanism. A satellite peak detected at 718.6 eV for the pristine Fe3O4 sample, which was a typical signal produced by the high spin Fe(III) in Fe2O3 or FeOOH [31, 32],

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indicated that ≡Fe(II) underwent oxidation when exposed to ambient air. However, for the magnetite samples collected from Fe3O4/AA/H2O2 and Fe3O4/H2O2 systems, the intensities of the satellite peaks

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decreased significantly, which suggested the regeneration of ≡Fe(II) via the reduction by AA or H2O2

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(Figure S6a). To quantify the surface Fe(II) fractions, the Fe 2p3/2 XPS spectra were fitted by deconvolution according to the Gupta and Sen (GS) multiple peak fitting method [33], with fitting

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parameters like BE (binding energy) and FWHM (full width at half maximum) listed in Table S3. The ≡Fe(II) proportion in the commercial Fe3O4 sample was calculated to be 22.8%, and the ratio was improved to 27.3% and 31.6% after heterogeneous Fenton reaction in the Fe3O4/H2O2 and Fe3O4/AA/H2O2 systems, respectively (Figure S6b), showing similar trends with those in Fig. 4a. The discrepancy between ≡Fe(II) fractions obtained by XPS analysis and those measured by the 1, 10-

phenanthraline extraction method could be explained as follows. XPS reflects the average signals of both surface exposed atoms and those in the bulk phase of Fe3O4 within a depth of ca. 0.5-10 nm. To study the homogeneous Fenton reaction in the Fe3O4/AA/H2O2 system, we also investigated the release of dissolved Fe2+ from Fe3O4 surface. In the Fe3O4/H2O2 system, dissolved Fe2+ or Fe3+ was not detected within 60 min. As expected, the concentration of dissolved Fe2+ in the Fe3O4/AA system increased steadily due to the reductive dissolution effect of AA, and dissolved Fe3+ was still not detected owing to the excess amount of AA and its much lower reduction potential (E0 = 0.06V) than

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Fe3+/Fe2+ (E0 = 0.77V) [34, 35]. In the Fe3O4/AA/H2O2 system, however, the release of Fe2+ was apparently inhibited by the addition of H2O2. Meanwhile, the concentration of dissolved Fe2+ decreased after the depletion of AA at 20 min of reaction, accompanied with the increased

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concentration of Fe3+ through the oxidation of Fe2+ by H2O2 and •OH (Fig. 4c). Therefore, AA could

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also promote the Fe3+/Fe2+ redox cycle of the homogeneous Fenton reaction. Although the concentrations of dissolved iron were low, the contribution of homogeneous Fenton in the

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Fe3O4/AA/H2O2 heterogeneous system could not be neglected because of the efficient Fe3+/Fe2+ cycle in the presence of AA. Obviously, it is not straight forward to assess the individual contributions of

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surface and homogeneous Fenton regarding the dynamic variations of iron, AA and H2O2

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concentrations during the reaction, which calls for quantitative description of the reaction kinetics in the Fe3O4/AA/H2O2 system as explored subsequently.

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3.4. Kinetic modeling

Taking into account the importance of surface Fenton in the Fe3O4/AA/H2O2 system, it is necessary

to figure out the initial concentrations of ≡Fe(II) and ≡Fe(III) for kinetic modeling. Unfortunately, the measurement of surface iron concentrations varied with the extraction time using the 1, 10phenanthraline extraction method, which made it challenging to quantify the exact concentrations of ≡Fe. Nevertheless, the fractions of ≡Fe(II) obtained by this method should be reliable since both ≡Fe(II)

and ≡Fe(III) were extracted from the surface of Fe3O4 simultaneously. Alternatively, the total concentration of ≡Fe was quantified by using the acid-base potentiometric titration method as described in Section 2.5. The titration curves indicated that the Fe3O4 sample had pH buffer capacity in the pH range of 3.5-10, demonstrating the protonation and deprotonation processes during titration (Fig. 5a, Eqs. 3-5). From the Gran plots, the equivalence points in the acidic (Vea) and alkaline (Veb) titration were determined to be 0.750 mL and 0.963 mL for the blank as well as 0.763 mL and 1.090 mL for the 1 g/L Fe3O4 sample (Fig. 5b, Table S3). Then the total concentration of ≡Fe on the Fe3O4

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surface was calculated to be 1.11 × 10-4 mol/L (Eqs. 7-9). According to the proportion (9.25%) of ≡Fe(II) in pristine Fe3O4 determined by the 1, 10-phenanthroline method (Fig. 4a), the initial ≡Fe(II) and ≡Fe(III) concentrations of 1.007 × 10-4 mol/L and 1.027 × 10-5 mol/L were employed in the kinetic

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modelling, respectively.

Table 1. Proposed kinetic model of the Fe3O4/AA/H2O2 heterogeneous Fenton reactions.

Fe(III)-Fe3 O4 + H2 O2 ↔ Fe(II)-Fe3 O4 + HO•2 + H +

H+

2

Fe(II)-Fe3 O4 + H2 O2 →

3

HO• + Alachlor → Alachlorprod

4

Fe(III)-Fe3 O4 + AA ↔ AA-Fe(III)-Fe3 O4

6

Fe(III)-Fe3 O4 + HO• + H2 O b

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5

a

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1

Reaction

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No.

c

H+

References

k1 = 2.5×10-3 M-1 s-1

[36, 37]

k-1 = 2.4×106 M-1 s-1 k2 = 0.38 M-1 s-1

this study

k3 = 7×109 M-1 s-1

this study

k4 = 15 M-1 s-1

this study

k-4 = 9.0×10-3 s-1

AA-Fe(III)-Fe3 O4 → Fe2+ + Fe(III)-Fe3 O4 + AAox AA-Fe(III)-Fe3 O4 + H2 O2 →

Rate constant

d

Fe(III)-Fe3 O4 + HO• +

k5 = 3×10-4 s-1

this study

k6 = 35 M-1 s-1

this study

k7 = 55 M-1 s-1

[36, 38]

k8 = 2.5×10-3 M-1 s-1

[36, 37]

H2 O + AAox H+

7

Fe2+ + H2 O2 →

8

Fe3+ + H2 O2 → Fe2+ + HO•2 + H +

Fe3+ + HO• + H2 O

k-8 = 2.4×106 M-1 s-1

9

H2 O2 + HO• → HO•2 + H2 O

k9 = 3.0×107 M-1 s-1

[36, 37]

10

Fe3+ + HO•2 → Fe2+ + O2 + H +

k10 = 7.7×106 M-1 s-1

[36, 37]

k-10 = 0.1 M-1 s-1 11

Fe2+ + HO• → Fe3+ + OH−

k11 = 3.2×108 M-1 s-1

[36, 37]

12

HO• + HO•2 → H2 O + O2

k12 = 7.5×109 M-1 s-1

[36, 37]

13

HO• + HO• → H2 O2

k13 = 5.2×109 M-1 s-1

[36, 37]

14

Fe3+ + AA → Fe2+ + AAox

k14 = 4.5×103 M-1 s-1

This study

15

HO• + AA → AAox

k15 = 1.0×1010 M-1 s-1

This study

a

Fe(II)-Fe3O4 and Fe(III)-Fe3O4 represent surficial Fe(II) and Fe(III) of Fe3O4, respectively; Alachlorprod represent the products following interactions of alachlor with HO•; c AA represents ascorbic acid, while AA-Fe(III)-Fe3O4 indicates the surface complex formed between AA and Fe(III)-Fe3O4; d AAox represents the oxidized forms of ascorbic acid following the interaction of ascorbic acid with Fe3O4;

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b

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Subsequently, a first principle kinetic model of the Fe3O4/AA/H2O2 heterogeneous Fenton system

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was developed by fitting the experimental data and the rate constants obtained from literature. Table 1 summarizes the elementary reactions involved. The key components of this model include: (i) The

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reactions in the Fe3O4/H2O2 system, mainly composed of the reduction of ≡Fe(III) by H2O2, the surface Fenton reaction and the oxidation of alachlor by •OH (Reaction 1-3 in Table 1). The dissolution of

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Fe2+ from the Fe3O4 surface and the subsequent homogeneous Fenton reactions were not included here

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since Fe2+ or Fe3+ concentrations did not significantly increase in the Fe3O4/H2O2 system (Fig. 4b); (ii) The surface reactions involved in the Fe3O4/AA system, like the formation of surface complex between

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AA and Fe3O4 (≡Fe(III)AA, Reaction 4 in Table 1), and the subsequent reductive dissolution of Fe2+ (Reaction 5, which was the overall reaction of Eqs. 12-13 for simplicity of the model); (iii) The surface Fenton reaction between ≡Fe(III)AA complex and H2O2 in the Fe3O4/AA/H2O2 system (Reaction 6 in Table 1); (iv) The classical homogeneous Fenton process in the Fe3O4/AA/H2O2 system (Reaction 713 in Table 1), the kinetic rate constants of which were well established in previous studies; and (v)

The homogeneous reactions of AA in the Fe3O4/AA/H2O2 Fenton system, i.e. the oxidation of AA by dissolved Fe3+ and •OH, respectively (Reaction 14 and 15 in Table 1). The rate constants of the elementary reactions in the proposed model were determined as follows. (i) It is difficult to figure out the exact rate constants of ≡Fe(III) reduction by H2O2 and its reverse reaction in the Fe3O4/H2O2 system (k1 and k-1). Instead, the rate constants of the corresponding homogeneous reactions (k8 and k-8) were adopted in our current model, which may contribute to the inaccuracy of the modelling results. However, the simulation results of the whole model did not change

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much when altering k1 and k-1 over a wide range (several magnitudes), indicating that the reduction of ≡Fe(III) by H2O2 was less important in the Fe3O4/AA/H2O2 system. The rate constant of Reaction 2 (k2 = 0.38 M-1 s-1) was deduced from the decomposition of H2O2 in the Fe3O4/H2O2 system (Fig. 2a).

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This value was much smaller than that of the homogeneous Fenton reaction, probably due to the

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limitation of diffusion and adsorption of H2O2 at the interface, as well as the desorption of H2O2 decomposition products (•OH, •HO2 and OH-) from the active sites. The rate constant of alachlor

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oxidation by •OH (k3 = 7 × 109 M-1 s-1) was acquired by fitting the alachlor degradation data in the Fe3O4/H2O2 system (Fig. 1a), and this value was in accordance with that reported by Haag and Yao

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[39], demonstrating the reliability of our fitting process. (ii) To obtain the kinetic rate constants of the

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surface complexation between AA and ≡Fe(III) (Reaction 4), the equilibrium constant (Kads) was first determined by conducting AA adsorption experiments with the literature method (Text S6) [29, 30],

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resulting in a Kads of ca. 1670 M-1 derived from the double reciprocal plots of Langmuir adsorption isotherm (Figure S7), and then the rate constants of AA adsorption (k4 = 15 M-1 s-1) on the Fe3O4 surface and AA desorption (k-4 = 9.0 × 10-3 s-1) were determined together with the rate constant of Fe2+ release from ≡Fe(III)AA complex (k5 = 3.0 × 10-4 s-1), based on the best fit to the experimental data of AA consumption and Fe2+ dissolution in the Fe3O4/AA system, respectively (Fig. 1b and Fig. 4c). (iii) Subsequently, the rate constant of the surface Fenton reaction between ≡Fe(III)AA complex and H2O2

(k6 = 35 M-1 s-1, Reaction 6) was obtained by fitting both the data of AA depletion and H2O2 decomposition in the Fe3O4/AA/H2O2 system, and this value was enhanced by ca. 92 times when compared with the rate constant of the surface Fenton reaction (k2) in the Fe3O4/H2O2 system, confirming that AA promoted the surface Fenton by increasing the concentration of active sites on the Fe3O4 surface, and improving the rate constant. (iv) The rate constants of the elementary reactions in classical homogeneous Fenton system (Reaction 7-13) were well established by previous studies and thus the literature reported values were directly adopted in our present kinetic model. (v) The rate

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constant (k14 = 4.5 × 103 M-1 s-1) of Fe3+ reduction by AA in the homogeneous solution was determined from the best fit of the dissolved Fe2+ and Fe3+ concentrations during reaction in the Fe3O4/AA/H2O2 system (Fig. 4c). In the Fe3O4/AA/H2O2 system, AA and alachlor would compete for the •OH produced

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by Fenton reaction. Therefore, when other rate constants involved were fixed, the rate constant of AA

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oxidation by •OH (k15 = 1.0 × 1010 M-1 s-1) could be deduced from the alachlor degradation data in the

× 109 – 1.3 × 1010 M-1 s-1) [27].

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Fe3O4/AA/H2O2 system (Fig. 1a), and this value is also consistent with those reported previously (4.1

The superoxide radical produced from Reaction 1 could subsequently reduce ≡Fe(III) to ≡Fe(II) at

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relatively higher rate constant than H2O2 to accelerate the iron cycle on the Fe3O4 surface (Eq 16).

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Besides, superoxide radical could also degrade alachlor in the Fe3O4/AA/H2O2 system (Eq 17). To evaluate the importance of these processes, we first compared the alachlor degradation profiles which

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were predicted by the kinetic models including and excluding Eq 16, and assuming their rate constants were identical to that of Reaction 10 in Table 1. However, the difference between the results simulated by the two models were negligible, both in the Fe3O4/H2O2 and Fe3O4/AA/H2O2 systems (Figure S8), suggesting that the reduction of ≡Fe(III) by superoxide radical (Eq 16) was not crucial in the Fe3O4/AA/H2O2 system. To confirm the contribution of superoxide radical for alachlor degradation in the Fe3O4/AA/H2O2 system (Eq 17), ROSs quenching experiments were conducted. It was found that

the addition of catalase and tert-butyl alcohol (TBA) completely inhibited the alachlor degradation, whereas superoxide dismutase (SOD) had little impact on the performance (Figure S9). These results suggested that •OH was the predominant ROSs for the alachlor removal, and the contribution of superoxide radical was negligible. Therefore, the reactions of Eqs 16-17 were not included in the current kinetic model. ≡Fe(III) + •HO2

→

≡Fe(II) + O2 + H+

(16)

alachlor + •HO2

→

product

(17)

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We could thus figure out the rate limiting steps of the heterogeneous Fenton systems based on the rate constants acquired in Table 1. In the Fe3O4/H2O2 system (Reaction 1-3), Reaction 1 is the rate limiting step because k1 is 9 magnitudes lower than k-1, resulting in the insufficient generation of ≡Fe(II)

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for the subsequent Fenton reactions. In the Fe3O4/AA/H2O2 system, the rate constants of Reactions 4-

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8 are much lower than the others, and might be potential rate determining elementary reactions. To quantify the roles of these reactions in determining the overall Fenton reaction rate, we increased k4-

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k8 by 2 times separately in the model, and compared the predicted alachlor degradation. The doubled k4 improved alachlor degradation the most, followed by k5 and k7, whereas the increased k6 and k8

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caused little change of alachlor removal (Figure S10). These results indicated that the adsorption of

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AA on the Fe3O4 surface to form electron transfer complex was the rate determining step in the Fe3O4/AA/H2O2 heterogeneous Fenton system.

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The resultant kinetic model could well describe the experimental data. To further check the

reliability of this model, we conducted the alachlor removal experiments with varying doses of the reactants such as Fe3O4, AA and H2O2, and the prediction by the model basically matched the corresponding experimental data (Fig. 6). Nevertheless, some discrepancies still existed between the experimental data and the modeling predicted values, which could be explained as follows. First, the reactions occurred at the surface of Fe3O4 are very complicated, including the diffusion of substrates

to the interfacial area, the adsorption on Fe3O4 surface, the complexation of substrates with surface active sites, the electron transfer processes, the generation of reactive intermediate species, the desorption of intermediates and products, and the dissolution of surface iron species, etc. Whereas in this study it was apparently unrealistic to include all the above-mentioned reactions since it would make the kinetic model too complicated to be addressed. Therefore, approximation was made to simplify the model, which would introduce uncertainties to the simulation results. Second, the heterogeneous Fe3O4 catalysts might aggregate during the reaction, resulting in less active sites

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available for AA complexation and surface Fenton reaction than those in the model. Third, some of the rate constants in our proposed kinetic model were directly adopted from literature values, which might bring about discrepancies due to the inconsistent experimental conditions between our case and

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kinetics of heterogeneous Fenton system in the future.

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literature. Therefore, scope of refinement to the model still exists for a more precise modeling of the

3.5. Contribution of surface and homogeneous Fenton reactions to alachlor degradation

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The individual contributions of surface and homogeneous Fenton reactions in the Fe3O4/AA/H2O2 system were then quantified by modifying the above kinetic model (Table S5). To discriminate the

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•OH generated by the surface and homogeneous Fenton processes, they were denoted as •OHsurface and

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•OHhomo, respectively. Consequently, the alachlor degraded by •OHsurface and •OHhomo were indicated as Alachlorsurface and Alachlorhomo, and identical rate constants were employed for the elementary

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reactions involved with •OHsurface and •OHhomo. Fig. 7 presents the model predicted time profile of Alachlorsurface and Alachlorhomo, as well as the contribution of surface Fenton to the overall alachlor removal (Alachlorsurface/(Alachlorsurface + Alachlorhomo)). Initially, the removal of alachlor was dominated by the surface Fenton reaction, taking up more than 85% before 10 min. With the increase of dissolved iron ions, the homogeneous Fenton reaction accelerated, corresponding to the gradual decrease of the surface Fenton fraction. Both surface and homogeneous Fenton reactions decelerated

sharply after 20 min owing to the depletion of AA. Nonetheless, surface Fenton still accounted for 62.6% of the overall alachlor removal at the end of the reaction (60 min), suggesting the main contribution of surface Fenton process in the Fe3O4/AA/H2O2 heterogeneous system. 3.6. Reusability of Fe3O4 To assess the reusability of the Fe3O4 sample as a heterogeneous Fenton catalyst in the Fe3O4/AA/H2O2 system, the Fe3O4 after reaction was separated and subjected to more cycles of test, fed with fresh AA, H2O2 and alachlor. The recycled Fe3O4 still exhibited high activity for the removal

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of alachlor after 4 successive runs (Fig. 8a). The stability of magnetite could be explained by the limited Fe leaching during the reaction of Fe3O4/AA/H2O2 system, since the ≡Fe(II) generated via AA reduction was compensated by the H2O2-mediated in situ oxidation. This coupled redox processes

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caused the iron cycle on the Fe3O4 surface, and thus helped maintain the surface structure and bulk

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composition of Fe3O4. 3.7. Transformation mechanism of alachlor and AA

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The dechlorination of alachlor in the Fe3O4/AA/H2O2 Fenton system was first studied by employing ion chromatography. The Cl- concentrations increased sharply to 0.054 mmol/L within 20 min, and

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leveled off thereafter, consistent with the removal trend of alachlor (Fig. 8b). The dechlorination ratio,

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defined as the ratio of released Cl- versus the total Cl contained in the removed alachlor, was 98.2% at 30 min. This suggested that dechlorination was an essential step during the transformation of

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alachlor in the Fe3O4/AA/H2O2 system. Moreover, considering that the substituted chlorine often contributes to the toxicity of pollutants, the nearly complete dechlorination indicated the detoxification of alachlor by Fe3O4/AA/H2O2 system. Possible degradation intermediates of alachlor in the Fe3O4/AA/H2O2 Fenton system were also identified using gas chromatograph-mass spectroscopy (GC-MS). The identified intermediates of alachlor included both the dealkylation and dechlorination products, such as 2-chloro-N-(2-ethyl-6-

methylphenyl)-N-(methoxymethyl) acetamide, 2-chloro-N-(2-ethylphenyl) acetamide, 2-chloro-N(2,6-diethylphenyl) acetamide, 1,3-diethyl-2-isocyanatobenzene, 2,6-diethyl benzenamine, and 1,3diethyl-2-nitrosobenzene, etc (Table S6). The possible degradation pathways of alachlor in the Fe3O4/AA/H2O2 Fenton system was proposed (Scheme 1). The first step of alachlor degradation mainly involved with the cleavage of the C-N single bond, resulting in the formation of intermediate 1, which could be further oxidized through dechlorination or dealkylation to generate compounds 2 and 3. The intermediate would undergo stepwise dealkylation of the ethyl group to give compounds 4

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and 5. The –NH2 group of compound 3 was oxidized by •OH into –N=O (compound 6) or –NO2, and the ethyl group transformed to –CHO (compound 7). Meanwhile, the ortho- ethyl group of alachlor could also be dealkylated at the initial stage of degradation, producing compound 8, which would

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further transform to 9 after the C–N bond broke up. The above intermediates could be eventually

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oxidized into smaller organic molecules like oxalic acid and formic acid after the cleavage of the benzene ring, or mineralized into carbon dioxide, nitrate, and chloride. The acute and chronic toxicity

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endpoint values of the intermediates showed similar trends, i.e., toxicity of the intermediates to fish or daphnia did not change much compared with alachlor, but the toxicity to green algae decreased

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significantly, indicating that alachlor can be gradually detoxified by Fe3O4/AA/H2O2 Fenton treatment

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(Figure S11).

Since AA was also degraded in the Fe3O4/AA/H2O2 Fenton system, the intermediates of AA

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transformation were also identified by HPLC-MS-MS (Table S7). The hydroxyl group adjacent to the five-membered ring of AA could be oxidized to carbonyl group either by •OH or Fe(III), resulting in dehydroascorbic acid (m/z = 173). •OH could also attack the side chain of AA molecules, producing 3,4-dihydroxy-5-(hydroxymethyl)furan-2(5H)-one (m/z = 144), which could be further oxidized to 5(hydroxymethyl)furan-2,3,4(5H)-trione (m/z = 143), 5-hydroxyfuran-2,3,4(5H)-trione (m/z = 129) and furan-2,3,4(5H)-trione (m/z = 113). The above intermediates would undergo further ring-opening

degradation, generating byproducts such as 2,3-diketogulonic acid (m/z = 191), 2, 3, 4trihydroxybutanoic acid (m/z = 135), oxalic acid, acetic acid and formic acid (Scheme 2). Ecotoxicity evaluation showed that most of the intermediates were nontoxic since their acute and chronic toxicity endpoint values were comparable with ascorbic acid. Though the toxicity of product 5 slightly increased, it would undergo further degradation and be detoxified (Figure S12). This implied that the application of AA to enhance the heterogeneous Fenton reaction was environmentally benign and would cause no secondary pollution.

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3.8. Removal of other pollutants and impacts of coexisting inorganic and organic matters We further employed the Fe3O4/AA/H2O2 system to remove other organic pollutants, including organic dyes (methylene blue), phenols (4-chloropenol and bisphenol A), and pharmaceuticals

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(ranitidine and propranolol). As expected, all these pollutants could be effectively degraded within 60

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min, indicating the wide applicability of the Fe3O4/AA/H2O2 Fenton system for organic pollution removal (Figure S13). Besides, we evaluated the impacts of coexisting inorganic and organic matters

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on the pollutant removal performance of the Fe3O4/AA/H2O2 system. Chloride had negligible impact on the alachlor removal (Figure S14a), and HCO3- showed slight inhibition effect only at high

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concentration (10 mM, Figure S14b). Interestingly, H2PO4- of low concentration (1 mM) promoted the

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degradation of alachlor, but 10 mM of H2PO4- decelerated alachlor removal, possibly due to the competitive adsorption of H2PO4- and AA on the Fe3O4 surface. Fortunately, good removal efficiency

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could still be achieved with prolonged reaction time (Figure S14c). The presence of 1 mg/L humic acid as a representative of natural organic matters had little impact on the alachlor removal. When the concentration of humic acid increased to 10 mg/L, slight inhibition of alachlor degradation was observed (Figure S14d). These results suggested that the Fe3O4/AA/H2O2 heterogeneous Fenton system was highly promising for natural water or waste water treatment. 4. Conclusions

In this study, we have demonstrated that the Fenton removal of alachlor in the Fe3O4/H2O2 system can be drastically accelerated by AA via the rapid iron redox cycle both on the surface of Fe3O4 and in the solution. To elucidate the individual contributions of surface and homogeneous Fenton processes to the alachlor removal, a conceptual kinetics model composed of (i) the surface complexation and redox reactions between AA and surface iron, (ii) the surface Fenton reactions, and (iii) the classical homogeneous Fenton reactions, was developed based on fitting of the experimental data, and the model prediction was in good accordance with experimental results. The individual contributions of

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surface and homogeneous Fenton to alachlor degradation were simulated with using this model, and the main role of surface Fenton was revealed. The kinetic modeling in this work could offer the

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understanding of the heterogeneous Fenton mechanism.

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quantitative description of the reactions in heterogeneous Fenton systems, contributing to a better

Declaration of interests

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☒ The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper.

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Acknowledgments

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☐The authors declare the following financial interests/personal relationships which may be considered as potential competing interests:

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This research was financially supported by The National Key Research and Development Program of China (2018YFC1800801 and 2018YFC1802003), Natural Science Funds for Distinguished Young Scholars (21425728), National Natural Science Foundation of China (41601543, 21936003 and 41807349), Project funded by China Postdoctoral Science Foundation (2017M620327 and 2018T110782), the Fundamental Research Funds for the Central Universities (CCNU17XJ005), the program of China Scholarship Council (201706775080), and Science Funds for Outstanding Postdocs of Hubei Province, China (Z13).

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Fig. 1. (a) Time profile of alachlor degradation in the Fe3O4/AA/H2O2 Fenton system and the control

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systems. (b) Depletion of ascorbic acid in the Fe3O4/AA/H2O2 and Fe3O4/H2O2 systems. The dots show the experimental results, and the thick lines represent the model predictions. The initial concentrations of alachlor, hydrogen peroxide, ascorbic acid and magnetite were 20 mg/L, 1 mmol/L,

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0.5 mmol/L, and 1 g/L, respectively; the initial pH was 4 without adjustment.

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Fig. 2. (a) Time profile of the decomposition of H2O2 in Fe3O4/AA/H2O2 and Fe3O4/H2O2 systems. (b) The corresponding pseudo-first order kinetic curves of H2O2 decomposition. The dots and the thick

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lines show the experimental results and model predictions, and thin lines indicate the linear regression.

Fig. 3. (a) The EPR signals of the DMPO adducts in different Fenton systems at 5 min of reaction. (b)

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Accumulative concentrations of hydroxyl radical (•OH) in Fe3O4/AA/H2O2 and Fe3O4/H2O2 systems,

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determined by using benzoic acid as a radical probe.

Fig. 4. (a) The surface Fe(II) fractions in different Fe3O4 systems. (b) The release of Fe2+ and Fe3+ into

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the bulk solution plotted as a function of reaction time in different systems. The dots and thin lines

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show the experimental results. The thick lines in panel c represent the model predictions.

Fig. 5. (a) The titration curves and (b) the Gran plots of 1 g/L Fe3O4 suspension and 0.01 M NaCl solution (control).

Fig. 6. The impacts on alachlor degradation of (a) magnetite dose, (b) AA concentration and (c) H2O2 dose. The dots represent experimental results and the thick lines indicate the model predictions.

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Conditions: (a) [alachlor]0 = 20 mg/L, [AA]0 = 0.5 mM, [H2O2]0 = 1 mM. (b) [alachlor]0 = 20 mg/L,

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[Fe3O4]0 = 1 g/L, [H2O2]0 = 1 mM. (c) [alachlor]0 = 20 mg/L, [Fe3O4]0 = 1 g/L, [AA]0 = 0.5 mM.

Fig. 7. The model predicted time profile of alachlor removed by surface Fenton (blue open square)

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and homogeneous Fenton (red open circle) in the Fe3O4/AA/H2O2 system, as well as the contribution

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of surface Fenton to the overall alachlor removal (green line). The initial concentrations of alachlor, H2O2, AA and Fe3O4 were 20 mg/L, 1 mmol/L, 0.5 mmol/L and 1 g/L, respectively.

Fig. 8. (a) Reusability of the Fe3O4 as Fenton catalyst for alachlor degradation. (b) The evolution of

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Cl- concentrations in the alachlor degradation process by Fe3O4/AA/H2O2 Fenton system.

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Scheme 1. Possible degradation pathways of alachlor in the Fe3O4/AA/H2O2 Fenton system.

Scheme 2. The possible degradation pathways of ascorbic acid in the Fe3O4/H2O2/AA system.

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