Ascorbate-Promoted Surface Iron Cycle for ... - ACS Publications

Subscriber access provided by UB + Fachbibliothek Chemie | (FU-Bibliothekssystem)

Article

Ascorbate Promoted Surface Iron Cycle for Efficient Heterogeneous Fenton Alachlor Degradation with Hematite Nanocrystals Xiaopeng Huang, Xiaojing Hou, Falong Jia, Fahui Song, Jincai Zhao, and Lizhi Zhang ACS Appl. Mater. Interfaces, Just Accepted Manuscript • DOI: 10.1021/acsami.6b16600 • Publication Date (Web): 27 Feb 2017 Downloaded from http://pubs.acs.org on February 28, 2017

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a free service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are accessible to all readers and citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

ACS Applied Materials & Interfaces is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ACS Applied Materials & Interfaces

Ascorbate Promoted Surface Iron Cycle for Efficient Heterogeneous Fenton Alachlor Degradation with Hematite Nanocrystals Xiaopeng Huang, Xiaojing Hou, Falong Jia, Fahui Song, Jincai Zhao, and Lizhi Zhang*

Key Laboratory of Pesticide & Chemical Biology of Ministry of Education, Institute of Environmental Chemistry, College of Chemistry, Central China Normal University, Wuhan 430079, People’s Republic of China

Keywords: Surface Iron Cycle, Heterogeneous Fenton Oxidation, Alachlor Degradation, Ascorbate Ions, Hematite Nanocrystals

*

To whom correspondence should be addressed. E-mail: [email protected]. Phone/Fax: +86-27-6786 7535 1

ACS Paragon Plus Environment


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

ABSTRACT This study reports the H2O2 activation with different hematite nanocrystals and ascorbate ions for the herbicide alachlor degradation at pH 5. We found that hematite nanoplates (HNPs) exposed with {001} facets exhibited better catalytic performance than hematite nanocubes (HNCs) exposed with {012} facets, which was attributed to the formation of inner−sphere iron−ascorbate complexes on the hematite facets. The 3−fold undercoordination Fe cations of {001} facet favors the formation of inner−sphere iron−ascorbate complexes, while the 5−fold undercoordination Fe cations of {012} facet has stereo−hindrance effect, disfavoring the complex formation. The surface area normalized alachlor degradation rate constant (23.3 × 10−4 min−1 L m−2) of HNPs−ascorbate Fenton system was about 2.6 times that (9.1 × 10−4 min−1 L m−2) of HNCs−ascorbate counterpart. Meanwhile, the 89.0% of dechlorination and 30.0% of denitrification in the HNPs−ascorbate Fenton system were also significantly higher than those (60.9% and 13.1%) of the HNCs−ascorbate one. More importantly, the reductive dissolution of hematite by ascorbate was strongly coupled with the subsequent H2O2 decomposition by surface bound ferrous ions through surface iron cycle on the hematite facets in the hematite−ascorbate Fenton systems. This coupling could significantly inhibit the conversion of surface bound ferrous ions to dissolved ones, and thus account for the stability of hematite nanocrystals. This work sheds light on the internal relationship between iron geochemical cycling and contaminants degradation, and also inspires us to utilize surface iron cycle of widely existent hematite for environmental remediation.

1. Introduction Fenton reaction is an efficient and environmentally friendly method for organic pollutants oxidation.1, 2

While, drawbacks such as low working pH and iron sludge restrict its wide applications. For 2


Page 2 of 26

Page 3 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


decades, great efforts have been made to design heterogeneous Fenton systems to overcome these shortcomings. For example, iron oxides have proven to be highly efficient to prevent ferric sludge formation and broaden working pH range.3-6 Among various iron oxide minerals, ubiquitous hematite is the most thermodynamically stable and environmentally benign one,7 thus a promising heterogeneous Fenton catalyst. Nevertheless, its specific surface area normalized hydroxyl radical generation rate constant of H2O2 decomposition was as low as 8 × 10−9 g s−1 m−2.8 Therefore, it is of great environmental significance to improve the H2O2 decomposition efficiency of hematite. Ascorbate is an antioxidant and commonly used for agricultural and food industry applications. The interaction of ascorbate ions and hematite can result in the reductive dissolution of hematite.9, 10 This reductive dissolution process of iron−bearing minerals enormously contributes to the iron geochemical cycling, and also strongly influences the elements geochemical cycling and the pollutants transformation.11 Hence, the combination of ascorbate and iron oxides was extensively utilized to remove environmental contaminates. For instance, ascorbate and iron oxides could efficiently reduce carbon tetrachloride.12 Meanwhile, ascorbate ions could accelerate iron cycling to enhance the contaminants degradation in homogeneous Fenton and Fenton–like systems.13-15 However, the interaction of ascorbate ions with well−defined hematite nanocrystals is not well understood, which is adverse to design highly efficient heterogeneous Fenton systems with hematite for the degradation of contaminants. It is widely accepted that the reactivity and catalytic activity of hematite might be strongly affected by its facet effects. Nevertheless, the detailed effects of different hematite facets on the surface iron cycle and the H2O2 decomposition as well as the pollutant degradation are seldom reported previously. In this study, we choose {001} and {012} facets of hematite nanocrystals to 3



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

investigate hematite surface iron cycle in the presence of ascorbate ions during the heterogonous Fenton processes, regarding that {001} and {012} facets are two dominant exposed surfaces of natural hematite mineral and possess relatively low surface energies,16-19 as well as are extensively studied in the fields of catalysis,20, 21 geochemistry,22, 23 and environmental science.24, 25 The binding mode of ascorbate ions on the surface of hematite nanocrystals are systematically investigated to clarify the effects of iron−ascorbate complexes on the H2O2 decomposition and the subsequent alachlor degradation. On the basis of experimental and theoretical results, we intend to elucidate the crucial roles of surface iron cycle for heterogeneous Fenton processes.

2. EXPERIMENTAL SECTION 2.1 Chemicals and Materials. Iron(III) chloride hexahydrate, H2O2 (30%), and sodium ascorbate were obtained from Sinopharm Chemical Reagent Co. Ltd., China. Alachlor was obtained from Sigma-Aldrich. Methanol and acetonitrile were of high-performance liquid chromatography grade and purchased from Fisher Scientific. All the other reagents were of analytical grade without purification. All the solutions were prepared with deionized (DI) water. All the stock solutions were freshly prepared before use. 2.2 Sample Preparation. The hematite nanoplates (HNPs) were prepared by a previously reported methods.26 Briefly, 1.09 g of iron(III) chloride hexahydrate and 3.20 g of sodium acetate were subsequently dissolved under magnetic stirring in the solution composed of ethanol (40.0 mL) and DI water (2.8 mL). The solution was then kept in a Teflon−lined stainless steel autoclave and treated in an electric oven at 180 oC for 12 h. Upon cooling to room temperature, the solid precipitate was collected and washed with DI water and ethanol several times before drying at 40 oC overnight. 4


Page 4 of 26

Page 5 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


According to previous methods to prepare hematite nanocubes (HNCs),27 2.08 g of iron(III) chloride hexahydrate and 6.93 g of sodium oleate were subsequently added to the solution composed of ethanol (35 mL) and oleic acid (4.3 mL) in a Teflon−lined stainless autoclave. The autoclave was sealed and treated at 180 oC for 12 h and then cooled to room temperature. The precipitate was collected by centrifugation and washed with cyclohexane, water and ethanol before drying at 40 oC overnight. 2.3 Materials Characterization. The powder X−ray diffraction (XRD, D8 Advance, Bruker) was performed on a D/Max−IIIA X−ray diffractometer, using a Cu Kα source (λ = 0.15418 nm). The morphologies and structures of products were examined by scanning electron microscopy (SEM, 6700-F, JEOL) and high−resolution transmission electron microscopy (HRTEM, JSM−2010, JEOL). The Brunauer−Emmett−Teller (BET) surface areas of two hematite samples were obtained using a Micromeritics Tristar 3000 instrument. 2.4 Degradation Experiments. Batch degradation experiments were conducted in 100 mL flasks which were covered with foil under argon gas at room temperature. The initial solution pH was adjusted to 5 by dilute sulfuric acid and sodium hydroxide aqueous solutions. The initial concentrations of sodium ascorbate and H2O2 aqueous solutions were 1.0 × 10−3 and 4.0 × 10−4 mol/L, respectively; the dosage of hematite nanocrystals was 0.4 g/L. Samples were taken out at predetermined time intervals from the flasks with a syringe and filtered with a membrane (0.22 µm). 2.5 Analytical Methods. The H2O2 concentration was determined according to a modified p−hydroxyphenylacetic acid fluorescent method.28 The benzoic acid and alachlor concentrations were analyzed by high performance liquid chromatography (HPLC, LC-20A, Shimadzu) with a Shimadzu SB−C18 reverse phase column. The alachlor degradation intermediates were determined 5



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 6 of 26

by gas chromatography−mass spectrometry (GC−MS, Trace 1300 equipped with ISQ, Thermo) with a DB−5 column and high performance liquid chromatography mass spectrometry (LC−MS, TSQ Quantum Access MAX, Thermo).29 The dissolved Fe(II) and total iron ions concentrations were determined colorimetrically with a previously reported 1,10−phenanthroline method.30,

31

The

generated surface bound Fe(II) of hematite were monitored by the hydrochloric extraction technique.31 The concentration of ascorbate ions was determined with a modified 2, 2’–bipyridine method.15, 32 An Orion model PHS−25m pH meter was employed to determine the solution pH. A JES FA 200 X−band spectrometer (JEOL) was used to obtain 5,5−dimethyl−1−pyrroline−N−oxide (DMPO)−trapped EPR spectra. The anions were analyzed with an ion chromatograph (IC, Dionex ICS−900, Thermo) of an AS23 column.

3. RESULTS AND DISCUSSION 3.1 Characterization of the As−Prepared Hematite Nanocrystals. The XRD results of as−synthesized samples coincided well with the standard XRD patterns of hematite (JCPDS card No. 33−0664), revealing their phase purity (Figure S1). The BET−specific surface areas of HNPs and HNCs were 21.3 and 20.4 m2 g-1 respectively, determined by their N2 adsorption and desorption isotherm curves (Figure S2). The SEM and TEM images demonstrated that the width and thickness of HNPs were about 82.2 and 13.1 nm, respectively (Figures 1a−1c). The representative HRTEM image (Figure 1e) and the fast Fourier transforms (FFT) pattern (Figure 1d) showed the lattice fringes of 0.25 nm, which agreed well with (−120), (−210) and (110) planes, respectively.24, 33 Hence, {001} was the dominant exposed facet of HNPs. The SEM and TEM images revealed that HNCs appeared to be uniform cubes with a mean size of 33.1 nm (Figures 1f−1h). The HRTEM images and 6


Page 7 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


FFT pattern of HNCs showed their lattice fringes of 0.37 nm (Figure 1i and 1j). The high magnification TEM images revealed that HNCs were actually pseudocubes with identically parallelepiped planes and the dihedral angles between lateral adjacent facets were 94° or 86° (Figure 1h). These structural features demonstrated that HNCs were completely exposed with {012} facets.34 Furthermore, the particle size distribution data of hematite nanocrystals confirmed the homogeneous particle size distribution of HNPs and HNCs (Table S1), which were consistent with the SEM and TEM characterization results. 3.2 The H2O2 Decomposition with Hematite−Ascorbate Catalysts. The as−prepared hematite nanocrystals and sodium ascorbate were used to decompose H2O2. Before evaluating the H2O2 decomposition efficiencies, we compared the concentrations of dissolved Fe(II) ([Fe(II)]dissolved) and total Fe(II) ([Total Fe(II)]) change as a function of time in the presence of ascorbate and hematite nanocrystals with/without H2O2 (Figure 2). Herein, the total Fe(II) concentration was the sum of dissolved and accumulated surface bound Fe(II) concentrations ([Total Fe(II)] = [Fe(II)]dissolved + [Fe(II)]surface). Although ascorbate could induce the reductive dissolution of hematite in the absence of H2O2, we interestingly found that the addition of H2O2 strongly inhibited the dissolved ferrous ions release (Figure 2a) and the total iron reduction (Figure 2b) of hematite in the two hematite−ascorbate Fenton systems. Meanwhile, the concentrations of dissolved ferrous ions increased within first 60 min and then slightly decreased along with time. These phenomena suggested that the reductive dissolution of hematite facets by ascorbate might be strongly coupled with the subsequent H2O2 decomposition by surface bound ferrous ions in the hematite−ascorbate Fenton system. This coupling could significantly inhibit the conversion of surface bound ferrous ions to dissolved ones, and thus account for the negligible reductive dissolution of hematite nanocrystals 7



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

after the addition of H2O2. Then, we could evaluate the H2O2 decomposition efficiencies of different hematite−ascorbate systems at pH 5. It was found that the pH did not change during the whole Fenton reaction (Figure S3). The H2O2 concentration was determined according to a modified p−hydroxyphenylacetic acid fluorescent method.28 From the curves of H2O2 concentration change as a function of time, we found that H2O2 could be effectively decomposed in the hematite−ascorbate systems (Figure 3a), and the corresponding kinetic curves obeyed pseudo−first−kinetic model (Figure 3b). The apparent H2O2 decomposition rate constants (k) were calculated to be 3.0 × 10−3 and 1.9 × 10−3 min−1 for the HNPs−ascorbate and HNCs−ascorbate systems, respectively, revealing that hematite nanoplates with exposed {001} facets exhibited better catalytic performance than hematite nanocubes with exposed {012} facets. The different catalytic decomposition efficiencies of H2O2 with hematite−ascorbate catalysts might be attributed to the binding modes of inner−sphere iron−ascorbate complexes formed on the hematite facets. To validate this assumption, we then compared the binding modes of ascorbate coordination on the two hematite facets (Figure 4). The {001} facets can form O− or Fe−terminated surfaces (Figure 4a), while the previous experimental and theoretical results demonstrated that Fe−terminated form was the stable one.35-37 The Fe terminated hematite {001} facets have 3−fold undercoordination (Fe3c) surface sites (Figure 4a) which could be coordinated with ascorbate ions to form inner−sphere iron−ascorbate complexes on the {001} facets. Differently, Fe−terminated form with five−fold undercoordinated Fe (Fe5c) sites are stable on hematite {012} facets (Figure 4b), in which the Fe5c sites could be coordinated with ascorbate ions to form inner−sphere iron−ascorbate coordination on the {012} facets. The Fe5c sites of the {012} facets were not favorable for ascorbate binding because higher number of undercoordination iron cations could bring a relatively strong 8


Page 8 of 26

Page 9 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


stereo−hindrance effect to disfavor the ascorbate complexes formation. Therefore, the {001} facets of 3−fold undercoordination Fe cations are more suitable for the formation of inner−sphere iron−ascorbate complexes than the {012} facets of 5−fold undercoordination Fe cations with stereo−hindrance effect. The undercoordinated Fe cations on the hematite surfaces can determine the ascorbate ions adsorption and complexation.38 When the hematite facets contacted with the ascorbate aqueous solution, the inner−sphere iron−ascorbate complexes would be formed on the hematite surfaces. Simultaneously, the electrons would be transferred from ascorbate to undercoordinated Fe cations (≡FeIIIOH2) within the surface complexes to produce surface bound ferrous (≡FeIIOH2), which would be attacked by H2O2 to form the intermediate complexes (≡FeII(H2O2)*), preventing ferrous ions from leaching into the solution. Owing to the strong reducibility of ≡FeIIOH2, ≡FeIIOH2 would donate its electrons to H2O2, producing hydroxyl radicals and surface ferric ions (≡FeIIIOH2) to realize surface iron cycle on the hematite facets. The accumuated hydroxyl radicals could degrade organic pollutants into CO2 and H2O finally, while the surface ferric ions (≡FeIIIOH2) would react with ascorbate to trigger another cycle of H2O2 decomposition. The effects of surface iron−ascorbate complexes on the H2O2 decomposition were further investigated. The DMPO−trapped electron paramagnetic resonance (EPR) spectra were employed to determine the hydroxyl radicals (•OH) accumulation. The DMPO−trapped EPR spectra demonstrated that •OH was produced in the hematite−ascorbate Fenton systems (Figure 5d). Obviously, the HNPs−ascorbate Fenton system showed a higher production of •OH than the HNCs−ascorbate counterpart at the first 2 minutes. The steady−state concentration of •OH in the HNPs−ascorbate Fenton system was also much higher than that of HNCs−ascorbate, which could be quantified by a 9



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

probe method with benzoic acid (BA).8, 39 We subsequently assessed the •OH generation rate (V•OH) in the hematite−ascorbate Fenton systems. We found that BA could not be oxidized in the presence of H2O2 and hematite (HNPs/H2O2 and HNCs/H2O2), or H2O2 and ascorbate (Asc/H2O2), because neither hematite nor ascorbate could induce the H2O2 decomposition to degrade BA in this study (Figure 5a). Interestingly, BA could be efficiently oxidized in the presence of H2O2, HNPs and ascorbate (Asc/HNPs/H2O2), or H2O2, HNCs and ascorbate (Asc/HNCs/H2O2), respectively (Figure 5a). From the BA oxidation and the corresponding kinetic curves (Figures 5a and 5b), the oxidation rates (kp) of HNPs−ascorbate and HNCs−ascorbate Fenton systems were found to be 1.22 × 10−3 and 6.33 × 10−4 s−1, respectively. Because the concentration of benzoic acid was much higher than that of ascorbate ions, the ascorbate oxidation by •OH could be ruled out during the hydroxyl radical measurement.8, 39 Then, the •OH generation rates (V•OH) were 5.55 × 10−7 and 3.37 × 10−7 M s−2 for the HNPs−ascorbate and HNCs−ascorbate Fenton systems, respectively. The corresponding •OH formation rate constants (k•OH) were 2.08 × 10−3 and 1.31 × 10−3 s−1 for the HNPs−ascorbate and HNCs−ascorbate Fenton systems, respectively (Figure 5c). Meanwhile, their specific surface area normalized k•OH (k•OH′) were 9.95 × 10−5 and 6.42 × 10−5 g s−1 m−2, respectively. According to the definition of turnover frequency (TOF),40 the TOF values of HNPs−ascorbate and HNCs−ascorbate Fenton systems were calculated to be 6.70 × 10−3 and 2.72 × 10−3 s−1, respectively (Table S2). The above kp, V•OH, TOF, k•OH and k•OH′ of hematite−ascorbate Fenton systems were summarized in Table S2. Obviously, the iron−ascorbate complexes on the {001} facets of hematite could more efficiently catalyze the H2O2 decomposition to produce •OH radicals than the counterparts on the {012} facets. 3.3 Contaminants Degradation with Hematite−Ascorbate Fenton systems. As expected, various organic contaminants could be efficiently degraded in the presence of H2O2, ascorbate and 10


Page 10 of 26

Page 11 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


HNPs at an initial pH of 5 (Table S3). Regarding the hardly biodegradable characteristic of alachlor, a widely used chloroacetanilide herbicide,41-44 we chose alachlor as a model pollutant to evaluate the performances of two hematite−ascorbate Fenton systems in detail. The degradation curves of alachlor in the presence of ascorbate and hematite nanocrystals with different H2O2 concentrations revealed that the optimal concentration of H2O2 was 4.0 × 10−4 mol/L in this study (Figure S4). A comparative experiment revealed that alachlor could not be degraded in the systems of HNPs/H2O2, HNCs/H2O2, or Asc/H2O2. As the reductive dissolution of hematite facets by ascorbate was strongly coupled with the subsequent H2O2 decomposition by surface bound ferrous ions to produce plenty of hydroxyl radicals, alachlor could be significantly removed in the Asc/HNPs/H2O2 and Asc/HNCs/H2O2 systems (Figure 6a). The apparent alachlor degradation rate constants (kalachlor) of HNPs−ascorbate and HNCs−ascorbate Fenton systems were 19.5 × 10−3 and 7.4 × 10−3 min−1, respectively (Figure 6b). Therefore, the surface area normalized alachlor degradation rate constants (kalachlor′) could be obtained from the kalachlor, specific surface area (SSA) and concentration of hematite nanocrystals, where kalachlor′ = kalachlor / (SSA × catalyst concentration). The surface area normalized alachlor degradation rate constant (23.3 × 10−4 min−1 L m−2) of HNPs−ascorbate Fenton system was even about 2.6 times that (9.1 × 10−4 min−1 L m−2) of HNCs−ascorbate counterpart, further confirming that the iron−ascorbate complexes on the {001} facets of hematite exhibited better contaminants degradation performance than the counterparts on the {012} facets. During the alachlor degradation process, ascorbate was also consumed simultaneously (Figure 6d). It was interesting to find that two kinds of hematite nanocrystals were reusable even after 5 cycles (Figure 7a). From Figure 2, ascorbate could significantly induce the reductive dissolution of hematite in the systems of ascorbate and hematite nanocrystals, but the addition of H2O2 strongly inhibited the 11



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

dissolved ferrous ions release of hematite in the hematite−ascorbate Fenton systems. As expected, the Fe leaching ratios of hematite nanocrystals were less than 1% after five cycles of use (0.9% for HNPs, 0.5% for HNCs) in the hematite−ascorbate Fenton systems (Figure 7b). This inhibition of reductive dissolution of hematite by ascorbate after the addition of H2O2 might be attributed to the coupling between the reductive dissolution of hematite induced by ascorbate and the subsequent H2O2 decomposition induced by surface bound ferrous ions through surface iron cycle on the hematite facets in the hematite−ascorbate Fenton systems. Although the reductive dissolution of hematite by ascorbate might affect the crystal structures, chemical composition, and morphologies of hematite nanocrystals,45 we interestingly found that the XRD patterns and the morphologies of hematite nanocrystals did not significantly change along with the Fenton reaction (Figures S5 and S6), confirming their catalytic stability. The stability of hematite nanocrystals in the hematite−ascorbate Fenton systems could be also explained as follows. Surface iron cycle on the hematite facets triggered a tight coupling between the reductive dissolution of hematite induced by ascorbate and the subsequent H2O2 decomposition induced by surface bound ferrous ions in the hematite−ascorbate Fenton systems. This coupling could significantly inhibit the conversion of surface bound ferrous ions to dissolved ones, and thus result in the stability of hematite nanocrystals. The alachlor mineralization in the hematite−ascorbate Fenton systems was also investigated by ion chromatography. The concentrations of Cl− ions after 5 hours' HNP−ascorbate and HNC−ascorbate Fenton degradation of alachlor were 2.34 and 1.60 ppm, respectively. As the theoretical concentration of Cl− was 2.63 ppm when all the chlorine of alachlor was converted into Cl− species, the HNPs−ascorbate and HNCs−ascorbate Fenton systems could respectively convert 89.0% and 12


Page 12 of 26

Page 13 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


60.9% of chlorine in alachlor to Cl− ions in 5 hours (Figure 6c). Simultaneously, 1.38 ppm and 0.60 ppm of NO3− were generated after 5 hours' alachlor degradation in the HNP−ascorbate and HNCs−ascorbate Fenton systems. Hence, the nitrogen removal ratios of alachlor in HNPs−ascorbate and HNCs−ascorbate Fenton systems were 30.0% and 13.1%, respectively (Figure 6c). The possible degradation intermediates of ascorbate and alachlor were analyzed by gas chromatograph−mass spectroscopy and high performance liquid chromatography mass spectrometry (Figures S7 and S8). The ascorbate degradation intermediates included dehydroascorbic acid, 2,3−diketogulonic acid, and 4,5,5,6−tetrahydroxy−2,3−diketohexanoic acid (Figure S7). The alachlor

degradation

intermediates

included

2,6−diethyl−N−(methoxymethyl)−benzenamine,

2−chloro−N−(2,6−diethylphenyl) 2,6−diethylbenzenamine,

acetamide,

1,3−diethylbenzene,

1,3−diethyl−2−nitrosobenzene, 3−ethylbenzaldehyde, N−(2−chloroethyl)−2,6−diethylbenzenamine, and 4−amino−3,5−diethylphenol, confirming that alachlor was oxidatively degraded (Figure S8 and Table S4), which could be involved in the first dealkylation and dechlorination, and the subsequent alkylic−oxidation and hydroxylation, as well as the final mineralization. On the basis of the above results, a possible alachlor degradation pathway in the hematite−ascorbate Fenton systems was proposed (Figure 8).

4. CONCLUSIONS In conclusion, the H2O2 decomposition with different hematite nanocrystals in the presence of ascorbate ions was thoroughly studied to clarify the critical roles of surface iron cycle in heterogeneous Fenton degradation process. We found that hematite nanoplates exposed with {001} facets exhibited better catalytic performance than hematite nanocubes exposed with {012} facets, 13



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

which was attributed to the different binding modes of inner−sphere iron−ascorbate complexes formed on the hematite facets. The 3−fold undercoordination Fe cations of {001} facet favors the formation of inner−sphere iron−ascorbate complexes, while 5−fold undercoordination Fe cations of {012} facet has stereo−hindrance effect, disfavoring the inner−sphere iron−ascorbate formation. The surface area normalized alachlor degradation rate constant (23.3 × 10−4 min−1 L m−2) of HNPs−ascorbate Fenton system was even about 2.6 times that (9.1 × 10−4 min−1 L m−2) of HNCs−ascorbate counterpart, along with 89.0% and 60.9% of dechlorination as well as 30.0% and 13.1% of denitrification for HNPs−ascorbate and HNCs−ascorbate Fenton systems, respectively. Meanwhile, the dealkylation, dechlorination, alkylic−oxidation, hydroxylation and mineralization could occur during the hematite−ascorbate Fenton alachlor oxidation processes. Meanwhile, the reductive dissolution of hematite by ascorbate was strongly coupled with the subsequent H2O2 decomposition by surface bound ferrous ions through surface iron cycle on the hematite facets in the hematite−ascorbate Fenton systems. This coupling could significantly inhibit the conversion of surface bound ferrous ions to dissolved ones, and thus account for the stability of hematite nanocrystals. This study sheds light on the internal relationship between iron geochemical cycling and contaminants degradation, and also enlightens us to utilize the surface iron cycle of the widely existent iron oxide minerals for the environmental remediation in the future.

ASSOCIATED CONTENT Supporting Information The powder XRD patterns; The N2 adsorption and desorption isotherm curves; The pH change as a function of time; The optimal of H2O2 concentrations; XRD patterns after the reaction; SEM images 14


Page 14 of 26

Page 15 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


before and after the reaction; The ascorbate degradation intermediates; The mass spectra of alachlor degradation intermediates; The particle size distribution data; A comparison of kp, V•OH, TOF, k•OH, and k•OH′; Degradation of various contaminants; Identification of major intermediate products derived from alachlor. This material is available free of charge via the Internet at http://pubs.acs.org.

ACKNOWLEDGEMENTS This work was supported by Natural Science Funds for Distinguished Young Scholars (Grant 21425728), National Science Foundation of China (Grant 21677059), 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 2015YBZD013 and 2016YBZZ036), and the CAS Interdisciplinary Innovation Team of the Chinese Academy of Sciences.

REFERENCES (1) Pignatello, J. J.; Oliveros, E.; MacKay, A. Advanced Oxidation Processes for Organic Contaminant Destruction Based on the Fenton Reaction and Related Chemistry. Crit. Rev. Environ. Sci. Technol. 2006, 36, 1-84. (2) Brillas, E.; Sires, I.; Oturan, M. A. Electro-Fenton Process and Related Electrochemical Technologies Based on Fenton's Reaction Chemistry. Chem. Rev. 2009, 109, 6570-6631. (3) Huang, H.; Chen, J.; Lu, M. Catalytic Decomposition of Hydrogen Peroxide and 2-Chlorophenol with Iron Oxides. Water Res. 2001, 35, 2291-2299. (4) Pham, A. L.; Lee, C.; Doyle, F. M.; Sedlak, D. L. A Silica-Supported Iron Oxide Catalyst Capable 15



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

of Activating Hydrogen Peroxide at Neutral pH Values. Environ. Sci. Technol. 2009, 43, 8930-8935. (5) Lin, S.-S.; Gurol, M. D. Catalytic Decomposition of Hydrogen Peroxide on Iron Oxide: Kinetics, Mechanism, and Implications. Environ. Sci. Technol. 1998, 32, 1417-1423. (6) Wang, L.; Yao, Y.; Sun, L.; Lu, W.; Chen, W. Activation of Hydrogen Peroxide by Activated Carbon Fibers Coupled with Fe(III)-Citrate for Degradation of Dyes at Neutral pH. Acta Chim. Sinica 2013, 71, 1633-1638. (7) Tartaj, P.; Morales, M. P.; Gonzalez-Carreño, T.; Veintemillas-Verdaguer, S.; Serna, C. J. The Iron Oxides Strike Back: From Biomedical Applications to Energy Storage Devices and Photoelectrochemical Water Splitting. Adv. Mater. 2011, 23, 5243-5249. (8) 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. (9) Echigo, T.; Aruguete, D. M.; Murayama, M.; Hochella, M. F. Influence of Size, Morphology, Surface Structure, and Aggregation State on Reductive Dissolution of Hematite Nanoparticles with Ascorbic Acid. Geochim. Cosmochim. Acta 2012, 90, 149-162. (10) Huang, X. P. ; Hou, X. J.; Song, F. H.; Zhao, J. C.; Zhang, L. Z. Ascorbate Induced Facet Dependent Reductive Dissolution of Hematite Nanocrystals. J. Phys. Chem. C 2017, 121, 1113-1121. (11) Borch, T.; Kretzschmar, R.; Kappler, A.; Cappellen, P. V.; Ginder-Vogel, M.; Voegelin, A.; Campbell, K. Biogeochemical Redox Processes and Their Impact on Contaminant Dynamics. Environ. Sci. Technol. 2010, 44, 15-23. (12) Lin, Y.-T.; Liang, C. Carbon Tetrachloride Degradation by Alkaline Ascorbic Acid Solution. Environ. Sci. Technol. 2013, 47, 3299-3307. (13) Zhang, Z.; Wang, X.; Zhao, M.; Qi, H. Free-Radical Degradation by Fe2+/Vc/H2O2 and 16


Page 16 of 26

Page 17 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Antioxidant Activity of Polysaccharide from Tremella Fuciformis. Carbohydr. Polym. 2014, 112, 578-582. (14) Bolobajev, J.; Trapido, M.; Goi, A. Improvement in Iron Activation Ability of Alachlor Fenton-Like Oxidation by Ascorbic Acid. Chem. Eng. J. 2015, 281, 566-574. (15) 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. (16) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrences and Uses, 2nd Ed. Wiley-VCH: Weinheim: 2003. (17) Mackrodt, W. Atomistic Simulation of Oxide Surfaces. Phys. Chem. Miner. 1988, 15, 228-237. (18) Hartman, P. The Effect of Surface Relaxation on Crystal Habit: Cases of Corundum (α-Al2O3) and Hematite (α-Fe2O3). J. Cryst. Growth 1989, 96, 667-672. (19) Mackrodt, W.; Davey, R.; Black, S.; Docherty, R. The Morphology of α-Al2O3 and α-Fe2O3: The Importance of Surface Relaxation. J. Cryst. Growth 1987, 80, 441-446. (20) Qiao, B.; Wang, A.; Yang, X.; Allard, L. F.; Jiang, Z.; Cui, Y.; Liu, J.; Li, J.; Zhang, T. Single-Atom Catalysis of CO Oxidation Using Pt1/FeOx. Nat. Chem. 2011, 3, 634-641. (21) Zhou, X.; Lan, J.; Liu, G.; Deng, K.; Yang, Y.; Nie, G.; Yu, J.; Zhi, L. Facet-Mediated Photodegradation of Organic Dye over Hematite Architectures by Visible Light. Angew. Chem., Int. Ed. 2012, 124, 182-186. (22) Yanina, S. V.; Rosso, K. M. Linked Reactivity at Mineral-Water Interfaces through Bulk Crystal Conduction. Science 2008, 320, 218-222. (23) Catalano, J. G.; Fenter, P.; Park, C.; Zhang, Z.; Rosso, K. M. Structure and Oxidation State of 17



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Hematite Surfaces Reacted with Aqueous Fe(II) at Acidic and Neutral pH. Geochim. Cosmochim. Acta 2010, 74, 1498-1512. (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. (25) Henderson, M. A. Surface Chemistry of Trimethyl Phosphate on α-Fe2O3. J. Phys. Chem. C 2011, 115, 23527-23534. (26) Chen, L.; Yang, X.; Chen, J.; Liu, J.; Wu, H.; Zhan, H.; Liang, C.; Wu, M. Continuous Shapeand Spectroscopy-Tuning of Hematite Nanocrystals. Inorg. Chem. 2010, 49, 8411-8420. (27) Wang, S. B.; Min, Y. L.; Yu, S. H. Synthesis and Magnetic Properties of Uniform Hematite Nanocubes. J. Phys. Chem. C 2007, 111, 3551-3554. (28) Tanner, P. A.; Wong, A. Y. S. Spectrophotometric Determination of Hydrogen Peroxide in Rainwater. Anal. Chim. Acta 1998, 370, 279-287. (29) 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. (30) Harvey Jr, A. E.; Smart, J. A.; Amis, E. Simultaneous Spectrophotometric Determination of Iron(II) and Total Iron with 1, 10-Phenanthroline. Anal. Chem. 1955, 27, 26-29. (31) 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. (32) Burman, D. M.; Shertzer, H. G.; Senft, A. P.; Dalton, T. P.; Genter, M. B. Antioxidant Perturbations in the Olfactory Mucosa of Alachlor-Treated Rats. Biochem. Pharmacol. 2003, 66, 1707-1715. (33) Huang, X. P.; Hou, X. J.; Zhao, J. C; Zhang, L. Z. Hematite Facet Confined Ferrous Ions as 18


Page 18 of 26

Page 19 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


High Efficient Fenton Catalysts to Degrade Organic Contaminants by Lowering H2O2 Decomposition Energetic Span. Appl. Catal., B 2016, 181, 127-137. (34) Ouyang, J.; Pei, J.; Kuang, Q.; Xie, Z.; Zheng, L. Supersaturation-Controlled Shape Evolution of α-Fe2o3 Nanocrystals and Their Facet-Dependent Catalytic and Sensing Properties. ACS Appl. Mater. Interfaces 2014, 6, 12505-12514. (35) Wasserman, E.; Rustad, J. R.; Felmy, A. R.; Hay, B. P.; Halley, J. W. Ewald Methods for Polarizable Surfaces with Application to Hydroxylation and Hydrogen Bonding on the (012) and (001) Surfaces of α-Fe2O3. Surf. Sci. 1997, 385, 217-239. (36) Lübbe, M.; Moritz, W. A Leed Analysis of the Clean Surfaces of α-Fe2O3 (0001) and α-Cr2O3 (0001) Bulk Single Crystals. J. Phys.: Condens. Matter 2009, 21, 134010-134010. (37) Wang, X. G.; Weiss, W.; Shaikhutdinov, S. K.; Ritter, M.; Petersen, M.; Wagner, F.; Schl; Ouml, R.; Scheffler, M. The Hematite (α-Fe2O3) (0001) Surface: Evidence for Domains of Distinct Chemistry. Phys. Rev. Lett. 1998, 81, 1038-1041. (38) 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. (39) Lindsey, M. E.; Tarr, M. A. Quantitation of Hydroxyl Radical During Fenton Oxidation Following a Single Addition of Iron and Peroxide. Chemosphere 2000, 41, 409-417. (40) Boudart, M. Turnover Rates in Heterogeneous Catalysis. Chem. Rev. 1995, 95, 661-666. (41) Gan, J.; Wang, Q.; Yates, S. R.; Koskinen, W. C.; Jury, W. A. Dechlorination of Chloroacetanilide Herbicides by Thiosulfate Salts. Proc. Natl. Acad. Sci. U.S.A. 2002, 99, 5189-5194. 19



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

(42) Qiang, Z.; Liu, C.; Dong, B.; Zhang, Y. Degradation Mechanism of Alachlor During Direct Ozonation and O3/H2O2 Advanced Oxidation Process. Chemosphere 2010, 78, 517-526. (43) Xin, Y.; Liu, H.; Han, L.; Zhou, Y. Comparative Study of Photocatalytic and Photoelectrocatalytic Properties of Alachlor Using Different Morphology TiO2/Ti Photoelectrodes. J. Hazard. Mater. 2011, 192, 1812-1818. (44) Wang, X.; Yong, Z. Degradation of Alachlor in Aqueous Solution by Using Hydrodynamic Cavitation. J. Hazard. Mater. 2009, 161, 202-207. (45) Tan, W.-F.; Yu, Y.-T.; Wang, M.-X.; Liu, F.; Koopal, L. K. Shape Evolution Synthesis of Monodisperse Spherical, Ellipsoidal, and Elongated Hematite (α-Fe2O3) Nanoparticles Using Ascorbic Acid. Cryst. Growth Des. 2014, 14, 157-164.

20


Page 20 of 26

Page 21 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure Captions

Figure 1. Representative morphologies and structures of HNPs and HNCs. (a) SEM image, (b) TEM image, (c) a single nanoplate, (d) FFT pattern and (e) HRTEM image of HNPs. (f) SEM image, (g) TEM image, (h) a single nanocube, (i) FFT pattern and (j) HRTEM image of HNCs.

Figure 2. A comparison of (a) dissolved Fe(II) and (b) total Fe(II) change as a function of time in the presence of ascorbate and hematite nanocrystals with/without H2O2. The initial ascorbate ions concentrations were 1.0 × 10-3 mol/L; the initial H2O2 concentration was 4.0 × 10-4 mol/L; the dosage of hematite nanocrystals was 0.4 g/L; the initial pH was 5. 21



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 3. (a) A comparison of change in H2O2 concentration as a function of time. (b) The corresponding pseudo-first kinetic curves. The initial ascorbate ions concentrations were 1.0 × 10-3 mol/L; the initial H2O2 concentration was 4.0 × 10-4 mol/L; the dosage of hematite nanocrystals was 0.4 g/L; the initial pH was 5.

Figure 4. Atomic arrangement of hematite exposed {001} and {012} facet. (a) {001} and (b) {012} facet. Exposed iron cations represent undercoordination iron cations on the surfaces.

22


Page 22 of 26

Page 23 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 5. (a) The oxidation of benzoic acid in different systems. (b) The kinetics curves of benzoic acid oxidation in the presence of H2O2, ascorbate and hematite nanocrystals. (c) A comparison of the •OH formation rates as a function of H2O2 concentration in the presence of H2O2, ascorbate and hematite nanocrystals. (d) DMPO−trapped EPR at 2 minutes in the presence of H2O2, ascorbate and hematite nanocrystals.

23



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Figure 6. (a) Time profile of alachlor degradation in different systems. (b) The degradation kinetics curves of alachlor degradation in the presence of H2O2, ascorbate and hematite nanocrystals. (c) A comparison of the nitrate and chlorine ions formation in the presence of H2O2, ascorbate and hematite nanocrystals. (d) Change in ascorbate concentration as a function of time in the presence of H2O2, ascorbate and hematite nanocrystals. The initial concentrations of alachlor, H2O2 and ascorbate were 20 mg/L, 4.0 × 10−4 mol/L and 1.0 × 10−3 mol/L; the dosage of hematite nanocrystals was 0.4 g/L; the initial pH was 5.

24


Page 24 of 26

Page 25 of 26

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 7. (a) The reusability of hematite nanocrystals for the Fenton alachlor degradation. (b) Fe leaching of hematite nanocrystals during five cycles of Fenton alachlor degradation.

Figure 8. The possible alachlor degradation pathway in the presence of H2O2, ascorbate and hematite nanocrystals.

25



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

TOC Art

26


Page 26 of 26

Ascorbate-Promoted Surface Iron Cycle for ... - ACS Publications

Recommend Documents