Simultaneous removal of Cr(VI) and triclosan from aqueous solutions through Fe3O4 magnetic nanoscale-activated persulfate oxidation

Chemical Engineering Journal 381 (2020) 122586

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Simultaneous removal of Cr(VI) and triclosan from aqueous solutions through Fe3O4 magnetic nanoscale-activated persulfate oxidation

T

⁎

Xiulan Song , Caiwen Ren, Qingyun Zhao, Bingqin Su College of Environmental Science and Engineering, Taiyuan University of Technology, Taiyuan, China

H I GH L IG H T S

G R A P H I C A L A B S T R A C T

removal of Cr(VI) and • Simultaneous Triclosan was achieved by Fe O ac3

• • •

4

tivated persulfate system. SO4∙− produced by Fe3O4 activated persulfate are the primary radical species responsible for TCS degradation. The mechanism of Cr(VI) removal by Fe3O4 activated persulfate was explained. Three possible degradation pathways were proposed, involving in the breakage of ether bond, dechlorination and hydroxylation.

A R T I C LE I N FO

A B S T R A C T

Keywords: Fe3O4 magnetic nanoscales Persulfate Cr(VI) Triclosan Simultaneous removal Sulfate radical oxidation

Triclosan (TCS) and Cr(VI), which are dangerous to human and ecological health, are frequently detected in surface waters. As a proposed solution, in this study, Cr(VI) and TCS were removed through an Fe3O4-activated persulfate (PS) oxidation process. At first, several influence factors on TCS degradation were investigated such as initial pH value, PS dose, temperature, Fe3O4 dosage and Cr(VI) concentration. 87.5% degradation of 5 mg L−1 TCS and 99.5% removal of 1 mg L−1 Cr(VI) were achieved within 120 min by using 2 g L−1 of Fe3O4 and 1 mM PS at 30 °C with a pH of 7. The degradation of TCS follows the pseudo-first-order kinetics under the given experimental conditions. Acid conditions were beneficial to TCS removal. TCS oxidation remained feasible at circumneutral pH. Moreover, CO32−, HPO42−, less than 10 mM of chloride ions and humic acid slightly inhibited the removal of Cr(VI) and TCS. Radical scavenger experiments indicated that sulfate radicals are the primary radical species responsible for TCS degradation. Then, three possible degradation pathways were proposed, including the breakage of ether bond, dechlorination and hydroxylation. This study indicates that the Fe3O4-activated PS oxidation process could be used as a novel treatment technique for the simultaneous removal of heavy metal and organic pollutants.

1. Introduction Recently, co-contamination by antibiotics and heavy metals has received more attentions [1]. Antibiotics and heavy metals commonly coexist in real environments such as agriculture production, different

⁎

water sources, bio-treated effluent and mixed industry wastewater effluent, causing more serious damage to environment and humans for their combined toxicity and relative mobility [2,3]. Combined pollution with heavy metals and antibiotics is more difficult to be treated [1,4], so it is necessary and appropriate to investigate the treatment of

Corresponding author. E-mail address: [email protected] (X. Song).

https://doi.org/10.1016/j.cej.2019.122586 Received 27 April 2019; Received in revised form 31 July 2019; Accepted 22 August 2019 Available online 26 August 2019 1385-8947/ © 2019 Elsevier B.V. All rights reserved.

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2.2. Preparation and characterization of Fe3O4

composite pollutants. Among all antibiotics, triclosan (5-chloro-2-(2,4-dichlorophenoxy) phenol, TCS) is a broad-spectrum antibacterial agent widely used at appropriate concentrations in personal and health care products, such as soap, toothpaste, and detergents [5]. The extensive use of TCS has resulted in its discharge into wastewater and subsequently into surface waters [6,7]. TCS has been detected in water bodies in China, and its concentration is 478 ng·L−1 in Zhujiang river [8]. Serious environmental concerns have been raised regarding TCS because of its toxicity to some aquatic organisms and the formation of toxic byproducts. TCS may react with free chlorine, which is used as a disinfectant in wastewater and drinking water, to generate chloroform and chlorinated phenols [9]. Furthermore, TCS has been detected in human samples and proved that it may cause skin irritation and endocrine disruption [10].On the other hand, drinking water bodies can be contaminated by heavy metals, such as Cr(VI), produced by textile processing, leather tanning, mining, electroplating, and pigment manufacturing [11,12]. The concentrations of Cr(VI) in sediments of Zhujiang river in South China is 5.869 – 66.559 mg/kg [13]. Cr(VI) have been extensively researched because it is highly mobile, toxic, mutagenic, and carcinogenic [14]. It is possible for TCS and Cr(VI) exists in real waters since organic pollutants and heavy metals could be found in aquatic environments [15,16].A new process must thus be developed to simultaneously remove TCS and Cr(VI) from drinking water bodies for water body pollution control and remediation. In situ chemical oxidation (ISCO) has been recognized in recent years for its potential in the degradation and mineralization of organic pollutants [17]. Among all ISCO oxidants, persulfate (PS) is generally used and has been extensively researched because of its excellent solubility, stable structure, relatively long half-life of SO4∙−, and high standard redox potential (E0 = 2.6 V)[18]. For SO4∙− generation, the activation of PS by using transition metal ions [19,20], heat [21,22], and ultraviolet (UV) rays [23,24] has been widely studied. Previous studies have reported that TCS can be degraded using transition metals (Fe2+, Co2+, Cu2+, and Ag+) and thermally activated PS [25–27]. However, the energy costs limit the application of thermally activated PS [26]. Among these transition metals, iron (other metal ions are toxic) [28] can efficiently activate PS for water treatment [29,30]. Accumulation of iron sludge, however, limits the application of this homogeneous system [31,32]. To overcome these problems, researchers are developing heterogeneous catalysts, such as iron minerals and nanoscale Fe3O4, in a Fenton-like system [33,34] Because of the relatively high availability and magnetic, specific structural, and catalytic properties [35,36] of Fe3O4 magnetic nanoparticles (MNPs), such particles have been employed as a satisfactory heterogeneous catalyst for activating PS to removal organic pollutants [33,34]. To date, the simultaneous removal of organic matters and heavy metals by using Fe3O4 and functional modified magnetic materials has been previously reported [3,37,38]. Based on the above, TCS and Cr(VI) were used as representative target pollutants in this study. Fe3O4 MNPs were used as a catalyst to activate PS for the simultaneous removal of TCS and Cr(VI). The primary objectives of this study were as follows: (1) to understand the effect of initial pH, PS concentration, temperature, initial Cr(VI) concentration and Fe3O4 dosage on TCS removal; (2) to determine the influence of a solution matrix on the removal of TCS and Cr(VI); (3) to identify the removal mechanism of TCS and Cr(VI); (4) to elucidate the stability and reusability of Fe3O4 MNPs. (5) to identify the reaction pathway of TCS degradation.

Details are shown in Supporting Information Text S1.

2.3. Experimental procedures Batch experiments of single TCS removal were performed in 150 mL conical flasks, which were placed in a thermostatic rotary shaker at 30 ± 1 °C and 200 rpm. An appropriate amount of Fe3O4 was rapidly added into the reaction solution with TCS and stirred for 30 min to achieve an adsorption–desorption equilibrium. Then PS was rapidly added into the reaction solution and carried out for 120 min and immediately quenched with appropriate MeOH. 2 mL samples were collected at designated time intervals. Samples were passed through 0.22 μm syringe filters and analyzed through high-performance liquid chromatography (HPLC). Batch experiments were conducted to study the removal efficiency of simultaneous removal of Cr(VI) and TCS in the Fe3O4/PS system under different conditions. Unless stated otherwise, the initial concentrations of PS and TCS were set at 1 mM and 5 mg L−1, respectively, and 2 g L−1 of Fe3O4 was used. The reaction was performed at room temperature (30 ± 1 °C), within a reaction time of 120 min. All experiments were performed without pH adjustment except the study of effects of pH(initial pH = 7 ± 0.3).All experiments were performed thrice, and the results are presented as mean values and the standard deviation of the results obtained from triplicates.

2.4. Analytical methods The concentration of TCS was quantified using an HPLC system (Agilent 1100, Agilent Technologies, CA, USA) equipped with a C18 column (Agilent XDB-C18 [4.6 mm × 250 mm, 5 μm], Agilent Technologies) with UV detection. The analysis was conducted using a mobile phase comprising methanol (90%) and water (10%) at a flow rate of 1.0 mL min−1, with an injection volume and a detection wavelength of 50 μL and 280 nm, respectively. In addition, the intermediate products of TCS degradation were detected by Thermo Scientific Q Exactive Orbitrap LC-MS/MS System (Details are shown in Supporting Information Text S2). Cr(VI) concentrations were measured by using a 1,5-diphenylcarbazide (DPC) colorimetric method through UV–vis spectrophotometry at 540 nm [39]. The total chromium concentrations were determined by atomic absorption spectrometer (AA-7020, EWAI, China). Cr(III) concentrations were determined by calculating the difference between Cr(VI) and the total chromium concentration. The total chromium on Fe3O4 surface was determined by atomic absorption spectrometer (AA-7020,EWAI,China) after digested [39]. Chloride concentrations were measured by Ion Chromatography (ICS-4000, Dionex, USA). The remaining concentration of PS was determined using a slightly modified spectrophotometer [40]. The concentrations of total iron and Fe(II) were measured using a ferrozine method through UV–vis spectrophotometry at 510 nm [41]. The pH was measured using a portable pH meter (Mettler Toledo FE-20, Shanghai Miqingke Industrial Co. Ltd, China). The mineralization of TCS was measured using a total organic carbon (TOC)-VCPH analyzer (TOC-VCPH, Shimadzu Corporation, Kyoto, Japan). In this study, the rate of TCS degradation was described by the pseudo-first-order kinetics, which is given in the following equation.

C ln ⎛ ⎞ = −kobs t ⎝ C0 ⎠ ⎜

2. Materials and methods

⎟

(1)

where C is the TCS concentration at any selected time (t); C0 is the concentration of TCS at initial time (t); t is the treatment time and kobs is the degradation rate constant of TCS (min−1). The removal efficiency of Cr(VI) was calculated from the following equation.

2.1. Materials Details are shown in Supporting Information Text S1. 2

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oxidized by the only-PS system. By comparison, the Fe3O4/PS system can induce TCS removal of 89.0% over a 120 min period. The kobs of TCS was 0.022 min−1in Fe3O4/PS system. The results showed that PS was activated by Fe3O4 to form activated substance, which was responsible for the TCS removal effectively. 3.2.2. Effect of PS concentration Experiments were conducted to study the effects of PS dose (0.3–2.0 mM) on TCS removal. At each PS dose, the removal of TCS followed a pseudo-first-order kinetic model (Table S1). The kobs of TCS increased from 0.006 to 0.022 min−1 as PS concentration increased from 0.3 to 1.0 mM. This phenomenon may be attributed to PS being a precursor in the reaction with Fe3O4, which generates SO4∙− that serves as driving forces for degrading TCS. More PS was activated using Fe3O4 to produce SO4∙− for TCS oxidization [17,44]. However, the TCS removal rate decreased when the PS concentration was higher than 1.0 mM, and the removal rate decreased to 75.0% when the PS concentration increased to 2.0 mM (Fig. 4a). The results could be explained by the following factors: when the PS concentration exceeded 1.0 mM, superfluous PS may cause the release of a large amount of SO4∙−, which result in the self-scavenging effect of SO4∙− or the consumption of the excessive PS (Eqs. (3–4)) [44]. The scavenging reactions may result in the decrease the TCS degradation efficiency. Therefore, 1 mM PS was used for the follow-up experiments.

Fig. 1. XRD pattern of Fe3O4 before and after usage;

η=

C0 − Ct × 100% C0

(2)

3. Results and discussion.

SO4∙− + S2 O82 − → SO4 2 − + S2 O8∙−

(3)

SO4∙− + SO4∙− → S2 O82 −

(4)

3.1. Characterization of Fe3O4 3.2.3. Effect of Fe3O4 dosage For evaluating the catalytic activity of Fe3O4 on PS for TCS degradation, a set of experiments was conducted with an Fe3O4 dose of 0.5–3.0 g L−1. In Fig. 4b, when the Fe3O4 dose increased, the TCS removal improved, Furthermore, the kobs of TCS degradation increased from 0.004 to 0.022 min−1 when the Fe3O4 dosage changed from 0.5 to 2.0 g/L (Table S1). When Fe3O4 dosage is beyond 2.0 g L−1, the kobs of TCS increased a little. The improve of kobs may be attributed to increasing Fe3O4 dosage, which improved the number of active sites, causing the generation of more SO4∙− [45]. The highest TCS removal rate was observed for an Fe3O4 dose of 3.0 g L−1. Similarly, at a lower concentration of Fe3O4 (2.0 g L−1), the TCS removal rate was close to that at 3.0 g L−1 of Fe3O4. Therefore, considering the cost and practical application of Fe3O4, Fe3O4 of 2.0 g L−1 was adopted for the subsequent experiments.

Fig. 1 shows a diffraction pattern with seven broad peaks at 18.3°, 30.1°, 35.5°, 43.1°, 53.5°, 57.0°, and 62.6° corresponding to (1 1 1), (2 2 0), (3 1 1), (4 0 0), (4 2 2), (5 1 1), and (4 4 0) of Fe3O4, respectively, which were consistent with the data from previous studies [42,43]. Furthermore, the XRD patterns of materials before and after usage showed that the XRD spectra did not change, which indicating good stability of Fe3O4 in this study. The diffraction peaks and positions indicated that all structures of the precipitated magnetite (Fe3O4) were inverse spinels. The particle size of the Fe3O4 particles calculated by the Debye–Sherrer formula D = Kλ/(β cos θ) is 16.7 nm. TEM images (Fig. 2) show that the Fe3O4 were spherical and exhibited an approximately 10–20 nm diameter, which is well in consistent with the value obtained from the XRD results. 3.2. TCS degradation by using the Fe3O4 /PS system

3.2.4. Effect of pH Experiments were conducted to study the effects of pH (3–11) on TCS removal by varying the pH. Moreover, the pH was previously adjusted to the target value using H2SO4 and NaOH (0.01 mM) to avoid interference from other anions. In (Fig. 4c), TCS oxidation was

3.2.1. TCS removal in different systems TCS removal was investigated by PS, Fe3O4 and Fe3O4/PS system. As shown in Fig. 3, only 13.2% of TCS was absorbed by the only-Fe3O4 system, whereas approximately only 11.2% of TCS degradation was

Fig. 2. TEM images of Fe3O4 before usage; 3

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Fig. 3. Effects of different systems on TCS degradation. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [PS]0 = 1.0 mM, [Cr(VI)]0 = 10 mg L−1, and T = 30 °C.

various temperatures indicated that the reaction was temperature-dependent (shown in Fig. 4d). The degradation of TCS (75.9%) at 25 °C was obviously lower than that at other temperatures (Fig. 4d). As shown in Table S1, the kobs of TCS were 0.015, 0.022, 0.025, 0.030, and 0.038 min−1 at 25, 30, 35, 40, and 45 °C, respectively. In Fig. S4, TCS degradation in the range of 25–45 °C was in line with an Arrhenius equation (R2 = 0.98). Based on Eq. (8), the average activation energy (Ea) was determined to be 34.80 kJ/mol (where R is the universal gas content (8.314 J mol−1 K−1)) [26,27]. Therefore, considering the costeffectiveness, 30 °C was adopted for the subsequent experiments [26,27].

depended on the pH conditions. The highest TCS removal was observed at a pH of 3; however, a dramatic decrease was observed at pH of 9–11. Approximately 61.4% of TCS was removed after 120 min at a pH of 11, whereas nearly 94.6% of TCS was removed at a pH of 3, which suggests that acidic conditions favor TCS removal. Meanwhile, the kobs of TCS decreased from 0.027 to 0.010 min−1 when the initial pH increased from 3.0 to 11.0 (Table S1). This phenomenon may be explained by the following aspects: First, SO4∙− was likely to be converted into ∙OH under alkaline conditions (Eqs. (5)) [46], resulting in the changes in the radical species of reaction system, and finally influencing TCS degradation. Second, the ionized form of TCS (pKa = 7.9–8.4) is the major speciation [47]. As pH increased, the electrostatic repulsive between TCS and SO4∙− was unfavorable to driving TCS being near SO4∙− [48]. Third, the acidic catalysis could boost the formation of SO4∙− (Eqs. (6–7) [49]. Furthermore, in Fig. S3(a), pH variation during 120 min treatment process indicates that pH of solution changed continuously during the whole reaction [50]. According to the reports [17,50,51], although stable pH value can be adjusted by buffers, a lot of anions will be added into solution, which causes some unknown influences on the system. Thus, without pH adjustment is more suitable for subsequent experiments. Acid conditions were beneficial to TCS removal. TCS oxidation remained feasible at pH 7. Thus, compared with the traditional Fenton process, Fe3O4/PS system can work within larger pH range.

SO4∙− + HO− → SO24− + ∙OH

(5)

S2 O82 − + H+ → HS2 O8−

(6)

HS2 O8− → SO4∙− + SO24− + H+

(7)

ln(K obs) = lnA −

Ea RT

(8)

3.3. Simultaneous removal of TCS and Cr(VI) in the Fe3O4/PS system The effects of Cr(VI) on TCS degradation in the Fe3O4/PS system were investigated in the concentration range of 0.0–2.0 mg L−1. Fig. 5 shows that Cr(VI) inhibited TCS degradation. The removal efficiency of TCS decreased with the increasing initial Cr(VI) concentration, and thus, although TCS removal was approximately 87.5% within 120 min at a Cr(VI) concentration of 1.0 mg L−1, it was only 66.0% within 120 min at a Cr(VI) concentration of 2.0 mg L−1. The plot appropriately fits the pseudo-first-order model (Tab.S1), with the observed the kobs of TCS significantly decreasing with the increasing initial Cr(VI) concentration. Fig. 5 shows that Cr(VI) removal was slightly decreased with the increasing initial Cr(VI) concentration. Cr(VI) is a recognized passivator and a strong oxidant, and therefore, the proximity between Cr (VI) and Fe3O4 positively corresponds with the oxidization and loss of activity of Fe3O4, which reduces the rate of TCS degradation [52]. It has been proposed that, during the redox reaction, Fe(II) is primarily oxidized to Fe(III) and Cr(VI) is reduced to Cr(III), and when Cr(III) and Fe

3.2.5. Effect of temperature In order to confirm the effects of temperature on the system, tests were conducted at various temperatures. The TCS degradation at 4

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Fig. 4. Effects of (a) PS dosage, (b) Fe3O4 dosage, (c) initial pH, and (d) temperature on TCS degradation by the Fe3O4/PS system. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C.

(III) are in close proximity, a thin film of Fe(III)-Cr(III) hydroxides forms on the surface of the Fe3O4 nanoparticles [53]. A thin film can increase the resistance for electron transfer from Fe3O4 to Cr(VI) and thus retard the degradation of TCS. A recent study has validated that the actual Fe–Cr surface structure was a mixture of CrxFe1−x(OH)3 or CrxFe1−x(OOH) [54].

SO4∙− + Cl− ↔ Cl∙ + SO24− , E0 (SO4∙− /SO24−) = 2.6V

Cl∙

Cl−

+

2Cl∙− 2

→

↔

Cl∙− 2 ,

2Cl−

E0 (Cl∙/Cl∙− 2 )

+ Cl2, E0 (Cl2

= 2.41V

/Cl−)

= 1.36V

Cl2 + H2 O ↔ HClO + H+ + Cl−, E0 (HClO/Cl−) = 1.48V

(9) (10) (11) (12)

Table 1 shows TCS degradation in the Fe3O4/PS system in the presence of carbonate. The results suggest that carbonates significantly inhibit TCS removal. The minimum removal efficiency of TCS in the presence of CO32− was 54.8%. For CO32−, in the first place, carbonate was adsorbed onto amorphous iron oxides. In the second place, the oxidation rate of TCS was retarded at an elevated pH because of CO32− addition, as argued in Section 3.2.4. Furthermore, CO32− may react with the catalyst to form Carbonate-Fe compounds (FeCO3), thereby inhibiting the catalytic activities of catalyst and blocking active sites [58]. CO32− can quench SO4∙− as radical scavengers to produce a carbonate radical (CO∙− 3 ) (Eqs. (13–14)) [27,59].

3.4. Effects of natural water constituents on simultaneous removal of TCS and Cr(VI) by using the Fe3O4/PS system 3.4.1. Effect of coexisting anions In the next experiment, coexisting anions (including Cl−, CO32−, HPO42− and SO42−) were added to examine their effects on Cr(VI) and TCS removal by using the Fe3O4/PS system. As shown in Table.1, the degradation of TCS in Fe3O4/PS system reduced slightly affected at low concentrations of Cl− (< 10 mM), which was comparable to the previous studies [55,56]. In contrast, as Cl− concentration increased from 20 to 30 mM, kobs of TCS appreciably increased from 0.022 to 0.025 min−1, suggesting that TCS removal was appreciably promoted after Cl− concentration increased (> 10 Mm). This can be explained by two aspects: 1) low concentrations of Cl− is used as SO4∙− scavenger to react with SO4∙−, which reduces TCS degradation (Eqs. (9)) [27,57]; 2) high concentrations of Cl− could be oxidized by SO4∙− to produce reactive chlorine species including reactive chlorine radicals and free chlorine (e.g., Cl∙, Cl∙− 2 , Cl2 and HClO) (Eqs. (9–12)), which could compensate the depletion of SO4∙− due to Cl− scavenging and enhance the TCS oxidation [56].

·OH + CO32 − → CO·3− + OH−

(13)

SO·4−

(14)

+

CO32 − 2−

→

SO24−

+

CO·3−

HPO4 had a negative effect on TCS degradation. Table1 shows that the removal efficiency decreased when HPO42− concentration increased. Furthermore, the kobs of TCS decreased with HPO42− concentration increasing. This phenomenon may be following reasons: 1) HPO42− can react with the catalyst to form phosphate-Fe compounds, thereby inhibit the catalytic activities of catalyst [56]; 2) HPO42− could react with SO4∙− as radical scavengers to generate a phosphate radical 5

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Fig. 5. Effects of Cr(VI) concentration on TCS degradation and Cr(VI) removal. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [Cr (VI)]0 = 1.0 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C.

(HPO4·) (Eqs. (15–16)) [60]

Table 1 Effects of coexisting substances on Cr(VI) removal and TCS degradation. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [Cr (VI)]0 = 1.0 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C. Coexisting substance

Removal of TCS (%)

kobs (min−1)

R2

Removal of Cr (VI) (%)

No addition Cl− dosage (mM) 0.01 0.1 1 10 20 30 CO32− dosage (mM) 0.01 0.2 0.1 1 10 HPO42−dosage (mM) 0.01 0.1 1 10 SO42− dosage (mM) (mg/L) 0.01 0.1 1 10 HA (mg/L) 2.5 5 7.5 10 20

87.5 – 87.4 87.4 86.8 86.9 90.8 93.5 – 86.9

0.022 – 0.020 0.020 0.020 0.020 0.022 0.025 – 0.019

0.99 – 0.93 0.93 0.93 0.93 0.97 0.98 – 0.98

99.5 – 99.5 99.5 99.2 98.1 97.0 96.9 – 99.5

80.1 64.9 54.8 – 87.0 86.7 77.5 64.5 –

0.017 0.012 0.007 – 0.020 0.019 0.013 0.009 –

0.91 0.94 0.99 – 0.96 0.97 0.97 0.97

99.2 95.7 67.2 – 96.6 95.6 93.7 68.3

87.4 87.4 87.4 87.4 – 86.3 84.4 81.3 78.0 71.2

0.021 0.020 0.019 0.019 – 0.020 0.017 0.015 0.013 0.011

0.92 0.95 0.95 0.96 – 0.93 0.93 0.97 0.97 0.99

98.8 97.8 97.4 97.1 – 98.9 97.1 97.0 96.9 95.1

HPO42 HPO42 -

+ SO·4− →

HPO4· + SO24−

+ ·OH → HPO4· +

OH−

(15) (16)

In addition, the addition of SO24− has no influences on TCS removal, indicating that SO24− is not scavengers for the SO4∙− produced in the Fe3O4/PS system. However, the presence of Cl− and SO24− in the Fe3O4/ PS system has no obvious inhibitory effect on Cr(VI) removal. Moreover, the removal of Cr(VI) was affected by HPO42- and CO32 − becuase they could react with the catalyst to form compounds, thereby inhibiting the catalytic activities of catalyst and blocking active sites. 3.4.2. Effect of humid acid In the next experiment, 0–20 mg L−1 of HA was used to examine the effects of natural water constituents in the Fe3O4/PS system. Table 1 shows the results, which reveal that HA had inhibitory effects on TCS removal in the Fe3O4/PS system. The removal efficiency of TCS was decreased from 86.3% to 71.2%. This inhibition can be attributed to two primary reasons. First, HA and TCS compete for surface active sites. HA has numerous functional groups. These acidic groups are used to easily combine HA with Fe(II) and Fe(III) for occupying reactive sites [53,58]. Second, HA might compete with TCS for the active species that affected the degradation efficiency, which was comparable to the previous studies [61]. However, the presence of HA in the Fe3O4/PS system has no evident inhibitory effects on Cr(VI) removal. This phenomenon may be attributed to the increased sorption and electron transfer between Fe3O4 and contaminants caused by organic metal chelates deposited on the Fe3O4 surfaces, which offset the negative effect of the competition between HA and Cr(VI) for surface active sites [53]. 3.5. Removal mechanism 3.5.1. Quenching test MeOH and TBA were used as quenching agents to identify major reactive radical species in the Fe3O4/PS system. In general, MeOH 6

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Fig. 6. (a) The variation of residual PS in different systems as a function of time; (b) Effect of MeOH and TBA as radical scavengers on removal of TCS and Cr(VI); (c) concentrations of TCS and 2,4-DCP as a function of time; and (d) TOC consumption in Fe3O4 /PS system. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [Cr(VI)]0 = 1.0 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C.

reacts at rapid rates with both ∙OH and SO4∙−. However, compared with ∙OH (rate constant = 3.8 × 108–7.6 × 108 Ms), the rate of reaction between SO4∙− (rate constant = 4.0 × 105–9.1 × 105 Ms) and TBA is slower [27,62]. Fig. 6b shows the results with the quenching agents. Without any radical scavenger, approximately 87.5% of TCS is removed in 120 min with kobs = 2 × 10−2 min−1 (Table S1). However, the rate of the TCS degradation is reduced to 0.9 × 10−2 and 1.6 × 10−2 min−1 (Table S1) with the addition of MeOH and TBA, respectively. Comparing with the control experiment, the TCS removal had a slightly decrease of 7.0% with the addition of TBA. This result indicated that ∙OH plays a less important role on TCS degradation. The decrease in kobs caused by MeOH is more significant than that by TBA, which indicates that sulfate radicals are the principal radical species generated during PS activation by Fe3O4. In addition, Fig. 4b shows that there is no significant difference in Cr(VI) removal with any scavenger compared with the control experiment.

which indicated that the PS-only system is not enough to degrade TCS. 3.5.3. Products and TOC analysis Fig. 6d presents the concentration changes of 2,4-dichlorophenol (2,4-DCP) and TCS as a function of time in the solution. The concentration of 2,4-DCP dramatically increased with the increasing reaction time in the first 10 min and then gradually started to decrease. This observation may be attributed to 2,4-DCP degradation by SO4∙−. However, the TCS removal dramatically increased in the initial 30 min and slightly increased with the increasing reaction time. TCS removal increased from 77.3% to 93.5% with the increasing reaction time from 30 to 180 min and became relatively stable. A slow increase in TCS removal was observed because of the self-scavenging of radicals and the competition between intermediates and TCS for SO4∙− [63]. In addition, the TOC rapidly decreased in the initial 90 min with the increasing reaction time and subsequently reached a relatively stable value. The mineralization efficiency of TCS was 56.8% at 300 min (Fig. 6c). The mineralization efficiency of TCS was 56.8% (Fig. 6c). The less accumulation of 2,4-DCP was also confirmed by the analysis of chloride ions (Fig. S3). It can be seen that the concentration of chloride ion was 0.99 mg L−1 at 300 min in the Fe3O4/PS system, indicating the dechlorination of TCS.

3.5.2. Residual PS Fig. 6a presents the variation of residual PS in different systems as a function of time. Approximately 47.0% of PS was rapidly decomposed in the first 10 min, and the decomposition rate gradually decreased before becoming relatively stable. This attributes to PS being activated by Fe3O4, which improve PS decompose rate. However, PS concentration in the PS-only system decomposed slightly in the first 10 min, and became relatively stable,

3.5.4. Characterizations of the Fe3O4 For an improved understanding of mechanism of PS activation and 7

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Fig. 7. XPS spectra of the fresh and reacted Fe3O4 MNPs catalyst: (a) Fe 2p, (b) O1s core level, (c) full-range scan of the samples, and (d) Cr 2p core level. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [Cr(VI)]0 = 1.0 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C.

Cr(VI) reduction, the XPS spectra of Fe3O4 were compared before and after the reaction (Fig. 7). Fig. 7a shows that the iron band at 710.6 and 713.0 eV in Fe3O4 before use was assigned to Fe (2p). The observed position is consistent with previous studies on iron assignments of magnetite [45]. Binding energy values in Fe3O4 after use were slightly increased, which suggests changes in the Fe(II) and Fe(III) fractions in Fe3O4. Based on the Fe (2p) envelope, the Fe(III) fraction increased to 9.1% after use, which suggests that electron capture occurred for part of the Fe(II) during treatment. In addition, the typical XPS spectrum of the O1s region was decomposed, as shown in Fig. 7b. Moreover, three peaks at 529.6, 530.6, and 532.3 eV validated that lattice oxygen (O2−), hydroxide (OH−), and H2O, respectively, coexisted in Fe3O4 before use, which is consistent with previous studies on iron oxide surfaces [45,64]. The fractions of O2−, OH−, and H2O in Fe3O4 after use were changed after activation. Based on deconvolution of the O2−, OH−, and H2O envelopes, the fraction of O2− in the catalyst was increased after the treatment, which was accompanied with the decreasing H2O and OH− fractions, thus indicating that H2O and OH− on the catalyst participated in the reactions and generated H+ and Cr–Fe (oxy) hydroxide [45,65]. To obtain insights into the Cr(VI) removal mechanism by using Fe3O4/PS, the variations in chromium concentrations in the liquid and solid phase were studied during the reaction. XPS analyses of Fe3O4 were performed before and after the reaction. Fig. 7c shows that compared with Fe3O4 before the reaction, evident Cr peaks emerged in the XPS analysis after the reaction, indicating that chromium was adsorbed by Fe3O4. The valence state of the Cr species on the surface of the used Fe3O4 was analyzed through XPS (Fig. 8d). The strong signals of Cr(III) at 576.3and 586.1 eV validate the reduction and collection of Cr(VI) in the form of Cr(III) on the Fe3O4 surface, which further indicates the reduction reaction effect on Cr(VI) removal. Fig. 8 illustrates

Fig. 8. Concentrations of chromium species as a function of time. Experimental conditions: [Fe3O4]0 = 2 g L−1, [TCS]0 = 5 mg L−1, [Cr(VI)]0 = 1.0 mg L−1, [PS]0 = 1.0 mM, and T = 30 °C.

the detailed reaction process. With the rapid decrease in Cr(VI) in the solution, Cr(III) was detected. The concentration of Cr(III) was increased to 0.41 mg L−1 at 60 min and then gradually decreased to 0.38 mg L−1 at 300 min, which may be attributed to co-precipitation with Fe ions. Moreover, the total chromium on Fe3O4 surface was 0.29 mg g−1 at 300 min. Based on the above analysis, a probable catalytic mechanism of PS activated by Fe3O4 for TCS degradation was put forward in Fig. 9. The Fe(II) on the surface of the Fe3O4 would activate PS to produce the SO4∙− (Eqs. (17)). ∙OH can be produced by the reaction of SO4∙− with 8

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System. The procedure of TCS and about nine kind of intermediates identification were summarized in Table S2 and Fig. S6–S16. According to the detected intermediates during reaction, three possible TCS degradation paths were proposed (Fig. 10). As for pathway A, hydroxylation was the degradation route of TCS, which could generate intermediates ·OH-Triclosan (P-2) and P-3 [67,68]. Formation of P-3 was the hydroxylated products of ·OH-Triclosan (P-2) either by molecular oxygen or disproportionation. ·OH-Triclosan (P-2) may be attacked by reactive species ∙OH , causing the formation of intermediates 2,4-Dichlorophenol(P-4) and Chlorohydroquinone (P-5). The formation of intermediates 2,4-Dichlorophenol(P-4) and may also be explained by the homolytic reaction of the C-O bond occurrence, while Chlorohydroquinone (P-5) is the result of a disproportionation process. The ether bond of P-3 was broken to generate P-5 [68]]. In pathway B, the breakage of ether bond of TCS is the main reaction route [67]. As C-O bond on the ether group was the relative weak site, and firstly attacked by ∙OH and SO4∙−, resulting to the breakage of ether bond and the generation of 2,4-Dichlorophenol (P-4) and 4-Chlorobenzene-1,2-diol (P-7) [69]. As one of the chlorine ion of 2,4-Dichlorophenol (P-4) benzene ring would be replaced by a hydroxyl group to produce Chlorohydroquinone (P-5), and Chlorohydroquinone (P-5) would be further oxidized to 2-Chloro-1,4-benzoquinone (P-6). In pathway C, direct dechlorination of TCS occurred, and to primarily produce intermediates P-8. The chlorine ion of P-8 would be replaced by a hydrogen to produce P-9. A similar phenomenon was reported in previous study [70]. As the breakage of ether bond of P-9 by attacking of excessive ∙OH and SO4∙−, p-chlorophenol(P-10) could be formed. Ultimately, intermediates P6,P7 and P10 were observed for ring-opening reaction, which could be decomposed to CO2, H2O and Cl−.

Fig. 9. Schematic of the mechanisms of Cr(VI) removal and TCS degradation from aqueous solution by using the Fe3O4/PS system.

H2O (Eqs. (18)). They will take part in TCS degradation process. On the other hand, Fig. 9 shows the stepwise removal mechanism of Cr(VI) in the Fe3O4/PS system. A four-step mechanism of Cr(VI) removal was proposed [40]. Initially, Cr(VI) was in contact with the reactive sites of Fe3O4 and absorbed on the Fe3O4 surface [66]. A reduction reaction was observed between Cr(VI) and Fe(II). Cr(VI) and Fe(II) were reduced to Cr(III) and oxidized to Fe(III), respectively, as expressed using Eq. (19). Finally, Cr(III) and Fe(III) were synchronously co-precipitated on the Fe3O4 surface, thus forming Cr–Fe (oxy) hydroxide (Eqs. (19–21)) [39].

≡ Fe(II) + S2 O82 − → ≡Fe(III) + SO·4− + SO24− H2 O+

SO·4−

→ H+ +

SO24−

(17)

+ ·OH

(18)

Cr2O72 − + 6Fe2 + + 14H+ → 2Cr 3 + + 6Fe3 + + 7H2 O

(19)

(1 − x)Fe3 + + xCr 3 + + 3H2 O → Crx Fe1 − x (OH)3

(20)

(1 − x)Fe3 + + xCr 3 + + 2H2 O → Crx Fe1 − x OOH + 3H+

(21)

3.7. Reusability of Fe3O4 To evaluate the reusability of the catalyst, Fe3O4 was used four times to activate PS. The TCS and Cr(VI) removal efficiency and the concentration of leached Fe ions were determined during the following catalytic experiments. After each run, the catalysts were collected using a magnet, rinsed with distilled water, and then recycled for the next batch. Fig. S4 shows that when the Fe3O4 cycle time increased, the degradation efficiency of TCS and Cr(VI) removal gradually decreased.

3.6. Degradation pathway TCS and its reaction intermediate products during reaction were analyzed by using Thermo Scientific Q Exactive Orbitrap LC–MS/MS

Fig. 10. Degradation pathways of TCS in Fe3O4/PS system. 9

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However, the removal efficiencies of TCS and Cr(VI) were 30.5% and 89.4%, respectively, when Fe3O4 was used until the fourth cycle. The decreasing removal efficiency indicated that Fe3O4 was slightly deactivated. This observation can be attributed to the following reasons: (1) Cr(III) adsorbed on the Fe3O4 surface inhibited the interaction between PS and Fe3O4; and (2) the concentration of leached total Fe ions was approximately 1.86 mg L−1 for the first run (Fig. S5), thus accounting for 0.12% of the total Fe ions content in the catalyst, respectively. Although the simultaneous removal of Cr(VI) and TCS was achieved in the Fe3O4/PS system, some common concerns were observed regarding Fe leaching, and reuse of this material for practical applications. Thus, further studies must be conducted in the future.

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4. Conclusion This study demonstrated that the Fe3O4/PS system could effectively remove Cr(VI) and TCS through the simultaneous involvement of adsorption and degradation. The presence of PS could not significantly reduce Fe3O4 reactivity toward Cr(VI) reduction; however, Cr(VI) had a slight effect on the TCS oxidation rate. The removal rate of Cr(VI) and TCS decreased with the presence of CO32−, HPO42− and HA because of the competition for active sites of Fe3O4 and scavenging of reactive species. In addition, a four-step mechanism of Cr(VI) removal was proposed, involving adsorption, reduction, immobilization, and conversion. Furthermore, nine kinds of main intermediates were identified for the degradation of TCS. These results suggest that the Fe3O4/PS system is a cost-effective technology that is favorable for the application of chromium and TCS removal in drinking water bodies. Acknowledgements This work was supported by the program of Natural Science Foundation of Shanxi Province , China (201801D121274), the Key Research and Development Projects of Shanxi province, China (201803D421098) and the Research Grants Council of the Hong Kong Special Administrative Region, China (project:UGC/FDS25/E16/17). This manuscript was edited by Wallace Academic Editing. Appendix A. Supplementary data Supplementary data to this article can be found online at https:// doi.org/10.1016/j.cej.2019.122586. References [1] Z. Yang, S. Jia, T. Zhang, N. Zhuo, Y. Dong, W. Yang, Y. Wang, How heavy metals impact on flocculation of combined pollution of heavy metals–antibiotics: A comparative study, Separ. Purif. Technol. 149 (2015) 398–406. [2] A. Chen, C. Shang, J. Shao, Y. Lin, S. Luo, J. Zhang, H. Huang, M. Lei, Q. Zeng, Carbon disulfide-modified magnetic ion-imprinted chitosan-Fe(III): A novel adsorbent for simultaneous removal of tetracycline and cadmium, Carbohydr. Polym. 155 (2017) 19–27. [3] Y. Zhou, F. Zhang, L. Tang, J. Zhang, G. Zeng, L. Luo, Y. Liu, P. Wang, B. Peng, X. Liu, Simultaneous removal of atrazine and copper using polyacrylic acid-functionalized magnetic ordered mesoporous carbon from water: adsorption mechanism, Sci. Rep. 7 (2017) 43831. [4] D. Wu, B. Pan, M. Wu, H. Peng, D. Zhang, B. Xing, Coadsorption of Cu and sulfamethoxazole on hydroxylized and graphitized carbon nanotubes, Sci. Total Environ. (2012) 427247–428252. [5] M. Adolfsson-Erici, M. Pettersson, J. Parkkonen, J. Sturve, Triclosan, a commonly used bactericide found in human milk and in the aquatic environment in Sweden, Chemosphere 46 (2002) 1485–1489. [6] D. Kolpin, E. Furlong, M. Meyer, E.M. Thurman, S. Zaugg, L. Barber, H. Buxton, Pharmaceuticals, Hormones, and Other Organic Wastewater Contaminants in U.S. Streams, 2005. [7] S. Heinz, M. Stephan, T. Céline, P. Laurent, Triclosan: occurrence and fate of a widely used biocide in the aquatic environment: field measurements in wastewater treatment plants, surface waters, and lake sediments, Environ. Sci. Technol. 36 (2002) 4998. [8] J.L. Zhao, G.G. Ying, Y.S. Liu, F. Chen, J.F. Yang, L. Wang, Occurrence and risks of triclosan and triclocarban in the Pearl River system, South China: from source to the receiving environment, J. Hazard Mater. 179 (2010) 215–222.

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