Production of Sulfate Radical from Peroxymonosulfate Induced by a

Feb 25, 2013 - react very rapidly with ethanol (k·OH = (1.2−2.8) × 109 M. −1. Figure 3. ..... (18) Brandt, C.; van Eldik, R. Transition metal-ca...
0 downloads 0 Views 361KB Size
Download PDF
Article pubs.acs.org/est

Production of Sulfate Radical from Peroxymonosulfate Induced by a Magnetically Separable CuFe2O4 Spinel in Water: Efficiency, Stability, and Mechanism Tao Zhang,† Haibo Zhu,‡ and Jean-Philippe Croué†,* †

Water Desalination and Reuse Center, King Abdullah University of Science and Technology, Thuwal 4700, Kingdom of Saudi Arabia KAUST Catalysis Center, King Abdullah University of Science and Technology, Thuwal 4700, Kingdom of Saudi Arabia

‡

S Supporting Information *

ABSTRACT: A simple, nonhazardous, efficient and low energyconsuming process is desirable to generate powerful radicals from peroxymonosulfate (PMS) for recalcitrant pollutant removal. In this work, the production of radical species from PMS induced by a magnetic CuFe2O4 spinel was studied. Iopromide, a recalcitrant model pollutant, was used to investigate the efficiency of this process. CuFe2O4 showed higher activity and 30 times lower Cu2+ leaching (1.5 μg L−1 per 100 mg L−1) than a well-crystallized CuO at the same dosage. CuFe2O4 maintained its activity and crystallinity during repeated batch experiments. In comparison, the activity of CuO declined significantly, which was ascribed to the deterioration in its degree of crystallinity. The efficiency of the PMS/CuFe2O4 was highest at neutral pH and decreased at acidic and alkaline pHs. Sulfate radical was the primary radical species responsible for the iopromide degradation. On the basis of the stoichiometry of oxalate degradation in the PMS/CuFe2O4, the radical production yield from PMS was determined to be near 1 mol/mol. The PMS decomposition involved an inner-sphere complexation with the oxide’s surface Cu(II) sites. In situ characterization of the oxide surface with ATR-FTIR and Raman during the PMS decomposition suggested that surface Cu(II)−Cu(III)−Cu(II) redox cycle was responsible for the efficient sulfate radical generation from PMS.

■

INTRODUCTION Efficient removal of organic contaminants is one of the major targeted objectives of water treatment and wastewater reclamation unit operators. Oxidation is an important technical option in pollutant removal. A large number of organic pollutants (e.g., molecules incorporating activated aromatic structures and double bonds) can be easily degraded by usual oxidants such as HOCl/OCl−, ClO2, permanganate, or ozone. However, many emerging contaminants are refractory to conventional oxidation process. Hydroxyl radical is highly reactive toward nearly all organic compounds. It can be generated from ozone (O3/H2O2, O3/ UV, O3/UV/H2O2, and catalytic ozonation) or H2O2 (Fenton and Fenton-like). These techniques also have their limitations in the application. The ozone based technique requires a sophisticated ozone generation system, that is gas pretreatment, ozone generator, cooling, and waste gas destruction. The Fenton process needs large quantity of acid and base for pH adjustment; it also leads to significant sludge production. Although supported iron oxides showed higher efficiency than the iron oxide alone near neutral pH, the Fenton-like process relying on solid catalysts still has relatively low efficiency in terms of both H2O2 utilization rate and reaction time. 1 During the past years, sulfate radical attracted increasing research interests as an alternative to hydroxyl radical for the degradation of recalcitrant organic pollutants.2−9 Sulfate radical © 2013 American Chemical Society

can be generated from peroxydisulfate (PDS) and peroxymonosulfate (PMS) with heating, UV irradiation, or alkaline activation. Because of the high energy input or high alkaline dosing, these sulfate radical generation techniques are not costeffective for water treatment. Co2+ was found to be very active in initiating sulfate radical generation from PMS for pollutant degradation. 2 Co2+ catalysis relies on the Co2+/Co3+ redox cycle. Because of the complexity to remove Co2+ from water, Co3O4 and supported cobalt oxide were further investigated.10−12 Although these oxides are still effective in initiating sulfate radical generation from PMS, concerns about leaching of potentially carcinogenic Co2+/Co3+ 13 from the solid phase is still a major limitation for practical application in water treatment. The search for alternative metal oxide to efficiently induce sulfate radical generation remains a priority for the development of this technology. Recently, a well-crystallized CuO was reported to be active in promoting PMS oxidation of phenol.14 In contrast, commercial CuO was much less active because of its poor crystalline structure. It was assumed that sulfate radical was generated from the reaction of HSO5− with Cu+, considering that the Received: Revised: Accepted: Published: 2784

November 19, 2012 February 20, 2013 February 25, 2013 February 25, 2013 dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

latter was produced from reducing Cu2+ by HSO5−. This hypothesis needs to be reevaluated, because (1) Cu+ reacts with HSO5− generating ·OH but not SO4·‑,15,16 and (2) the reduction of Cu2+ (E0Cu(II)/Cu(I) = 0.15 V) 17 by HSO5− (E0HSO5‑/HSO4‑ = 1.8 V)18 is thermodynamically unfavorable. Recycle of this crystallized CuO is not favorable because of its relatively high Cu2+ leaching (0.78 mg L−1 at the CuO dose of 1 g L−1and pH 7)14 as well as the loading difficulty caused by the low-temperature preparation condition to get high crystallinity (hydrothermal synthesis at 180 °C). However, because Cu2+ is presently not considered a potential carcinogen, 19 the copperbased oxide would be acceptable for PMS activation in water treatment, if it is highly stable, efficient, and recyclable. During the synthesis of spinel ferrite with Fe3+, namely the ferrite process, heavy metals can be removed from water. 20 Because the spinel structure is relatively stable, the leaching of heavy metals from the solid phase can be significantly reduced. 21 The ferrite process was applied for the removal of heavy metals from mining and industrial wastewaters.22,23 Another advantage of the spinel ferrite is its magnetic property, which makes its separation from water very easy. Thus, the spinel oxides can be prepared and used as powders, which can provide relatively large surface area assisting the reactants’ diffusion onto the active sites. In this work, we investigated the activity and stability of a spinel copper ferrite (CuFe2O4) as compared with the wellcrystallized CuO in the catalytic PMS oxidation, its efficiency in terms of radical yield, and the radical generation mechanism. Iopromide (molecular structure shown in Scheme S1 of the Supporting Information), an iodinated X-ray contrast media detected at elevated concentrations these years in surface water, groundwater, and potable water, 24 was selected as a model pollutant. This compound is very stable and hydrophilic. It cannot be efficiently degraded during water treatment with ozonation (kO3 < 0.8 M−1 s−1)25 or O3/H2O2 oxidation.26 In situ characterization with ATR-FTIR and confocal Raman was applied to analyze the catalytic PMS decomposition process occurring on the oxide surface. Results indicate that the PMS/ CuFe2O4 couple is an attractive process providing efficient sulfate radical generation for recalcitrant pollutant degradation.

■

measured on a Mastersizer 2000 laser particle size analyzer. pHpzc (pH at which the surface is zero-charged) was determined with acid−base titration. The field dependent magnetization (M−H) at room temperature was measured with a superconductor quantum interference magnetometer SQUID-VSM. All results are listed in Table S1 of the Supporting Information. The morphology of the oxide particles was examined from SEM pictures taken on a Quanta 200 (FEG) scanning electronic microscopy (Figure S1 of the Supporting Information). XRD was used to determine the modification of the crystalline structure of the particles after the reaction. Metal oxidation states of CuFe2O4 and CuO particles before and after reaction were characterized with X-ray photoelectron spectroscopy (XPS) (Kratos AMICUS/ESCA 3400 spectrometer) (Figure S2 of the Supporting Information) indicating that they were only Fe(III) and Cu(II). Experimental Procedure. Batch Reaction. Experiments were conducted in a brown glass bottle. Predetermined volumes of iopromide (Schering AG) and PMS (Aldrich Oxone, KHSO5·0.5KHSO4·0.5K2SO4) stock solutions were injected into 200 mL Milli-Q water to get the desired initial concentrations. Tetra-borate rather than phosphate was used as a buffer in most of the reactions, because phosphate usually is a strong coordinate for transition metals. Mechanical stirring was applied during the reaction at a rotary speed of 850 rpm and room temperature (20 °C). The reaction was started by introducing the oxide particles (usually 100 mg L−1) into the solution. Samples taken at different time intervals were filtered through 0.2 μm acetate-fiber syringe filters. The filtration had no impact on the iopromide concentration. Sodium nitrite solution was used to quench HSO5− residual. Analysis. The iopromide used in this study presented in two isomers. They were separated on a Waters HPLC equipped with an Agilent Eclipse XDB C-18 column (150 × 4.6 mm) at a UV wavelength of 238 nm. The separation was performed using an isocratic mobile phase composed of phosphoric acid buffer (0.3% in Milli-Q water) and acetonitrile at a volume ratio of 96/4 and a flow rate of 1.0 mL min−1. Because the peak area of the two isomers declined in nearly equal proportions during the reactions, the sum of the peak areas were used to establish the calibration curve and determine the iopromide concentration for each water sample. The iopromide detection limit was about 10−2 μM at the injection volume of 50 μL. The concentration of HSO5− was analyzed on the same HPLC unit at 260 nm using a Varian Polaris 3 C-18 A column (250 × 3.0 mm) and a mixture of methanol (0.01 mL min−1) and phosphoric acid buffer (0.24 mL min−1 of 0.3% H3PO4 in Milli-Q water) as mobile phase. The detection limit of HSO5− was about 0.5 μM at the injection volume of 50 μL. Oxalate, the probe compound used to quantify the radical yield, was analyzed on a Dionex ICS-3000 IC equipped with an AS-15 column (mobile phase, 30 mM KOH; flow rate, 0.35 mL min−1). TOC was measured on a Teledyne Tekmar TOC Fusion analyzer. Dissolved copper was determined on an ICP-MS (Agilent 7500) with detection limit of 0.1 μg L−1. The procedures of in situ characterization of the oxide surface during catalytic PMS decomposition with ATR-FTIR and confocal Raman were described in the Text S1 of the Supporting Information.

EXPERIMENTAL SECTION

Metal Oxides Preparation. The spinel CuFe2O4 was prepared with a citrate combustion method27 successfully used in the synthesis of a variety of spinel oxides. Copper nitrate, ferric nitrate, and citric acid were dosed in a molar ratio of 1:2:3.6 into 200 mL Milli-Q water. The mixture was stirred for 4 h and then subjected to vacuum evaporation to remove water. The sticky gel was calcined at 300 °C. After calcination, the particles were grinded, washed several times with Milli-Q water, reclaimed with a magnet, and further dried at 120 °C for 12 h. The spinel crystalline in cubic phase was confirmed with X-ray diffraction (XRD) using a Bruker D8 Advanced A25 diffractometer (Cu Kα radiation, operated at 40 kV and 40 mA). Fe2O3 and MnFe2O4 particles were prepared with the same method to test the contribution of Fe(III) in nonspinel and spinel phases to iopromide removal. The well-crystallized CuO, synthesized following the procedure of Ji et al.,14 was used as a comparison for the activity and stability study. Metal Oxides Characterization. BET surface area and average pore size of the oxides were determined on a Micromeritics ASAP 2420 analyzer. Average particle size was

■

RESULTS AND DISCUSSION Activity and Stability of the Spinel CuFe2O4. The hydrothermally synthesized CuO was reported to be able to improve PMS oxidation of phenol in water because of its high 2785

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

degree of crystallinity.14 It was used here as a reference to evaluate the activity and the stability of CuFe2O4 in improving iopromide degradation. Fe2O3 and MnFe2O4 were also used to test the activity of Fe(III) in nonspinel and spinel structures. Figure 1 shows the iopromide concentration decline in PMS

Figure 2. Variation of the simulated pseudofirst-order degradation rate of iopromide with oxide dose in catalytic PMS oxidation (simulated with the mean values of triplicate experiments shown in Figure S3 of the Supporting Information). Conditions: initial iopromide concentration = 1 μM, PMS dose = 20 μM, 10 mM tetraborate buffered pH 6.0, T = 20 °C.

Figure 1. Iopromide degradation in PMS oxidation and catalytic PMS oxidation. Conditions: oxide dose = 100 mg L−1, initial iopromide concentration = 1 μM, PMS dose = 20 μM, 10 mM tetraborate buffered pH 6.0, T = 20 °C. Error bars represent standard deviations from triplicate experiments.

highest reaction rate constant of the former was 1.7 times of that of the latter. The result means that: 1) the increase of oxide dose below its optimal value led to more efficient iopromide degradation with PMS/CuFe2O4 than with PMS/CuO, and 2) CuFe2O4 at the optimal dose was more efficient than synthetic CuO no matter how much was used for the reaction. This phenomenon can be related to the smaller particle size and larger pore size of CuFe2O4 than CuO (Table S1 of the Supporting Information), thus providing more accessible active sites. The decrease of the reaction rate constant at catalyst dose over the optimal value can be interpreted as a diffusion limitation phenomenon in heterogeneous reactions.28 PMS (ionic, Mw = 113 Da) could diffuse much faster toward the active sites of the oxide particles than iopromide (nonionic, Mw = 791 Da) because of its negative charge (contrary to the positive surface charge of the oxide particles in the reaction) and much smaller size. In this case, side-reactions (e.g., SO4·− and HSO5− reaction generating less oxidative SO5·−, and SO4·− combination reaction forming S2O82−)16 might occur leading to ineffective oxidant consumption on the oxide surface before iopromide arriving. To confirm the catalytic property and stability of the oxides in the PMS oxidation system, the particles were recovered (i.e., centrifugation and drying) and reused (after compensation of the mass loss, i.e., 2−3 mg, mainly caused by sampling and transfer, with virgin oxide). This procedure was repeated 7 times (Figure 3). The iopromide removal by PMS/CuO decreased with the number of reaction cycles, whereas the removal by PMS/CuFe2O4 remained almost constant. There was no decrease in specific surface area of CuO after 4 cycles of reaction (virgin, 0.9 m2 g−1; used, 1.0 m2 g−1). Therefore, the activity decline of CuO cannot be attributed to the change of surface area. The crystallinity variation of the CuO and the CuFe2O4 after 4 successive reactions was qualitatively characterized by XRD (Figure S5 of the Supporting Information). Considering that lower diffraction intensity is usually ascribed to a lower crystallinity of the oxide,29,30 the result clearly showed that the crystallinity of the used CuO decreased as compared with the virgin one, whereas for CuFe2O4 no apparent change was noticed. The stability of CuFe2O4 was also supported by its low copper leaching (1.5 ± 0.2 μg L−1 at the oxide dosage of 100 mg L−1) in the repeated

solution alone and PMS coupled with CuO, CuFe2O4, Fe2O3, MnFe2O4, and Cu2+, respectively. Nearly no iopromide degradation was noticed during PMS oxidation alone. A rapid decrease of the iopromide concentration was observed for the PMS/CuO mixture and even faster for the PMS/CuFe2O4 couple. CuO and CuFe2O4 particles exerted negligible iopromide adsorption (Figure S3 of the Supporting Information) finding that confirmed iopromide degradation. Fe2O3 and MnFe2O4 improved only slightly the iopromide degradation. Less than 8% iopromide was reduced in the PMS/Fe2O3 mixture. The removal rate in the PMS/MnFe2O4 mixture was even lower. The result indicates that Cu(II) is the active metal site of CuFe2O4, whereas Fe(III), whether in spinel structure or not, is not active. Copper leached from CuO and CuFe2O4 was detected to be 46 ± 3 and 1.5 ± 0.1 μg L−1, respectively. Under the same reaction condition, Cu2+ ion (1 μM) showed very low activity with less than 5% of iopromide being degraded. It can be concluded that the iopromide degradation in PMS/CuO and PMS/CuFe2O4 followed a heterogeneous catalytic reaction (absence of homogeneous reaction catalyzed by Cu2+). Iopromide degradation data of PMS/CuO and PMS/ CuFe2O4 (up to 70% iopromide degradation) can be fitted well with pseudofirst-order kinetics (R2 > 0.98). Reaction rate constant got from the pseudofirst-order kinetic fitting was thus used as a simple index to compare iopromide degradation efficiency for the catalytic oxidations. The degradation rate constant with PMS/CuFe2O4 was found to be 2.4 times higher than PMS/CuO for a dose of copper about 3 times lower. Because CuFe2O4 is characterized by a higher specific surface area than CuO (Table S1 of the Supporting Information), this finding does not mean that the Cu(II) sites of CuFe2O4 are more active than those of CuO. It is interesting that the increase of total surface area through increasing oxide dose does not necessarily improve the reaction rate. Figure 2 (original data shown in Figure S4 of the Supporting Information) shows optimal oxide doses for the two oxidation systems. Linear relationships could be established between the simulated rate constant and the oxide dose below its optimal value. The slope for PMS/CuFe2O4 was 2.5 times of that for PMS/CuO. The 2786

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

surface charge and also the proportion of SO52−, which possibly inhibited the static interaction between the oxide surface and the PMS species. The strong inverse H+-dependency was also reported in homogeneous PMS activation systems,8,31 which was attributed to the stabilization effect of H+ on the HSO5−. It is suspected that H-bond formation between H+ and the O−O group of HSO5− would be more significant at acidic pH, attaching positive charge to HSO5− and hindering its interaction with the positively charged oxide surface. The high efficiency of PMS/CuFe2O4 at neutral pH supports again that this process is suitable for the treatment of surface and ground waters. Influence of Ionic Strength and Radical Scavengers. The increase of the ionic strength is known to reduce the zeta potential of colloid particles in water. The increase of ionic strength significantly influences outer-sphere interactions (electrostatic bonding) between the solute and the particle’s surface in both equilibrium and kinetics. However, inner-sphere complexation (covalent bonding or a combination of covalent and ionic bonding) is not affected.32 Figure 5 shows that

Figure 3. Iopromide removed in repeated batch catalytic reactions with the same oxide particles. Conditions: oxide dose = 100 mg L−1, initial iopromide concentration in each cycle = 1 μM, PMS dose in each cycle = 20 μM, 10 mM tetraborate buffered pH 6.0, T = 20 °C, reaction for 10 min.

reactions. Ji et al. observed that the commercial CuO had low activity in promoting the PMS oxidation of phenol,14 which was ascribed to its poor crystallization. Consequently, the activity decline of the synthetic CuO in the repeated reactions of this study is largely due to the decrease of its degree of crystallinity. The stable activity of the CuFe2O4 is related to the high stability of the spinel structure. The deterioration of the CuO structure in the reaction also suggests that PMS forms complexes with the surface Cu (II) reducing its stability in the crystalline and thus promoting its release from the solid phase. This assumption is supported by the fact that CuO particles were completely dissolved when immersed into a high concentrated PMS solution (Figure S6 of the Supporting Information). At the same PMS concentration, CuFe2O4 showed no obvious dissolution and still can be separated magnetically meaning that its spinel ferrite structure is quite stable. Because of the high stability and the magnetic property, CuFe2O4 is suitable for water treatment application. Influence of pH. In the pH range of 3.8 to 9.8, the PMS/ CuFe2O4 couple achieved highest iopromide degradation at pH 7.2 and 5.7 (Figure 4). Its efficiency decreased significantly when the solution became acidic or alkaline. pKa1 of H2SO5 is less than 0.16 Its pKa2 is 9.4.8 It can be assumed that HSO5− was the only PMS species present in solution at acid and neutral pHs, whereas a small fraction of SO52− existed at pH 9.2. The pHpzc of CuFe2O4 was determined to be around 7.9. The increase of pH above its pHpzc raised the amount of negative

Figure 5. Influence of ionic strength (represented as NaClO4 concentration) on iopromide degradation during PMS/CuFe2O4 oxidation. Conditions: oxide dose = 100 mg L−1, initial iopromide concentration = 1 μM, PMS dose = 20 μM, initial pH 6.0 without buffer, T = 20 °C.

iopromide degradation was nearly not influenced when the concentration of NaClO4, which was used to adjust the ionic strength, increased from 10 to 200 mM in the reaction solution. The result indicates that there is a strong inner-sphere interaction between the HSO5− and the active Cu(II) on the oxide surface, that is the catalytic oxidation arises from a strong HSO5− binding or complexation with the surface Cu(II) sites. It confirms the assumption made from CuO dissolution (Figure S6 of the Supporting Information). The oxide surface was positively charged at the reaction pH of 6.0. The result also shows that the intensity of electrostatic attraction between the positively charged surface and the negatively changed HSO5− has no significant influence on the reaction rate. Then, the uncharged O in the HSO5− (molecular structure shown in Scheme S1 of the Supporting Information) is likely the coordination atom for the surface Cu(II) sites leading to the formation of powerful radicals. Hydroxyl radical, sulfate radical, and peroxymonosulfate radical are considered possible radical species in activated PMS system.2 Hydroxyl radical reaction can be scavenged by tertbutanol (t-BuOH) (k·OH = (3.8−7.6) × 108 M−1 s−1). Sulfate radical reacts with t-BuOH much more slowly (kSO4·− = (4−9.1) × 105 M−1 s−1). 33 Both hydroxyl radical and sulfate radical react very rapidly with ethanol (k·OH = (1.2−2.8) × 109 M−1

Figure 4. Influence of pH on iopromide degradation during PMS/ CuFe2O4 oxidation. Conditions: oxide dose = 100 mg L−1, initial iopromide concentration = 1 μM, PMS dose = 20 μM, 10 mM tetraborate adjusted pH, T = 20 °C. 2787

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

s−1; kSO4·− = (1.6−7.8) × 107 M−1 s−1). Peroxymonosulfate radical is relatively inert toward alcohols. The two radical scavengers were used to differentiate the contribution of the three radical species in pollutant removal. 2 It was noticed that the two scavengers had nearly no negative effect on the HSO5− degradation (part A of Figure S7 of the Supporting Information) indicating that they did not retard the interactions between HSO5− and the oxide surface. About 74% iopromide were degraded in the presence of 10 mM t-BuOH in 20 min reaction as compared with nearly 100% iopromide degradation in the absence of radical scavengers (part B of Figure S7 of the Supporting Information). The presence of 10 mM of ethanol nearly completely inhibited iopromide degradation. These results suggest that sulfate radical is the major radical species responsible for the iopromide degradation. A small amount of hydroxyl radical was also generated in the catalytic oxidation, which possibly comes from the reaction of sulfate radical with OH− in water at the experimental pH 7.5. 8 No contribution of peroxymonosulfate radical was observed in the iopromide degradation, which might be due to its negligible production or low oxidation potential (1.1 V).15 Chlorine radical (Cl2·−) can be produced from sulfate radical oxidation of chloride (kSO4·− = 2.0 × 108 M−1 s−1) followed by rapid combination with another chloride in water.34 Because the oxidation potential of chlorine radical (EoCl2·−/2Cl− = 2.09 V) is lower than sulfate radical (EoSO4·−/SO42− = 2.5−3.1 V), the presence of chloride in water can reduce the oxidation efficiency of sulfate radical. 35 It was observed that the iopromide degradation rate was reduced by about 50% in the presence of 2 mM chloride (Figure S8 of the Supporting Information). As the chloride concentration increased from 0.5 to 4 mM, the iopromide degradation rate varied slightly. The result indicates that chlorine radical may acts as the main oxidant in the iopromide degradation in the normal chloride concentration range of surface waters. At 2.0 mM chloride, the increase of PMS dose from 20 to 100 μM improved the iopromide degradation significantly (Figure S8 of the Supporting Information). Therefore, the negative effect of chloride on iopromide degradation can be compensated by increasing the PMS dosage. Natural organic matter (NOM) in water can consume a large fraction of sulfate radical due to its relatively high concentration and reactivity toward the oxidant compared with the trace recalcitrant pollutant. It was shown that, in the PMS/CuFe2O4 system, the increase of the dosage of the Suwannee River hydrophobic NOM fraction (i.e., humic substance) significantly reduced the iopromide degradation rate (Figure S9 of the Supporting Information). Higher PMS dosage compensated the negative effect of the NOM on the iorpomide degradation rate. To reduce PMS dosage and to improve pollutant degradation efficiency during the treatment of surface water or reclaimed water, the PMS/CuFe2O4 process cannot be applied alone. It should be applied as a polishing treatment (i.e., after partly NOM removal using enhanced coagulation, resin or activated carbon adsorption, or after reducing oxidant demand of the NOM using preoxidation with other oxidants). Quantification of Radical Yield. To quantify the radical yield from PMS in the catalytic oxidation, oxalate was used as another probe compound, because (1) the initiation of oxalate degradation needs strong oxidant species, (2) oxalate is hydrophilic and has a small molecular size facilitating the diffusion into the pores of the catalyst, thus reducing the side

reactions of sulfate radicals that caused ineffective oxidant consumption. The oxidation of oxalate requires the abstraction of two electrons from the molecule. The initial step is to produce an oxalate radical. Because it can be oxidized to CO2 even by O2 with an unknown rate constant, the oxalate radical was reported to be reductive,36 thus being more oxidizable than oxalate. For an oxalate/PMS molar ratio of 1:1, about 51% of oxalate was finally removed in the PMS/CuFe2O4 reaction system (Figure 6). The DOC removal rate was similar (Figure

Figure 6. Oxalate removal in the PMS/CuFe2O4 reaction system under ambient condition, nitrogen gas bubbling, and oxygen gas bubbling, respectively. Conditions: oxide dose = 100 mg L−1; PMS dose = 0.1 mM, initial oxalic acid concentration = 0.1 mM; initial pH 3.8 without buffer, T = 20 °C.

S10 of the Supporting Information). The oxalate removal by CuFe2O4 adsorption alone was only 4%. Therefore, one mole of PMS can mineralize about 0.5 mol of oxalate during the catalytic oxidation. The oxalate removal was also tested in the presence of O2 and N2 gas bubbling, respectively. The oxalate degradation rates (Figure 6) and mineralization rates (Figure S10 of the Supporting Information) for the two conditions equaled the one observed without gas introduction. Therefore, O2 does not contribute significantly to the oxidation of oxalate radical in this system. Sulfate radical as well as hydroxyl radical are much more reactive toward the oxalate radical through one electron transfer. The result indicates that the radical yield ratio from PMS is about 1 mol/mol, a relatively high radical generation ratio as compared to Fenton, Fenton-like and baseactivated peroxodisulfate reactions.1,6 Sulfate Radical Generation Mechanism. It was suggested that the feasibility of the catalytic PMS oxidation with free metal ions (M n+ ) was determined by the thermodynamic favorability of reducing the M(n + 1)+ to Mn+ by PMS, i.e. the regeneration of Mn+.33 It is difficult to directly identify the existence of transient high-valent metal in the oxide surface during PMS-oxide interaction. However, M(n+1)+ has a smaller ionic radii than Mn+ and thus a stronger binding capacity with ligands.37 The binding of surface hydroxyl groups (−OH), universally formed through dissociative adsorption of water molecules on transition metal oxides, are expected to change with the oxidation state of the surface metal. Because PMS forms inner-sphere complexation with the surface Cu(II) of the oxides, some intermediates originated from PMS decomposition could be detected with in situ spectroscopic analysis. In situ characterization of the oxides’ surface during catalytic decomposition of PMS was conducted with ATR-FTIR (Figure 7) and confocal Raman (Figure S11 of the Supporting 2788

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

surface −OH partly shifted to the surface metal sites. It could be due to the increase of electron withdrawing power of the bonded surface metal, possibly arising from a higher transient valence. Therefore, the ATR-FTIR results indicate the formation of surface Cu(II)-(HO)OSO3− complex (depicted in eq 1) and surface −OH group bonded to a possible higher transient valence. It is supposed that there could be one electron transfer inside the surface Cu(II)−(HO)OSO3− complex. In this way, the new surface −OH group can be formed, and sulfate radical can be generated (eq 2). This electron transfer is similar to a Cu(II)−H2O2 reaction (eq 3) proposed in literature.41,42 ≡Cu(II) − −OH + HSO5− → ≡Cu(II)−(HO)OSO3− + OH−

(1)

≡Cu(II)−(HO)OSO3− → ≡ Cu(III)−−OH + SO4•− (2) −

•

Cu(II) + H 2O2 → Cu(III) − OH + OH

(3)

The in situ Raman spectra (Figure S11 of the Supporting Information) show that a new small and broad peak appeared at around 831 cm−1 during the catalytic decomposition of HSO5−. According to a number of studies,43−45 this peak can be ascribed to the formation of a peroxo species bond to surface metal sites which will finally detach from the surface as O2. This explanation is supported by the production of small gas bubbles from the oxide surface during the catalytic decomposition of HSO5− observed at high concentration (Figure S6 of the Supporting Information). It can be assumed that higher valence copper (i.e., Cu(III)) oxidized HSO5− to SO5·‑ (eq 4), adjacent SO5·− on the surface further combined with each other forming the surface peroxo species (finally O2) and the sulfate radicals (eq 5). On the basis of this series of surface reactions, one sulfate radical can be generated from one HSO5−, which is consistent with the radical yield stoichiometry drawn from Figure 6.

Figure 7. ATR-FTIR spectra of the PMS solution alone, the oxide particles in water, and the oxide particles in PMS solution. Conditions: initial pH of the pure water was 6.0; initial pH of the PMS solution was adjusted to pH 6.0 with NaOH.

Information). According to Gonzalez et al.,38 the IR bands at 1103 cm−1 and 1249 cm−1 can come from the S−O stretching vibration of either HSO5− or SO42−. The intensity of the 1249 cm−1 band was significantly reduced relative to the neighboring band (1103 cm−1) in presence of the metal oxides. Considering the significant PMS decomposition in presence of the oxide (Figure S7 of the Supporting Information), this band can be assigned to HSO5−. Moreover, there was a blue-shift of 28 and 7 cm−1 of this band in presence of CuO and CuFe2O4 respectively as compared to the PMS solution alone. The result indicates that the OH in the S−O−(OH) structure of HSO5−, when bonded to the oxide surfaces, draw less electron density from the neighboring S−O bond leaving it stronger. It is reasonable to assume that the surface Cu(II) donated electron to the OH of HSO5− inhibiting the electron attraction from the neighboring S−O. ATR-FTIR spectra of aqueous slurries of metal oxides, when water spectrum was subtracted as background, usually show stretching vibrations of surface −OH around 3100 cm−1 and surface H-bonded water molecules in the range of 3200−3500 cm−1.39,40 The spectra of the two oxides in water showed broad peaks with maximum intensity at 3076 and 3135 cm−1 respectively indicating the existence of surface −OH groups. In the presence of HSO5−, these peaks became narrow and redshifted by 36 and 70 cm−1 for CuO and CuFe2O4, respectively. This result indicates that HSO5− replaced some surface −OH to bond with the metal sites of the oxides during the catalytic decomposition. Because the inner-sphere complexation on the oxide surface occurs through the replacement of the surface −OH groups, 32 this result is consistent with the conclusion drawn from Figure 5. The red-shift of the surface −OH band during the reaction suggests that the electron density of the

≡Cu(III)−−OH + HSO5− → ≡Cu(II)−•OOSO3− + H 2O (4)

2≡Cu(II)−•OOSO3− + 2H 2O → 2≡Cu(II)−−OH + O2 + 2SO4•− + 2H+

(5)

In a latest study, Ding et al. tested tetragonal CuFe2O4 (a different crystalline from and 10 times lower of saturated magnetization than the catalyst of this study) to improve PMS oxidation of tetrabromobisphenol A. 46 They proposed that sulfate radical could be produced from PMS oxidation of Cu(I) and Fe(II), which were supposed to be prereduced from Cu(II) and Fe(III) by PMS. These reactions might occur in some cases, but they are not dominant reactions in our system considering the following facts: (1) the reduction of either Cu (II) or Fe(III) by PMS is quite thermodynamically unfavorable because of much higher reducing potential of HSO5− (1.8 V) than Cu(II) (0.15 V) and Fe(III) (0.77 V) leading to inefficient reduction for the two metals (e.g., a first-order rate constant of only 9 × 10−13 s−1 was reported for Cu(II) reduction by acidic H2O2)16 and consequently inefficient sulfate radical generation; (2) XPS results (ex situ and indirect characterization) showed no discernible metal valence change for the catalyst reused for several cycles (Figure S2 of the Supporting Information); (3) PMS−Cu(I) reaction produces hydroxyl radical but not sulfate 2789

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

radical;15 (4) Fe2O3 and MnFe2O4 showed negligible activity (Figure 1) indicating that Fe(III) both in spinel and in nonspinel structures are not active enough; (5) according to their proposed mechanism, the molar yield ratio of sulfate radical from PMS is 0.5 which is only half of that in our system (i.e., 1). Anipsitakis and Dionysiou tested the activity of a variety of free metal cations to generate sulfate radical from the catalytic decomposition of HSO5− in water.33 They found that only Co2+ had high efficiency, which was attributed to the high reduction potential of Co3+/Co2+ (1.92 V) making the regeneration of Co2+ with HSO5− thermodynamically feasible. For metal ions with lower M(n + 1)+/Mn+ reduction potential such as Ce4+/Ce3+ (1.72 V) and Mn3+/Mn2+ (1.54 V), their activity were much lower. The reduction potential of Cu(III)/ Cu(II) was reported to be 1.57 V in ionized form and 2.3 V in solid phase, respectively.47 Free Cu2+ showed very low activity in this process (Figure 1) possibly due to the low reduction potential of Cu3+/Cu2+. In contrast, the high activity of the Cu(II) in the solid phase is likely related to the higher reduction potential of Cu(III)/Cu(II). A simple, nonhazardous, low energy input, and high efficient oxidation process is always desirable for the degradation of recalcitrant organic pollutants in case of small scale water treatment. The results of this study showed that PMS oxidation catalyzed with magnetic spinel CuFe2O4 might be a feasible choice to meet with this requirement. CuFe2O4 was found to be stable, efficient, and magnetically separable for the catalytic PMS oxidation solution. Both NOM and chloride in water can compete with the targeted pollutants toward the highly powerful sulfate radical. Considering the regulation of maximum sulfate concentration in potable water (usually 2.5 mM), it is necessary to reduce, as much as possible, the concentration of NOM or its oxidant demand prior to sulfate radical oxidation with the objective to minimize PMS dose. Low-DOC and low-salinity water is more desirable for this oxidation process. Additional research is needed to identify oxidation products and to evaluate selectivity and limitation of this process for a variety of recalcitrant pollutants under various water matrixes. The optimal preparation conditions and the influence of any other metal doping also need to be studied to further improve the stability, the activity, and the recyclability of the spinel CuFe2O4.

■

WDRC for SEM and ICP-MS analysis. We also appreciate the anonymous reviewers’ valuable revision suggestions.

■

(1) Pham, A. L.; Lee, C.; Doyle, F. M.; Sedlak, D. L. A silicasupported iron oxide catalyst capable of activating hydrogen peroxide at neutral pH. Environ. Sci. Technol. 2009, 43, 8930−8935. (2) Anipsitakis, G. P.; Dionysiou, D. D. Degradation of organic contaminants in water with sulfate radical generated by the conjunction of peroxymonosulfate with cobalt. Environ. Sci. Technol. 2003, 37, 4790−4797. (3) Hori, H.; Yamamoto, A.; Hayakawa, E.; Taniyasu, S.; Yamashita, N.; Kutsuna, S. Efficient decomposition of environmentally persistent perfluorocarboxylic acids by use of persulfate as a photochemical oxidant. Environ. Sci. Technol. 2005, 39, 2383−2388. (4) Lau, T. K.; Chu, W.; Graham, N. J. D. The aqueous degradation of butylated hydroxyanisole by UV/S2O82‑: Study of reaction mechanism via dimerization and mineralization. Environ. Sci. Technol. 2007, 41, 613−619. (5) Waldemer, R. H.; Tratnyek, P. G.; Johnson, R. L.; Nurmi, J. T. Oxidation of chlorinated ethenes by heat-activated persulfate: Kinetics and products. Environ. Sci. Technol. 2007, 41, 1010−1015. (6) Furman, Q. S.; Teel, A. L.; Watts, R. J. Mechanism of base activation of persulfate. Environ. Sci. Technol. 2010, 44, 6423−6428. (7) Antoniou, M. G.; De La Cruz, A. M.; Dionysiou, D. D. Intermediates and reaction pathways from the degradation of microcystin-LR with sulfate radicals. Environ. Sci. Technol. 2010, 44, 7238−7244. (8) Guan, Y.; Ma, J.; Li, X.; Fang, J.; Chen, L. Influence of pH on the formation of sulfate and hydroxyl radicals in the UV/peroxymonosulfate system. Environ. Sci. Technol. 2011, 45, 9308−9314. (9) Drzewicz, P.; Perez-Estrada, L.; Alpatova, A.; Martin, J. W.; Gamal El-Din, M. Impact of peroxydisulfate in the presence of zero valent iron on the oxidation of cyclohexanoic acid and naphthenic acids from oil sands process-affected water. Environ. Sci. Technol. 2012, 46, 8984−8991. (10) Anipsitakis, G. P.; Stathatos, E.; Dionysiou, D. D. Heterogeneous activation of oxone using Co3O4. J. Phys. Chem. B 2005, 109, 13052−13055. (11) Yang, Q.; Choi, H.; Chen, Y.; Dionysiou, D. D. Heterogeneous activation of peroxymonosulfate by supported cobalt catalysts for the degradation of 2,4-dichlorophenol in water: The effect of support, cobalt precursor, and UV irradiation. Appl. Catal., B 2008, 77, 300− 307. (12) Yang, Q.; Choi, H.; Al-Abed, S.; Dionysiou, D. D. Iron-cobalt mixed oxide nanocatalysts: Heterogeneous peroxymonosulfate activation, cobalt leaching, and ferromagnetic properties for environmental applications. Appl. Catal., B 2009, 88, 462−469. (13) Hayes, R. B. The carcinogenicity of metals in humans. Cancer Causes Control 1997, 8, 371−385. (14) Ji, F.; Li, C.; Deng, L. Performance of CuO/Oxone system: Heterogeneous catalytic oxidation of phenol. Chem. Eng. J. 2011, 178, 239−243. (15) Gilbert, B. C.; Stell, J. K. Mechanisms of peroxide decomposition: An electron paramagnetic resonance study of the reaction of the peroxomonosulphate anion with CuI. A marked contrast in behavior with that of TiIII and FeII. J. Chem. Soc., Faraday Trans. 1990, 86, 3261−3266. (16) Negri, A. R.; Jimenez, G.; Hill, R. T.; Francis, R. C. Caroate delignification Part 4: The generation and role of hydroxyl radicals. Tappi J. 1998, 81 (5), 241−246. (17) Handbook of Chemistry and Physics, 77th ed.; Lide, D. R., Eds.; CRC Press: Boca Raton, Fl, 1996. (18) Brandt, C.; van Eldik, R. Transition metal-catalyzed oxidation of sulfur(IV) oxides. Atmospheric-relevant processes and mechanisms. Chem. Rev. 1995, 95, 119−190. (19) Toxicological Profile of Copper; U.S. Department of Health and Human Health Services, Public Health Services, Agency for Toxic Substances and Disease Registry: Atlanta, 2004.

ASSOCIATED CONTENT

S Supporting Information *

Procedures of ATR-FTIR and Raman characterization. This material is available free of charge via the Internet at http:// pubs.acs.org.

■

REFERENCES

AUTHOR INFORMATION

Corresponding Author

*Tel.: + 966 (0) 2 808 2984, e-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS The authors want to thank Dr. Yang Yang and Mr. Bei Zhang of Imaging and Characterization Laboratory of KAUST for their help in performing Raman and saturated magnetization analysis, Dr. Hua Tan of Analytical Chemistry Lab of KAUST for BET analysis, and Dr. Cyril Aubry and Ms. Tong Zhan of 2790

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791

Environmental Science & Technology

Article

(20) Wang, W.; Xu, Z.; Finch, J. Fundamental study of ambient temperature process in the treatment of acid mine drainage. Environ. Sci. Technol. 1996, 30, 2604−2608. (21) Shih, K.; White, T.; Leckie, J. O. Nickel stabilization efficiency of aluminate and ferrite spinels and their leaching behavior. Environ. Sci. Technol. 2006, 40, 5520−5526. (22) Mckinnon, W.; Choung, J. W.; Xu, Z.; Finch, J. A. Magnetic seed in ambient temperature ferrite process applied to acid mine drainage treatment. Environ. Sci. Technol. 2000, 34, 2576−2581. (23) Morgan, B. E.; Lahav, O.; Loewenthal, R. E. Advances in seeded ambient temperature ferrite formation for treatment of acid mine drainage. Environ. Sci. Technol. 2005, 39, 7678−7683. (24) Kormos, J. L.; Schulz, M.; Ternes, T. A. Occurrence of iodinated X-ray contrast media and their biotransformation products in the urban water cycles. Environ. Sci. Technol. 2011, 45, 8723−8732. (25) Huber, H. H.; Canonica, S.; Park, G. Y.; von Gunten, U. Oxidation of pharmaceuticals during ozonation and advanced oxidation process. Environ. Sci. Technol. 2003, 37, 1016−1024. (26) Westerhoff, P.; Yoon, Y.; Snyder, S.; Wert, E. Fate of endocrine -disruptor, pharmaceutical, and personal care product chemicals during simulated drinking water treatment processes. Environ. Sci. Technol. 2005, 39, 6649−6663. (27) Li, W.; Li, J.; Guo, J. Synthesis and characterization of nanocrystalline CoAl 2 O4 spinel powder by low temperature combustion. J. Eur. Ceram. Soc. 2003, 23, 2289−2295. (28) Najjar, W.; Azabou, S.; Sayadi, S.; Ghorbel, A. Catalytic wet peroxide photo-oxidation of phenolic olive oil mill wastewater contaminants. Appl. Catal., B 2007, 74, 11−18. (29) Faungnawakij, K.; Tanaka, Y.; Shimoda, N.; Fukunaga, T.; Kikuchi, R.; Eguchi, K. Hydrogen production from dimethyl ether steam reforming over composite catalysts of copper ferrite spinel and alumina. Appl. Catal., B 2007, 74, 144−151. (30) Faour, A.; Mousty, C.; Prevot, V.; Devouard, B.; De Roy, A.; Bordet, P.; Elkaim, E.; Taviot-Gueho, C. Correlation among structure, microstructure, and electrochemical properties of NiAl-CO3 layered double hydroxide thin film. J. Phys. Chem. C 2012, 116, 15646−15659. (31) Thompson, R. C. Catalytic decomposition of peroxymonosulfate in aqueous perchloric acid by the dual catalysts Ag+ and S2O82‑ and by Co2+. Inorg. Chem. 1981, 20, 1005−1010. (32) Stumm, W. Chemistry of the Solid-Water Interface: Process at the Mineral-Water and Particle-Water Interface in Natural Systems; John Wiley & Sons: New York, 1992. (33) Anipsitakis, G. P.; Dionysiou, D. D. Radical generation by the interaction of transition metals with common oxidants. Environ. Sci. Technol. 2004, 38, 3705−3712. (34) Heta, P.; Huie, R. E.; Ross, A. B. Rate constants for reactions of inorganic radicals in aqueous solution. J. Phys. Chem. Ref. Data 1988, 17, 1027−1284. (35) Anipsitakis, G. P.; Dionysiou, D. D.; Gonzalez, M. A. Colbaltmediated activation of peroxymonosulfate and sulfate radical attack on phenolic compounds: Implication of chloride ions. Environ. Sci. Technol. 2006, 40, 1000−1007. (36) Zuo, Y.; Hoigné, J. Formation of hydrogen peroxide and depletion of oxalic acid in atmospheric water by photolysis of iron (III)-oxalato complexes. Environ. Sci. Technol. 1992, 26, 1014−1022. (37) Ahrland, S.; et al. Solution chemistry. In Comprehensive Inorganic Chemistry; Bailar, J. C., Emeleus, H. J., Nyholm, R., TrotmanDickenson, A. F., Eds.; Pergamon Press: Oxford, 1973. (38) Gonzalez, J.; Torrent-Sucarrat, M.; Anglada, J. M. The reactions of SO3 with HO2 radical and H2O...HO2 radical complex. Theoretical study on the atmospheric formation of HSO5 and H2SO4. Phys. Chem. Chem. Phys. 2010, 12, 2116−2125. (39) Tejedor-Tejedor, M. I.; Anderson, M. A. “In Situ” attenuated total reflectance Fourier transformation infrared studies of the goethite (α-FeOOH)-aqueous solution interface. Langmuir 1986, 2, 203−210. (40) Zhang, T.; Li, C.; Ma, J.; Tian, H.; Qiang, Z. Surface hydroxyl groups of synthetic a-FeOOH in promoting •OH generation from aqueous ozone: Property and activity relationship. Appl. Catal., B 2008, 82, 131−137.

(41) Amine, A.; Atmani, Z.; El Hallaoui, A.; Giorgi, M. Pierrot, M.; Réglier, M. Copper(II)/H2O2-mediated DNA cleavage: Improvement of a copper(III) species in H-atom abstraction of deoxyribose units. Bioorg. Med. Chem. Lett. 2002, 12, 57−60. (42) Hatcher, L. Q.; Karlin, K. D. Oxidant types in copper-dioxygen chemistry: The ligand coordination defines the Cun-O2 structure and subsequent reactivity. J. Biol. Inorg. Chem. 2004, 9, 669−683. (43) Pushkarev, V. V.; Kovalchuk, V. I.; d’Itri, J. L. Probing defect sites on the CeO2 surface with dioxygen. J. Phys. Chem. B 2004, 108, 5341−5348. (44) Shimizu, K.; Murata, Y.; Satsuma, A. Dicopper-dioxygen complexes in Y zeolite for selective oxidation of cyclohexane under photoirradiation. J. Phys. Chem. C 2007, 111, 19043−19051. (45) Zhang, T.; Li, W.; Croué, J. P. Catalytic ozonation of oxalate with a cerium supported palladium oxide: An efficient degradation not relying on hydroxyl radical oxidation. Environ. Sci. Technol. 2011, 45, 9339−9346. (46) Ding, Y.; Zhu, L.; Wang, N.; Tang, H. Sulfate radicals induced degradation of tetrabromobisphenol A with nanoscaled magnetic CuFe2O4 as a heterogeneous catalyst of peroxymonosulfate. Appl. Catal., B 2013, 129, 153−162. (47) Popova, T. V.; Aksenova, N. V. Complexes of copper in unstable oxidation states. Russ. J. Coord. Chem. 2003, 29 (11), 743−76.

2791

dx.doi.org/10.1021/es304721g | Environ. Sci. Technol. 2013, 47, 2784−2791