Efficient pyrite activating persulfate process for degradation of p-chloroaniline in aqueous systems: A mechanistic study

Accepted Manuscript Efficient pyrite activating persulfate process for degradation of p-chloroaniline in aqueous systems: a mechanistic study Yongqing Zhang, Hien Phuong Tran, Xiaodong Du, Imtyaz Hussain, Shaobin Huang, Shaoqi Zhou, William Wen PII: DOI: Reference:

S1385-8947(16)31348-1 http://dx.doi.org/10.1016/j.cej.2016.09.104 CEJ 15814

To appear in:

Chemical Engineering Journal

Received Date: Revised Date: Accepted Date:

26 May 2016 20 September 2016 21 September 2016

Please cite this article as: Y. Zhang, H. Phuong Tran, X. Du, I. Hussain, S. Huang, S. Zhou, W. Wen, Efficient pyrite activating persulfate process for degradation of p-chloroaniline in aqueous systems: a mechanistic study, Chemical Engineering Journal (2016), doi: http://dx.doi.org/10.1016/j.cej.2016.09.104

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Efficient pyrite activating persulfate process for degradation of pchloroaniline in aqueous systems: a mechanistic study Yongqing Zhanga,b,c,e*, Hien Phuong Trana, Xiaodong Dua, Imtyaz Hussaina, Shaobin

Huanga,b,c, Shaoqi Zhoua, d, e, William Wenf

a

School of Environment and Energy, Guangdong Provincial Key Laboratory of Atmospheric

Environment and Pollution Control, South China University of Technology, Guangzhou 510006, PR China b

The Key Lab of Pollution Control and Ecosystem Restoration in Industry Clusters, Ministry of Education, Guangzhou 510006, PR China

c

State Key Laboratory of Pulp and Paper Engineering, South China University of Technology, Guangzhou 510640, PR China d e

Guizhou Academy of Sciences, Shanxi Road 1, Guiyang 550001, PR China

State Key Laboratory of Subtropical Building Sciences, South China University of Technology, Guangzhou 510641, PR China

f

Centre for Clean Environment and Energy, Environmental Futures Research Institute, Griffith School of Environment, Griffith University, Gold Coast campus, QLD 4222, Australia

* Corresponding author: Yongqing Zhang Tel: +86-20-39380569 Fax: +86-20-39380508 E-mail address: [email protected]

1

ABSTRACT In recent years, persulfate activation systems have received increasing attention due to their high oxidation reactivity when removing environmental pollutants. Pyrite, the most common metal sulfide on Earth’s surface, can supply abundant Fe2+ for persulfate activation. The role of the generated reactive oxygen species (ROS) in persulfate-pyrite systems however, is not fully understood. In this study, batch experiments were used to investigate p-chloroaniline (PCA) degradation by a pyrite-persulfate system. The effects of pyrite dosage, pH, temperature, air conditions (aerobic vs. anaerobic) and pyrite particle size on PCA degradation were examined. Radical detection was conducted using electron paramagnetic resonance (EPR) methods. Results from the EPR spectra indicated that PCA degradation was achieved by sulfate radical and hydroxyl radical oxidation. Aerobic conditions were more beneficial to PCA degradation than anaerobic conditions due to the generated superoxide radicals (O2•-) that activated the persulfate to produce more sulfate radicals (SO4•-). PCA degradation also increased with higher pyrite doses and under acidic conditions (pH 3.0 and 5.0). PCA was removed completely at pH 3.0 after 60 min. Temperature increase from 10 to 50 ◦C significantly promoted PCA degradation. These findings provide new understanding of the mechanism involved in pyrite activation of persulfate which can be used to improve PCA degradation by pyrite-persulfate systems. Key words: p-chloroaniline; Pyrite; Persulfate; Sulfate radical; Superoxide radical; Hydroxyl radical

2

1. Introduction In recent years, advanced oxidation processes (AOPs) have become a promising technology for the degradation of recalcitrant organic pollutants in aqueous solution [1]. In AOPs, persulfate (PS) activation technology has received increasing attention due to its higher redox potential (E0 = 2.01 V). Persulfate is a suitable oxidant source because of its stability whether in solution or as a solid, transport-convenience and cost-efficiency [2]. More importantly, a stronger oxidant (E0 = 2.6 V), sulfate free radical SO4•-, can be generated from PS activation [3]. In aqueous solution, hydroxyl radicals can be produced by sulfate radicals and may be partly responsible for the oxidation of contaminants [3]. Iron has been widely used as an activator for persulfate as it is nontoxic, cheap and naturally abundant [4-8]. However, SO4•- can be rapidly scavenged by Fe2+ in Fe(II) - persulfate systems (Eqs. 1-3), which results in decreasing efficiency of pollutant degradation [8]. S 2 O 82 − + 2Fe 2 + → 2Fe 3+ + 2SO 24 −

（1）

•− S2O82− + Fe 2+ → Fe3+ + SO 2− 4 + SO 4

（2）

SO•4− + Fe2 + → Fe3+ + SO24 −

(3)

Persulfate could be effectively activated by heat, UV-Visible light, base, carbon materials and transition metal. While heat and light require considerable energy input, base needs chemical invest, carbon materials are restricted in lab studies and transition metals are related to second pollution of heavy metal. Therefore all these problems should be taken into account for persulfate application of in situ chemical oxidation (ISCO). Pyrite is natural occurring mineral. As an abundant iron-containing species, pyrite shows the potential to activate persulfate in a high-efficiency and cost-benefit way by not only avoiding large energy and chemical input but also providing enough environment-friendly activator.

3

High content of iron in pyrite is an attractive reason to research on this material. Iron minerals have attracted attention as alternative activators in Fenton-like chemical oxidation processes [9]. Teel et al. [10] reported that 13 natural minerals can potentially facilitate persulfate decomposition and produce reactive oxygen species (ROS). Pyrite (FeS2), the most common metal sulfide on the Earth’s surface, can activate persulfate under both aerobic [9, 11] and anaerobic [12] conditions (Eqs. 4-9) due to the abundant Fe2+ supply. Therefore pyrite is a promising persulfate activator and the activation mechanism is remained to be confirmed.

2FeS2 + 7O2 + H 2 O → 2Fe2+ 4SO42- + 4H+

（4）

2Fe2 + + O2 + 4H+ → 2Fe3 + + 2H2O

（5）

2FeS2 + 14Fe3+ + 8H 2O → 15Fe2+ + 2SO42- + 16H +

（6）

FeS2 + 8H2O → Fe2 + + 2SO24 - + 16H+ + 14e-

（7）

2FeS2 + 15S2O82- + 16H2O → 2Fe3+ + 34SO24- + 32H+

（8）

2FeS 2 + 2S2O82 - → Fe 2 + + 2SO •4- + 2S

（9）

As one of common pollutants, chlorinated anilines present in wastewater, sludge and agricultural soils; and are intermediates of synthetic organic chemical and polymer production [13]. In a previous study [14], we investigated the degradation of p-chloroaniline (PCA) and the ROS in a pyrite system, including O2•-, H2O2 and •OH radicals. Although the oxidation of PCA by ZVI (zero-valent iron) has been thoroughly investigated [15], PCA degradation details and the effects of the generated ROS in persulfate-pyrite systems are not fully understood. In this study, batch experiments were used to investigate p-chloroaniline (PCA) degradation in a persulfatepyrite system. Quenching studies and Electron Paramagnetic Resonance (EPR) techniques were applied to identify the radical species generated in this process. Solution pH, temperature, pyrite 4

concentration and air conditions were controlled to examine their effects on PCA degradation. The main objectives of this work were to (i) determine the mechanism involved in pyrite activation of persulfate, and (ii) elucidate the main reaction oxidation species in the degradation of pollutants. 2. Experimental 2.1. Materials All chemicals used in this study were reagent grade. p-Chloroaniline (C6H6ClN, > 99.5% purity) and methanol were purchased from Sigma-Aldrich (USA). 2,9-dimethyl-1, 10phenanthroline (C14H12N2, DMP), 1,4-benzoquinone (C6H4O2, BQ) and 1,10-phenanthroline monohydrate (C12H8N2.H2O) were purchased from Aladdin Reagents Co. Ltd. (Shanghai, China). Ethanol (EtOH) was purchased from Shanghai Reagents Co. Ltd. (Shanghai, China). Sodium thiosulfate pentahydrate (Na2S2O3·5H2O, > 99.0%) was obtained from Guangzhou Chemical Reagents Factory (Guangzhou, China). Concentrated sulfuric acid (H2SO4, > 98.0%) and sodium hydroxide (NaOH, > 96.0%) were purchased from Tianjin Baishi Chemical Reagents Co. Ltd. (Tianjin, China). Pyrite samples were collected from the Dabaoshan sulfur-polymetallic mines in the north of Guangdong Province, China. The main chemical composition of the pyrite samples is described in our previous paper [14]. Water used in this study was deionized distilled water produced by a Millipore Milli-Q system (USA). 2.2. Experiments Batch experiments were performed in 250 mL Erlenmeyer flasks at 25 ◦C in a rotary shaker. PCA and persulfate stock solutions with concentrations of 10 mM and 50 mM, respectively, were prepared in ultrapure water before each experiment. Reaction system pH values were 5

adjusted using 1 M NaOH and H2SO4 solutions. The reaction solutions were purged with N2 for 20 min to remove dissolved oxygen before the addition of PCA. A volume of 0.5 mL PCA stock solution was added to 94.5 mL of deionized water to prepare reaction mixtures which were fixed on a rotary shaker at 125 rpm and 25◦C. After 10 min, 0.5 mL of the mixture was collected from the reactor to determine the initial PCA concentration ([PCA]0). The reaction was initiated by adding 5 ml sodium persulfate (50 mM) and a predetermined amount (0.05 g, 0.5 g/L) of pyrite particles to the reactor. At regular time intervals, a 0.5 ml aliquot of solution was added to a 2 mL vial with 0.5 mL of 2 M Na2S2O3 to quench the reaction. 1 ml of 0.1 M DMPO was added to 10 ml of the reaction mixture to react for 20 min before EPR analysis. 2.3. Analytical methods Pristine pyrite and used pyrite (pyrite after reaction with water, and pyrite after reaction with persulfate solution) were characterized by X-ray diffraction (XRD, Bruker advance 8), Scanning Electron Microscope (SEM, EVO 18) and X-ray photoelectron spectroscopy (XPS, PHI X-tool). A Shimadzu LC-20A high performance liquid chromatography (HPLC) system was used to measure the concentration of PCA, with the UV detector set at 254 nm. A reversed-phase C18 column (25 cm x 4.6 mm, 5 µm) was used. Methanol-water (30:70, v/v) was used as the mobile phase and the flow rate was 0.5 ml/min. External PCA standards were run with the samples. An UV-vis spectrophotometer at 510 nm was used to colorimetrically determine Fe2+ concentration by forming a colored complex with the addition of 1,10-phenanthroline [16]. The formation of hydrogen peroxide was analyzed using the copper(II) and DMP method [17,18] where the optical absorbance was measured at 551 nm. pH values were determined by a Shanghai Leici E-201-C pH meter. An EPR (Electron Paramagnetic Resonance) Spectrometer Bruker A300 was used to identify the radicals, and was set up as follows: microwave frequency at 9.8752 GHz, sweep 6

width at 100 G, center field at 3504.12 G, modulation amplitude at 1 G, time constant at 327.68 ms, and scan time of 40 s. 3. Results and discussion 3.1. Characterization Pyrite crystal structure, morphology, composition and surface modification after reaction were tested by XRD, SEM, XRF and XPS, and all the details of result were stored in support information(XRD, SEM, XRF and XPS). The XRD result indicated that the pristine pyrite is composed of FeS2 crystal (PDF#71-1680) and FeSO4(H2O) crystal (PDF#81-0019), and when it is introduced in water or persulfate solution, the FeSO4(H2O) crystal is dissolved and pure FeS2 crystal remained. The triclinic FeS2 crystal were characterized by 14 relatively high intensity peaks (correspond to 14 lattice planes). The monoclinic FeSO4(H2O) crystal were characterized by 11 relatively high intensity peaks (correspond to 11 lattice planes), which scattered between FeS2 characterized peaks and the intensity is much weaker (correspond to much lower content compared with FeS2). From the images obtained from SEM, the pristine pyrite was the mixture of FeS2 and FeSO4(H2O) crystals which tangled together. When it is introduced in water or persulfate solution, FeSO4(H2O) dissolved and the pure FeS2 crystal appear with glossy planes with clear edges. When magnification increases, the glossy planes become bumpy and cracked, indicating the dissolution of adherent FeSO4(H2O) and FeS2 as well. Compared with water, Persulfate leads to more and larger cracks, which is ascribed to greater FeS2 dissolution. XPS test was conducted with pristine pyrite, pyrite after reaction with water, and pyrite after reaction with persulfate solution. After the fitting of Fe, S and O elements peaks, it was concluded that pristine pyrite surface was mainly covered by ferrous and ferric sulfate (the characterized Fe and S peaks of ferrous and ferric sulfate were dominant) and bulk FeS2 (the characterized Fe and S peaks of 7

FeS2 were dominated) was underneath. After reaction with water and persulfate solution, sulfate species dissolved and FeS2 came out (the characterized Fe and S peaks of FeS2 were dominant), and FeOOH deposit from dissolved iron ions precipitation appeared (the characterized Fe and S peaks of ferrous and ferric sulfate degraded, the characterized Fe peak of FeOOH appeared). The ferrous sulfate dissolution supported by above evidences indicated its contribution to initial Fe2+ concentration. The XRF result showed that Fe and S elements took up 93.33% mass fraction of pristine pyrite, which indicated the high content of FeS2. Apart from Fe and S, Si, Al, K, Ca, Mg, Zn and other metal elements also existed. 3.2. Degradation of PCA by pyrite activated persulfate system More PCA was degraded oxidatively in the pyrite-persulfate system than in the pyrite or persulfate only systems (Fig. 1). Only about 10.31% PCA was removed from solution over a 240 min period by pyrite alone. This result indicated that most PCA removal in the pyrite system was due to oxidative degradation [15,19-22] or sorption [20] on the pyrite surface. In the control with persulfate but no pyrite, PCA removal was 30.63% over the same time period. This result was consistent with observations by Hussain et al. [15]. In contrast, the addition of 0.5 g/L of pyrite to the persulfate solution resulted in much more rapid and prominent PCA degradation, with 91.02% removal in 240 min. This result indicated that in the pyrite-persulfate system, the pyrite was able to activate persulfate to degrade organic pollutants. Liang et al. [11] reported that pyrite could activate persulfate to degrade methyl tert-butyl ether (MTBE). They suggested that sulfate free radicals generated from persulfate activation by the pyrite and Fe(II) from pyrite played a key role in pollutant degradation. The sulfate and hydroxyl radicals were identified by radical scavengers namely Phenol, methanol and tert-butyl alcohol (TBA). During the experiments, scavengers amounted to 1000 8

times of PCA molar concentration were added in FeS2/PS/PCA system respectively. The Fig. 2 showed that PCA removal rate was decreased to 70.44%, 45.66% and 30.99% with the addition of TBA, methanol and phenol respectively, compared with 91.02% without scavengers addition. The inhibition effect of TBA indicated the hydroxyl radical generation, while that of methanol and phenol delivered that sulfate radical was dominant. Therefore, sulfate and hydroxyl radicals were reactive oxygen species that led to PCA degradation. 3.3. Effect of pyrite concentration on PCA degradation To explore the effect of pyrite dose on PCA degradation in the presence of persulfate, experiments were conducted by adding 0.1 g/L, 0.3 g/L, 0.5 g/L, 0.7 g/L and 1.0 g/L of pyrite to solution. The initial PCA concentration was fixed at 0.1 mM, persulfate concentration was 2.5 mM and initial pH was 7.0. The degradation of PCA was 100% at 0.7 g/L and 1.0 g/L of pyrite after 180 and 120 min, respectively. When the pyrite dose was 0.1 g/L, 0.3 g/L and 0.5 g/L, the degradation of PCA after 240 min was 35.0%, 65.0% and 91.02%, respectively. In addition, persulfate decomposition rapidly increased with increasing pyrite dosage (Fig. S1). These results showed that increasing pyrite dosage improved both PCA degradation efficiency and PS decomposition efficiency. To clarify the role of ferrous ion in the system, the ferrous ion concentration evolution in the pyrite-persulfate process was also determined. The ferrous ion concentration increased with increasing pyrite dosage, indicating that pyrite was a source of ferrous ion (Fig. 3). These results were in agreement with previous studies [15,30,31] ascribing persulfate activation by pyrite to ferrous ion released from pyrite in the persulfate–pyrite system. Ferrous ion is one of the main species that can efficiently activate persulfate to produce sulfate radicals ( SO •4- ).

9

3.4. Effect of initial pH on PCA degradation Previous studies have shown that solution pH has a strong influence on sulfate radical oxidative degradation of organic pollutants [3,23]. Acidic conditions favored oxidations of TCE in a ferrous iron activated persulfate system [24]. pH higher than 4.0 led to the precipitation of Fe3+ ions to form oxyhydroxides [26], which have low efficiency for activating persulfate to form sulfate radicals [27,28]. The following equations (Eqs. 10-12) demonstrate the processes of Fe3+ oxyhydroxide formation [28].

Fe3+ + H2O → FeOH2+ + H+ (k = 2.3 x10-7s-1)

（10）

Fe3+ + 2H 2O → Fe(OH)22 + + 2H + (k = 4.7 x10-7s-1)

（11）

2Fe3+ + 2H2O → Fe2 (OH)22 + + 2H + (k = 1.1 x10-7s-1)

（12）

In acidic pH conditions, persulfate generated more sulfate radicals [29] through the following equations. S2O82- + H + → SO •4- + HSO 4

（13）

SO•4- + SO•4- → S2O82-

（14）

SO •4- + S2O 82 - → SO 24 - + S2O8• -

（15）

Therefore, to explore the effect of pH on PCA degradation in pyrite-persulfate systems, the batch experiments were conducted with initial pH of 3.0, 5.0, 7.0, 9.0 or 11.0. The initial PCA concentration was fixed at 0.1 mM, persulfate concentration was 2.5 mM and pyrite concentration was 0.5 g/L. PCA degradation rate declined with increasing solution pH (Fig. 4). The elimination of PCA was 100% at pH 3.0 and pH 5.0 after 60 min and 180 min, respectively. At pH 9.0 and 11.0, PCA removal after 240 min reaction time was 73.49% and 44.08%, respectively. In the study of pH and Fe2+ concentration evolution in pyrite-deionized water 10

system at five initial pH, all the pH was decreased to a specific value respectively and Fe2+ concentration was higher at lower initial pH (Fig. S2, 5). Much lower Fe2+ was detected when the initial pH was 9 and 11 since the pH during the reaction would be more than initial pH 4 though with the trend of decrease. This indicated that acidic pH was more beneficial for PCA degradation as pyrite released more Fe2+ under these conditions [10,25,26]. The higher released Fe2+ would enhance the persulfate activation and finally promoted PCA removal. 3.5. Effect of pyrite particle size on PCA degradation Li et al. [32] reported that ZVI of various sizes exerted different levels of ability to activate persulfate to remove acid orange 7, nano-ZVI (approx. 50 nm) > micro-ZVI (approx. 150 µm) > milli-ZVI (approx. 1 mm). Therefore, experiments were designed to investigate the effect of pyrite particle size on PCA degradation in the pyrite-persulfate system. Pyrite particle size was set as 20 mesh, 40 mesh, 60 mesh, 100 mesh or 200 mesh. The initial PCA concentration was fixed at 0.1 mM, persulfate concentration was 2.5 mM and initial pH was 7.0. Pyrite particle size had a remarkable impact on the degradation rate of PCA (Fig. 6). PCA removal rate after 240 min was 38.99%, 44.01%, 56.0%, 77.0% and 91.02% with pyrite particle size of 20 mesh, 40 mesh, 60 mesh, 100 mesh and 200 mesh, respectively. The ferrous ion concentration in reaction mixtures increased with reduced pyrite particle size (Fig. 7). This indicated that smaller pyrite particles with higher total surface area, released more dissolved ferrous ions. The decomposition of persulfate rapidly increased with reduction in pyrite particle size (Fig. S3). To distinguish PCA adsorption on pyrite which would contribute to PCA removal, we performed the experiment of PCA absorption on pyrite with 20 mesh, 40 mesh, 60 mesh, 100 11

mesh and 200 mesh particle sizes in the absence of persulfate. PCA degradation rate for these pyrite particle sizes was 5.42%, 6.74%, 7.49%, 8.3% and 10.31%, respectively (Fig. S4). This result confirmed that oxidation was the main PCA degradation pathway in the pyrite-persulfate system. 3.6. Effect of temperature on PCA degradation To explore the effect of temperature on PCA degradation, experiments were conducted at 10◦C, 20◦C, 30◦C, 40◦C and 50◦C. The results demonstrated that PCA degradation increased with temperature (Fig. 8). This was due to greater sulfate free radical generation under higher temperature, which was responsible for the increased PCA oxidation. The degradation of PCA after 240 min was 79.0% and 91.02% at 10◦C and 20◦C, respectively, while 100% degradation was achieved at 30◦C, 40◦C and 50◦C after 240 min, 180 min and 120 min, respectively. In contrast, our previous work showed that PCA removal rate increased from 20.93% to 58.16% when the reaction temperature increased from 15◦C to 50◦C in a persulfate-only system [15]. Based on the PCA degradation in these two systems, the persulfate-pyrite system is better in removing organic pollutants in subsurface environments since heat activation needs more energy. 3.7. Effect of dissolved oxygen on PCA degradation Dissolved oxygen has an impact on PCA degradation as oxygen is an electron acceptor. The PCA removal rate was 100% after 180 min with air purging and after 240 min without air purging, while only 89.02% of PCA was degraded under N2 purging (Fig. 9). These results indicated that increasing dissolved oxygen concentration increased the PCA degradation rate in the reaction solution. Furthermore, the decomposition of persulfate also increased with increasing O2 concentration (Fig. S5(a)). The decomposition rate of persulfate under air purging, 12

without purging and with N2 purging was 51.87%, 40.91% and 31.27%, respectively. Increasing O2 concentration could lead to the generation of more superoxide radicals (O2•-), which would activate the persulfate to produce SO•4- [33] (shown in Eq. 16). To test the effects of dissolved oxygen on the formation of O2•-, generated hydrogen peroxide concentration was determined and 1,4-benzoquinone (BQ) was selected as an O2•- quencher in the reaction system [34]. 1,4benzoquinone clearly exhibited the ability to inhibit PCA degradation under air purging conditions (Fig. 10). Furthermore, the persulfate decomposition rate was also lower (Fig.S5(b)). These results present the delayed formation of O2•- during the reaction. Results in Fig. S6 showed that hydrogen peroxide concentration increased in the order of N2 purging-pyrite system, no purging-pyrite system and air purging-pyrite system, it is evidence that dissolved oxygen would be reduced to O 2•- which further transferred to hydrogen peroxide. It is noted that molecular oxygen is paramagnetic with two unpaired electrons, making it unlikely to react directly with diamagnetic species such as pyrite due to spin restriction [35], therefore other reactive species were formed to reacted with oxygen. When persulfate was introduced into the system, the pyrite was oxidized to form ferric ions which would further oxidize pyrite to produce ferrous ions. Because absorption of ferrous ions on the pyrite surface changes the sulfur combined iron magnetism, it reacts with oxygen more easily. Ferrous oxyhydroxides would form, which are better reactants with oxygen [36]. The ferrous species could donate one electron to oxygen to generate superoxide radicals which would subsequently react with persulfate to form sulfate radicals.

Fe 2 + + H 2 O → Fe(OH) + + H +

(16)

Fe(OH) + + H 2 O → Fe(OH) 2 + H +

(17) 13

Fe(II) + O 2 → O •2- + Fe(III)

(18)

S 2O 82 - + O •2- → SO 42 - + SO •4- + O 2

(19)

3.8. Electron Paramagnetic Resonance study In this study, electron paramagnetic resonance (EPR) was used to study the mechanism of persulfate activation by directly detecting radicals. Effects of solution pH and oxygen content on the persulfate activation mechanism were investigated in the pyrite-persulfate system. Three experiments at pH 3, 7 and 11 without air purging and two experiments at pH 3 and 7 with air purging were conducted. Sulfate and hydroxyl radicals were observed in the pyrite-persulfate system (Fig. 11). Results also showed that acidic and aerobic conditions favored the generation of radicals. In the pH 7 system without air purging however, the radical signal was too weak to observe. In contrast, the radical signal increased in the pH 7 system with air purging (shown in Fig. 17b). It confirmed that increasing dissolved oxygen concentration had a positive impact on the production of radicals. It has been reported that SO4•- generated from persulfate activation was the dominant reactive oxygen species under acidic conditions [11]. There were hydroxyl radicals mainly in the pH 11 system due to the base activation of persulfate (Fig. 11a). In addition, when the pH was 3 or after air purging for 20 min at both pH 3 and 7, sulfite radicals were clearly identified according to the hyperfine coupling constants reported by Mottly and Ranguelova [37,38], where aβH and aN were 16.0G and 14.7G, respectively. Pyrite could be oxidized by dissolved oxygen and ferric ions to form sulfoxy intermediates such as sulfite (SO32), thiosulfate (S2O32-) and polythionate (SnO62-, n = 4,5) [35]. These sulfur species could then react with sulfate radicals or superoxide radicals to form sulfite radicals. The sulfite radicals would react preferentially with unsaturated bonds since they are sensitive to steric effects [39]. 14

The increased production of sulfate radicals or superoxide radicals and higher sulfite ions supply would result in more sulfite radicals due to the reaction equilibrium between sulfate radicals and sulfite radicals. Sulfite radicals were observed under lower pH and high oxygen conditions (Fig. 11) because of DMPO with C=N bonds. Both sulfate and hydroxyl radicals were the main reactive oxidative species in the system with target pollutants and played the leading role in PCA degradation. Relevant reactions are typified by equations 20-27: (FeS 2 ) + (O 2 ) + Fe(III) → (SO 42 - + SO 32 - + S 2 O 32 - + S n O 62 - )

(20)

S 2 O 32 - + S n O 62 - → S n +1O 62 - + SO 32 -

(21)

8S 2 O 32 - + H + → S8 + 7SO 32 - + HSO -3

(22)

2− •− 2− SO•− 4 + SO 3 → SO3 + SO 4

(k>2×109M-1S-1)

+ •− SO32− + O•− 2 + 2H → SO3 + H 2 O 2

(23)[40]

(24)

•− SO•− 3 + O 2 → SO 5

(25)[41]

2− •− 2− SO•− 5 + SO 3 → SO 4 + SO 4

(26)

•− SO•− 3 + O 2 → SO3 + O 2

(27)

4. Conclusions

This study suggests that the persulfate-pyrite system is an effective oxidant process for PCA degradation. In the persulfate-pyrite system, Fe2+ released from the pyrite facilitated persulfate decomposition to form sulfate radicals. Increased pyrite dosages and dissolved oxygen concentration resulted in enhanced PCA degradation. The optimum pH for PCA degradation in 15

the persulfate-pyrite system was 3.0 where 100% removal rate was achieved at 60 min. PCA degradation was enhanced under aerobic conditions when compared with anaerobic conditions. The superoxide radical (O2•-) generated in the persulfate-pyrite system could react with persulfate to form more sulfate radicals (SO4•-), which promoted the oxidation of PCA. The finding of sulfite radicals (SO3•-) indirectly proved the generation of sulfate radicals (SO4•-) and superoxide radicals (O2•-). Both sulfate radicals and hydroxyl radicals played a significant role in PCA degradation.

Acknowledgements

This work was financially supported by the National Natural Science Foundation of China (51572089), the National Natural Science Foundation of Guangdong Province, China (2015A030313232), the Foundation of Science and Technology Planning Project of Guangdong Province (No.2016A050502007)，the Research Funds of State Key Lab of Subtropical Building Science, South China University of Technology (2015ZB25) and the Funds of the State Key Laboratory of Pulp and Paper Engineering, China (201477) .

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18

Figure caption

Fig.1. Degradation of PCA by pyrite, persulfate and pyrite/persulfate systems. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite] 0 0.5 g/L, [PS] 0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.2. Effect of scavenger on the degradation of PCA. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.3. The concentration of ferrous ion in the pyrite-persulfate system at different pyrite dosages. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.4. Effect of different initial pH values on PCA degradation in pyrite-persulfate system. Experiment conditions: 20 oC, [PCA]0 0.1 mM, [pyrite] 0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.5. The concentration of Fe

2+

in the degradation of PCA by pyrite at different initial

pH values. Experiment conditions: 20 oC, [PCA]0 0.1 mM, [pyrite] 0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.6. Effect of pyrite particle size on PCA degradation rate. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM and reaction time 4 h Fig.7. Effect of pyrite particle size on the concentration of (a) ferrous and (b) total dissolved iron. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM and reaction time 4 h Fig.8. Effect of temperature on PCA degradation. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.9. Effects of dissolved oxygen condition on pyrite activated persulfate PCA degradation. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.10. Effect of 1, 4-benzoquinone on PCA degradation by pyrite activated persulfate in the presence of dissolved oxygen. Experiment conditions: pHinitial 7.0, at 20 oC, [PCA]0 0.1 mM, [pyrite]0 0.5 g/L, [PS]0 0.5 mM, pyrite particle size 200 mesh and reaction time 4 h Fig.11. EPR signals of different reaction systems; a: effect of pH on radical generation in PS/pyrite system; b:effect of oxygen on radical generation in PS/pyrite system. Green shape stands for sulfite radical, red shape stands for sulfate radical and blue shape stands for hydroxyl radical. Experiment conditions: pHinitial 7.0, at 20 oC, [pyrite]0 0.5 g/L, [PS]0 0.5 mM pyrite particle size 200 mesh and reaction time 4 h

1.0

PCA C/C0

0.8

0.6

0.4 FeS2/PCA

0.2

PS/PCA FeS2/PS/PCA

0.0 0

50

100

150

200

250

200

250

t/min

Fig.1

1.0

PCA C/C0

0.8 0.6 0.4 Phenol MA TBA Without radical scavenger

0.2 0.0 0

50

100

150 t/min

Fig.2

1.2

1.0 gL 0.7 gL 0.5 gL

1.0

0.3 gL 0.1 gL

(a)

2+

-1

Fe (mgL )

0.8 0.6 0.4 0.2 0.0 0

50

100

150 t/min

Fig.3

200

250

1.0 pH:3.0 pH:5.0

0.8

pH:7.0 pH:9.0 pH:11

PCA C/C0

0.6

0.4

0.2

0.0 0

50

100

150

200

250

t/min

Fig.4

8

pH 3.0 pH 5.0

4

2+

Fe (mg/L)

6

pH 7.0 pH 9.0 pH 11.0

2

0 0 Fig.5

100

200 t/min

300

400

1.0

0.8

PCA C/C0

0.6

0.4

200 mesh 100 mesh 60 mesh 40 mesh 20 mesh

0.2

0.0 0

50

100

150 t/min

Fig.6

200

250

0.8

200 mesh 100 mesh 60 mesh 40 mesh (a) 20 mesh

0.4

2+

Fe (mg/L)

0.6

0.2

0.0 0

50

100

150

200

250

t/min

Total dissolved iron (mg/L)

5 200 mesh 100 mesh 60 mesh

4

40 mesh 20 mesh

3 (b) 2 1 0 0

50

100

150 t/min

Fig.7

200

250

1.0 0

10 C 0 20 C 0 30 C 0 40 C 0 50 C

0.8

PCA C/C0

0.6

0.4

0.2

0.0 0

50

100

t/min

150

200

250

Fig.8

1.0

PCA/FeS2/PS/N2

0.8

PCA C/C0

PCA/FeS2/PS PCA/FeS2/Air

0.6

0.4

0.2

0.0 0

50

100

150 t/min

Fig.9

200

250

1.0 PCA/FeS2/PS/Air

PCA C/C0

0.8

PCA/FeS2/PS/Air/BQ

0.6

0.4

0.2

0.0 0

50

100

150 t/min

Fig.10

Fig. 11

200

250

Graphical abstract

19

Highlights


PCA degradation was enhanced by pyrite in persulfate solution


The mechanism of the generation of reactive oxygen species (ROS) was explored.
The role of the ROS for the degradation of PCA was elucidated.
Superoxide radical generated from oxygen could transfer into sulfate radical.
Sulfate radical, hydroxyl radical and sulfite radical were detected by EPR.

20