Hydroxylamine Promoted Goethite Surface Fenton Degradation of

Mar 30, 2017 - ... Fe(III)/Fe(II) Cycle on Ilmenite Surface to Enhance Persulfate Catalytic Activation and Aqueous Pharmaceutical Ibuprofen Degradatio...
1 downloads 0 Views 2MB Size
Download PDF
Article pubs.acs.org/est

Hydroxylamine Promoted Goethite Surface Fenton Degradation of Organic Pollutants Xiaojing Hou, Xiaopeng Huang, Falong Jia, Zhihui Ai, Jincai Zhao, and Lizhi Zhang* Key Laboratory of Pesticide & Chemical Biology of Ministry of Education, Institute of Environmental & Applied Chemistry, Central China Normal University, Wuhan 430079, P. R. China

Downloaded via TULANE UNIV on January 22, 2019 at 02:19:11 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

S Supporting Information *

ABSTRACT: In this study, we construct a surface Fenton system with hydroxylamine (NH2OH), goethite (α-FeOOH), and H2O2 (α-FeOOH−HA/H2O2) to degrade various organic pollutants including dyes (methyl orange, methylene blue, and rhodamine B), pesticides (pentachlorophenol, alachlor, and atrazine), and antibiotics (tetracycline, chloramphenicol, and lincomycin) at pH 5.0. In this surface Fenton system, the presence of NH2OH could greatly promote the H2O2 decomposition on the α-FeOOH surface to produce ·OH without releasing any detectable iron ions during the alachlor degradation, which was different from some previously reported heterogeneous Fenton counterparts. Moreover, the ·OH generation rate constant of this surface Fenton system was 102−104 times those of previous heterogeneous Fenton processes. The interaction between α-FeOOH and NH2OH was investigated with using attenuated total reflectance Fourier transform infrared spectroscopy and density functional theory calculations. The effective degradation of organic pollutants in this surface Fenton system was ascribed to the efficient Fe(III)/Fe(II) cycle on the α-FeOOH surface promoted by NH2OH, which was confirmed by X-ray photoelectron spectroscopy analysis. The degradation intermediates and mineralization of alachlor in this surface Fenton system were then systematically investigated using total organic carbon and ion chromatography, liquid chromatography−mass spectrometry, and gas chromatography−mass spectrometry. This study offers a new strategy to degrade organic pollutants and also sheds light on the environmental effects of goethite.

■

diaminetetraacetic acid (EDTA), and oxalate.7 They attributed the enhanced efficiency of the heterogeneous Fenton system to iron ions absorbed on the catalyst surface, rather than more dissolved iron ions in the solution. For example, EDTA and CMCD could more strongly complex with iron ions than oxalate and thus promote the dissolution of iron bearing minerals, increasing the dissolved iron ion concentrations to the range of 2.8−14 mg/L during the heterogeneous Fenton processes, but oxalate was found to be more efficient to improve the efficiency of heterogeneous Fenton system than EDTA and CMCD. Interestingly, Wang et al. found that mesoporous copper ferrite could effectively catalyze the H2O2 decomposition without significantly releasing iron ions, as the concentration of iron ions during the heterogeneous Fenton processes was less than 1 mg/L even under acidic condition. More importantly, this catalyst retained its high catalytic activity even after 5 cycles of use.8 According to the concentration of iron ions released during the heterogeneous Fenton processes, herein we intend to classify heterogeneous Fenton systems into heterogeneous Fenton-like and surface Fenton ones. In the heterogeneous Fenton-like systems of considerable dissolved iron ions in the solution, the Fe(III)/Fe(II) cycle, the generation of reactive oxygen species (ROS), and the pollutant

INTRODUCTION To avoid the drawbacks of traditional homogeneous Fenton (Fe2+/H2O2) systems, such as ferric oxide sludge generation, the straight working pH value of 2.0−3.5, and damage to the catalyst, heterogeneous Fenton systems are designed with iron bearing catalysts for pollutant control and environmental remediation.1−4 However, the Fe(III)/Fe(II) cycle of most iron bearing catalysts during the heterogeneous Fenton process is not as effective as that of traditional homogeneous Fenton counterpart, lowering the efficiencies of H2O2 decomposition and hydroxyl radical (·OH) generation, and thus restricting the application of heterogeneous Fenton process. For instance, the ·OH generation rate constants of heterogeneous Fenton reagents such as goethite, hematite, and ferrihydrite were as low as 4.00 × 10−7, 4.25 × 10−5, and 2.00 × 10−5 s−1, respectively.5 In order to promote the efficiency of heterogeneous Fenton process, chelating or reducing agents were often employed to enhance the Fe(III)/Fe(II) cycle of iron oxide in heterogeneous Fenton reactions. For instance, the addition of carboxymethyl-β-cyclodextrin (CMCD) could improve the degradation of 2,4,6-trinitrophenolin by a factor of 3 in the magnetitebearing mineral heterogeneous Fenton system, because CMCD could favor the dissolution of iron ions from mineral surface.6 Moreover, Xue et al. reported that the pentachlorophenol degradation rate in the magnetite heterogeneous Fenton process increased by 1.7−5.7 times after adding chelating agents such as succinate, citrate, tartrate, CMCD, ethylene© 2017 American Chemical Society

Received: Revised: Accepted: Published: 5118

November 22, 2016 March 29, 2017 March 30, 2017 March 30, 2017 DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126

Article

Environmental Science & Technology

solution was produced by dissolving HA in deionized and oxygen-free water. All the concentrations of organic pollutant solutions were 20 mg/L. The NaOH and HCl solutions were utilized to adjust the pH values of degradation solutions. All the stock solutions were freshly prepared before use and covered with foil paper to avoid light. Degradation Procedures. Batch trials were carried out in 100 mL conical flasks under magnetic stirring at 25 ± 2 °C. Briefly, the reactions were triggered by adding 5 mg of αFeOOH, 250 μL of 0.1 mol/L HA solution, and 500 μL of 0.1 mol/L H2O2 solution into 50 mL of alachlor solution (20 mg/ L) in sequence. After that, 900 μL of the reaction solution was sampled at predetermined intervals, following with the filtration through 0.22 μm filter membranes. Subsequently, ethanol (100 μL) was added immediately into the sampled solutions to stop the reaction for the alachlor concentration measurement. The initial pH value of the α-FeOOH−HA/H2O2 system was 5.0 without adjustment. HCl and NaOH were used to adjust the initial pH values of the reaction solutions. The anaerobic experiments were conducted by bubbling Ar gas to remove dissolved oxygen. Briefly, conical flasks sealed with the rubber stoppers were pumped with high purity Ar gas, along with needles being inserted in the rubber stoppers for gas escaping. To assess the stability of α-FeOOH, the catalyst after reaction was separated from the solution, then washed with deionized water and ethanol thoroughly, and finally vacuum-dried for the reuse. Besides alachlor, the degradation experiments of other organic pollutants, including rhodamine B, methyl orange, atrazine, pentachlorophenol, lincomycin, tetracycline, and chloramphenicol, were also conducted. All the other experimental procedures were the same as those of alachlor degradation. Analytical Methods. The analysis procedures of organic pollutants were described in the Supporting Information. The concentrations of total dissolved irons and surface Fe(II) were determined by a 1,10-phenanthroline method with a UV−vis spectrophotometer (UV-2550, Shimadzu, Japan) unless it was specially stated.18 Hydroxylamine hydrochloride was used for the measurement of total iron concentration. Briefly, 0.5 mL of the resulting solution was taken, following with filtration through 0.22 μm filter membranes. A 0.5 mL portion of hydroxylamine hydrochloride solution (100 g/L) was then added. After 5 min of reduction, 0.5 mL of 1,10-phenanthroline solution (2 g/L) and 0.5 mL of sodium acetate solution (10%, w:w) were added in sequence. After 30 min of chromogenic reaction, the total dissolved irons amount was then analyzed. For the dissolved Fe(II) detection, 0.5 mL of hydroxylamine hydrochloride solution (100 g/L) was replaced with 0.5 mL of H2O. Then, the Fe(II)−1,10-phenanthroline complex was measured at λ = 510 nm by the UV−vis spectrophotometer with the detection limit of 0.47 μmol/L (0.026 mg/L). Surface Fe(II) was determined with a modified 10-phenanthroline method after careful pretreatment.19,20 Typically, the αFeOOH slurry after reaction with HA for 1 h was deaerated by bubbling Ar gas for 10 min before being transferred to the centrifuge tube in a glovebox under anaerobic conditions. Subsequently, the supernatant was discarded and the solid sample in the centrifuge tube was rinsed with neutral oxygenfree water twice. All the operations were done in the anaerobic glovebox except the centrifugation process. As HA are easily dissolved in water and iron ions on the α-FeOOH surface are not soluble in water under neutral conditions, HA in the αFeOOH slurry and on the α-FeOOH surface could be removed

oxidation might be similar to those in homogeneous Fenton systems.6 However, in the surface Fenton systems with limited dissolved iron ions (90%) in the α-FeOOH−HA/H2O2 system. The low degradation efficiency of α-FeOOH/H 2 O 2 system was attributed to the poor surface Fe(III)/Fe(II) cycle of αFeOOH. Although HA of high concentration (10 mmol/L) could react with H2O2 to produce ·OH,16 only about 5% of alachlor was degraded within 60 min during the HA/H2O2 process because the concentration of HA was as low as 0.5 mmol/L in this α-FeOOH−HA/H2O2 system. The influence of HA concentration on the alachlor degradation in the αFeOOH−HA/H2O2 system was therefore investigated at the initial pH values of about 5.0. Along with increasing the HA concentration from 0.0 to 0.75 mmol/L, the alachlor degradation rate first increased but then decreased slightly when the HA concentration was changed from 0.75 to 1.0 mmol/L (SI Figure S5). The slight decrease of alachlor degradation rate at higher HA concentration might be attributed to the competitive consumption of reactive oxygen species by excess HA. As the degradation percentage of alachlor (92%) in case of 0.5 mmol/L HA was close to that (94%) of 0.75 mmol/L HA, 0.5 mmol/L was chosen as the HA concentration for the subsequent investigation. The generation of ·OH in the α-FeOOH−HA/H2O2 system was first investigated by adding ·OH scavenger. As shown in Figure 3a, the addition of iso-propanol (IPA) significantly decreased the alachlor degradation efficiency of the αFeOOH−HA/H2O2 system, suggesting that ·OH was the major reactive species in the α-FeOOH−HA/H2O2 system. Meanwhile, the alachlor degradation curve did not change when we bubbled Ar gas to remove dissolved oxygen from the α-FeOOH−HA/H2O2 system (Figure 3b), ruling out the participation of molecular oxygen in the alachlor degradation. 5122

DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126

Article

Environmental Science & Technology

Figure 5. (a) DMPO trapped ESR spectra in the α-FeOOH−HA/H2O2 and α-FeOOH−HA/H2O2/F− systems. (b) Time profiles of the ESR signal of DMPO−·OH generation. The initial concentrations of HA, H2O2, and F− were 5 × 10−4, 1 × 10−3, and 2 × 10−3 mol/L, respectively; the dosage of α-FeOOH was 0.1 g/L; the initial pH was 5.0.

cost a longer time under these conditions. Interestingly, the surface Fe(II) chelation with 1,10-phenanthroline or bipyridine completely inhibited the alachlor degradation and the H2O2 decomposition in the α-FeOOH−HA/H2O2 system, confirming the indispensable roles of surface Fe(II) in the alachlor removal with the α-FeOOH−HA/H2O2 system (SI Figure S12). Since Fe(II) on the α-FeOOH surface was the main species to catalyze the H2O2 decomposition, H2O2 might be decomposed on the α-FeOOH surface and the surface-bound · OH was generated in the α-FeOOH−HA/H2O2 system. To confirm the decomposition of H2O2 on the α-FeOOH surface, the changes of surface hydroxyl groups (−OH) and surface peroxide groups (−O22−) of α-FeOOH in the αFeOOH/H2O2 and α-FeOOH−HA/H2O2 systems were monitored with O 1s XPS spectra analysis (Figure 4b). The O 1s peaks at 530.0 and 531.0 eV were attributed to the lattice oxygen (Fe−O) and lattice hydroxyl groups (Fe−OH lattice) of α-FeOOH, respectively.31 Differently, the O 1s XPS spectra of α-FeOOH separated from the α-FeOOH/H2O2 system possessed another shoulder at 532.1 eV, which was ascribed to the −O22− groups.35 As expected, this shoulder from −O22− groups on the α-FeOOH surface decreased and the peak of surface hydroxyl groups at 533.3 eV increased obviously for αFeOOH in the α-FeOOH−HA/H2 O2 system.31 These phenomena strongly confirmed the generation of surfacebound ·OH via the H2O2 decomposition catalyzed by Fe(II) in situ generated on the α-FeOOH surface in the presence of HA. To further confirm the existence of surface bound ·OH generated by the surface bound Fe(II) on α-FeOOH, we compared the amounts of ·OH generated in the supernatants of the α-FeOOH−HA/H 2O 2 and α-FeOOH−HA/H2 O2 /F− systems with DMPO trapped ESR technique (Figure 5), regarding that F− in the solution could desorb ·OH bound on α-FeOOH surface by forming strong hydrogen bond (surfacebound ·OH···F-FeOOH).36,37 As expected, the addition of F− improved the intensity of DMPO trapped ·OH ESR signal by 30% in the α-FeOOH−HA/H2O2 system after 120 s of reaction, well validating that the surface-bound ·OH was generated by the surface bound Fe(II) on α-FeOOH in the αFeOOH−HA/H2O2 system. Subsequently, we compared the α-FeOOH−HA/H2O2 system with the corresponding homogeneous Fenton ones. Interestingly, the alachlor degradation percentages in the αFeOOH−HA/H2O2 (92%) and HA/Fe(II)/H2O2 (27%) systems were far higher than that (8%) in the Fe(II)/H2O2 system when the soluble Fe(II) concentrations in these homogeneous Fenton ones were equal to the surface Fe(II) concentration in the α-FeOOH−HA/H2O2 system (SI Figure

IPA were conducted (SI Figure S9). The addition of IPA completely inhibited the alachlor degradation in the αFeOOH−HA/H2O2 systems at different pH, confirming that ·OH mainly contributed to the alachlor removal in the αFeOOH−HA/H2O2 system. As mentioned before, surface Fe(III) of α-FeOOH could be reduced by HA. Regarding that dissolved iron ions were not detected in the α-FeOOH−HA system or in the α-FeOOH− HA/H2O2 system in the presence/absence of alachlor at pH 5.0 with ICP-OES (SI Figures S1 and S10), we concluded the H2O2 decomposition should mainly occur on the α-FeOOH surface. To understand the Fe(III)/Fe(II) cycle on the αFeOOH surface in the α-FeOOH−HA/H2O2 system, we employed high-resolution X-ray photoelectron spectroscopy (HRXPS) to check the generation of Fe(II) on the α-FeOOH surface during this Fenton process (Figure 4a). Previous study reported that the binding energies of Fe 2p3/2 and Fe 2p1/2 XPS peaks of α-FeOOH were located at 711.0 and 724.9 eV, respectively.31 In this study, the Fe 2p3/2 and Fe 2p1/2 peaks of the α-FeOOH samples treated with H2O or H2O2 were also located at 711.0 and 724.9 eV, respectively. This consistence suggested that H2O and H2O2 did not affect the Fe speciation on α-FeOOH surface. However, the Fe 2p1/2 peak of the αFeOOH sample separated from the α-FeOOH−HA/H2O2 system appeared at 724.3 eV, with a binding energy shift of −0.6 eV, suggesting that Fe(II) was generated on the αFeOOH surface.32 To further identify the generation of Fe(II) on the α-FeOOH surface separated from the α-FeOOH−HA/ H2O2 system, XPS depth profiling technique was employed (Figure 4a). As expected, the binding energies of Fe 2p1/2 peaks were located at 724.3, 724.4, 724.5, and 724.9 eV for the depths of 0, 0.5, 1.0, 1.5 nm, respectively. Therefore, Fe(II) existed at least between 0 and 1.0 nm depth of the α-FeOOH sample separated from the α-FeOOH−HA/H2O2 system. These results strongly confirmed that Fe(II) was generated inside or on the surface of α-FeOOH in the α-FeOOH−HA/H2O2 system. Subsequently, we checked the contribution of surface Fe(II) bound on α-FeOOH to the alachlor degradation by chelating the surface Fe(II) with 1,10-phenanthroline or bipyridine during the alachlor degradation.33,34 Meanwhile, we checked the iron dissolution from α-FeOOH in the presence of phenanthroline or bipyridine and HA (SI Figure S11) and did not detect any dissolved iron ions in the α-FeOOH/ phenanthroline or bipyridine, α-FeOOH−HA/phenanthroline or bipyridine, and α-FeOOH−HA/H2O2/phenanthroline or bipyridine systems at pH 5.0 within 1 h. These results were reasonable because the reductive dissolution of goethite would 5123

DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126

Article

Environmental Science & Technology

Scheme 1. Possible Alachlor Degradation Pathway in the αFeOOH−HA/H2O2 System

S13), confirming that HA could promote Fe(III)/Fe(II) cycle either in the solution or on the surface and the α-FeOOH surface bound Fe(II) was more active than the dissolved Fe(II) on the H2O2 decomposition, consistent with the previous reports.38,39 As expected, the α-FeOOH−HA/H2O2 system could generate more ·OH than these two homogeneous Fenton ones (SI Figure S14). Although the alachlor degradation rate in the α-FeOOH−HA/H2O2 system was lower than that in the Fe(II)/H2O2 one after increasing the initial concentration of Fe(II) to 1.1 mmol/L, the same as the concentration of αFeOOH (SI Figure S15), the α-FeOOH−HA/H2O2 system still succeeded with the reusability of α-FeOOH (SI Figure S16). Surprisingly, the reactivity of α-FeOOH even slightly increased along with repeated use, further confirming the indispensible roles of surface Fe(II) in the surface Fenton alachlor degradation. The variations of nitrate ions, chloride ions, and small molecule acid concentrations were then monitored to evaluate the alachlor degradation in the α-FeOOH−HA/H2O2 system. Given that both alachlor and HA contributed to the NO3− generation during the alachlor degradation in the α-FeOOH− HA/H2O2 system, we therefore compared the differences of NO3− concentration changes in the absence or presence of alachlor, which were 0.42 and 0.47 mmol/L, respectively. Their difference (0.05 mmol/L) suggested about 68% of nitrogen mineralization of alachlor (SI Figure S17a). During the alachlor degradation in the α-FeOOH−HA/H2O2 system, the Cl− concentration increased to 0.072 mmol/L within 60 min, corresponding to 98% of dechlorination. Meanwhile, the concentrations of HCOOH, CH3COOH, and HOOCCOOH, respectively, increased to 37.3, 96.2, and 108.8 μmol/L after 60 min, revealing that about 43% of alachlor was degraded to small molecule acids (SI Figure S17b). Moreover, the mineralization of alachlor in the α-FeOOH−HA/H2O2 system was analyzed by total organic carbon (TOC) analysis (SI Figure S18a). After 5 h of reaction, about 35% of TOC was removed. When the H2O2 concentration was increased from 1 to 10 mmol/L, the TOC removal percentage improved to 84% within 5 h. To eliminate the interference effect of H2O2, the TOC content changes of alachlor solution in the presence and absence of H2O2 were also monitored (SI Figure S18b). We found that the TOC content of alachlor solution was nearly unchanged in the presence of H2O2, suggesting the TOC analysis of alachlor solution was not affected by H2O2. The degradation intermediates of alachlor (1) were then detected with LC-MS and GC-MS. Eight degradation intermediates, including 2-chloro-N-(2,6-diethylphenyl)-Nmethylacetamide (2), 2-chloro-N-(2,6-diethylphenyl)acetamide (3), 2-chloro-N-(2-ethylphenyl)acetamide (4), N-(2-ethyl-6methylphenyl)acetamide (5), 2,6-diethylbenzenamine (6), 2ethyl-6-methylbenzenamine (7), 1,3-diethyl-2-nitrosobenzene (8), and 2-ethyl-6-methylphenol (9), were identified by GCMS (SI Figure S19), while two other intermediates, 2,6-diethylN-methylbenzenamine (10) and N-(2-acetyl-6-ethylphenyl)-2chloroacetamide (11), were checked out by LC-MS (SI Figure S20). The mass spectra of all these detected intermediates of alachlor were consistent with those reported in the literature,40−42 which were summarized in Table S4 (Supporting Information). The appearance of these degradation intermediates suggested a possible alachlor degradation pathway involving dechlorination, alkylic oxidation, hydroxylation, and dealkylation process in the α-FeOOH−HA/H2O2 system, as shown in Scheme 1. First, the C−Cl bond and C−N bond

were likely broken up to generate the N-dealkylation intermediates (compounds 2, 3, 4) and the dechlorination intermediate (compound 5). The compound 3 could be further oxidized by ·OH to generate compound 11. Meanwhile, paraposition carbon or N-adjacent atom was attacked by ·OH to generate the compounds 6, 7, and 10. Subsequently, the amino side chain on the benzene ring would be oxidized to generate the compounds 8 and 9. Eventually, ·OH could attack the benzene ring to generate small-molecule organic acids and finally mineralize alachlor. Meanwhile, it was interesting to find that HA could be degraded along with alachlor in the αFeOOH−HA/H2O2 system (SI Figure S21a), accompanying with 84% of HA mineralization into NO3− (SI Figure S21b). Therefore, we conclude that this surface Fenton system is very environmentally friendly. According to the above results, we would like to account for the high pollutant removal performance of α-FeOOH−HA/ H2O2 surface Fenton system as follows (Scheme 2). First, iron−HA complexes (Fe(III)−HA) formed on the αFeOOH surface. Then, the electron transfer from HA to surface ferric iron (Fe(III)) would produce surface bound ferrous species (Fe(II)) to decompose H2O2 for the generation of abundant ·OH to degrade alachlor and HA, Scheme 2. Schematic Illustration for the Possible Hydroxyl Radical Generation Mechanism in the α-FeOOH−HA/H2O2 System

5124

DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126

Article

Environmental Science & Technology accompanied by the formation of surface Fe(III)−HA complex. Subsequently, the formed Fe(III)−HA complex would be reduced by HA to trigger another cycle of surface Fenton alachlor degradation. Environmental Implications. In geochemical processes, the contaminants transformation dynamics and the geochemical cycling of other redox−active elements are strongly affected by iron redox cycling. Therefore, the understanding of interaction between goethite and reductants is of important environmental implications. In this study, the effective degradation of alachlor in the surface Fenton system (αFeOOH−HA/H2O2) was ascribed to the efficient Fe(III)/ Fe(II) cycle over the α-FeOOH surface promoted by NH2OH. Moreover, the reactivity of α-FeOOH even slightly increased along with repeated use, further confirming the indispensable roles of surface Fe(II) in the surface Fenton degradation of alachlor. As goethite, reductants, and hydrogen peroxide widely exit in nature, similar surface Fenton reaction will occur in the environment, affecting the transformation and conversion of organic contaminants. Thus, the mechanistic investigation of the surface Fenton system in this study can provide new insights into the intrinsic relationship between iron cycling and the geochemical cycling of other redox−active elements and also offer a new strategy with iron cycling to realize pollutant control and environmental remediation.

■

(4) Ma, J.; Song, W.; Chen, C.; Ma, W.; Zhao, J.; Tang, Y. Fenton degradation of organic compounds promoted by dyes under visible irradiation. Environ. Sci. Technol. 2005, 39, 5810−5815. (5) Kwan, W. P.; Voelker, B. M. Rates of hydroxyl radical generation and organic compound oxidation in mineral-catalyzed Fenton-like systems. Environ. Sci. Technol. 2003, 37, 1150−1158. (6) Matta, R.; Hanna, K.; Kone, T.; Chiron, S. Oxidation of 2,4,6trinitrotoluene in the presence of different iron-bearing minerals at neutral pH. Chem. Eng. J. 2008, 144, 453−458. (7) Xue, X.; Hanna, K.; Despas, C.; Wu, F.; Deng, N. Effect of chelating agent on the oxidation rate of PCP in the magnetite/H2O2 system at neutral pH. J. Mol. Catal. A: Chem. 2009, 311, 29−35. (8) Wang, Y.; Zhao, H.; Li, M.; Fan, J.; Zhao, G. Magnetic ordered mesoporous copper ferrite as a heterogeneous Fenton catalyst for the degradation of imidacloprid. Appl. Catal., B 2014, 147, 534−545. (9) Klausen, J.; Troeber, S. P.; Haderlein, S. B.; Schwarzenbach, R. P. Reduction of substituted nitrobenzenes by Fe(II) in aqueous mineral suspensions. Environ. Sci. Technol. 1995, 29, 2396−2404. (10) Hofstetter, T. B.; Schwarzenbach, R. P.; Haderlein, S. B. Reactivity of Fe(II) species associated with clay minerals. Environ. Sci. Technol. 2003, 37, 519−528. (11) Huang, X. P.; Hou, X. J.; Zhao, J. C.; Zhang, L. Z. Hematite facet confined ferrous ions as high efficient Fenton catalysts to degrade organic contaminants by lowering H2O2 decomposition energetic span. Appl. Catal., B 2016, 181, 127−137. (12) Viollier, E.; Inglett, P. W.; Hunter, K.; Roychoudhury, A. N.; Van Cappellen, P. The ferrozine method revisited: Fe(II)/Fe(III) determination in natural waters. Appl. Geochem. 2000, 15, 785−790. (13) Dain, R. P.; Gresham, G.; Groenewold, G. S.; Steill, J. D.; Oomens, J.; Van Stipdonk, M. J. Infrared multiple photon dissociation spectroscopy of group I and group II metal complexes with Bochydroxylamine. Rapid Commun. Mass Spectrom. 2013, 27, 1867−1872. (14) Zou, J.; Ma, J.; Chen, L.; Li, X.; Guan, Y.; Xie, P.; Pan, C. Rapid acceleration of ferrous iron/peroxymonosulfate oxidation of organic pollutants by promoting Fe(III)/Fe(II) cycle with hydroxylamine. Environ. Sci. Technol. 2013, 47, 11685−11691. (15) Chen, L.; Ma, J.; Li, X.; Zhang, J.; Fang, J.; Guan, Y.; Xie, P. Strong enhancement on Fenton oxidation by addition of hydroxylamine to accelerate the ferric and ferrous iron cycles. Environ. Sci. Technol. 2011, 45, 3925−3930. (16) Chen, L.; Li, X.; Zhang, J.; Fang, J.; Huang, Y.; Wang, P.; Ma, J. Production of Hydroxyl Radical via the Activation of Hydrogen Peroxide by Hydroxylamine. Environ. Sci. Technol. 2015, 49, 10373− 10379. (17) Fukuchi, S.; Nishimoto, R.; Fukushima, M.; Zhu, Q. Effects of reducing agents on the degradation of 2,4,6-tribromophenol in a heterogeneous Fenton-like system with an iron-loaded natural zeolite. Appl. Catal., B 2014, 147, 411−419. (18) Harvey, A. E.; Smart, J. A.; Amis, E. S. Simultaneous spectrophotometric determination of iron(II) and total iron with 1,10-phenanthroline. Anal. Chem. 1955, 27, 26−29. (19) Cooper, D. C.; Picardal, F.; Rivera, J.; Talbot, C. Zinc immobilization and magnetite formation via ferric oxide reduction by Shewanella putrefaciens 200. Environ. Sci. Technol. 2000, 34, 100−106. (20) Jeon, B. H.; Dempsey, B. A.; Burgos, W. D. Kinetics and mechanisms for reactions of Fe(II) with iron(III) oxides. Environ. Sci. Technol. 2003, 37, 3309−3315. (21) Lazrus, A. L.; Kok, G. L.; Gitlin, S. N.; Lind, J. A.; Mclaren, S. E. Automated fluorimetric method for hydrogen peroxide in atmospheric precipitation. Anal. Chem. 1985, 57, 917−922. (22) Frear, D.; Burrell, R. Spectrophotometric method for determining hydroxylamine reductase activity in higher plants. Anal. Chem. 1955, 27, 1664−1665. (23) Qin, Y. X.; Song, F. H.; Ai, Z. H.; Zhang, P. P.; Zhang, L. Z. Protocatechuic acid promoted alachlor degradation in Fe(III)/H2O2 Fenton system. Environ. Sci. Technol. 2015, 49, 7948−7956. (24) Huang, X. P.; Hou, X. J.; Song, F. H.; Zhao, J. C.; Zhang, L. Z. Facet-dependent Cr(VI) adsorption of hematite nanocrystals. Environ. Sci. Technol. 2016, 50, 1964−1972.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.6b05906. Additional descriptions, figures, and tables as mentioned in the text (PDF)

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone/Fax: +86-27-6786 7535. ORCID

Lizhi Zhang: 0000-0002-6842-9167 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported by Natural Science Funds for Distinguished Young Scholars (Grant 21425728), National Science Foundation of China (Grant 51472100), the 111 Project (Grant B17019), Self−Determined Research Funds of CCNU from the Colleges’ Basic Research and Operation of MOE (Grant CCNU14Z01001), excellent doctorial dissertation cultivation grant from Central China Normal University (Grant 2016YBZZ037), and the CAS Interdisciplinary Innovation Team of the Chinese Academy of Sciences.

■

REFERENCES

(1) Beltran, F. J.; Tarr, M. Chemical degradation methods for wastes and pollutants environmental and industrial applications; Marcel Dekker Inc.: New York, USA, 2003. (2) Haber, F.; Weiss, J. The catalytic decomposition of hydrogen peroxide by iron salts. Proc. R. Soc. London, Ser. A 1934, 147, 332−351. (3) Neyens, E.; Baeyens, J. A review of classic Fenton’s peroxidation as an advanced oxidation technique. J. Hazard. Mater. 2003, 98, 33− 50. 5125

DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126

Article

Environmental Science & Technology (25) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; et al. Gaussian 09, Revision A. L.; Gaussian, Inc.: Wallingford, CT, 2009. (26) Yeo, G. A.; Ford, T. A. Ab initio molecular orbital calculations of the infrared spectra of hydrogen-bonded complexes of water, ammonia and hydroxylamine. Theor. Chim. Acta. 1992, 81, 255−268. (27) Johnson, M. D.; Hornstein, B. J. The kinetics and mechanism of the ferrate (VI) oxidation of hydroxylamines. Inorg. Chem. 2003, 42, 6923−6928. (28) Keenan, C. R.; Sedlak, D. L. Factors affecting the yield of oxidants from the reaction of nanoparticulate zero-valent iron and oxygen. Environ. Sci. Technol. 2008, 42, 1262−1267. (29) Yang, X. J.; Xu, X. M.; Xu, J.; Han, Y. F. Iron oxychloride (FeOCl): an efficient Fenton-like catalyst for producing hydroxyl radicals in degradation of organic contaminants. J. Am. Chem. Soc. 2013, 135, 16058−16061. (30) Hug, S. J.; Leupin, O. Iron-catalyzed oxidation of Arsenic(III) by oxygen and by hydrogen peroxide: pH-dependent formation of oxidants in the Fenton reaction. Environ. Sci. Technol. 2003, 37, 2734− 2742. (31) Baltrusaitis, J.; Cwiertny, D. M.; Grassian, V. H. Adsorption of sulfur dioxide on hematite and goethite particle surfaces. Phys. Chem. Chem. Phys. 2007, 9, 5542−5554. (32) Fan, J. X.; Wang, Y. J.; Fan, T. T.; Dang, F.; Zhou, D. M. Effect of aqueous Fe(II) on Sb(V) sorption on soil and goethite. Chemosphere 2016, 147, 44−51. (33) Zhou, H.; He, Y.; Lan, Y.; Mao, J.; Chen, S. Influence of complex reagents on removal of chromium(VI) by zero-valent iron. Chemosphere 2008, 72, 870−874. (34) Carter, M. T.; Rodriguez, M.; Bard, A. J. Voltammetric studies of the interaction of metal chelates with DNA. 2. Tris-chelated complexes of cobalt(III) and iron(II) with 1,10-phenanthroline and 2,2′bipyridine. J. Am. Chem. Soc. 1989, 111, 8901−8911. (35) Lee, H. Y.; Wu, B. K.; Chern, M. Y. Study on the formation of zinc peroxide on zinc oxide with hydrogen peroxide treatment using xray photoelectron spectroscopy (XPS). Electron. Mater. Lett. 2014, 10, 51−55. (36) Sheng, H.; Li, Q.; Ma, W.; Ji, H.; Chen, C.; Zhao, J. Photocatalytic degradation of organic pollutants on surface anionized TiO2: Common effect of anions for high hole-availability by water. Appl. Catal., B 2013, 138−139, 212−218. (37) Xu, Y.; Lv, K.; Xiong, Z.; Leng, W.; Du, W.; Liu, D.; Xue, X. Rate enhancement and rate inhibition of phenol degradation over irradiated Anatase and Rutile TiO2 on the addition of NaF: New Insight into the Mechanism. J. Phys. Chem. C 2007, 111, 19024− 19032. (38) Klausen, J.; Troeber, S. P.; Haderlein, S. B.; Schwarzenbach, R. P. Reduction of substituted nitrobenzenes by Fe(II) in aqueous mineral Suspensions. Environ. Sci. Technol. 1995, 29, 2396−2404. (39) Liger, E.; Charlet, L.; Van Cappellen, P. Surface catalysis of uranium (VI) reduction by Iron (II). Geochim. Cosmochim. Acta 1999, 63, 2939−2955. (40) Hou, X. J.; Huang, X. P.; Ai, Z. H.; Zhao, J. C.; Zhang, L. Z. Ascorbic acid/Fe@Fe2O3: A highly efficient combined Fenton reagent to remove organic contaminants. J. Hazard. Mater. 2016, 310, 170− 178. (41) Wei, J.; Feng, Y.; Sun, X.; Liu, J.; Zhu, L. Effectiveness and pathways of electrochemical degradation of pretilachlor herbicides. J. Hazard. Mater. 2011, 189, 84−91. (42) Moza, P. N.; Hustert, K.; Pal, S.; Sukul, P. Photocatalytic decomposition of pendimethalin and alachlor. Chemosphere 1992, 25, 1675−1682.

5126

DOI: 10.1021/acs.est.6b05906 Environ. Sci. Technol. 2017, 51, 5118−5126