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Applied Catalysis B: Environmental

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

Hongwei Suna,b, Guihong Xiea, Di Hec,*, Lizhi Zhanga,*a Key Laboratory of Pesticide & Chemical Biology of Ministry of Education, Institute of Environmental & Applied Chemistry, College of Chemistry, Central China NormalUniversity, Wuhan 430079, ChinabDepartment of Chemical Engineering and Division of Environmental Science and Engineering, Pohang University of Science and Technology (POSTECH), Pohang 37673,Republic of KoreacGuangdong Key Laboratory of Environmental Catalysis and Health Risk Control, Institute of Environmental and Ecological Engineering, Guangdong University ofTechnology, Guangzhou, 510006, China

A R T I C L E I N F O

Keywords:Ascorbic acidSurface FentonMagnetiteIron redox cycleKinetics

A B S T R A C T

In this study we constructed a heterogeneous Fenton system with ascorbic acid (AA), magnetite (Fe3O4) andH2O2 for the alachlor degradation, aiming to clarify the heterogeneous Fenton mechanism. The addition of AAcould significantly accelerate the Fenton reaction by promoting the surface Fe(III)/Fe(II) redox cycle (iron cycle)of Fe3O4. A kinetic model was successfully developed to quantitatively describe the complicated reactions in theFe3O4/AA/H2O2 system. We thus employed this model to identify the individual contributions of surface andhomogeneous Fenton reactions to the overall alachlor degradation in the Fe3O4/AA/H2O2 system, and found thesurface Fenton reaction was mainly responsible for the alachlor degradation with more than 62.6% of con-tribution. This work offers a new strategy to improve the heterogeneous Fenton activity via promoting surfaceFenton reaction, and sheds light on the possibility to quantitatively describe and predict the heterogeneousFenton processes with first principle kinetic models.

1. Introduction

Fenton reaction is an efficient advanced oxidation process (AOP) forthe remediation of organic pollutants [1]. However, the wide applica-tion of traditional homogeneous Fenton system is still restricted bydrawbacks like narrow working pH (2.0–3.5) and mass iron sludge [2].One solution to overcome these shortcomings is to develop hetero-geneous Fenton systems [3,4]. Iron bearing minerals such as hematite(α-Fe2O3) and goethite (α-FeOOH) are abundant in earth crust, in-expensive and environmentally benign, thus often serve as hetero-geneous Fenton catalysts [5,6]. In comparison with hematite and goe-thite, magnetite (Fe3O4) might be more attractive because of itsintrinsic Fe(II) sites which act as active centers for H2O2 decomposition,and better electron transfer capacity resulting from its inverse spinelstructure [6]. Nonetheless, the Fenton efficiency of Fe3O4 is still un-satisfactory for practical application and need further improvement.

Fe(III)+H2O2→ Fe(II)+ %HO2+H+ (1)

Fe(II) +H2O2→ Fe(III)+ %OH+OH− (2)

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 to Fe(II)by H2O2 (Eq. (1)) is approximately 4 orders of magnitude lower thanthat of Fe(II) oxidation by H2O2 (Eq. (2)), resulting in rapid depletion ofFe(II) during the reaction [6,8]. To enhance the efficiency of hetero-geneous Fenton, researchers employed chelating or reducing reagentsto promote the Fe(III)/Fe(II) redox cycle (iron cycle). For example, ourgroup found that hydroxylamine could improve the surface Fentondegradation rate of alachlor by 46 times in goethite/H2O2 systemwithout any release of iron ions into solution, via accelerating the ironcycle on the goethite surface [7]. In contrast, other agents like succi-nate, citrate, oxalate and ethylene-diaminetetraacetic acid (EDTA) in-creased the heterogeneous Fenton performance along with significantrelease of iron ions [9], which thus involved both surface and homo-geneous Fenton processes. However, the individual contributions ofsurface and homogeneous Fenton reactions are not clear yet, goingagainst a better understanding of heterogeneous Fenton mechanism.

Normally, the contribution of homogeneous Fenton process wasassessed by conducting control experiments with dissolved Fe2+ orFe3+ of the same concentrations as those released by heterogeneous

https://doi.org/10.1016/j.apcatb.2019.118383Received 27 September 2019; Received in revised form 23 October 2019; Accepted 2 November 2019

⁎ Corresponding authors.E-mail addresses: di.he@gdut.edu.cn (D. He), zhanglz@mail.ccnu.edu.cn (L. Zhang).

Applied Catalysis B: Environmental 267 (2020) 118383

Available online 05 November 20190926-3373/ © 2019 Elsevier B.V. All rights reserved.

T

Fenton catalysts. Obviously, this approach could only offer an ap-proximated assessment of real cases, since the concentrations of ironions, hydrogen peroxide, and chelating/reducing reagents changeddynamically during the heterogeneous Fenton processes. So it is ne-cessary to establish quantitative description of the kinetics in the het-erogeneous Fenton system for more precise assessment. Moreover, thedevelopment of a first principle kinetic model, which is composed of theessential elementary reactions to simulate the heterogeneous Fentonsystem, will enable more detailed understanding about the thermo-dynamics and kinetics, such as the rate constants of surface andhomogeneous Fenton reactions, the rate-limiting steps, and the equili-brium constants of surface complex. In addition, such a model is able topredict the consumption of reactants, the production of reactive speciesand the degradation of pollutants [10], and thus could benefit the op-timization of experimental parameters in the heterogeneous Fentonsystems. Unfortunately, the mechanism or kinetics of the heterogeneousFenton reactions were often experimentally and qualitatively in-vestigated [7,11,12], with very limited quantitative comprehension ofthe detailed mechanistic information of heterogeneous Fenton systems.

Different from chelating/reducing reagents such as EDTA or hy-droxylamine which are toxic and may cause secondary pollution, as-corbic acid (AA), also known as vitamin C, is an eco-friendly reducingand chelating agent widely found in plants, animals and human, andalso widely used as dietary supplement or medicine [11]. Our grouppreviously demonstrated that the addition of AA significantly ac-celerated the iron cycle in hematite and Fe@Fe2O3 nanowires hetero-geneous Fenton systems [11,13], and AA could also induce reductivedissolution of iron oxides [14]. Considering that magnetite possessesmore surface Fe(II) than hematite, and is easily available and moreinexpensive than Fe@Fe2O3 nanowires, herein we constructed theFe3O4/AA/H2O2 system to unveil the contributions of surface andhomogeneous Fenton to the removal of alachlor, a recalcitrant herbi-cide [15,16]. The mechanisms involved with the consumption of AAand H2O2, the generation of reactive species, the iron cycle, and therelease of iron ions were investigated in detail. On the basis of ex-perimental data, we developed a kinetic model to quantitatively de-scribe the Fe3O4/AA/H2O2 system. The stability of magnetite, the de-gradation pathways of alachlor and AA, as well as their eco-safety werealso studied carefully.

2. Materials and methods

2.1. Chemicals and materials

Magnetite (Fe3O4) was purchased from Shanghai Aladdin Bio-ChemTechnology Co., Ltd. H2O2 (30 wt %), L-ascorbic acid (AA), sodiumacetate, benzoic acid, 4-hydroxyl benzoic acid, ethyl alcohol, sodiumhydroxide, tert-butyl alcohol, iso-propanol, 1,10-phenanthroline, and2,2′-bipyridine were of analytical grade and bought from SinopharmChemical Reagent Co. Ltd., China. Dichloromethane, methanol andacetonitrile were of HPLC grade and supplied by Merck KGaA.Hydroxylamine hydrochloride (99.9%), formic acid, acetic acid, andalachlor were purchased from Sigma-Aldrich. All the solutions wereprepared in deionized water throughout the experiments.

2.2. Materials characterization

The powder X-ray diffraction (XRD, D8 Advance, Bruker) analysiswas performed with a D/Max-IIIA X-ray diffractometer, using a Cu Kαsource (λ=0.15418 nm). The morphology of the as-received com-mercial magnetite sample was examined by scanning electron micro-scopy (SEM, 6700-F, JEOL). The nitrogen adsorption-desorption iso-therm of the Fe3O4 was obtained using a Micromeritics Tristar 3000instrument at 77 K, and the specific surface area (SSA) was calculatedbased on the Brunauer–Emmett–Teller (BET) model.

2.3. Alachlor degradation experiments

Batch trials of alachlor degradation were conducted in 100mLconical flasks fixed on an orbital shaker at 250 rpm and ambient tem-perature. Briefly, the reaction systems consisted of 50mL of alachlorsolution (20mg/L), 0.05 g of Fe3O4, 0.5mL of 50mmol/L AA solution,and the reaction was triggered by adding 0.5 ml of 100mmol/L H2O2

stock solution. At predetermined intervals, 900 μL of the reaction so-lution was withdrawn and mixed with 100 μL of ethanol to stop thereaction, and the mixture was filtered through 0.22 μm syringe filtermembranes to remove the suspended solids before high performanceliquid chromatography (HPLC) analysis. The initial pH value of theFe3O4/AA/H2O2 system was 4.0 without adjustment. The catalyst afterreaction was collected by magnetic separation, carefully washed withdeionized water and ethanol, and finally dried in a vacuum ovenovernight to test the reusability of Fe3O4.

2.4. Analytical methods

The concentration of alachlor was determined with the methoddescribed in our previous reports and the details described in the sup-plementary material (SM Text S1) [13]. H2O2 concentration was mea-sured using the iodide colorimetric method [17], while AA concentra-tions were determined using a modified 2,2′-bipyridine method with aUV–vis spectrometer [18]. Notably, H2O2 and AA showed mutual in-terference of the colorimetric processes in the mixed Fe3O4/H2O2/AAsystem, but fortunately, the interference was dose-dependent and thusthe exact concentrations of H2O2 and AA can be obtained throughnormalization and calibration, the details of which were provided in SMText S2, Tables S1 and S2. Electron spin resonance (ESR) spectra wererecorded with a JES FA 200 X-band spectrometer (JEOL, Japan). Ac-cumulative %OH was probed by the well-defined reaction betweenbenzoic acid (BA) and %OH to produce 4-hydroxybenzoic acid (4-HBA),which could be easily quantified by the HPLC method [19]. The ele-mental chemical states on magnetite surface before and after the Fentonreaction 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-phenan-throline colorimetric method [20]. For the measurement of dissolvedFe3+, hydroxylamine hydrochloride was used for the reduction of Fe3+

into Fe2+ for subsequent 1,10-phenanthroline measurement (SM TextS4). Surface Fe(II) and Fe(III) concentrations were measured by mod-ifying the method described by He et al. [21]. Briefly, 10mL of thereaction suspension was sampled at fixed intervals and saturated withargon gas for 5min, then the magnetite was collected by magnetic se-paration and washed with 5mL of oxygen-free deionized water to re-move the possible residues of AA or H2O2 from the surface of magnetite.After the addition of 5mL 1, 10-phenanthroline reagent, which con-sisted of 1.5 mL water, 2 mL 2 g/L 1,10-phenanthroline and 1.5mLacetate buffer, degassed with argon for 30min before use, the sampletube was shaken on an orbital shaker for 30min. Finally, 3 mL of thesuspension was filtered and measured for the absorbance at 510 nm(Abs510nm) by the UV–vis spectrophotometer. Subsequently, 100 μL of100 g/L hydroxylamine hydrochloride solution was added into 3mL ofthe filtered suspension to reduce Fe3+ into Fe2+, and Abs510nm wasmeasured again after another 30min. The procedures described abovewere conducted under argon gas protection with great care to avoid theoxidation of Fe(II). The possible degradation intermediates of alachlorwere probed by gas chromatography–mass spectrometry (GCeMS,Trace 1300, ISQ, Thermo) with a DB–5 column(30m×0.25mm×0.25 μm). Meanwhile, the AA degradation by-products were determined using liquid chromatography–mass spectro-metry (LCeMS, TSQ Quantum MAX, Thermo). The sample pretreat-ment procedures and instrumental parameters for GCeMS and LCeMSanalysis were described in the Text S5 of SM [22].

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2.5. Quantification of Fe surface site concentration

The total surface iron concentrations were determined by measuringthe density of surface hydroxyls since the surface iron would react withwater molecules and become coordinated with hydroxyl groups inaqueous solution (Eq. (3)). In this study, the stoichiometric ratio be-tween surface Fe site and hydroxyl group was fixed at 1:1, assumingthat the surface hydroxyl groups were singly coordinated with Fe.Notably, doubly or triply coordinated hydroxyl groups also exist, but itis still a challenge to experimentally quantify their proportions and thuswas not considered here [23]. This approximation would contribute tothe uncertainty of kinetic modeling. The surface hydroxyl groups getprotonated in acidic pH (Eq. (4)) and deprotonated under alkalinecondition (Eq. (5)), therefore could be quantified by an acid-base po-tentiometric titration method [24].

^Fe(II,III)+H2O→^Fe(II,III)OH+H+ (3)

^Fe(II,III)OH+H+→^Fe(II,III)OH2+ (4)

^Fe(II,III)OH+OH−→^Fe(II,III)O- (5)

1 g/L of Fe3O4 was dispersed in 100mL 0.01M NaCl and stirred for3 h at ambient condition to reach equilibrium. The pH of the suspensionwas first adjusted to 3 with 0.1 M HCl, and the volume of 0.1 M HClneeded was recorded as Va. Then the suspension was titrated by using0.1 M NaOH in dropwise manner until pH=10.5, with the stepwise pHvalues and volumes of base (Vb) recorded. 100mL of 0.01M NaClwithout Fe3O4 was tested as control, following the same procedure.Nitrogen gas was constantly bubbled to eliminate the possible inter-ference of dissolved CO2 gas. The total concentration of protons (Ht)introduced into the titration system was calculated by Eq. (6):

Ht= (Ca×Va− CbVb)/(V0+Va+Vb) (6)

where Ca, Cb and V0 refer to the concentration of acid, base and theinitial volume of Fe3O4 suspension, respectively. The pH was plotted asa function of Ht to indicate the titration curves of Fe3O4 and the controlsample. The Gran functions (Eqs. (7) and (8)) of the titration curveswere plotted versus Vb, then the linear regression analysis on bothacidic and alkaline sides were performed, with the intercepts at x-axisindicating 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 un-dertaken using the KinTeck Explorer 8.0 software [26], with the rate

constants of the elementary reactions within the kinetic model eitherobtained from the literature or by fitting the experimental data.

2.7. Ecotoxicity assessment

The ecotoxicity of the detected intermediates toward typical speciesat three different trophic levels (fish, daphnia, and green algae) wascalculated using ECOSAR program version 1.11 developed by USEPA.For acute toxicity, the endpoints were EC50 (the concentration for 50%growth inhibition in 96 h) for green algae, and LC50 (the concentrationinducing 50% death of the organism, 96 h for fish and 48 h for daphnia).For chronic toxicity, the endpoint chronic value (ChV) is defined as thegeometric mean of the no observed effect concentration (NOEC) and thelowest observed effect concentration (LOEC).

3. Results and discussion

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

The commercially available magnetite (Fe3O4) was used as hetero-geneous Fenton catalyst in this study. The XRD results of the as-re-ceived magnetite sample coincided well with the standard XRD patternof Fe3O4 (JCPDS card No. 88-315), showing the chemical purity of themagnetite sample (Fig. S1a). The SEM images revealed that the mag-netite particles were in the shapes of either octahedrons or spheres,with diameters in the range of 100–200 nm (Fig. S1b). According to theN2 adsorption and desorption isotherms, the specific surface area of as-received magnetite was calculated to be 5.5 m2/g by using theBrunauer–Emmett–Teller (BET) model (Fig. S2).

The heterogeneous Fenton activity of the magnetite sample wasevaluated with the alachlor removal. The alachlor removal efficiency ofthe Fe3O4/H2O2 system was less than 10% within 1 h, whereas the in-troduction of ascorbic acid greatly enhanced the removal ratio of ala-chlor to 69.5% within only 20min (Fig. 1a). As either the Fe3O4/AA orthe H2O2/AA systems could hardly degrade alachlor, AA might enhancethe degradation of alachlor in the Fe3O4/AA/H2O2 system by accel-erating the Fenton reaction. As a weak acid, AA could lower the pH ofthe system, which might induce iron dissolution and benefit the Fentonreaction. Lower than the initial pH (6.1) of the Fe3O4/H2O2 system, theinitial pH of the Fe3O4/AA/H2O2 system was 4 ± 0.1 without anyadjustment with this pH constant throughout the reaction (Fig. S3). Toinvestigate the role of lower pH on Fenton activity, we used H2SO4

instead of AA to obtain Fe3O4/H2O2/H2SO4 system with initial pH of4.0. However, the degradation of alachlor in the Fe3O4/H2O2/H2SO4

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

system (Fig. S4), suggesting that the decreased pH caused by AA did notcontribute to the better alachlor removal performance of the Fe3O4/AA/H2O2 system.

In the Fe3O4/AA/H2O2 system, the degradation of alachlor sig-nificantly decelerated after 20min. To check if the depletion of re-actants was responsible for this deceleration, we conducted experi-ments by replenishing different reactants at 60min. The addition of

Fig. 1. (a) Time profile of alachlor degradationin the Fe3O4/AA/H2O2 Fenton system and thecontrol systems. (b) Depletion of ascorbic acidin the Fe3O4/AA/H2O2 and Fe3O4/H2O2 sys-tems. The dots show the experimental results,and the thick lines represent the model pre-dictions. The initial concentrations of alachlor,hydrogen peroxide, ascorbic acid and magne-tite were 20mg/L, 1mmol/L, 0.5 mmol/L, and1 g/L, respectively; the initial pH was 4without adjustment.

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Fe3O4 or H2O2 did not cause any further removal of alachlor, whereasthe supplement of AA triggered further removal of alachlor (Fig. S5).Therefore, the depletion of AA might be responsible for the decelerationof alachlor degradation after 20min. To confirm this assumption, wefurther monitored the concentration profile of AA during the reaction,and found the complete consumption of AA within 20min in the Fe3O4/AA/H2O2 system (Fig. 1b). Therefore, AA played the indispensable rolein the enhanced Fenton degradation of alachlor. It was reported that AAcould reductively dissolve iron minerals like hematite and magnetite,resulting in its oxidation [14]. However, the concentration of AA in theFe3O4/AA/H2O2 system decreased much more quickly than that in theFe3O4/AA system, and slightly declined in the AA/H2O2 system. Thefast depletion of AA in the Fe3O4/AA/H2O2 system could be attributedto 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 byFe3O4 or H2O2 [27].

3.2. AA promoted H2O2 decomposition and %OH production

To further confirm the accelerated heterogeneous Fenton reactionby AA, we monitored the decomposition of H2O2 and the production of•OH. In the Fe3O4/H2O2 system, less than 2% of the H2O2 was con-sumed within 60min, whereas the introduction of AA greatly increasedthe decomposition percentage of H2O2 to ca. 26% within 20min, andthe decomposition of H2O2 slowed down thereafter due to the ex-haustion of AA (Fig. 2a). The decomposition of H2O2 was fitted bypseudo-first order kinetics (Fig. 2b), and the apparent rate constant (k)of H2O2 decomposition increased by 174 times from9.56×10−5 min−1 in Fe3O4/H2O2 system to 1.66× 10-2 min−1 inFe3O4/AA/H2O2 system.

Consequently, the production of •OH was also enhanced by AA.Electron paramagnetic resonance (EPR) spectra were first employed toconfirm the yield of %OH by using 5,5-dimethyl-L-pyrroline-N-oxide(DMPO) as a spin trapper. As depicted by Fig. 3a, typical signals ofDMPO-•OH spin adduct, quartet peaks with relative intensity ratios of1:2:2:1, were distinguished from the ESR spectra in both Fe3O4/H2O2

and Fe3O4/H2O2/AA systems [28]. Obviously, the addition of AA sig-nificantly promoted the generation of •OH. We then measured the ac-cumulative concentrations of %OH in both systems by monitoring theproduction of p-hydroxyl benzoic acid from the reaction between %OHand benzoic acid. In the Fe3O4/H2O2 system, the production of %OH wasvery limited, with approximate 10.5 μmol/L accumulated in 60min. Inthe presence of AA, the accumulation of %OH was greatly acceleratedespecially within the initial 20min, with the peak value of 241.6 μmol/L obtained in the subsequent 60min (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, regardingthe surface and homogeneous Fenton processes, respectively. In a ty-pical surface Fenton system, the generation of %OH originated from thereaction between active sites on the surface of catalysts and H2O2, thusthe generation rate of %OH (VOH) could be defined as:

= = kV d[ OH]dt

[active site][H O ]OH

•

active site, H O 2 2•2 2 (10)

where ktive site, H2O2 refers to the rate constants between the surficialactive sites and H2O2, while [active site] and [H2O2] represent theirconcentrations. Thus, the rise of VOH could be attributed to the pro-motion of kactive site, H2O2, and/or the increased concentration of activesites which referred to faction of Fe(II) species on the surface of themagnetite in the Fe3O4/AA/H2O2 system. AA was able to induce re-ductive dissolution of various iron oxides like hematite, goethite andmagnetite, to produce surface Fe(II), described by the following steps[14,29,30]:

^Fe(III)+AA→^Fe(III)AA (11)

^Fe(III)AA→^Fe(II)+AAox (12)

^Fe(II)→^Fe(III)+ Fe2+ (13)

First, ascorbic acid formed surface complexes (^Fe(III)AA) with theiron atoms on the surface of iron oxide (^Fe(III)), followed by theelectron transfer process from AA ligands to iron atoms, generatingsurface ferrous species (^Fe(II)) and oxidized AA molecules (AAox).The ^Fe(II) will subsequently release dissolved ferrous ion (Fe2+) intothe bulk solution phase slowly. Therefore, AA could increase the densityof ^Fe(II), the active center for the decomposition of H2O2 in surfaceFenton reaction. To confirm this mechanism, we monitored the frac-tions of ^Fe(II) to total surface iron species (sum of ^Fe(II) and ^Fe(III)) on Fe3O4 by a 1, 10-phenanthroline dissolution method underanoxic condition. In pristine Fe3O4, the initial ^Fe(II) fraction wasdetermined to be ca. 10%, lower than the theoretical ferrous fraction(33%) in Fe3O4, possibly due to the oxidation of the surface Fe(II) by airduring storage (Fig. 4a). In the Fe3O4/H2O2 system, the^Fe(II) fractiongradually increased to ca. 30% after 25min and then decreased, in-dicating that the ^Fe(III) could be slowly reduced by H2O2 to ^Fe(II)(Eq. (14)), which subsequently underwent re-oxidation during theFenton process (Eq. (15)). In the Fe3O4/AA system, the ^Fe(II) fractionincreased drastically to nearly 100% within only 5min, and maintainedat this high level thereafter. As for the Fe3O4/AA/H2O2 system, the^Fe(II) proportion showed similar sharp rise to 97% within the initial5 min, and decreased gradually after 15min, coinciding with the de-pletion of AA. These results suggested that AA was able to promote therapid reduction of ^Fe(III) to produce more ^Fe(II) that acted as theactive center for H2O2 decomposition, improving surface Fenton effi-ciency of Fe3O4.

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

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^Fe(III)+H2O2→^Fe(II)+ %HO2+H+ (14)

^Fe(II)+H2O2→^Fe(III)+ %OH+OH− (15)

X-ray photoelectron spectroscopy (XPS) analysis of the reactedFe3O4 samples in different systems also supports the proposed me-chanism. 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) inFe2O3 or FeOOH [31,32], indicated that ^Fe(II) underwent oxidationwhen exposed to ambient air. However, for the magnetite samplescollected from Fe3O4/AA/H2O2 and Fe3O4/H2O2 systems, the in-tensities of the satellite peaks decreased significantly, which suggestedthe regeneration of ^Fe(II) via the reduction by AA or H2O2 (Fig. S6a).To quantify the surface Fe(II) fractions, the Fe 2p3/2 XPS spectra werefitted by deconvolution according to the Gupta and Sen (GS) multiplepeak fitting method [33], with fitting parameters like BE (binding en-ergy) and FWHM (full width at half maximum) listed in Table S3. The^Fe(II) proportion in the commercial Fe3O4 sample was calculated tobe 22.8%, and the ratio was improved to 27.3% and 31.6% after het-erogeneous Fenton reaction in the Fe3O4/H2O2 and Fe3O4/AA/H2O2

systems, respectively (Fig. S6b), showing similar trends with those inFig. 4a. The discrepancy between ^Fe(II) fractions obtained by XPSanalysis and those measured by the 1, 10-phenanthraline extractionmethod could be explained as follows. XPS reflects the average signalsof 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 notdetected within 60min. As expected, the concentration of dissolvedFe2+ in the Fe3O4/AA system increased steadily due to the reductivedissolution effect of AA, and dissolved Fe3+ was still not detectedowing to the excess amount of AA and its much lower reduction

potential (E0= 0.06 V) than Fe3+/Fe2+ (E0=0.77 V) [34,35]. In theFe3O4/AA/H2O2 system, however, the release of Fe2+ was apparentlyinhibited by the addition of H2O2. Meanwhile, the concentration ofdissolved Fe2+ decreased after the depletion of AA at 20min of reac-tion, accompanied with the increased concentration of Fe3+ throughthe oxidation of Fe2+ by H2O2 and %OH (Fig. 4c). Therefore, AA couldalso promote the Fe3+/Fe2+ redox cycle of the homogeneous Fentonreaction. Although the concentrations of dissolved iron were low, thecontribution of homogeneous Fenton in the Fe3O4/AA/H2O2 hetero-geneous system could not be neglected because of the efficient Fe3+/Fe2+ cycle in the presence of AA. Obviously, it is not straight forward toassess the individual contributions of surface and homogeneous Fentonregarding the dynamic variations of iron, AA and H2O2 concentrationsduring the reaction, which calls for quantitative description of the re-action kinetics in the Fe3O4/AA/H2O2 system as explored subsequently.

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 concentrationsof ^Fe(II) and ^Fe(III) for kinetic modeling. Unfortunately, the mea-surement of surface iron concentrations varied with the extraction timeusing the 1, 10-phenanthraline extraction method, which made itchallenging to quantify the exact concentrations of ^Fe. Nevertheless,the fractions of ^Fe(II) obtained by this method should be reliablesince both ^Fe(II) and ^Fe(III) were extracted from the surface ofFe3O4 simultaneously. Alternatively, the total concentration of ≡Fewas quantified by using the acid-base potentiometric titration methodas described in Section 2.5. The titration curves indicated that theFe3O4 sample had pH buffer capacity in the pH range of 3.5–10, de-monstrating the protonation and deprotonation processes during titra-tion (Fig. 5a, Eqs. (3)–(5)). From the Gran plots, the equivalence points

Fig. 3. (a) The EPR signals of the DMPO adducts in different Fenton systems at 5min of reaction. (b) Accumulative concentrations of hydroxyl radical (%OH) inFe3O4/AA/H2O2 and Fe3O4/H2O2 systems, 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 the bulk solution plotted as a function of reaction time indifferent systems. The dots and thin lines show the experimental results. The thick lines in panel c represent the model predictions.

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in the acidic (Vea) and alkaline (Veb) titration were determined to be0.750mL and 0.963mL for the blank as well as 0.763mL and 1.090mLfor the 1 g/L Fe3O4 sample (Fig. 5b, Table S3). Then the total con-centration of ≡Fe on the Fe3O4 surface was calculated to be1.11×10−4 mol/L (Eqs. (7)–(9)). According to the proportion (9.25%)of ^Fe(II) in pristine Fe3O4 determined by the 1, 10-phenanthrolinemethod (Fig. 4a), the initial ^Fe(II) and ^Fe(III) concentrations of1.007× 10−4 mol/L and 1.027×10−5 mol/L were employed in thekinetic modelling, respectively.

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

heterogeneous Fenton system was developed by fitting the experi-mental data and the rate constants obtained from literature. Table 1summarizes the elementary reactions involved. The key components ofthis model include: (i) The reactions in the Fe3O4/H2O2 system, mainlycomposed of the reduction of ^Fe(III) by H2O2, the surface Fentonreaction and the oxidation of alachlor by %OH (Reaction 1–3 in Table 1).The dissolution of Fe2+ from the Fe3O4 surface and the subsequenthomogeneous Fenton reactions were not included here since Fe2+ orFe3+ concentrations did not significantly increase in the Fe3O4/H2O2

system (Fig. 4b); (ii) The surface reactions involved in the Fe3O4/AAsystem, like the formation of surface complex between 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) and (13) for simplicity of the model); (iii) The surface Fenton re-action between ^Fe(III)AA complex and H2O2 in the Fe3O4/AA/H2O2

system (Reaction 6 in Table 1); (iv) The classical homogeneous Fentonprocess in the Fe3O4/AA/H2O2 system (Reaction 7–13 in Table 1), thekinetic rate constants of which were well established in previous stu-dies; 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 proposedmodel were determined as follows. (i) It is difficult to figure out theexact rate constants of ^Fe(III) reduction by H2O2 and its reverse re-action in the Fe3O4/H2O2 system (k1 and k-1). Instead, the rate constantsof the corresponding homogeneous reactions (k8 and k-8) were adoptedin our current model, which may contribute to the inaccuracy of themodelling results. However, the simulation results of the whole modeldid not change 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 ofReaction 2 (k2= 0.38 M−1 s-1) was deduced from the decomposition ofH2O2 in the Fe3O4/H2O2 system (Fig. 2a). This value was much smallerthan that of the homogeneous Fenton reaction, probably due to the

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

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

No. Reaction Rate constant References

1 − + ↔ − + ++Fe III Fe O H O Fe II Fe O HO H( ) ( )3 4 2 2 3 4 2

• a k1 = 2.5× 10−3 M-1 s-1 [36,37]k-1 = 2.4× 106 M−1 s-1

2− + → − + +

+

Fe II Fe O H O Fe III Fe O HO H O( ) ( )H

3 4 2 2 3 4 • 2k2= 0.38 M−1 s-1 this study

3 + →HO Alachlor Alachlorprod• b k3= 7×109 M−1 s-1 this study4 − + ↔ − −Fe III Fe O AA AA Fe III Fe O( ) ( )3 4 3 4

c k4= 15 M−1 s-1 this studyk-4 = 9.0× 10−3 s-1

5 − − → + − ++AA Fe III Fe O Fe Fe III Fe O AA( ) ( ) ox3 4 2 3 4

d k5= 3×10−4 s-1 this study6

− − + → − + + +

+

AA Fe III Fe O H O Fe III Fe O HO H O AA( ) ( )H

ox3 4 2 2 3 4 • 2k6= 35 M−1 s-1 this study

7+ → + +

+

+

+Fe H O Fe HO H OH2 2 2 3 • 2

k7= 55 M−1 s-1 [36,38]

8 + → + ++ + +Fe H O Fe HO H3 2 2 2

2• k8=2.5× 10−3 M-1 s-1 [36,37]

k-8 = 2.4× 106 M−1 s-1

9 + → +H O HO HO H O2 2 •2•

2 k9= 3.0× 107 M−1 s-1 [36,37]10 + → + +

+ + +Fe HO Fe O H32• 2 2 k10=7.7× 106 M−1 s-1 [36,37]

k-10=0.1M−1 s-1

11 + → ++ + −Fe HO Fe OH2 • 3 k11= 3.2×108 M−1 s-1 [36,37]

12 + → +HO HO H O O•2•

2 2 k12= 7.5×109 M−1 s-1 [36,37]13 + →HO HO H O• • 2 2 k13= 5.2×109 M−1 s-1 [36,37]14 + → +

+ +Fe AA Fe AAox3 2 k14 = 4.5×103 M−1 s-1 This study15 + →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.b 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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limitation of diffusion and adsorption of H2O2 at the interface, as wellas the desorption of H2O2 decomposition products (%OH, %HO2 and OH-)from the active sites. The rate constant of alachlor oxidation by %OH(k3=7×109M−1 s−1) was acquired by fitting the alachlor degrada-tion data in the Fe3O4/H2O2 system (Fig. 1a), and this value was inaccordance with that reported by Haag and Yao [39], demonstratingthe reliability of our fitting process. (ii) To obtain the kinetic rateconstants of the surface complexation between AA and ^Fe(III) (Re-action 4), the equilibrium constant (Kads) was first determined byconducting AA adsorption experiments with the literature method (TextS6) [29,30], resulting in a Kads of ca. 1670 M-1 derived from the doublereciprocal plots of Langmuir adsorption isotherm (Fig. S7), and then therate constants of AA adsorption (k4=15M−1 s−1) on the Fe3O4 surfaceand AA desorption (k-4=9.0×10−3 s−1) were determined togetherwith 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 AAconsumption and Fe2+ dissolution in the Fe3O4/AA system, respec-tively (Figs. 1b and 4c). (iii) Subsequently, the rate constant of thesurface Fenton reaction between ≡Fe(III)AA complex and H2O2

(k6=35M-1 s-1, Reaction 6) was obtained by fitting both the data of AAdepletion and H2O2 decomposition in the Fe3O4/AA/H2O2 system, andthis value was enhanced by ca. 92 times when compared with the rateconstant of the surface Fenton reaction (k2) in the Fe3O4/H2O2 system,confirming that AA promoted the surface Fenton by increasing theconcentration of active sites on the Fe3O4 surface, and improving therate constant. (iv) The rate constants of the elementary reactions inclassical homogeneous Fenton system (Reaction 7–13) were well es-tablished by previous studies and thus the literature reported valueswere directly adopted in our present kinetic model. (v) The rate con-stant (k14=4.5×103M−1 s−1) of Fe3+ reduction by AA in thehomogeneous solution was determined from the best fit of the dissolvedFe2+ and Fe3+ concentrations during reaction in the Fe3O4/AA/H2O2

system (Fig. 4c). In the Fe3O4/AA/H2O2 system, AA and alachlor wouldcompete for the %OH produced by Fenton reaction. Therefore, whenother rate constants involved were fixed, the rate constant of AA oxi-dation by %OH (k15=1.0×1010 M-1 s-1) could be deduced from thealachlor degradation data in the Fe3O4/AA/H2O2 system (Fig. 1a), andthis value is also consistent with those reported previously (4.1× 109 –1.3×1010 M-1 s-1) [27].

The superoxide radical produced from Reaction 1 could subse-quently reduce ^Fe(III) to ^Fe(II) at relatively higher rate constantthan H2O2 to accelerate the iron cycle on the Fe3O4 surface (Eq. (16)).Besides, superoxide radical could also degrade alachlor in the Fe3O4/AA/H2O2 system (Eq. (17)). To evaluate the importance of these pro-cesses, we first compared the alachlor degradation profiles which werepredicted by the kinetic models including and excluding Eq. (16), andassuming their rate constants were identical to that of Reaction 10 inTable 1. However, the difference between the results simulated by thetwo models were negligible, both in the Fe3O4/H2O2 and Fe3O4/AA/

H2O2 systems (Fig. S8), suggesting that the reduction of ^Fe(III) bysuperoxide radical (Eq. (16)) was not crucial in the Fe3O4/AA/H2O2

system. To confirm the contribution of superoxide radical for alachlordegradation in the Fe3O4/AA/H2O2 system (Eq. (17)), ROSs quenchingexperiments were conducted. It was found that the addition of catalaseand tert-butyl alcohol (TBA) completely inhibited the alachlor de-gradation, whereas superoxide dismutase (SOD) had little impact on theperformance (Fig. S9). These results suggested that %OH was the pre-dominant ROSs for the alachlor removal, and the contribution of su-peroxide radical was negligible. Therefore, the reactions of Eqs. (16)and (17) were not included in the current kinetic model.

^Fe(III)+ %HO2→^Fe(II)+O2+H+ (16)

alachlor+ %HO2→ product (17)

We could thus figure out the rate limiting steps of the heterogeneousFenton systems based on the rate constants acquired in Table 1. In theFe3O4/H2O2 system (Reaction 1–3), Reaction 1 is the rate limiting stepbecause k1 is 9 magnitudes lower than k-1, resulting in the insufficientgeneration of ≡Fe(II) for the subsequent Fenton reactions. In theFe3O4/AA/H2O2 system, the rate constants of Reactions 4–8 are muchlower than the others, and might be potential rate determining ele-mentary reactions. To quantify the roles of these reactions in de-termining the overall Fenton reaction rate, we increased k4-k8 by 2times separately in the model, and compared the predicted alachlordegradation. The doubled k4 improved alachlor degradation the most,followed by k5 and k7, whereas the increased k6 and k8 caused littlechange of alachlor removal (Fig. S10). These results indicated that theadsorption of AA on the Fe3O4 surface to form electron transfer complexwas the rate determining step in the Fe3O4/AA/H2O2 heterogeneousFenton system.

The resultant kinetic model could well describe the experimentaldata. To further check the reliability of this model, we conducted thealachlor removal experiments with varying doses of the reactants suchas Fe3O4, AA and H2O2, and the prediction by the model basicallymatched the corresponding experimental data (Fig. 6). Nevertheless,some discrepancies still existed between the experimental data and themodeling 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 ad-sorption on Fe3O4 surface, the complexation of substrates with surfaceactive sites, the electron transfer processes, the generation of reactiveintermediate species, the desorption of intermediates and products, andthe dissolution of surface iron species, etc. Whereas in this study it wasapparently unrealistic to include all the above-mentioned reactionssince it would make the kinetic model too complicated to be addressed.Therefore, approximation was made to simplify the model, whichwould introduce uncertainties to the simulation results. Second, theheterogeneous Fe3O4 catalysts might aggregate during the reaction,

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 thicklines indicate the model predictions. Conditions: (a) [alachlor]0= 20mg/L, [AA]0= 0.5mM, [H2O2]0= 1mM. (b) [alachlor]0= 20mg/L, [Fe3O4]0= 1 g/L,[H2O2]0= 1mM. (c) [alachlor]0= 20mg/L, [Fe3O4]0= 1 g/L, [AA]0=0.5mM.

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resulting in less active sites available for AA complexation and surfaceFenton reaction than those in the model. Third, some of the rate con-stants in our proposed kinetic model were directly adopted from lit-erature values, which might bring about discrepancies due to the in-consistent experimental conditions between our case and literature.Therefore, scope of refinement to the model still exists for a moreprecise modeling of the kinetics of heterogeneous Fenton system in thefuture.

3.5. Contribution of surface and homogeneous Fenton reactions to alachlordegradation

The individual contributions of surface and homogeneous Fentonreactions in the Fe3O4/AA/H2O2 system were then quantified bymodifying the above kinetic model (Table S5). To discriminate the %OHgenerated by the surface and homogeneous Fenton processes, they weredenoted as •OHsurface and %OHhomo, respectively. Consequently, thealachlor degraded by %OHsurface and %OHhomo were indicated asAlachlorsurface and Alachlorhomo, and identical rate constants were em-ployed for the elementary reactions involved with %OHsurface and%OHhomo. Fig. 7 presents the model predicted time profile of Ala-chlorsurface and Alachlorhomo, as well as the contribution of surfaceFenton to the overall alachlor removal (Alachlorsurface/(Ala-chlorsurface+Alachlorhomo)). Initially, the removal of alachlor wasdominated by the surface Fenton reaction, taking up more than 85%before 10min. With the increase of dissolved iron ions, the homo-geneous Fenton reaction accelerated, corresponding to the gradualdecrease of the surface Fenton fraction. Both surface and homogeneousFenton reactions decelerated sharply after 20min owing to the deple-tion of AA. Nonetheless, surface Fenton still accounted for 62.6% of theoverall alachlor removal at the end of the reaction (60min), suggestingthe 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 heterogeneousFenton catalyst in the Fe3O4/AA/H2O2 system, the Fe3O4 after reactionwas separated and subjected to more cycles of test, fed with fresh AA,H2O2 and alachlor. The recycled Fe3O4 still exhibited high activity forthe removal of alachlor after 4 successive runs (Fig. 8a). The stability ofmagnetite could be explained by the limited Fe leaching during thereaction 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 caused the iron cycle on the Fe3O4 sur-face, and thus helped maintain the surface structure and bulk compo-sition of Fe3O4.

3.7. Transformation mechanism of alachlor and AA

The dechlorination of alachlor in the Fe3O4/AA/H2O2 Fentonsystem was first studied by employing ion chromatography. The Cl-concentrations increased sharply to 0.054mmol/L within 20min, andleveled off thereafter, consistent with the removal trend of alachlor(Fig. 8b). The dechlorination ratio, defined as the ratio of released Cl−

versus the total Cl contained in the removed alachlor, was 98.2% at30min. This suggested that dechlorination was an essential step duringthe transformation of alachlor in the Fe3O4/AA/H2O2 system. More-over, considering that the substituted chlorine often contributes to thetoxicity of pollutants, the nearly complete dechlorination indicated thedetoxification 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-massspectroscopy (GCeMS). The identified intermediates of alachlor in-cluded 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 1,3-diethyl-2-nitrosobenzene, etc. (Table S6). The possible de-gradation pathways of alachlor in the Fe3O4/AA/H2O2 Fenton systemwas proposed (Scheme 1). The first step of alachlor degradation mainlyinvolved with the cleavage of the C–N single bond, resulting in theformation of intermediate 1, which could be further oxidized throughdechlorination or dealkylation to generate compounds 2 and 3. Theintermediate would undergo stepwise dealkylation of the ethyl group togive compounds 4 and 5. The –NH2 group of compound 3 was oxidizedby %OH into –N]O (compound 6) or –NO2, and the ethyl grouptransformed to −CHO (compound 7). Meanwhile, the ortho- ethylgroup of alachlor could also be dealkylated at the initial stage of de-gradation, producing compound 8, which would further transform to 9after the CeN bond broke up. The above intermediates could beeventually oxidized into smaller organic molecules like oxalic acid andformic acid after the cleavage of the benzene ring, or mineralized intocarbon dioxide, nitrate, and chloride. The acute and chronic toxicityendpoint values of the intermediates showed similar trends, i.e., toxi-city of the intermediates to fish or daphnia did not change much com-pared with alachlor, but the toxicity to green algae decreased sig-nificantly, indicating that alachlor can be gradually detoxified byFe3O4/AA/H2O2 Fenton treatment (Fig. S11).

Since AA was also degraded in the Fe3O4/AA/H2O2 Fenton system,the intermediates of AA transformation were also identified by HPLC-MS-MS (Table S7). The hydroxyl group adjacent to the five-memberedring of AA could be oxidized to carbonyl group either by %OH or Fe(III),resulting in dehydroascorbic acid (m/z=173). •OH could also attackthe side chain of AA molecules, producing 3,4-dihydroxy-5-(hydro-xymethyl)furan-2(5 H)-one (m/z=144), which could be further oxi-dized to 5-(hydroxymethyl)furan-2,3,4(5 H)-trione (m/z=143), 5-hy-droxyfuran-2,3,4(5 H)-trione (m/z=129) and furan-2,3,4(5 H)-trione(m/z=113). The above intermediates would undergo further ring-opening degradation, generating byproducts such as 2,3-diketogulonicacid (m/z=191), 2, 3, 4-trihydroxybutanoic acid (m/z=135), oxalicacid, acetic acid and formic acid (Scheme 2). Ecotoxicity evaluationshowed that most of the intermediates were nontoxic since their acuteand chronic toxicity endpoint values were comparable with ascorbicacid. Though the toxicity of product 5 slightly increased, it would un-dergo further degradation and be detoxified (Fig. S12). This impliedthat the application of AA to enhance the heterogeneous Fenton reac-tion was environmentally benign and would cause no secondary pol-lution.

Fig. 7. The model predicted time profile of alachlor removed by surface Fenton(blue open square) and homogeneous Fenton (red open circle) in the Fe3O4/AA/H2O2 system, as well as the contribution of surface Fenton to the overallalachlor removal (green line). The initial concentrations of alachlor, H2O2, AAand Fe3O4 were 20mg/L, 1mmol/L, 0.5 mmol/L and 1 g/L, respectively. (Forinterpretation of the references to colour in this figure legend, the reader isreferred to the web version of this article.)

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3.8. Removal of other pollutants and impacts of coexisting inorganic andorganic matters

We further employed the Fe3O4/AA/H2O2 system to remove otherorganic pollutants, including organic dyes (methylene blue), phenols(4-chloropenol and bisphenol A), and pharmaceuticals (ranitidine andpropranolol). As expected, all these pollutants could be effectively de-graded within 60min, indicating the wide applicability of the Fe3O4/AA/H2O2 Fenton system for organic pollution removal (Fig. S13).Besides, we evaluated the impacts of coexisting inorganic and organicmatters on the pollutant removal performance of the Fe3O4/AA/H2O2

system. Chloride had negligible impact on the alachlor removal (Fig.S14a), and HCO3

− showed slight inhibition effect only at high con-centration (10mM, Fig. S14b). Interestingly, H2PO4

− of low con-centration (1mM) promoted the degradation of alachlor, but 10mM ofH2PO4- decelerated alachlor removal, possibly due to the competitiveadsorption of H2PO4

− and AA on the Fe3O4 surface. Fortunately, goodremoval efficiency could still be achieved with prolonged reaction time(Fig. S14c). The presence of 1mg/L humic acid as a representative ofnatural organic matters had little impact on the alachlor removal. Whenthe concentration of humic acid increased to 10mg/L, slight inhibitionof alachlor degradation was observed (Fig. S14d). These results sug-gested that the Fe3O4/AA/H2O2 heterogeneous Fenton system washighly promising for natural water or waste water treatment.

4. Conclusions

In this study, we have demonstrated that the Fenton removal ofalachlor in the Fe3O4/H2O2 system can be drastically accelerated by AAvia the rapid iron redox cycle both on the surface of Fe3O4 and in thesolution. To elucidate the individual contributions of surface and

homogeneous Fenton processes to the alachlor removal, a conceptualkinetics model composed of (i) the surface complexation and redoxreactions between AA and surface iron, (ii) the surface Fenton reac-tions, and (iii) the classical homogeneous Fenton reactions, was de-veloped based on fitting of the experimental data, and the model pre-diction was in good accordance with experimental results. Theindividual contributions of surface and homogeneous Fenton to ala-chlor degradation were simulated with using this model, and the mainrole of surface Fenton was revealed. The kinetic modeling in this workcould offer the quantitative description of the reactions in hetero-geneous Fenton systems, contributing to a better understanding of theheterogeneous Fenton mechanism.

Conflict of interests

The authors declare that they have no known competing financialinterests or personal relationships that could have appeared to influ-ence the work reported in this paper.

Acknowledgments

This research was financially supported by The National KeyResearch and Development Program of China (2018YFC1800801 and2018YFC1802003), Natural Science Funds for Distinguished YoungScholars (21425728), National Natural Science Foundation of China(41601543, 21936003 and 41807349), Project funded by ChinaPostdoctoral 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 HubeiProvince, China (Z13).

Fig. 8. (a) Reusability of the Fe3O4 as Fenton catalyst for alachlor degradation. (b) The evolution of Cl− concentrations in the alachlor degradation process by Fe3O4/AA/H2O2 Fenton system.

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

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Appendix A. Supplementary data

Supplementary material related to this article can be found, in theonline version, at doi:https://doi.org/10.1016/j.apcatb.2019.118383.

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

H. Sun, et al. Applied Catalysis B: Environmental 267 (2020) 118383

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