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Activation of Persulfates Using Siderite as a Source of Ferrous Ions: Sulfate Radical Production, Stoichiometric Efficiency, and Implications Yong Feng, Deli Wu, Hailong Li, Jianfeng Bai, Yi-bo Hu, Chang-Zhong Liao, Xiao-yan Li, and Kaimin Shih ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.7b03948 • Publication Date (Web): 16 Jan 2018 Downloaded from http://pubs.acs.org on January 16, 2018

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ACS Sustainable Chemistry & Engineering

Activation of Persulfates Using Siderite as a Source of Ferrous Ions: Sulfate Radical Production, Stoichiometric Efficiency, and Implications Yong Fenga, Deli Wub, Hailong Lia,c, Jianfeng Baid, Yibo Hua, Changzhong Liaoa,e, Xiao-yan Lia and Kaimin Shiha,* a

Department of Civil Engineering, The University of Hong Kong, Pokfulam, HASAR, China

b

State Key Laboratory of Pollution Control and Resources Reuse, School of Environmental Science &

Engineering, Tongji University, Shanghai 200092, People’s Republic of China c

School of Energy Science and Engineering, Central South University, Changsha 410083, People’s

Republic of China d

School of Urban Development and Environmental Engineering, Shanghai Second Polytechnic University,

Shanghai 201209, People's Republic of China e

Guangdong Key Laboratory of Integrated Agro-environmental Pollution Control and Management,

Guangdong Institute of Eco-Environmental Science & Technology, Guangzhou 510650, People’s Republic of China

Contact Information: Yong Feng ([email protected]); Deli Wu ([email protected]) Hailong Li ([email protected]); Jianfeng Bai ([email protected]) Yibo Hu ([email protected]); Changzhong Liao ([email protected]) Xiaoyan Li ([email protected]); Kaimin Shih ([email protected]) Manuscript submitted to ACS Sustainable Chemistry & Engineering *Corresponding Author: Dr. Kaimin Shih Phone: +852-2859-1973 Fax: +852-2559-5337 E-mail: [email protected]

Word Counts: (Abstract, Manuscript Body, Description of Supporting Information, Acknowledgements, 6 Figures, and 2 Tables) = 6832 words

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ABSTRACT Ferrous ions (Fe2+) rapidly activate persulfates to produce sulfate radicals. However, the high reactivity of Fe2+ towards sulfate radicals means that they are easily scavenged, which reduces the stoichiometric efficiency of persulfates. To improve the stoichiometric efficiency, siderite was used to activate peroxydisulfate (PDS) and peroxymonosulfate (PMS), with phenol as a model contaminant. Near-100% degradation of phenol was achieved by siderite-activated PDS or PMS. In contrast, only 34% and 25% of the phenol was degraded by Fe2+- and nanoscalemagnetite-activated persulfates, respectively. The stoichiometric efficiencies of PMS and PDS activated by siderite were more than 4.4 and 3.6 times higher, respectively, than those activated by Fe2+. Electron paramagnetic resonance recorded both sulfate radicals and hydroxyl radicals. The effects of pH, iron dissolution, and scavenging were characterized, and the results indicated that siderite mainly activated persulfates by acting as a source of Fe2+, and that sulfate radicals were the major active species. The release of Fe2+ and the production of sulfate radicals were controllable via the pH of the solution. No deactivation occurred when the siderite was reused, because the acidic environment partially dissolved the surface. These findings may facilitate the application of iron-bearing materials for sulfate radical production. KEYWORDS: Persulfates; Siderite; Sulfate radicals; Ferrous Ions; Controlled release


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INTRODUCTION Chemical treatments based on advanced oxidation processes (AOPs) are of great interest to environmental scientists because of their use in the removal of refractory contaminants; such as the in situ remediation of groundwater and soil.1 The roles of hydroxyl radicals (·OH) and sulfate 2, 3 radicals (SO∙ Although ) in AOPs have been thoroughly investigated during the past decade.

both of these species are capable of mineralizing most known organic contaminants, SO∙ is generally more selective and therefore is less likely to be scavenged by common constituents of aqueous matrices.4 In addition, persulfates used to produce SO∙ , such as peroxydisulfate (PDS) and peroxymonosulfate (PMS), are solid under standard conditions, and therefore are more stable than liquid hydrogen peroxide (H2O2), which is commonly used to produce ·OH. Iron is a readily available and environmentally friendly material. Fenton reagents and Fenton-like reactions based on iron-containing materials (such as magnetite and ferrihydrite) are known to produce ·OH.5, 6 Because of the similarities between the structures of persulfates and H2O2, iron-containing materials have also been investigated for activating persulfates.7-9 Unfortunately, dissolved iron (Fe2+ and Fe3+) and iron oxides exhibit low reactivity toward persulfate activation,10,

11

mainly because of the strong scavenging effect exhibited by Fe2+

3+ 2+ 3+ toward SO∙ and the difficulty in reducing Fe or Fe(III). In this paper, Fe and Fe represent

dissolved iron species, and Fe(II) and Fe(III) represent the solid iron species. To alleviate the scavenging effect of Fe2+, Rastogi et al.12 used sodium citrate to control the quantity of Fe2+ in solution via complexation; this increased the degradation of 2-chlorobiphenyl by Fe2+-PMS from 44% to 74%. Oh et al.13 used zero-valent iron (ZVI) as an iron source to activate PDS, and found that ZVI-PDS completely degraded 2,4-dinitrotoluene. In contrast, Fe2+PDS only degraded ≤ 20% of the 2,4-dinitrotoluene. Wang et al.14 used carbon-sphere-



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encapsulated ZVI to activate PMS, which resulted in nearly complete removal of phenol within 20 min. However, because of its high surface energy, the stability of ZVI in air is poor.15 In addition, Kambhu et al.16 reported that serious scavenging effects were exhibited by Fe2+ produced by the reaction of ZVI and PDS. To promote the reduction of Fe3+, Zou et al.17 added hydroxylamine during Fe3+-PMS oxidation. As a result, the degradation of benzoic acid by hydroxylamine-Fe3+-PMS was > 80%; Fe3+-PMS alone exhibited benzoic acid degradation of < 10%. Wang et al.18 investigated the effect of ultraviolet radiation on the degradation of 2,4,5trichlorophenoxyacetic acid by Fe2+-PMS, and found that pollutant degradation was greatly improved because the radiation accelerated the regeneration of Fe2+. Fe(II)-bearing minerals are promising sources of Fe2+. The active sites on Fe(II)-bearing minerals do not depend on the reduction of Fe(III), which may improve the reactivity of iron materials toward persulfate activation. Siderite (FeCO3), which contains Fe(II) and carbonates, is an economically important iron ore. It is also a common product of the corrosion of iron materials in waterlogged environments, such as permeable reactive barriers,19,

20

and an

important component of many sediments.21-23 Importantly, under acidic conditions, the carbonates in siderite react with H+ ions to release Fe2+, which may activate either PDS or PMS to generate radicals. However, the application of siderite to persulfate activation using such a strategy has not previously been reported. In the present study, we investigated the activation of PDS and PMS by siderite under various acidic conditions and characterized the activation mechanism using phenol as a model contaminant. Because the activation of PMS by siderite has not been reported previously, we studied the parameters that we expected to influence the reactivity of siderite-PMS oxidation and explored the effects of common water components, including Cl− and bicarbonate ions.


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MATERIALS AND METHODS Chemicals. Oxone (the commercial form of PMS; KHSO5·1/2KHSO4·1/2K2SO4), sodium PDS (≥ 98.0%), and nanoscale magnetite (Fe3O4, 97% trace metals basis) were purchased from Sigma-Aldrich (St. Louis, MO, USA). Before use, the weight ratio of PMS in Oxone was quantified using a thiosulfate standard. A complete list of chemicals is listed in Note S1.

Material Synthesis and Characterization. Siderite was synthesized using a modified hydrothermal approach with urea as the carbonate source.24 Detailed steps are described in Note S2. Iron quantification indicated that the synthesized siderite had an iron content of around 46.8%. Zeta potential was measured using a Coulter Delsa 440SX zeta-potential analyzer (Beckman, Fullerton, CA, USA); the point of zero charge (pHpzc) was determined to be around 4.3 (Fig. S1). The crystal structures and purities of the synthesized and commercial materials were determined using a D8 Advance X-ray diffractometer (Bruker, Karlsruhe, Germany) with a Cu X-ray tube at 40 mA and 40 kV (Fig. S2). The morphology of siderite was analyzed using a field emission scanning microscope (LEO 1530); piled-up microspheres were observed, and these microspheres consisted of siderite particles with diameters of in the range of 100 to 200 nm (Fig. S3). The Brunauer-Emmett-Teller specific surface area of the synthesized siderite (101.1 ± 3.2 m2/g) and commercially obtained magnetite (89.0 ± 1.8 m2/g) were determined by measuring nitrogen adsorption using a micromeritics 3Flex surface characterization analyzer (Norcross, GA, USA). The valence of iron on the surface of the siderite was examined using an X-ray photoelectron spectrometer (ESCALAB 250XI, Thermo, Waltham, MA, USA) with Al-Kα radiation.



Chemical Analysis. An Aquity ultra performance liquid chromatography (UPLC) system (Waters, Milford, MA, USA) equipped with a photodiode array detector (PDA) and a Waters BEH C18 column (50 × 2.1 mm, 1.7 µm) was used to analyze the organic compounds. Detailed UPLC-PDA conditions, analysis of inorganic chemicals, and detection of degradation intermediates of phenol are listed in Note S3. The electron paramagnetic resonance investigation was conducted on a Bruker EMX 10/12 spectrometer (Karlsruhe, Germany) and DMPO was used as a spin-trapping agent. Detailed parameters and procedures can be found elsewhere.25

Batch Experiments. Unless otherwise specified, all of the experiments were conducted at ambient temperature (24 ± 1 °C) in 250-mL conical flasks. A stock solution of phenol (0.1 M) was prepared by dissolving an exact amount of phenol in ultrapure water. Working solutions of phenol (0.1 mM) for the degradation experiments were prepared by diluting the stock solution with ultrapure water; the pH was adjusted to 7.0 using 1 M NaOH. Typically, 100 mL of working solution was transferred to a conical flask, followed by the addition of Oxone (sodium PDS) powder. The release of H+ from the dissolution of Oxone reduced the pH of the working solution to 2.6. To examine the effects of pH, the pH of the resulting solution was adjusted to the target value with 5 M NaOH. Activators, such as siderite or Fe2+ (as aqueous iron(II) sulfate), were then added to initiate the degradation. A magnetic stirrer (RT 10, IKA, Guangzhou, China) operated at 600 rpm was used to keep the solid activators suspended. In the case of using magnetite as the activator, mechanical agitation was used. For LC-PDA analysis, aliquots (0.4 mL) were withdrawn using 1-mL syringes, filtered with 0.22-µm polytetrafluoroethylene (PTFE) syringe filters (Millipore, Darmstadt, Germany), diluted with ultrapure water, and transferred to 2-mL LC autosampler vials. LC-mass spectrometry (MS)-grade methanol (100 µL) was


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immediately added to the vials to prevent the further degradation of phenol by radicals. The pH was measured using an Orion 2-Star benchtop pH meter (Thermo Scientific, Waltham, MA, USA), which was calibrated with standard pH buffers (VWR, Radnor, PA, USA) before use. To evaluate the degradation of siderite, the siderite reactivity was measured over multiple activation cycles according to the procedures described in Note S4.

RESULTS AND DISCUSSION Reactivity of Different Activators. As shown in Fig. S4, neither persulfates (PDS and PMS) nor activators (Fe2+, magnetite, and siderite) alone had any significant effects on the concentration of phenol. Before investigating the reactivity of siderite, the activation of persulfates by Fe2+ and magnetite was examined. The concentration of Fe2+ used in these studies was higher than that leached from siderite (Fig. 4). Approximately 20 and 38% of the phenol was degraded by Fe2+-PMS and Fe2+-PDS, respectively, after a reaction time of 180 min in the presence of Fe2+ (50 mg/L) and persulfates (5 mM) (Fig. 1a). Because SO∙ was the major active species produced by Fe2+-activated persulfates,12 the low efficiencies of Fe2+-PMS and Fe2+-PDS 2+ 9 −1 −1 26 may be attributed to the competitive consumption of SO∙ by Fe (4.6 × 10 M s ) and the

thermodynamically unfavorable regeneration of Fe2+. Under identical conditions, around 9 and 28% of the phenol was degraded by magnetite-PMS and magnetite-PDS, respectively; 2.5 g/L of magnetite was used as the activator. As can be seen from these results, activated PDS exhibited higher oxidation performance than activated PMS, which can be explained by the significantly lower reaction rate of PDS with Fe2+ (27-32 M−1 s−1)7 than that of PMS [(3.0-3.56) × 104 M−1 s−1].27 The slower the generation of SO∙ , the lower the scavenging rate.



Figure 1. Degradation of phenol by (a) Fe2+- and magnetite-activated persulfates and (b) sideriteactivated persulfates. Conditions: [Fe2+] = 50 mg/L, [magnetite] = [siderite] = 2.5 g/L, [PMS] = [PDS] = 5 mM, [phenol] = 0.1 mM. The pH values of the phenol solutions containing PDS were adjusted to 2.6 before addition of the activators.

Near-complete degradation of phenol by persulfates occurred after 120 min when siderite was used as the activator (Fig. 1b). The efficient degradation of phenol by siderite-activated persulfates suggests that siderite had significantly higher reactivity for persulfate activation than either Fe2+ or magnetite. As with persulfates activated by Fe2+ and magnetite, more efficient degradation was observed with PDS than with PMS. The degradation of phenol by sideriteactivated persulfates followed pseudo-first-order kinetics. The kinetic constant (kobs) calculated for the siderite-activated reaction was around 10 times higher than that for the magnetiteactivated system (Table 1); this was also the case for the surface area- and metal-loadingnormalized rate constants. The half-lives of phenol with siderite-PMS and siderite-PDS were around 44 and 22 min, respectively. To quantify the reduced scavenging effects in the siderite-activated system, we calculated the stoichiometric efficiency of each system, defined as the amount of phenol degraded per mole 8 ACS Paragon Plus Environment

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of persulfate (Eq. 1). The results show that the stoichiometric efficiencies of the siderite systems were more than 4.4 and 3.6 times that those of the Fe2+-activated PMS and PDS systems, respectively. Stoichiometric efficiency = ∆[phenol]⁄∆[persulfate]

(1)

Table 1. Phenol degradation by different combinations a

Fe2+-PMS

Pseudo-firstorder constant (kobs × 103/min−1) -

Fe2+-PDS

Active species

R2

Surface areanormalized constant (× 104/min−1 m−2)

Metal-loadingnormalized constant (× 102/min−1 g−1)

Half-life (t1/2, min)

Stoichiom etric efficiency b

-

-

-

2.9 ± 0.2

-

-

-

-

4.5 ± 0.7

siderite-PMS

15.6 ± 0.6

0.992

6.2 ± 0.2

13.3 ± 0.5

44

12.9 ± 0.5

siderite -PDS

31.6 ± 0.2

0.925

12.5 ± 0.1

27.0 ± 0.2

22

16.4 ± 1.1

magnetite-PMS

1.2 ± 0.1

0.951

0.5 ± 0.0

0.67 ± 0.1

578

3.8 ± 0.4

magnetite-PDS

3.2 ± 0.2

0.927

1.4 ± 0.1

1.8 ± 0.1

217

21.4 ± 0.8

a

2+

Conditions: [persulfate] = 5 mM, [siderite] = [magnetite] = 2.5 g/L, Fe = 50 mg/L, [phenol] = 0.1 mM.

Reaction time = 60 min.

Effects of Experimental Parameters on Siderite-PMS Oxidation. To determine the parameters that influence the degradation of phenol by siderite-persulfates, the effects of siderite loading ([siderite]/[PMS] ratio), ionic strength, and the pH values of pollutant solutions were investigated. As shown in Fig. 2a, a significant increase in the degradation of phenol was observed when siderite loading was increased from 0.5 to 2.5 g/L. Although the near-complete degradation of phenol was realized after 120 min with 2.5 g/L of siderite, a further increase in the siderite loading did not result in observable scavenging effects; a slight increase in the degradation rate (from 87 to 97%) occurred after 90 min when siderite leading was increased from 2.5 to 4.0 g/L.


b


Figure 2. Effects of (a) siderite loading, (b) ionic strength, and (c) initial pH (phenol solution containing PMS) on the degradation of phenol by siderite-PMS. (d) Degradation of phenol by PMS alone at different initial pH values. Conditions: [PMS] = 5 mM, (b, c) [siderite] = 2.5 g/L, (a, b) initial pH of phenol solution = 7.0, [phenol] = 0.1 mM.

Because of the relatively low solubility of KClO4, Na2SO4 was used to increase background ionic strengths. Increasing the concentration of Na2SO4 from 5 to 50 mM inhibited phenol degradation, which decreased from nearly 100 to 82%. However, further increasing the concentration of Na2SO4 to 100 mM resulted in a slight promoting effect. These observations were attributed to two competing effects of Na2SO4: initially, the release of SO from the dissolution of Na2SO4 inhibited the decomposition (activation) of PMS; at higher concentrations, the increased ionic strength promoted the decomposition (activation) of PMS.28 10 ACS Paragon Plus Environment

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Under acidic conditions (pH ≤ 6.5), the reactivity of siderite-PMS was significantly affected by the initial pH of the reaction mixture. As shown in Fig. 2c, the degradation rate of phenol decreased from near-100% at pH 2.6, to 67, 18, 21, and 25% at pH 3.5, 4.5, 5.5, and 6.5, respectively. The release of Fe2+ from siderite under mildly acidic conditions was limited (Fig. 4b); the low efficiency of siderite activation at pH 4.5 and 5.5 suggests that the activation of PMS by the Fe(II) on the surface of siderite was weak, and that siderite functioned more efficiently as a source of Fe2+ under strongly acidic conditions (such as pH 2.6). The low reactivity of siderite under mildly acidic conditions was also reported by Huang et al.29 and Yan et al.;30 only 20% of trichloroethylene was degraded by siderite-PDS after reacting for 24 h. Under basic conditions (pH ≥ 7.5), the rate of degradation increased with increasing pH. Almost complete degradation of phenol was observed at pH 8.0 and 10.0 after reacting for only 15 min (Fig. 2c). PMS has a second ionization constant (pK2) of 9.4 ± 0.1,28 which suggests that a majority of the PMS was present in its deprotonated form (SO ) at pH 10.0, and that a small proportion was completely deprotonated at pH 8.5. Previous studies have shown that SO can be readily activated by benzoquinone to produce singlet oxygen (1O2).31, 32 Phenol may be oxidized 1 to quinones in air, and these quinones may have activated SO to produce O2. Moreover, under

basic conditions, PMS undergoes self-decomposition to form 1O2 (Eq. 2).33 1O2 reacts rapidly with phenol at a rate constant of (1.95-6.13) × 107 M−1 s−1.34 Hence, it may contribute significantly to the degradation of phenol at pH 8.0 and 10.0. To verify this hypothesis, we examined the degradation of phenol by PMS alone under various pH conditions (Fig. 2d). In each experiment, the pH of the PMS solution was adjusted to the desired value before the addition of phenol to initiate the reaction. As expected, rapid degradation occurred at pH 10.0, and the degradation was increasingly inhibited with decreasing pH (Fig. 2d). In addition to



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reacting with phenol, the 1O2 readily decayed to produce triplet oxygen, as evidenced by the production of gas bubbles during the degradation of phenol by PMS at pH 8.0 and 10.0 (Fig. S5). To further characterize the dominant role of 1O2 at pH 10, we examined the effects of adding methanol and sodium azide. Azide reacts with 1O2 with kinetics (2.2 × 108 M−1 s−1)35 that are more than 7 × 104 times higher than methanol (3.0 × 103 M−1 s−1);36 hence, these compounds were suitable for detecting 1O2. Near-complete degradation of phenol occurred in the presence of methanol (500 mM). However, only around 8.9% of the phenol was degraded in the presence of azide (1 mM) (Fig. S6). These results confirm that 1O2 was responsible for the degradation of phenol under basic conditions. HSO + SO → H + 2SO + 1O

(2)

Mechanisms of Degradation and Activation. As siderite-PMS has not been previously studied, investigation in the presence of DMPO as a spin trapper was conducted to reveal the production of radicals by this oxidation. As shown in Fig. 3a, radical adducts were recorded after 30 and 60 min of reaction. The hyperfine splitting constants of these adducts obtained by simulation were in agreement with previous reports37 and indicated that both ·OH (αN = αβ-H = 14.9) and SO∙ (αN = 13.8, αβ-H = 9.5, αγ1-H = 1.44, and αγ2-H = 0.79) were probably produced by siderite-PMS. Although the intensity of DMPO-OH adduct was significantly stronger than that of DMPO-SO4, it was hard to conclude that ·OH was the dominant species because of the probable 38, transformation of DMPO- SO to DMPO-OH via nucleophilic substitution.

39

To further

examine the role of these radicals in the degradation, we conducted classical scavenging experiments using different alcohols.10, 40, 41 Ethanol and tert-butanol were chosen because of the significant differences in their reactivities toward ·OH and SO∙ . Ethanol is relatively reactive


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7 −1 −1 42 toward both ·OH (2.2 × 109 M−1 s−1) and SO∙ s ). tert-Butanol is reactive (5.6 × 10 M 5 −1 toward ·OH (3.3 × 109 M−1 s−1),43 but exhibits low reactivity toward SO∙ (4.1-9.0 × 10 M

s−1).44 The results in Fig. 3b show that the degradation of phenol by siderite-PMS was significantly inhibited by ethanol (0.5 M). In contrast, the presence of tert-butanol (0.5 M) did not have a significant effect on phenol degradation. In addition, 1,4-dioxane has a higher affinity for activator surfaces than ethanol,45 and exhibited a significantly stronger inhibitory effect. 7 −1 −1 46 However, its rate of reaction with SO∙ (1.6 × 10 M s ) is slower than that of ethanol (5.6 ×

107 M−1 s−1). This increased inhibition of degradation by 1,4-dioxane indicated that active 4 ∙ species such as 1O2 (7.0 × 102 M−1 s−1)36 and SO∙ ( E (SO /SO ) = 1.1 V) were not

responsible for phenol degradation under these conditions. Similar alcohol inhibitory phenomena also occurred in phenol degradation by siderite-PDS (Fig. 3c). The scavenging effects of ethanol and 1,4-dioxane, and the lack of scavenging by tert-butanol, demonstrate that radicals were generated by siderite-persulfates and that SO∙ was the dominant species responsible for the degradation of phenol.



Figure 3. (a) Electron paramagnetic resonance spectra of siderite-PMS oxidation after reacting for 30 and 60 min. Effects of radical scavengers on the degradation of phenol by (b) sideritePMS and (c) siderite-PDS. (d) Production of HCHO by siderite-activated persulfates. Conditions: [PMS] = [PDS] = 5 mM, [siderite] = 2.5 g/L, (b, c) [phenol] = 0.1 mM, (d) [methanol] = 2 M, initial pH of phenol solution = 7.0. The pH of the phenol solution containing PDS was adjusted to 2.6 before the addition of siderite.

The radicals produced during the decomposition of persulfates were quantified using the formaldehyde (HCHO) approach with methanol as a substrate.45, 47 The results show that the production of HCHO was increased with increasing siderite loading, which is consistent with the increased degradation of phenol as a result of increased radical production (Fig. 3d). At a siderite loading of 2.5 g/L, approximately 740 and 610 µM of HCHO were generated by siderite-PMS and siderite-PDS, respectively, after reacting for 180 min. After determining the consumption of


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persulfates during the activation processes, the conversion rates of PMS and PDS to radicals (amount of HCHO produced per mole of persulfate) were calculated to be 52.1 ± 5.1 and 65.3 ± 1.1%, respectively. These conversion rates were comparable with the conversion rates (around 62%) achieved when Co2+ was used for PMS activation.45

Figure 4. Effect of PMS on (a) the dissolution of iron from siderite and the decomposition of PMS in the presence of siderite and (b) the dissolution of iron from siderite in the presence of PMS under different pH conditions. Conditions: [PMS] = 5 mM, [siderite] = 2.5 g/L, [phenol] = 0.1 mM. The pH of phenol solution (a) was decreased to 2.6 after addition of PMS. The pH of the phenol solution containing PMS (b) was adjusted to the desired value before the addition of siderite. As discussed above, siderite mainly activated PMS via the controlled release of Fe2+. Therefore, we determined the dissolution of iron from siderite. Because of the rapid reaction of Fe2+ with PMS, only Fe3+ was detected during siderite-PMS oxidation. As shown in Fig. 4, the concentration of dissolved iron increased with increasing reaction time. After reacting for 180 min, around 35 mg/L of dissolved iron was present in the reaction solution. This corresponds to a CO $ concentration of 0.63 mM, which was produced from the dissolution of siderite (FeCO3). Interestingly, the presence of PMS had an inhibitory effect on the dissolution of iron; under 15 ACS Paragon Plus Environment


identical conditions, around 54 mg/L of dissolved iron was determined in the absence of PMS. The pHpzc of siderite is around 4.3 (Fig. S1), which suggests that the surface of the siderite was positively charged at pH 2.6. Under these conditions, PMS existed mainly as HSO . Hence, the reduced dissolution of iron in the presence of PMS was attributed to the adsorption of PMS onto the surface of siderite, which hindered the interaction of H+ ions with siderite. The inhibition effect of PMS on the dissolution of iron explains the promotion effect of ionic strength, as an increase in the ionic strength is known to reduce the zeta potential of colloid particles in aqueous solutions.

High levels of iron dissolved from siderite can be a concern and may require

additional treatment procedures. Under mildly acidic conditions (pH 3.5 and 4.5), the dissolution of iron from siderite was significantly decreased with decreasing acidity. Only 6.4 and 4.4 mg/L of dissolved iron were detected after 180 min at pH 3.5 and 4.5, respectively. These relatively low concentrations of dissolved iron were consistent with the low reactivity of siderite-PMS at pH 3.5 and 4.5 (Fig. 2c). These results also confirmed our hypothesis that the reactivity of siderite toward PMS under strongly acidic conditions mainly resulted from the release of Fe2+.

Stability and Reutilization of Siderite. The reutilization of catalysts is a crucial factor in determining the overall cost of treatment processes. To evaluate the reutilization of siderite, we evaluated the degradation of phenol by siderite-PMS over multiple activation cycles (Table 2). Approximately 80% of the phenol was degraded during the first cycle. Degradation rates of 6770% were achieved in the second to fifth cycles. Correspondingly, the consumption of PMS in the first cycle (12.6%) was higher than in the subsequent cycles (6.9-7.6%). Iron quantification showed that iron dissolution from siderite was nearly constant; dissolved iron concentrations


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ranged from 20 to 24 mg/L for all five cycles. The stoichiometric efficiency of PMS was lower in the first cycle (12.7%) than in later cycles (17.7-20.8%). The higher consumption and degradation observed during the first cycle results may be explained by the existence of specific surface features on fresh siderite (Fig. 5b) that interacted actively with PMS and phenol; however, further investigations are required to confirm this hypothesis and characterize these features. Because the synthesized siderite was dried in air at an elevated temperature (105 °C), the majority of the iron on its surface was oxidized to Fe(III), as shown by its X-ray photoelectron spectrum (Fig. 5a). The Fe(III) on the siderite surface was dissolved during the catalytic degradation, and the surface Fe(II) was readily oxidized to Fe(III) in the presence of PMS and SO∙ (Fig. 5a). However, as shown by the X-ray diffraction patterns, no structural variations or new phases were formed as a result of the activation process (Fig. 5b). This was consistent with our conclusion that, under strongly acidic conditions, siderite mainly activated PMS via the release of Fe2+. Table 2. Results of siderite-PMS oxidation cycling experiments a

a

PMS consumption

Total dissolved iron

Stoichiometric

(%)

(mg/L)

efficiency (%)

80.2 ± 0.8

12.6 ± 0.9

21.5 ± 0.76

12.7

2

70.5 ± 1.6

6.9 ± 0.5

23.6 ± 0.27

20.4

3

68.7 ± 1.1

6.6 ± 0.5

22.6 ± 0.43

20.8

4

67.2 ± 1.9

7.3 ± 0.4

20.2 ± 2.35

18.4

5

67.4 ± 0.5

7.6 ± 1.5

22.6 ± 0.67

17.7

Cycle

Degradation (%)

1

Conditions: [PMS] = 5 mM, [siderite] = 2.5 g/L, [phenol] = 0.1 mM, initial pH of phenol solution =

7.0, reaction time = 120 min.



Figure 5. (a) High-resolution X-ray photoelectron spectrum of Fe2p region and (b) X-ray diffraction patterns of fresh siderite and siderite after different activation cycles.


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Effects of Carbonate and Chloride Ions. To evaluate the practical applicability of sideritePMS oxidation, the effect of bicarbonate (HCO $ ) on phenol oxidation was investigated. Na2SO4 (10 mM) was added to eliminate the potential influence of HCO $ on background ionic strengths. HCO $ (1-20 mM) had no significant effect on the degradation of phenol; almost complete degradation was achieved after 120 min under all of the conditions investigated (Fig. 6a). During the dissolution of iron from siderite, CO $ was also generated and released into the reaction solutions. The degradation of phenol in the presence of 1-20 mM HCO $ (Fig. 5a) indicates that the release of CO $ (about 0.63 mM from 2.5 g/L siderite) from siderite will have no adverse effects on the degradation of pollutants.

− Figure 6. Effects of (a) HCO $ and (b) Cl on the degradation of phenol by siderite-activated PMS.

Conditions: [PMS] = 5 mM, [siderite] = 2.5 g/L, [phenol] = 0.1 mM, [Na2SO4] = 100 mM. The pH values of the phenol solutions were adjusted to 2.6 after the addition of PMS and HCO $ and before the addition of siderite. Insets in (a) and (b) show the pseudo-first-order kinetic constants for various concentrations of HCO$ and Cl−, respectively.

− In addition to HCO $ , Cl is ubiquitous in aqueous environments and has been reported to

have inhibitory effects on the degradation of trichloroethylene by PDS48 and 1,4-dioxane by



PMS.45 Degradation experiments with Cl− showed that it exhibited promoting effects on sideritePMS oxidation, and that the promotion effect increased with increasing Cl− concentration (Fig. 6b). Because SO∙ is the major active species in siderite-PMS oxidation, the promoting effect − ∙ may be related to the generation of Cl∙ via the one-electron reaction of excess Cl with SO (3.1

∙ × 108 M−1 s−1).46 Although Cl∙ has a lower standard reduction potential than SO∙ [E (Cl / ∙ 2Cl ) = 2.09 V4 and E (SO∙ /SO ) = 2.5-3.1 V], Cl readily reacts with nucleophiles via one-

electron oxidation. A similar promoting effect was reported by Anipsitakis et al.,49 who examined the degradation of phenol by Co2+-PMS in the presence of Cl−. In addition, our previous study showed that phenol can be degraded by the reaction of PMS and Cl−.50 Hence, the accelerated degradation of phenol shown in Fig. 6b could be partially related to the active species generated by direct interactions between PMS and Cl−. However, based on this promotion mechanism, chlorinated products were likely produced. To further clarify this mechanism, we investigated the degradation intermediates of phenol in the presence of Cl− (5 mM) using GCMS (Note S3). The results show that 2-chlorophenol accumulated during the degradation process (Fig. S7). Moreover, 2-chlorophenol and 4-chlorophenol were detected in the phenol solution containing both PMS and Cl− after 120 min (Fig. S8). Because these chlorinated phenols are generally more toxic than phenol,51, 52 the generation of these species should be considered when PMS is used in systems containing elevated levels of Cl−.

ASSOCIATED CONTENT Supporting Information Note S1–Note S4: Chemicals, synthesis of siderite, chemical analysis, and reutilization test of siderite; Figure S1–Figure S3: Zeta potential, X-ray diffraction pattern, and scanning electron


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micrograph of siderite; Figure S4: Removal of phenol by persulfates or siderite alone; Figure S5: Images of phenol solutions containing PMS at different pH values; Figure S6: Effects of different quenching agents; Figure S7: Accumulation of 2-chlorophenol; Figure S8: Mass spectra of 2-chlorophenol and 4-chlorophenol.

AUTHOR INFORMATION Corresponding Author *Phone: +852-2859-1973; fax: +852-2559-5337; e-mail: [email protected]. Notes The authors declare no competing financial interest.

ACKNOWLEDGEMENTS This work was supported by the Research Grants Council of Hong Kong (17212015, 17257616, C7044-14G, and T21-771/16R). The authors thank Ms. Vicky Fung for her technical assistance.



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SYNOPSIS: Siderite activated persulfates effectively by acting as a source of ferrous ions for the controlled generation of sulfate radicals.


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