Synthesis of Mesoporous Iron Oxides by an Inverse Micelle Method

Article pubs.acs.org/JPCC

Synthesis of Mesoporous Iron Oxides by an Inverse Micelle Method and Their Application in the Degradation of Orange II under Visible Light at Neutral pH Ting Jiang,† Altug S. Poyraz,‡ Aparna Iyer,‡,§ Yashan Zhang,‡ Zhu Luo,∥ Wei Zhong,∥ Ran Miao,‡ Abdelhamid M. El-Sawy,‡ Curtis J. Guild,‡ Yu Sun,∥ David A. Kriz,‡ and Steven L. Suib*,†,‡,∥ †

Department of Chemical and Biomolecular Engineering, University of Connecticut, U-3222, 191 Auditorium Road, Storrs, Connecticut 06269-3222, United States ‡ Department of Chemistry, University of Connecticut, U-3060, 55 North Eagleville Road, Storrs, Connecticut 06269-3060, United States § Department of Chemical Engineering and Material Science, University of Minnesota, 421 Washington Ave. SE, Minneapolis, Minnesota 55455-0132, United States ∥ Institute of Materials Science, University of Connecticut, U-3060, 97 North Eagleville Road, Storrs, Connecticut 06269-3136, United States S Supporting Information *

ABSTRACT: Mesoporous iron oxides (2-line ferrihydrite, α-Fe2O3, γ-Fe2O3, and Fe3O4) are successfully synthesized by modifying the reaction temperatures and calcination atmospheres of the sol−gel-based inverse micelle method. Different characterization techniques, such as PXRD, N2 sorption, SEM, HRTEM, Raman spectroscopy, and XANES, are performed to determine the properties of the catalysts. Larger pore sizes can be obtained in mesoporous γ-Fe2O3 and Fe3O4 compared with mesoporous 2-line ferrihydrite and α-Fe2O3. The catalytic performance of mesoporous iron oxides are examined as Fenton catalysts in orange II degradation in the presence of oxidant H2O2 at neutral pH under visible light. Adsorption capacities of mesoporous iron oxides on orange II are greater than that of commercial Fe2O3. The greatest adsorption capacity is found to be 49.3 mg/g with mesoporous 2-line ferrihydrite. In addition, the degradation efficiency of orange II is found to be markedly improved by mesoporous iron oxides compared with the commercial catalyst. In the best case scenario, 2-line ferrihydrite shows the highest degradation rate constant (0.0258 min−1) among all the catalysts tested. The excellent performance of 2-line ferrihydrite is mainly attributed to the larger surface area but also related to surface hydroxyl groups, acidic products, and possible additional adsorption sites. The recyclability of mesoporous 2-line ferrihydrite catalyst can be achieved up to 3 times without performance decay. At last, a discussion regarding the possible mechanisms of degradation of orange II over mesoporous 2-line ferrihydrite is proposed, based on the previous literature work and the observed reaction intermediates monitored by ESI/MS in this study.

1. INTRODUCTION Mesoporous materials have been of great interest in the past decades in many fields because of their benefits such as large specific surface areas, large pore sizes, and volumes.1−8 Among the mesoporous materials, mesoporous iron oxides are widely used in photocatalytic applications.9−14 In the photo-Fenton process catalyzed by mesoporous iron oxides, diffusion can be © 2015 American Chemical Society

promoted by large pore sizes and pore volumes of the catalysts.15,16 At the same time, the number of active sites can also be increased by the large surface areas. Generally, iron Received: March 2, 2015 Revised: April 21, 2015 Published: April 22, 2015 10454

DOI: 10.1021/acs.jpcc.5b02057 J. Phys. Chem. C 2015, 119, 10454−10468

Article

The Journal of Physical Chemistry C

above, an ultraviolet light source was used for the Fenton reaction. Other examples with both visible light and neutral pH come from Cheng et al.37 and Du et al.38 Cheng et al. claimed that 90% of methyl orange was degraded by Fe(III)-loaded resin within 500 min at nearly a neutral pH condition (pH = 6.0) in the presence of H2O2 under visible irradiation. Du et al. reported that clay-supported iron oxide was able to degrade 65% of anionic orange II in 150 min and >90% of cationic dye methyl orange within 60 min at near neutral pH conditions (pH = 6.5) under visible light and H2O2. Recently, researchers from our group39 developed a unique sol−gel-based inverse micelle method to synthesize UCT materials (University of Connecticut mesoporous materials), including manganese oxide, iron oxide, cobalt oxide, nickel oxide, titanium oxide, cerium oxide, zirconium oxide, etc. Surface areas, pore volumes and sizes, and nanocrystallinity of these materials are tunable. UCT materials have shown excellent activities in many applications.40−42 Based on this synthesis method, mesoporous amorphous iron oxide (proved to be 2-line ferrihydrite in this study) and α-Fe2O3 have been successfully synthesized.39 If as-made iron oxides are heated directly at 250 °C without removing carboxyl groups and nitrate ions at 150 °C, thermally unstable Fe3O4 with nonporous features can be synthesized.39 In this study, we modified the sol−gel-based inverse micelle methods published before39 and successfully synthesized mesoporous 2-line ferrihydrite (UCT-5), α-Fe2O3 (UCT-6), γ-Fe2O3 (UCT-64), and Fe3O4 (UCT-65) and by changing the reaction temperature and calcination atmosphere. Various characterization techniques have been performed in order to understand the properties of synthesized materials. The synthesized mesoporous materials have been applied to azo dye (orange II) degradation in the presence of oxidant H2O2 at neutral pH under visible light. The possibility of recycling catalysts offered by this study serves a further advantage of using these catalysts for industrial applications. At the end of the study, the key factors on the activities of mesoporous iron oxides are discussed and the azo dye degradation mechanism is proposed based on reaction intermediates and surface properties of mesoporous iron oxides.

oxides can be synthesized by a hard template method or by a soft template method. Jiao et al.17 synthesized ordered mesoporous α-Fe2O3 with crystalline walls using silica KIT-6 as a hard template. Mesoporous γ-Fe2O3 and Fe3O4 could be further synthesized based on prepared mesoporous α-Fe2O3 by a post-template oxidation/reduction.18 Mesoporous iron oxide can also be synthesized by a soft template method. Jiao and Bruce 19 synthesized mesoporous iron oxide with two dimensions and three dimensions using decylamine as a soft template. Mitra et al.20 synthesized mesoporous iron oxides including α-Fe2O3 and γ-Fe2O3 using only sodium dodecyl sulfate or a mixture of sodium dodecyl sulfate and benzyl alcohol as soft templates. Utilization of mesoporous iron oxide catalysts in dye degradation can be found in the literature. Kim et al.21 synthesized mesoporous iron oxide-layered tinanate nanohybrids with pore sizes of 3.7−12 nm. The prepared material could degrade 100% methylene blue within 120 min under visible light. Panda et al.22 synthesized a mesoporous Fe2O3− SiO2 composite by using CTAB as a structure directing agent. The mesoporous Fe2O3−SiO2 composite was able to discolor 98.5% methyl orange within 20 min in the presence of H2O2. Moreover, after three cycles, the catalyst was very stable and only showed 1.5% loss of activity in discoloration and mineralization of methyl orange. However, their catalysts only work well under acid conditions (pH < 4). When iron oxides are used as photo-Fenton catalysts, usually there are several drawbacks which limit their effectiveness: (1) ultraviolet light, which only accounts for 2% of sunlight, is not as widely available as visible light; (2) acid reaction conditions, which have been proved to offer the best reaction conditions for Fenton catalysts,23,24 also needs further processing to avoid second-time contamination in the processing of wastewater treatment; (3) the large volume of Fe(III) sludge after a Fenton reaction needs to be removed at the end of treatment by precipitation, which causes additional expense of the wastewater treatment. Therefore, the development of active heterogeneous systems to enable the photo-Fenton reaction to operate with (1) visible light, (2) neutral or near-neutral pH conditions, and (3) no Fe ion sludge generation is of great significance. In addition, the use of heterogeneous catalysts can prevent generation of iron cement and subsequent catalyst deactivation, which usually happen in the traditional photoFenton process. So focus has been given to research in heterogeneous catalysis to allow the photo-Fenton process to work at the above-mentioned conditions without a trade-off of efficiency. For the Fenton process under visible light, significant contribution can be found from work in the literature.25−28 However, most of the work is limited by acidic conditions.29−31 This is because formation of the key oxidant •OH radicals can be accelerated under acid conditions (best at pH = 2−4) in the Fenton reaction.25,32−35 As an improvement, in the application of azo dye degradation, a few researchers have studied the catalyst performance under neutral or near neutral pH. Feng et al.36 studied the effectiveness of four different catalysts on discoloration and mineralization of orange II at pH = 3 and pH = 6.6 conditions. They observed bentonite clay-based Fe nanocomposite was the catalyst with the best performance in both pH conditions. Specifically, at pH = 3 conditions, Fe nanocomposite was able to degrade 90% of orange II within 10 min whereas at pH = 6.6 condition Fe nanocomposite needed 30 min to degrade 90% of orange II. However, in the case

2. EXPERIMENTAL SECTION 2.1. Synthesis. Iron nitrate nonahydrate (≥98%), poly(ethylene glycol)-block-poly(propylene glycol)-block-poly(ethylene glycol), P123 (Mn ∼ 5800), nitric acid (≥70%), 1butanol (≥99.4%), and iron(III) oxide (99.98%, metal basis) were purchased from Sigma-Aldrich; H2O2 solution (∼30%) was from Fisher Scientific. All chemicals were used as received without further purification. Iron nitrate, 0.01 mol, was dissolved in a solution containing 0.12 mol of 1-butanol, 0.019 mol of nitric acid, and 2.04 × 10−4 mol of P123 in a 150 mL beaker at room temperature with magnetic stirring.39 Then the clear gel was put in the oven under air at 95 °C for 2−3 h to synthesize mesoporous 2-line ferrihydrite and mesoporous αFe2O3. The clear gel was placed in the oven under air at 100 °C to synthesize mesoporous γ-Fe2O3 and Fe3O4. The obtained powders were washed twice with ethanol and then centrifuged. Finally, the powders were dried in a vacuum oven overnight. To synthesize mesoporous 2-line ferrihydrite and γ-Fe2O3, the dried powders (as-made samples) were heated at 150 °C for 12 h and then heated at 250 °C 4 h under an air atmosphere. To synthesize mesoporous α-Fe2O3, the dried powders were heated at 150 °C for 12 h, then heated at 250 °C for 4 h, 10455


Article

The Journal of Physical Chemistry C and finally heated at 350 °C for 3 h under an air atmosphere. To synthesize mesoporous Fe3O4, the dried powders were heated at 150 °C for 12 h and then heated at 250 °C for 4 h under a nitrogen atmosphere. The different parameters which were used in the synthesis method are summarized in Table 1. The synthesis processes of each phase of iron oxide were repeated at least five times for accuracy and repeatability.

TG-MS analyses were performed with a TG 209 F1 Libra thermogravimetric analyzer coupled to a QMS 403C quadrupole mass spectrometer. The samples (20 mg) were kept at 120 °C to removal adsorbed water and then heated to 600 °C with a 5 °C/min heating rate under 20 mL/min argon. A scanned bargraph mode was used for MS measurements. The dc magnetic measurements were carried out as a function of applied magnetic field (−10 to 10 kOe at 300 K) with a vibrating sample magnetometer attached to the Evercool Physical Property Measurement System (PPMS) from Quantum Design. The dye solution was separated from iron oxides by filtration of a microfilter with 0.45 μm pore size (Pall Corporation). The concentrations of dyes were quantified from their adsorbance intensities measured by a Shimadzu UV-2450 UV−vis spectrophotometer (Shimadzu Scientific Instruments, Tokyo, Japan). The decrease in absorption intensity at 485 nm for orange II was monitored to calculate the concentration and determine % photodegradation (C/C0 × 100%, where C is the final concentration and C0 is the initial concentration which here refers to concentrations after adsorption in the dark). A series of external standard dye solutions were used for calibration. An API 2000 mass spectrometer (Applied Biosystem/MDS SCIEX, CA) using an electrospray ionization (ESI) source in negative mode was used to analyze the dye degradation products under different conditions. A 1:1 ratio of HPLC-grade acetonitrile to dye sample solution was used. The solvent flow rate was set at 10 μL min−1 while the injection volume was 50 μL. The needle voltage was −4500 V. The desolvation temperature was 250 °C. The scan range was m/z 50−400. 2.3. Catalyst Performance Test. For the dye adsorption experiment, typically, 50 mg of iron oxide and 200 mL of 0.1 mM aqueous orange II solution at various initial concentrations C0 was mixed in a beaker, followed by stirring at a constant rate in the dark overnight. After equilibrium was reached, the suspension was filtered through a microfilter with a 0.45 μm PVDF membrane (Pall Corporation), and the substrate Ce concentration remaining in the filtrate was analyzed. The decreased concentration (C0 − Ce) was then used to calculate the amount of adsorption, in units of milligrams per gram of iron oxide. Reactor lamps RPR-5750A (λpeak = 575 nm, > 97% in the visible light range), 8 W, were used as a constant source of visible light radiation. Sixteen visible lamps were installed equally spaced on a Rayonet RPR-200 photochemical chamber reactor. Both the reactor lamps and the reactor were purchased from Southern New England Ultraviolet Co., CT. These types of lamps and photoreactor have been recently utilized to mimic sunlight for pollutant degradation.46 Loss of irradiance was minor as no water cooling jacket was installed around the quartz beaker. Dye solutions (0.1 mM) were prepared by dissolving relative amounts of dyes into DDW. Before the degradation reaction, dye solution and catalysts (0.25 g/L) were stirred vigorously in the dark overnight to eliminate the adsorption effect to reach equilibrium. Then the degradation reaction was started by adding H2O2 solution to the suspension (to reach 12.5 mM concentration in solution) under visible light irradiation in a 600 mL quartz beaker as the photoreactor. The pH of solution was the intrinsic pH and no more addictive was added to adjust the pH value. The solution was stirred by a magnetic stir bar. All the catalysts were suspended in the solution during the dye

Table 1. Experimental Parameters for Various Mesoporous Iron Oxides

2-line ferrihydrite α-Fe2O3 (hematite) γ-Fe2O3 (maghemite) Fe3O4 (magnetite)

reaction temp (°C)

heating cycles

heating atmosphere

95

150 °C (12 h)−250 °C (4 h)

air

95

air

100

150 °C (12 h)−250 °C (4 h) −350 °C (3 h) 150 °C (12 h)−250 °C (4 h)

100

150 °C (12 h)−250 °C (4 h)

nitrogen

air

2.2. Characterizations. Powder XRD data were collected on a Rigaku UltimaIV instrument using Cu Kα (1.54 Å) radiation at a beam voltage of 40 kV and a 45 mA beam current. High- and low-angle patterns were obtained by continuous scans in a 2θ range of 5−90° and 0.5−5° with a scan rate of 2°/ min and 0.5°/min, respectively. The surface area, pore size distribution, and pore volume of iron oxides were determined in a Micrometitics ASAP 2020 accelerated surface area and porosimetry system. The isotherms of N2 at 77 K were obtained from physisorption. Before analysis, the samples were degassed at 120 °C under vacuum for 12 h to remove the surface contaminants. The pore size distributions of the iron oxides were determined from the N2 desorption isotherms at 77 K, using the Barrett−Joyner−Halenda (BJH) method. Highresolution scanning electron microscope (HRSEM) photographs were taken on a Zeiss DSM 982 Gemini FESEM with a Schottky emitter at an accelerating voltage of 2.0 kV and a beam current of 1.0 mA. A high-resolution transmission electron microscope (HRTEM) was used to characterize the crystallized structure and the morphology of the samples using a JEOL 2010 instrument with an accelerating voltage of 200 kV. Raman spectra were obtained using a Renishaw 2000 Ramanscope which has an optical microscope (0.024 in. focus length), graphic grating (1800 mm−1), and CCD detector. A laser excitation source with 514 nm wavelength was used for testing mesoporous α-Fe2O3, γ-Fe2O3, and Fe3O4; a laser excitation source with 633 nm was used for testing mesoporous 2-line ferrihydrite. The laser focus was set to 40% to prevent local damage.43−45 For each sample, three different locations were analyzed to verify the spectra. The X-ray absorption spectroscopic (XAS) measurements were conducted at the National Synchrotron Light Source (NSLS) at Brookhaven National Laboratory using beamline X18A. The synchrotron radiation was monochromatized by a silicon (111) double crystal monochromator. The incident and transmitted beam intensities were monitored using ionization chambers filled with mixtures of He and N2 gases. The samples were diluted by h-BN with a ratio of (1:8), then pressed into pellets, and mounted between the ionization chambers of Io and It. A thin iron foil reference was used for energy calibration. The XANES data were analyzed using Athena software where background and post- and pre-edge corrections were made. 10456


Article


Figure 1. (a) High-angle and (b) low-angle PXRD patterns of mesoporous iron oxides.

Figure 2. (a) N2 sorption isotherms and (b) BJH pore size distributions for mesoporous iron oxides.

(Figure S1a,f), the weight loss takes place at 150 and 300 °C, caused by losing both H2O and CO2. The temperature of releasing H2O overlaps with that of losing CO2, which makes the attempt to calculate the formula of 2-line ferrihydrite unsuccessful. The crystallinity of mesoporous α-Fe2O3 was obtained by further heating mesoporous 2-line ferrihydrite at 350 °C. Recent studies from our group39 found that transition metal oxides are amorphous when calcined at a low temperature while they become crystalline with different phases when calcined at a higher temperature. Patterns of γ-Fe2O3 and Fe3O4 only show differences below 30°. Both γ-Fe2O3 and Fe3O4 are inverse spinel structures, and the differences between the crystal structures of them are the iron valence and the vacancies. Fe3O4 can be formulated as (FeIII)A[FeIIFeIII]BO4, and γ-Fe2O3 can be formulated as (FeIII8)A[FeIII40/3−8/3]BO32, where A represents tetrahedral sites and B represents octahedral sites. Figure 1b shows the low-angle diffraction lines for mesoporous iron oxide samples. The low-angle diffraction line position is 1.35° (6.5 nm) for 2-line ferrihydrite. Since mesoporous α-Fe2O3 was obtained by further heating 2-line ferrihydrite to 350 °C, the shift of low-angle diffraction peak positions of mesoporous α-Fe2O3 to 0.94° (9.4 nm) is attributed to the sintering nature of small crystalline nanosize particles.39 Mesoporous γ-Fe2O3 has a similar low-angle diffraction line position to 2-line ferrihydrite which is at 1.35° (6.5 nm). The meso structures (low-angle diffraction line) of

degradation reactions. The beaker was covered with a watch glass to minimize water evaporation and interference from the atmosphere.

3. RESULTS 3.1. Structure of Mesoporous Catalysts. High-angle PXRD is used to identify phases and crystallinity of the synthesized iron oxides samples (Figure 1a). PXRD patterns of mesoporous α-Fe2O3, γ-Fe2O3, and Fe3O4 match with standard PDF cards of hematite (33-0664), maghemite (39-1346), and magnetite (19-0629), respectively. In comparison with the literature, XRD patterns of α-Fe2O3, γ-Fe2O3, and Fe3O4 show low peak intensities. This indicates low crystallinity of these samples. The PXRD pattern of 2-line ferrihydrite (Fe2O3· nH2O) contains two peaks of low intensity appearing at dspacings of 2.53 Å (35.4°) and 1.47 Å (63.0°), which match patterns of 2-line ferrihydrite very well in the literature.47 For ferrihydrite, a range of compounds with different degrees of structural order exist; these compounds are generally named according to the number of broad X-ray peaks which they exhibit, e.g., 2-line and 6-line ferrihydrite.48 The formula of ferrihydrite has not been fully established.48 Fe5HO8·4H2O and 5Fe2O3·9H2O have been suggested, but the amounts of structural water may in fact be less than those indicated in these formulas.48,49 It was difficult to obtain the actual formula of 2-line ferrihydrite in this study. As shown in the TG-MS plot 10457


Article


because further heat treatment at 350 °C will cause more sintering of nanoparticles. SEM and TEM images of mesoporous iron oxides are shown in Figure 3. From SEM, iron oxide nanoparticles have the morphology of spherical, hedgehog-like, or pin-cushion aggregates that consist of radially oriented nanoparticles. From TEM images, the thorns are formed by small crystalline nanosize particles and pores are formed among these nanosize particles. All TEM images are collected at the same magnification. Hence, the effects of experimental parameters on pore sizes and aggregation of nanoparticles can be observed. The order of pore sizes of mesoporous iron oxides (Fe3O4 > γFe2O3 > α-Fe2O3 > 2-line ferrihydrite) matches that of pore sizes obtained from BET pore size distributions. Raman spectra of mesoporous 2-line ferrihydrite, α-Fe2O3, γFe2O3, and Fe3O4 are shown in Figure 4. Raman can detect phase information which can help us confirm the phases of iron oxides. Raman spectra of iron oxides match literature spectra43,50,51 very well. Mesoporous 2-line ferrihydrite matches 2-line ferrihydrite in the literature. Peaks in the spectrum of mesoporous α-Fe2O3 are 228, 246, 293, 411, 497, 616, and 1320 cm−1. Peaks in the spectrum of mesoporous γ-Fe2O3 are 352, 504, 719, and 1409 cm−1. Peaks in the spectrum of mesoporous Fe3O4 are 352, 528, and 682 cm−1. The peak shift (676 to 682 cm−1)43 of Fe3O4 is attributed to nonequivalent sites or internal strain in the structure.43,50 Figure 5 shows the XANES data of mesoporous 2-line ferrihydrite, α-Fe2O3, γ-Fe2O3, and Fe3O4 compared with standard iron oxides. From the XANES data, the mesoporous 2-line ferrihydrite, α-Fe2O3 match standard α-Fe2O3 very well. Mesoporous γ-Fe2O3 and Fe3O4 match standard γ-Fe2O3 and Fe3O4 very well. The average oxidation states of mesoporous 2line ferrihydrite, α-Fe2O3, γ-Fe2O3, and Fe3O4 are 2.84, 2.77, 2.68, and 2.56, respectively. From the average oxidation state, all the mesoporous iron oxides have some oxygen vacancies. Figure 6 presents field-dependent magnetism by VSM of mesoporous γ-Fe2O3 and Fe3O4 catalysts at 300 K. The mesoporous γ-Fe2O3 and Fe3O4 catalysts exhibit superparamagnetic behavior52,53 with saturated magnetization values of 39.8 and 45.7 emu/g, respectively. Mesoporous γ-Fe2O3 and Fe3O4 catalysts can be easily recycled by external magnets

mesoporous 2-line ferrihydrite, α-Fe2O3, and γ-Fe2O3 are all very clear. In the case of Fe3O4, we can only see a shoulder because the position of the line, 0.58° (15.2 nm) is too low to get the full peak. The mesoporous structures of the iron oxides are further confirmed by nitrogen sorption. Nitrogen sorptions and BJH pore size distributions are shown in Figure 2a,b. All iron oxides show characteristic type IV adsorption isotherms, suggesting that the materials preserve the mesoporous structure with different reaction temperatures, heating cycles, and heating atmosphere. In addition, these different experimental parameters in the synthesis of mesoporous iron oxide show excellent control in the development of pores and phases. Compared with 2-line ferrihydrite, mesoporous α-Fe2O3 shows pore expansion, increase in dspacing, and decrease in surface area (Figure 1b, Figure 2b, and Table 2) caused by nanoparticle sintering upon thermal Table 2. Structure Parameters of Mesoporous Iron Oxides surface area (m2/g) pore size (nm) pore volume (cm3/g)

2-line ferrihydrite

α-Fe2O3

γ-Fe2O3

Fe3O4

258 2.2 0.20

137 3.4 0.19

100 5.1 0.15

75 7.1 0.18

treatment at a higher calcination temperature (350 °C). At a reaction temperature of 100 °C, γ-Fe2O3 with a final calcination temperature of 250 °C has a similar pore size distribution as αFe2O3 with a final calcination temperature of 350 °C (Figure 2b). Furthermore, mesoporous Fe3O4 has much larger pore size than the other mesoporous iron oxides, including mesoporous α-Fe2O3 which is further treated at a higher calcination temperature. The detailed structural parameters of mesoporous iron oxides are listed in Table 2. UCT materials have a typical unit-cell expansion behavior upon heat treatment caused by nanoparticle sintering.39,42 The explanation of pore expansion could be that different phases of materials have different tendency of sintering on heat treatment. So even after the same calcination cycle at 150 °C for 12 h and 250 °C for 4 h, mesoporous 2-line ferrihydrite, γ-Fe2O3, and Fe3O4 nanocrystals have different extents of pore expansion. Mesoporous αFe2O3 has larger pore size than mesoporous 2-line ferrihydrite

Figure 3. SEM images (top row) and TEM images (bottom row) of mesoporous (a) 2-line ferrihydrite, (b) α-Fe2O3, (c) γ-Fe2O3, and (d) Fe3O4 (scale bars are 200 nm in SEM images and scale bars are 20 nm in TEM images). 10458


Article


Figure 4. Raman spectra of mesoporous (a) 2-line ferrihydrite, (b) α-Fe2O3, (c) γ-Fe2O3, and (d) Fe3O4.

Figure 5. XANES data of mesopporous (a) 2-line ferrihydrite, (b) α-Fe2O3, (c) γ-Fe2O3, and (d) Fe3O4 (blue solid lines) compared with standard samples (red dashed lines) (a) α-Fe2O3, (b) α-Fe2O3, (c) γ-Fe2O3, and (d) Fe3O4.

because of their superparamagnetic properties. This is an additional advantage of mesoporous iron oxide catalysts. The picture of the samples (Figure S5) shows that the mesoporous 2-line ferrihydrite, α-Fe2O3, and γ-Fe2O3 are dark brown and the mesoporous Fe3O4 is black. The colors of the samples indicate mesoporous iron oxides are able to absorb visible light. 3.2. Catalyst Performance Test. Figure 7 shows the adsorption capabilities of mesoporous iron oxides, compared

with that of commercial Fe2O3. The average adsorption capability of commercial Fe2O3 is 21.1 mg/g. The greatest adsorption capability (49.3 mg/g) is found on mesoporous 2line ferrihydrite. All of the mesoporous iron oxides have greater adsorption capability than the commercial iron oxide because of the large surface area of mesoporous Fe2O3. The adsorption capability of γ-Fe2O3 is similar to that of commercial Fe2O3. Furthermore, the adsorption capability of α-Fe2O3 and Fe3O4 is 50% more than that of commercial Fe2O3. In addition, the 10459


Article


ferrihydrite, α-Fe2O3, γ-Fe2O3, and Fe3O4 have higher rate constant than commercial Fe2O3. The commercial Fe2O3 catalysts can only degrade about 70% of the dye after 8 h reaction. With mesoporous 2-line ferrihydrite, α-Fe2O3, and γFe2O3 catalysts, 100% of dyes can be degraded within 8 h of reaction time. Among them, the mesoporous 2-line ferrihydrite was the most efficient catalyst. This catalyst can achieve 100% degradation of dyes within 3 h. Mesoporous γ-Fe2O3 can degrade the dyes completely within 5 h. The half-life of the dyes which are catalyzed by commercial Fe2O3 is 3 h whereas the half-life of the dyes which were catalyzed by mesoporous αFe2O3 is 2 h. In the case of the mesoporous Fe3O4 catalyst, there is a small platform at the beginning of the reaction, as shown in Figure 8. In the first 2 h of the reaction, the dye contents change very little with only 5% degraded. After an 8 h reaction, mesoporous Fe3O4 is able to degrade 80% of orange II dyes. Figure 9 shows the pseudo-first-order kinetic plot of dye

Figure 6. Vibrating sample magnetometer curves of mesoporous γFe2O3 and Fe3O4 at 300 K.

Figure 7. Adsorption capabilities of mesoporous 2-line ferrihydrite, αFe2O3, γ-Fe2O3, Fe3O4, and commercial Fe2O3.

adsorption capability of 2-line ferrihydrite is more than twice that of the commercial Fe2O3. All these results prove that the mesoporous iron oxides synthesized by the sol−gel-based inverse micelle method have greater adsorption capability for orange II than commercial Fe2O3. Figure 8 shows the dye degradation efficiency with different mesoporous iron oxide compared with commercial iron oxide

Figure 9. Pseudo-first-order kinetics plot of degradation of orange II by mesoporous iron oxides.

degradation reactions by mesoporous iron oxides. The corresponding apparent reaction rate constants of dye degradation reactions with iron oxides catalysts are listed in Table 3. The apparent rate constant of α-Fe2O3 is twice as Table 3. First-Order Kinetic Parameters for Degradation Reactions by Mesoporous Iron Oxides and Turnover Frequencies of Catalysts catalysts 2-line ferrihydrite α-Fe2O3 γ-Fe2O3 Fe3O4 commercial Fe2O3 catalyst free

ka (min−1)

R2

turnover freqb (×10−7 mol/(m2/(g h))

0.0258 0.0082 0.0137 0.0037 0.0040

0.930 0.992 0.961 0.952 0.998

1.260 0.888 1.960 1.375 4.236

0.0006

0.970

N/A

Figure 8. Dye degradation effects by different mesoporous iron oxides.

a

and catalyst free conditions under visible light in the presence of H2O2. Without catalysts, the dyes can be slightly degraded by H2O2 under irradiation of visible light. Without H2O2, the dyes can only be adsorbed rather than degraded by the catalysts under irradiation of visible light (Figure S6). After 8 h, no more than 30% of dyes were degraded. In Figure 8, the rate constant of reaction without iron oxide catalysts was very slow compared to that with catalytic reactions. All the mesoporous 2-line

much as that of commercial Fe2O3, and the apparent rate constant of γ-Fe2O3 is 4 times as much as that of commercial Fe2O3. Moreover, the apparent rate constant of 2-line ferrihydrite is 6 times more than that of commercial Fe2O3. These catalysts are much more efficient than commercial Fe2O3 in catalyzing dye degradation of orange II. The turnover frequency (TOF) is defined as the number of molecules of a

Apparent rate constant. bTurnover frequency is expressed as (no. dye reacted)/(unit surface area)(unit time).

10460


Article


decay of the catalyst appears at the end of the dye degradation reaction, the overall performance of the mesoporous 2-line ferrihydrite catalyst is good in terms of stability during the first three cycle, which is further confirmed by HRTEM (Figure S4). 3.4. pH Value and Intermediates of Reaction. The initial pH of solution is the intrinsic pH of orange II solution (8.4). However, the pH of the reaction solution is changed during both adsorption in the dark and the reaction. From Figure 12, after adsorption in the dark overnight, the pH values

given product per catalytic site per unit time. Unit surface area of catalyst is used to represent the number of active sites for the calculation of the TOF.54 Calculated TOF’s of iron oxides catalysts are also listed in Table 3. 3.3. Stability and Recyclability of Mesoporous Iron Oxide Catalyst. The performance of recycled mesoporous 2line ferrihydrite was tested by repeating the catalytic degradation of orange II over mesoporous 2-line ferrihydrite/ visible light/H2O2 system for four times, as shown in Figure 10.

Figure 12. pH values of reaction solutions on different mesoporous iron oxides.

Figure 10. Performance of recycled mesoporous 2-line ferrihydrite catalysts.

of solutions with Fe3O4 and 2-line ferrihydrite decrease significantly to 6.1 and 6.0, respectively, whereas the pH values of solutions with α-Fe2O3, γ-Fe2O3, and commercial Fe2O3 decrease only a little to 8.2, 8.2, and 8.0, respectively. Overall, during the degradation reaction of orange II, the pH values of reaction solutions have a decreasing tendency regardless of iron oxide catalyst types and properties. For experiments with 2-line ferrihydrite, the pH of the solution drops dramatically to 4.4 as soon as the degradation reaction time reaches 0.5 h. During the process of degradation, the pH of the solution with 2-line ferrihydrite is always the lowest. After the reactions, the pH of mesoporous Fe3O4 changes from 6.1 to 4.4, while the pH values of solutions with α-Fe2O3, γ-Fe2O3, and commercial Fe2O3 drop to 6.1, 5.1, and 5.3, respectively. The drop of pH caused by degradation products is consistent with the literature.38

The recycle experiments show that the performance of the mesoporous 2-line ferrihydrite catalysts can remain almost unchanged after at least three cycles of degradation reactions. High-angle PXRD pattern of used catalyst shows amorphous nature of catalysts. The appearance of the peak in the low-angle PXRD pattern of the used catalyst indicates that the used catalyst still maintains the meso structure after the fourth run, as shown in Figure 11b. However, the peak in the low-angle PXRD patterns of the used catalyst shifts to lower angles, which means the d-spacing of the used catalyst is larger than the fresh catalyst. The structural change of the catalyst could be possibly caused by the residue of carboxyl groups on the surface39 (Figure S1). For the high angle PXRD patterns, the used catalyst maintains the amorphous nature after the fourth run of the reaction (Figure 11a). Combining the results from Figures 10 and 11, although, in the fourth cycle, some degree of activity

Figure 11. (a) High-angle and (b) low-angle PXRD patterns of recycled mesoporous 2-line ferrihydrite. 10461


Article


Figure 13. ESI/MS plots of reaction solutions over (a) mesoporous 2-line ferrihydrite and (b) commercial Fe2O3. The intensity of each plot has been adjusted to the same scale.

4. DISCUSSION 4.1. Structure of Mesoporous Catalysts. During the synthesis of mesoporous iron oxides, reaction temperature is a critical parameter, since inverse micelle formed by surfactant species serves as a nanoreactor in our sol−gel-based inverse micelle method.39 At a higher reaction temperature, the formation of carboxyl groups from the surfactant has the possibility of resulting in partially deforming the inverse micelles. According to this, when higher reaction temperature (100 °C) is used for synthesizing mesoporous γ-Fe2O3, the pore volume of γ-Fe2O3 is smaller than that of 2-line ferrihydrite due to decomposition of surfactants. This explains the result of Figure 2a, where γ-Fe2O3 shows only a small hysteresis loop. The uniformity of the nanocrystalline distribution is achieved through a short nucleation period that produces all the particles obtained at the end of the reaction. During the synthesis procedure, for the late transition metal oxides such as Mn, Fe, and Co, the adsorbed NOx and carboxyl species on the materials are concerns for these systems because they have the possibility of forming multiple oxidation states.39 In our iron oxide case, the as-synthesized iron oxides possibly have an in situ redox reaction with NOx and carboxyl species at a higher reaction temperature (100 °C) because of increased oxidation potential of NOx and carboxyl species when the temperature is increased. On the other hand, at a lower reaction temperature (95 °C), there would be less chance for the redox reaction to take place which could form iron(III) oxide with higher oxidation states. Higher reaction temperatures such as 105 and 110 °C have also been tried during the experiment. However, these temperatures were not adopted because there was greater chance for the reaction gel to dangerously splash out when the reaction almost reached the end. When heated in a nitrogen atmosphere, Fe(II) species could be kept to form Fe3O4 while Fe(II) species were oxidized back to a higher valence when calcined at 250 °C in an air atmosphere. The formation of different phases of iron oxide nanoparticles has been observed in the literature55,56 during low temperature

ESI/MS characterization of the reaction solutions catalyzed by 2-line ferrihydrite and commercial Fe2O3 are shown in Figures 13a and 13b, respectively. The degradation reaction of orange II usually starts from hydroxyl radicals attacking −N N− functional groups. So the intermediate products after the first step dye degradation reaction are fragments with m/z 187 and 157. A large amount of 50−150 m/z fragments appearing right after reaction begins (0.5 h) with 2-line ferrihydrite while with commercial Fe2O3, they appear at the half time of the reaction (4 h). In addition, in Figure 13a, the initial fragment from orange II (m/z = 327) disappears after 1 h reaction time, and the fragment (m/z = 187, contains a naphthalene ring) after the first step degradation disappears after 2.5 h reaction time. However, in Figure 13b, with commercial Fe2O3, the initial fragment from orange II (m/z = 327) does not disappear until the end of the reaction. When the dye degradation reaction is catalyzed by mesoporous 2-line ferrihydrite, large m/ z fragments with the aromatic rings after the first-step degradation are quickly reacted to produce other small m/z compounds (appears from the very beginning of the degradation reaction) by possible reactions of being attacked by hydroxyl radicals. In summary, mesoporous iron oxides show better performance than commercial Fe2O3, not only on the aspect of higher reaction rate constant but also being able to quickly degrade dyes into smaller species by breaking aromatic rings. The intensity of orange II (m/z = 327) in the original solution was around 1.5 × 106. After 3 h of reaction over mesoporous 2-line ferrihydrite, the intensity of orange II (m/z = 327) dropped to 8.3 × 103 (∼0.5% of the original intensity). The complete degradation of the dye is more easily observed in the UV−vis results shown in Figure 8. The complete discoloration of the dye is also consistently observed after 3 h. In ESI/MS, some intermediates can still be observed after 3 h of reaction. Further degradation of the intermediates to gases or other smaller compounds through mineralization is not a focus in this study. 10462


Article


between commercial Fe2O3 and orange II. This causes a preferable adsorption of orange II on mesoporous iron oxides compared with the commercial Fe2O3. Another adsorption mechanism of dyes on iron oxides was proposed by Bandara et al. in their studies.61,62 They showed that sulfonic groups of orange II precipitated the adsorption on iron oxide surface by the formation of an unidentate complex with oxide surface. They observed an SO2 symmetric peak at 1190 cm−1 and an asymmetric peak 1305 cm−1 disappeared after adsorption. Similar observations were reported in the literature. Figure S3a shows infrared spectra of orange II adsorbed on mesoporous 2-line ferrihydrite, mesoporous αFe2O3, and commercial Fe2O3 compared with free orange II. After being adsorbed on catalysts, the asymmetric stretching peak (1191 cm−1) of −SO2 is missing since −SO2 groups are attached on the catalysts after adsorption. Peaks at 1405 and 1318 cm−1 on the spectra of mesoporous 2-line ferrihydrite also disappear. These two peaks represent the bending −OH deformation. This means that adsorption of orange II on mesoporous 2-line ferrihydrite may restrict the −OH bending deformation which leads to a possibility that −OH is on another binding site when adsorbed on mesoporous 2-line ferrihydrite. On one hand, additional sites of adsorption may make the position of the adsorbed molecules “more fixed”. On the other hand, additional sites of adsorption may give the adsorbed molecules extra strain by possibly changing the bond angle and bond length. Both effects caused by additional sites of adsorption will cause the adsorbed molecules to be attacked more easily by the hydroxyl radical. So with an increase in overall energy, adsorbed molecules will become much more unstable. Figure S3b shows infrared spectra of orange II adsorbed on mesoporous γ-Fe2O3 and Fe3O4 compared with free orange II. The peak at 1160 cm−1 and broad peak at the 1230 cm−1 are due to aromatic C−H bending. This is because crystal structures of γ-Fe2O3 and Fe3O4 are both cubic close packed while the crystal structures of mesoporous 2-line ferrihydrite and γ-Fe2O3 are both hexagonal close packed. 4.3. Catalyst Performance. The commonly used photoFenton process conditions for dye removal were reviewed by Soon and Hameed.63 As a summary, most researchers used 1.0 g/L catalyst, 0.1−0.2 mM dye, around 10−20 mM H2O2, mostly pH < 7, and ultraviolet light (∼8 W) as experimental conditions. Compared with these common conditions, this study used a relatively small amount of catalyst (0.25 g/L), neutral pH, and visible light. As for visible light power, Du et al.38 and Dong et al.28 used a 500 W halogen lamp and a 400 W mercury lamp with a cutoff filter to generate visible light, respectively. Compared with their study, the lamp power used in this study (16 × 8 W = 128 W) is relatively low. In addition, the electrical energy efficiency is calculated to be 0.0006% for the best catalyst (mesoporous 2-line ferrihydrite). The detailed calculation process is shown in the Supporting Information. For the catalytic degradation of orange II under visible light in the presence of H2O2, all the mesoporous iron oxides have higher rate constants than commercial iron oxide as Fenton catalysts. Furthermore, the reaction catalyzed by mesoporous 2line ferrihydrite has the highest rate constant. The major reason for the excellent performance of the mesoporous iron oxides is that they contain large surface area (Table 2), which will provide more available sites for orange II to adsorb on and for H2O2 to produce highly reactive hydroxyl radicals. The TOFs of mesoporous iron oxides (Table 3) are smaller than that of commercial Fe2O3. Since TOF is calculated by the number of

sol−gel syntheses. In order to obtain more information about the surface hydroxyl and carboxyl groups, TG-MS analyses of all the synthesized mesoporous materials and commercial Fe2O3 were performed and are shown in Figure S1. Although carboxyl groups on mesoporous iron oxides can be mostly removed by heating at 150 °C for 12 h and followed by heating at 250 °C for 4 h,39 some residues of carboxyl groups are observed on mesoporous 2-line ferrihydrite and Fe3O4. These residues of carboxyl groups may cause the structure of mesoporous 2-line ferrihydrite and Fe3O4 materials to be unstable while the structure of mesoporous α-Fe2O3, γ-Fe2O3, and commercial Fe2O3 may not be affected by carboxyl group residues. 4.2. Adsorption Properties. The adsorption capacities (mg dye/g catalyst) of mesoporous iron oxides in this study are greater than commercial Fe2O3 (shown in Figure 7) mainly due to the large surface area of mesoporous iron oxide. The mesoporous iron oxides synthesized in this study with small crystalline nanostructure could have more than 3 times adsorption capability of orange II compared with those (∼17.5 mg L−1) synthesized using conventional precipitation methods, when the orange II concentration was around 0.1 mM.57 Furthermore, the adsorption capability of orange II would become even higher when the pH (pH = 8.4 in this study) is adjusted to be as low as what is used in the literature (pH = 6.5).58 Other competitive catalysts come from flowerlike α-Fe2O3, γ-Fe2O3, and Fe3O4 with 3D ordered nanostructures, which were also synthesized by a self-assembly process. They could perform an average removal capacity of around 43.5 mg orange II g−1.55 The experiments in this study show that the average adsorption capability of mesoporous α-Fe2O3 is 37.5 mg orange II g−1; that of mesoporous 2-line ferrihydrite is 49.3 mg/g. Both of them are comparable with literature values. But the adsorption capability of γ-Fe2O3 and Fe3O4 are lower than the average removal capacity from the literature. The adsorption capability of iron oxides synthesized in this study is roughly in agreement with and can be generally reflected by the trend in surface area as reported in the literature. The adsorption mechanism of dyes on iron oxides is proved by Saha et al. in a recent report.59 They reported that preferable and enhanced adsorption phenomena of the dyes contains hydroxyl (−OH) groups on iron oxide nanoparticles with hydroxyl groups on the surface. The formation of hydrogen bonds will shift the frequency of hydroxyl vibrations (in the 3000−3500 cm−1 range) to lower wavenumbers after dye absorption on the iron oxide nanoparticle surface. The IR shifts due to hydrogen bond formation are also observed on our mesoporous iron oxides. In Figure S2, the vibrational band of hydroxyl group of orange II dye occurs at 3440 cm−1. After orange II adsorbed on the iron oxide catalysts, the peaks of hydroxyl groups of orange II adsorbed on mesoporous 2-line ferrihydrite, α-Fe2O3, γ-Fe2O3, Fe3O4, and commercial iron oxide shifted from 3440 to 3435, 3428, 3435, 3419, and 3438 cm−1, respectively. This is clear evidence of direct hydrogen bonding formation between hydroxyl groups of orange II and the surfaces −OH group of iron oxides. Moreover, comparing with the red-shifts on commercial Fe2O3, there are larger redshifts of −OH bonds adsorbed on mesoporous 2-line ferrihydrite, α-Fe2O3, γ-Fe2O3, and Fe3O4. The longer redshifts indicate stronger elongation of O−H bonds which is caused by stronger attractive interactions between positive H and iron oxide surfaces.60 This means the bonding between mesoporous iron oxides and orange II is stronger than that 10463


Article

The Journal of Physical Chemistry C Table 4. Effectiveness of Iron Oxide Catalysts on Azo Dye Degradation in the Literature dye

dye conc (mM)

36

orange II

0.2

57

orange II

0.1

37

orange II

0.1

38 this study

orange II orange II

0.2 0.1

ref

a

catalyst Fe-B α-Fe2O3 α-Fe2O3 γ-Fe2O3 Fe3O4 Fe(III)-loaded resin Fe2O3/clay 2-line ferrihydrite

catalyst conc (g/L)

light source

1

UV

0.5

UV

0.02 0.5 0.25

H2O2 conc (mM) 10

pH value

time for total conv (min)

kapp (min−1) × 10−3 a

6.6

120

1.2

6.5

b

visible

2.5

6.0

≫700

77 63 8 14 11 b

visible visible

2.0 12.5

6.5 neutral

b 180

4.9 26

Best reported kapp is used in this table. bNot available.

150 min. In comparison, mesoporous 2-line ferrihydrite catalyst can degrade 90% orange II in 150 min. Furthermore, as shown in the literature,25,65 increasing the amount of catalyst amount can largely increase the rate of dye degradation. The catalyst amount used in their study is twice as large as the catalyst amount used in this study. In Table 4, the refs 36 and 57 use a UV light as the light source, which is much stronger than visible light. Thus, higher rate constants and lower degradation times are observed in the literature compared with the current study. However, we believe that the effective catalysts which are active under visible light will be much more preferred for industrial applications in terms of utilization of solar light, since the content of visible light (∼47% of the solar light in energy distribution) is much greater than UV light (∼2% of the solar light in energy distribution). If a catalyst is active under visible light, it can utilize not only the UV part of sunlight but also the visible sunlight, which will greatly increase the energy efficiency. As far as the effectiveness of the catalyst used in this study is concerned, if comparison is performed between this study and those using visible light, the apparent rate constant in this study (26 × 10−3 min−1) is larger than that in ref 38 (4.9 × 10−3 min−1). The total time for complete conversion (180 min) in this study is smaller than that in ref 37 (≫700 min). 4.4. Mechanism Discussion. In order to distinguish the photo-Fenton process and the semiconductor-mediated photocatalytic reactions, we proposed an experiment which uses mesoporous 2-line ferrihydrite as the catalyst to degrade orange II without the presence of H2O2 under visible light in ambient air. The other conditions, such as catalyst concentration, dye concentration, and power of visible light, are kept the same as the reaction conditions used in this paper. The result is shown in Figure S6. The dye content only drops a little bit after 4 h irradiation with visible light. Further comparisons of the UV− vis spectra of the original dye, degraded dye in the presence of H2O2, and the dye after experiments without H2O2 are shown in Figure S6b. The dye in the experiment without H2O2 has not been degraded, showing different characteristics in UV−vis spectra (∼200 nm wavenumber) from the typical degraded dye. This means that the decrease of dye content is mainly caused by adsorption rather than degradation. The reason why this happens is that the mesoporous iron oxide catalyst is not able to generate O2• radicals under ambient air conditions under visible light exposure with TiO2. Therefore, the degradation of dyes on mesoporous iron oxides is contributed by the photoFenton reaction. The mechanism33 of Fenton reagents in aqueous solutions is shown as follows:

moles of dye degraded per unit surface area per hour, the calculated TOF number of mesoporous iron oxides is neutralized by the large surface areas of mesoporous iron oxides. This indicates that the major advantage of mesoporous iron oxides is their large surface area. Although the mesoporous structure in the catalysts could help to promote the diffusion of reactants, the transportation of reactants to the active sites may be still limited by the size of the dyes and morphology of the catalysts. The amorphous transition metal oxide has markedly different catalytic activity in photochemical reactions compared with highly crystallized transition metal oxides.64 The high photochemical activity of amorphous transition metal oxides could be due to highly disordered packing, oxygen vacancies, small particle sizes, and high surface areas. In this study, the high activity of mesoporous 2-line ferrihydrite (amorphous in high-angle PXRD) in dye degradation could be attributed to a variety of factors, such as large surface area, surface hydroxyl groups, acidic products, or possible additional adsorption sites. The rate constants of mesoporous iron oxides catalysts are comparable with those of iron oxides in the literature under mild reaction conditions, such as visible light and neutral pH. Table 4 summarizes the previous work regarding degradation of orange II using different iron oxides catalysts. All the literature data discussed here had the assumption of pseudo-first-order kinetics for the dye degradation process. Feng et al.36 reported degradation of orange II using different catalysts, including bentonite-clay-based Fe nanocomposite (Fe-B) and α-Fe2O3 in the presence of H2O2 at pH = 6.6 under UV light. Taking advantage of UV light, they obtained faster reaction kinetics, compared to this study, with their catalysts, for example, kapp = 0.077 with Fe-B nanocomposite and kapp = 0.063 with α-Fe2O3. The research from Du et al.57 included the effect of H2O2 on photodegradation orange II by different iron oxides, α-Fe2O3, γFe2O3, and Fe3O4, under UV light at a neutral pH = 6.5. Apparent rate constants on different iron oxides in their study were at least 50% smaller than the apparent rate constant shown in this study. As far as visible light is concerned, Cheng et al.37 synthesized Fe(III) iron exchange resin as a catalyst for the degradation of various dyes including orange II in the presence of H2O2 at pH = 6.0 under visible light. The catalyst degraded less than 5% orange II after 700 min irradiation whereas the mesoporous 2-line ferrihydrite in this study can degrade 100% orange II within 180 min. Du et al.38 synthesized clay-supported iron oxide with different ratios of clay and Fe2O3 as catalysts for cationic and anionic dyes including orange II at pH = 6.5 under visible light. Very good performance is observed by samples containing 20% Fe2O3 (20FeC in the literature). The 20FeC catalysts can degrade 65% orange II in

Fe2 + + H 2O2 → Fe3 + + OH• + OH− 10464

(chain initiation)


Article


Figure 14. Proposed mechanism of Fenton reactions on mesoporous 2-line ferrihydrite catalyst.

Fe3 + + H 2O2 ↔ Fe−OOH2 + + H+

at one position can form 2-diazenlyphenol and but-2-enedioic acid. Breakage of the naphthalene ring at another position can form benzoic acid and 2-hydroxy-3-(2-hydroxydiazen-1-yl)prop-2-enal. The possible breakage position is based on possible appearance of organic acid which causes the drop of pH value of the solution. The formation of organic acid can help the catalyst perform better by adjusting the pH value of the solution to around 4. This will further accelerate mesoporous 2-line ferrihydrite on discoloration and mineralization of orange II. The fragments with low m/z are observed in ESI/MS spectra.

Fe−OOH2 + → HO2• + Fe 2 + OH• + Fe2 + → OH− + Fe3 +

(chain termination)

In the heterogeneous photo-Fenton reaction, we believe that a similar mechanism also takes place on the surface of our mesoporous iron oxides. The hydroxyl radicals (•OH) and the hydroperoxyl radicals (•OOH) will react with dyes in solution. Bandara et al.66 showed that orange II could form Fe complexes with iron species in combination with H2O2, which is not preferred in basic media. During the degradation of azo dyes, the azo bonds −NN− are usually first broken down.67−69 The splitting of −NN− bonds will produce monosubstituted benzene and substituted naphthalene compounds, and then the opening of an aromatic ring of substituted naphthalene will form substituted benzene and other acid products.67,70−72 Similar results were reported in the literature. In Figure 12a, the fragments of orange II with m/z 327 starts to break down from 0.5 h. At the same time, the fragments with m/z 187 appears, which are products from splitting −NN− bonds. The fragments with m/z 187 are substituted naphthalene compounds. After 1 h, only weak peaks at m/z 327 are observed, and the fragments with m/z 187, which are the substituted naphthalene compounds, start to break down. A lot of fragments with m/z lower than 150 appear. Table S1 lists observed major peak positions and possible corresponding structures. The appearance of fragments with m/z 115 and 121 indicates that substituted naphthalene has been broken down by opening an aromatic ring. This step could form some organic acids. In addition, the pH drop of the reaction solutions also indicates the formation of acid products. So based on ESI/ MS and pH analyses, a proposed reaction pathway for orange II degradation over mesoporous 2-line ferrihydrite catalyst is shown in Figure 14. Degradation of orange II starts with azo bonds attacked by hydroxyl radicals to form benzenesolanate and (2-hydroxynaphthalen-1-yl)diazenolate. Then the naphthalene ring in the (2-hydroxynaphthalen-1-yl)diazenolate was attacked at different positions. Breakage of the naphthalene ring

5. CONCLUSIONS Mesoporous γ-Fe2O3 and Fe3O4 have been successfully synthesized by changing the reaction temperature of a sol− gel based on the inverse micelle method to a higher temperature (100 °C) and heating at 150 °C for 12 h and 250 °C for 4 h under an air atmosphere and nitrogen atmosphere, respectively. The mesoporous structures of mesoporous iron oxides are confirmed by low-angle XRD, BET, and TEM. Phases of mesoporous iron oxides are confirmed by high-angle XRD, Raman spectroscopy, and XANES. All the synthesized materials show vacancies in their own structure. The higher reaction temperature (100 °C) forms Fe(II) species. If as-made samples are heated under a nitrogen atmosphere, Fe(II) species can be preserved to form the Fe3O4 phase. If as-made samples are heated under air atmosphere, Fe(II) species will be oxidized to form the γ-Fe2O3 phase. Mesoporous γ-Fe2O3 and Fe3O4 have larger pore sizes and smaller surface areas compared with mesoporous 2-line ferrihydrite and α-Fe2O3. The sintering of nanoparticles will cause the expansion in pore size and decrease the surface area. The greatest adsorption capability of orange II is found in 2-line ferrihydrite (49.3 mg orange II/g cat.), which is 2 times larger than that of commercial Fe2O3. The trend of adsorption capacities of orange II on mesoporous iron oxides roughly follows the trend of surface area. The photocatalytic performances of mesoporous iron oxides are all better than that of commercial Fe2O3. The apparent rate constants clearly show that the best catalyst 2-line ferrihydrite has 6 times better 10465


Article


(7) Qiu, B.; Zhong, C.; Xing, M.; Zhang, J. Facile Preparation of CModified TiO2 Supported on MCF for High Visible-Light-Driven Photocatalysis. RSC Adv. 2015, 5, 17802−17808. (8) Qiu, B.; Xing, M.; Zhang, J. Mesoporous TiO2 Nanocrystals Grown in Situ on Graphene Aerogels for High Photocatalysis and Lithium-Ion Batteries. J. Am. Chem. Soc. 2014, 136, 5852−5855. (9) Zhang, Q.; Lu, X.; Chen, L.; Shi, Y.; Xu, T.; Liu, M. Mesoporous Flower-like a-Fe2O3 Nanoarchitectures: Facile Synthesis and Their Magnetic and Photocatalytic Properties. Mater. Lett. 2013, 106, 447− 451. (10) Deng, Y.; Cai, Y.; Sun, Z.; Liu, J.; Liu, C.; Wei, J.; Li, W.; Liu, C.; Wang, Y.; Zhao, D. Multifunctional Mesoporous Composite Microspheres with Well-Designed Nanostructure: A Highly Integrated Catalyst System. J. Am. Chem. Soc. 2010, 132, 8466−8473. (11) Pradhan, A. C.; Parida, K. M. Facile Synthesis of Mesoporous Composite Fe/Al2O3−MCM-41: An Efficient Adsorbent/Catalyst for Swift Removal of Methylene Blue and Mixed Dyes. J. Mater. Chem. 2012, 22, 7567. (12) Liu, Y. M.; Xu, J.; He, L.; Cao, Y.; He, H. Y.; Zhao, D. Y.; Zhuang, J. H.; Fan, K. N. Facile Synthesis of Fe-Loaded Mesoporous Silica by a Combined Detemplation-Incorporation Process through Fenton’s Chemistry. J. Phys. Chem. C 2008, 112, 16575−16583. (13) Zhong, X.; Royer, S.; Zhang, H.; Huang, Q.; Xiang, L.; Valange, S.; Barrault, J. Mesoporous Silica Iron-Doped as Stable and Efficient Heterogeneous Catalyst for the Degradation of C.I. Acid Orange 7 Using Sono-Photo-Fenton Process. Sep. Purif. Technol. 2011, 80, 163− 171. (14) Wang, M.; Shu, Z.; Zhang, L.; Fan, X.; Tao, G.; Wang, Y.; Chen, L.; Wu, M.; Shi, J. Amorphous Fe2+-Rich FeOx Loaded in Mesoporous Silica as a Highly Efficient Heterogeneous Fenton Catalyst. Dalton Trans. 2014, 43, 9234−9241. (15) Kim, S.; Durand, P.; Roques-Carmes, T.; Eastoe, J.; Pasc, A. Metallo-Solid Lipid Nanoparticles as Colloidal Tools for Meso− Macroporous Supported Catalysts. Langmuir 2015, 31, 1842−1849. (16) Zhou, X.; Cui, X.; Chen, H.; Zhu, Y.; Song, Y.; Shi, J. A Facile Synthesis of Iron Functionalized Hierarchically Porous ZSM-5 and Its Visible-Light Photocatalytic Degradation of Organic Pollutants. Dalton Trans. 2013, 42, 890−893. (17) Jiao, F.; Harrison, A.; Jumas, J. C.; Chadwick, A. V.; Kockelmann, W.; Bruce, P. G. Ordered Mesoporous Fe2O3 with Crystalline Walls. J. Am. Chem. Soc. 2006, 128, 5468−5474. (18) Jiao, F.; Jumas, J. C.; Womes, M.; Chadwick, A. V.; Harrison, A.; Bruce, P. G. Synthesis of Ordered Mesoporous Fe3O4 and γ-Fe2O3 with Crystalline Walls Using Post-Template Reduction/Oxidation. J. Am. Chem. Soc. 2006, 128, 12905−12909. (19) Jiao, F.; Bruce, P. G. Two- and Three-Dimensional Mesoporous Iron Oxides with Microporous Walls. Angew. Chem., Int. Ed. 2004, 43, 5958−5961. (20) Mitra, A.; Vázquez-Vázquez, C.; López-Quintela, M. A.; Paul, B. K.; Bhaumik, A. Soft-Templating Approach for the Synthesis of High Surface Area and Superparamagnetic Mesoporous Iron Oxide Materials. Microporous Mesoporous Mater. 2010, 131, 373−377. (21) Kim, T. W.; Ha, H. W.; Paek, M. J.; Hyun, S. H.; Baek, I. H.; Choy, J. H.; Hwang, S. J. Mesoporous Iron Oxide-Layered Titanate Nanohybrids: Soft-Chemical Synthesis, Characterization, and Photocatalyst Application. J. Phys. Chem. C 2008, 112, 14853−14862. (22) Panda, N.; Sahoo, H.; Mohapatra, S. Decolourization of Methyl Orange Using Fenton-like Mesoporous Fe2O3-SiO2 Composite. J. Hazard. Mater. 2011, 185, 359−365. (23) Lucas, M. S.; Peres, J. A. Decolorization of the Azo Dye Reactive Black 5 by Fenton and Photo-Fenton Oxidation. Dyes Pigm. 2006, 71, 236−244. (24) Hsueh, C. L.; Huang, Y. H.; Wang, C. C.; Chen, C. Y. Degradation of Azo Dyes Using Low Iron Concentration of Fenton and Fenton-like System. Chemosphere 2005, 58, 1409−1414. (25) Herney-Ramirez, J.; Vicente, M. A.; Madeira, L. M. Heterogeneous Photo-Fenton Oxidation with Pillared Clay-Based Catalysts for Wastewater Treatment: A Review. Appl. Catal., B 2010, 98, 10−26.

activity than that of commercial Fe2O3. All the reaction solutions catalyzed by mesoporous iron oxides have pH drops during the reactions. In particular, the pH of the solutions catalyzed by 2-line ferrihydrite drops to 4.4, the largest drop among mesoporous iron oxides. The stability test shows the mesoporous iron oxides can maintain the performance until the fourth cycle. The meso structure of the catalyst is reserved after the fourth run of reaction. According to reaction intermediates monitored by ESI/MS, 2-line ferrihydrite catalyst can start degrading aromatic rings to smaller species at the beginning of the reaction whereas in the solution catalyzed by commercial Fe2O3 a large amount of aromatic species remain until the end of reaction.

■

ASSOCIATED CONTENT

S Supporting Information *

TG-MS spectra of mesoporous 2-line ferrihydrite, α-Fe2O3, γFe2O3, Fe3O4, and commercial Fe2O3, IR spectra of orange II absorbed on mesoporous 2-line ferrihydrite, α-Fe2O3, γ-Fe2O3, Fe3O4, and commercial Fe2O3 at the range 3000−4000 and 500−1800 cm−1, major peaks on ESI/MS spectra and possible structures, the HRTEM image of mesoporous 2-line ferrihydrite after the fourth cycle of reaction, picture of mesoporous iron oxides, mesoporous 2-line ferrihydrite catalyzed dye degradation without H2O2 under visible light, calculation of electrical energy efficiency. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.5b02057.

■

AUTHOR INFORMATION

Corresponding Author

*(S.L.S.) Phone (860)-486-2797; Fax (860)-486-2981; e-mail [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was funded by the US Department of Energy. The authors thank Dr. Frank Galasso for correction of manuscript and thank Dr. Menka Jain for magnetic measurements.

■

REFERENCES

(1) Wen, Z.; Zhang, Y.; Dai, C.; Chen, B.; Guo, S.; Yu, H.; Wu, D. Synthesis of Ordered Mesoporous Iron Manganese Bimetal Oxides for Arsenic Removal from Aqueous Solutions. Microporous Mesoporous Mater. 2014, 200, 235−244. (2) Zhang, Y.; Li, L.; Ma, W.; Zhang, Y.; Yu, M.; Guo, J.; Lu, H.; Wang, C. Two-in-One Strategy for Effective Enrichment of Phosphopeptides Using Magnetic Mesoporous γ-Fe2O3 Nanocrystal Clusters. ACS Appl. Mater. Interfaces 2013, 5, 614−621. (3) Wen, Z.; Dai, C.; Zhu, Y.; Zhang, Y. Arsenate Removal from Aqueous Solutions Using Magnetic Mesoporous Iron Manganese Bimetal Oxides. RSC Adv. 2015, 5, 4058−4068. (4) Ivanova, A.; Fominykh, K.; Fattakhova-Rohlfing, D.; Zeller, P.; Döblinger, M.; Bein, T. Nanocellulose-Assisted Formation of Porous Hematite Nanostructures. Inorg. Chem. 2015, 54, 1129−1135. (5) Yi, D. K.; Lee, S. S.; Papaefthymiou, G. C.; Ying, J. Y. Nanoparticle Architectures Templated by SiO2/Fe2O3 Nanocomposites. Chem. Mater. 2006, 18, 614−619. (6) Wu, Z.; Li, W.; Webley, P. A.; Zhao, D. General and Controllable Synthesis of Novel Mesoporous Magnetic Iron Oxide@carbon Encapsulates for Efficient Arsenic Removal. Adv. Mater. 2012, 24, 485−491. 10466


Article

The Journal of Physical Chemistry C (26) 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. (27) 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, 51, 178−182. (28) Dong, Y.; Dong, W.; Cao, Y.; Han, Z.; Ding, Z. Preparation and Catalytic Activity of Fe Alginate Gel Beads for Oxidative Degradation of Azo Dyes under Visible Light Irradiation. Catal. Today 2011, 175, 346−355. (29) Wu, K.; Xie, Y.; Zhao, J.; Hidaka, H. Photo-Fenton Degradation of a Dye under Visible Light Irradiation. J. Mol. Catal. A: Chem. 1999, 144, 77−84. (30) Park, H.; Choi, W. Visible Light and Fe(III)-Mediated Degradation of Acid Orange 7 in the Absence of H2O2. J. Photochem. Photobiol. A: Chem. 2003, 159, 241−247. (31) Fernandez, J.; Dhananjeyan, M. R.; Kiwi, J.; Senuma, Y.; Hilborn, J. Evidence for Fenton Photoassisted Processes Mediated by Encapsulated Fe Ions at Biocompatible pH Values. J. Phys. Chem. B 2000, 104, 5298−5301. (32) Yuranova, T.; Enea, O.; Mielczarski, E.; Mielczarski, J.; Albers, P.; Kiwi, J. Fenton Immobilized Photo-Assisted Catalysis through a Fe/C Structured Fabric. Appl. Catal. B: Environ. 2004, 49, 39−50. (33) Neyens, E.; Baeyens, J. A Review of Classic Fenton’s Peroxidation as an Advanced Oxidation Technique. J. Hazard. Mater. 2003, 98, 33−50. (34) 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. (35) Kremer, M. L. The Fenton Reaction. Dependence of the Rate on pH. J. Phys. Chem. A 2003, 107, 1734−1741. (36) Feng, J.; Hu, X.; Po, L. Y. Discoloration and Mineralization of Orange II Using Different Heterogeneous Catalysts Containing Fe: A Comparative Study. Environ. Sci. Technol. 2004, 38, 5773−5778. (37) Cheng, M.; Ma, W.; Li, J.; Huang, Y.; Zhao, J.; Wen, Y. X.; Xu, Y. Visible-Light-Assisted Degradation of Dye Pollutants over Fe(III)Loaded Resin in the Presence of H2O2 at Neutral pH Values. Environ. Sci. Technol. 2004, 38, 1569−1575. (38) Du, W.; Sun, Q.; Lv, X.; Xu, Y. Enhanced Activity of Iron Oxide Dispersed on Bentonite for the Catalytic Degradation of Organic Dye under Visible Light. Catal. Commun. 2009, 10, 1854−1858. (39) Poyraz, A. S.; Kuo, C.-H.; Biswas, S.; King’ondu, C. K.; Suib, S. L. A General Approach to Crystalline and Monomodal Pore Size Mesoporous Materials. Nat. Commun. 2013, 4, 2952. (40) Luo, Z.; Poyraz, A. S.; Kuo, C.; Miao, R.; Meng, Y.; Chen, S.-Y.; Jiang, T.; Wenos, C.; Suib, S. L. Crystalline Mixed Phase (anatase/ rutile) Mesoporous Titanium Dioxides for Visible Light Photocatalytic Activity. Chem. Mater. 2015, 27, 6−17. (41) Song, W.; Poyraz, A. S.; Meng, Y.; Ren, Z.; Chen, S. Y.; Suib, S. L. Mesoporous Co3O4 with Controlled Porosity: Inverse Micelle Synthesis and High-Performance Catalytic Co Oxidation at −60 °C. Chem. Mater. 2014, 26, 4629−4639. (42) Poyraz, A. S.; Song, W.; Kriz, D.; Kuo, C. H.; Seraji, M. S.; Suib, S. L. Crystalline Mesoporous K2‑xMn8O16 and E-MnO2 by Mild Transformations of Amorphous Mesoporous Manganese Oxides and Their Enhanced Redox Properties. ACS Appl. Mater. Interfaces 2014, 6, 10986−10991. (43) Faria, D. L. A. De; Silva, S. V.; de Oliveira, M. T. Raman Microspectroscopy of Some Iron Oxides and Oxyhydroxides. J. Raman Spectrosc. 1997, 28, 873−878. (44) Shebanova, O. N.; Lazor, P. Raman Study of Magnetite (Fe3O4): Laser-Induced Thermal Effects and Oxidation. J. Raman Spectrosc. 2003, 34, 845−852. (45) Du, S.; Valla, J. A.; Bollas, G. M. Characteristics and Origin of Char and Coke from Fast and Slow, Catalytic and Thermal Pyrolysis of Biomass and Relevant Model Compounds. Green Chem. 2013, 15, 3214.

(46) Genuino, H. C.; Hamal, D. B.; Fu, Y. J.; Suib, S. L. Synergetic Effects of Ultraviolet and Microwave Radiation for Enhanced Activity of TiO2 Nanoparticles in Degrading Organic Dyes Using a Continuous-Flow Reactor. J. Phys. Chem. C 2012, 116, 14040−14051. (47) Drits, V. A.; Sakharov, B. A.; Salyn, A. L.; Manceau, A. Structural Model for Ferrihydrite. Clay Miner. 1993, 28, 185−207. (48) Schwertmann, U.; Cornell, R. M. Iron Oxides in the Laboratory. Preparation and Characterization, 2000. (49) Michel, F. M.; Ehm, L.; Liu, G.; Han, W. Q.; Antao, S. M.; Chupas, P. J.; Lee, P. L.; Knorr, K.; Eulert, H.; Kim, J.; et al. Similarities in 2- and 6-Line Ferrihydrite Based on Pair Distribution Function Analysis of X-Ray Total Scattering. Chem. Mater. 2007, 19, 1489− 1496. (50) Lübbe, M.; Gigler, A. M.; Stark, R. W.; Moritz, W. Identification of Iron Oxide Phases in Thin Films Grown on Al2O3(0001) by Raman Spectroscopy and X-Ray Diffraction. Surf. Sci. 2010, 604, 679−685. (51) Bellot-Gurlet, L.; Neff, D.; Réguer, S.; Monnier, J.; Saheb, M.; Dillmann, P. Raman Studies of Corrosion Layers Formed on Archaeological Irons in Various Media. J. Nano Res. 2009, 8, 147−156. (52) Zhang, L.; Wang, W.; Zhou, L.; Shang, M.; Sun, S. Fe3O4 Coupled BiOCl: A Highly Efficient Magnetic Photocatalyst. Appl. Catal. B: Environ 2009, 90, 458−462. (53) Wang, W.; Xu, Y.; Wang, D. I. C.; Li, Z. Recyclable Nanobiocatalyst for Enantioselective Sulfoxidation: Facile Fabrication and High Performance of Chloroperoxidase-Coated Magnetic Nanoparticles with Iron Oxide Core and Polymer Shell. J. Am. Chem. Soc. 2009, 131, 12892−12893. (54) Thomas, J. M.; Thomas, W. J. Principles and Practice of Heterogeneous Catalysis; Wiley: New York, 1997. (55) Zhong, L.-S.; Hu, J.-S.; Liang, H.-P.; Cao, A.-M.; Song, W.-G.; Wan, L.-J. Self-Assembled 3D Flowerlike Iron Oxide Nanostructures and Their Application in Water Treatment. Adv. Mater. 2006, 18, 2426−2431. (56) Cui, H.; Liu, Y.; Ren, W. Structure Switch between α-Fe2O3, γFe2O3 and Fe3O4 during the Large Scale and Low Temperature SolGel Synthesis of Nearly Monodispersed Iron Oxide Nanoparticles. Adv. Powder Technol. 2013, 24, 93−97. (57) Du, W.; Xu, Y.; Wang, Y. Photoinduced Degradation of Orange II on Different Iron (Hydr)oxides in Aqueous Suspension: Rate Enhancement on Addition of Hydrogen Peroxide, Silver Nitrate, and Sodium Fluoride. Langmuir 2008, 24, 175−181. (58) Herrera, F.; Lopez, A.; Mascolo, G.; Albers, P.; Kiwi, J. Catalytic Combustion of Orange II on Hematite Surface Species Responsible for the Dye Degradation. Appl. Catal. B: Environ. 2001, 29, 147−162. (59) Saha, B.; Das, S.; Saikia, J.; Das, G. Preferential and Enhanced Adsorption of Different Dyes on Iron Oxide Nanoparticles: A Comparative Study. J. Phys. Chem. C 2011, 115, 8024−8033. (60) Joseph, J.; Jemmis, E. D. Red-, Blue-, or No-Shift in Hydrogen Bonds: A Unified Explanation. J. Am. Chem. Soc. 2007, 129, 4620− 4632. (61) Bandara, J.; Mielczarski, J. A.; Kiwi, J. 1. Molecular Mechanism of Surface Recognition. Azo Dyes Degradation on Fe, Ti, and Al Oxides through Metal Sulfonate Complexes. Langmuir 1999, 15, 7670−7679. (62) Bandara, J.; Mielczarski, J. A.; Kiwi, J. 2. Photosensitized Degradation of Azo Dyes on Fe, Ti, and Al Oxides. Mechanism of Charge Transfer during the Degradation. Langmuir 1999, 15, 7680− 7687. (63) Soon, A. N.; Hameed, B. H. Heterogeneous Catalytic Treatment of Synthetic Dyes in Aqueous Media Using Fenton and Photo-Assisted Fenton Process. Desalination 2011, 269, 1−16. (64) Iyer, A.; Del-Pilar, J.; King’ondu, C. K.; Kissel, E.; Garces, H. F.; Huang, H.; El-Sawy, A. M.; Dutta, P. K.; Suib, S. L. Water Oxidation Catalysis Using Amorphous Manganese Oxides, Octahedral Molecular Sieves (OMS-2), and Octahedral Layered (OL-1) Manganese Oxide Structures. J. Phys. Chem. C 2012, 116, 6474−6483. (65) Perez, M.; Torrades, F.; Garcia-Hortal, J. A.; Domenech, X.; Peral, J. Removal of Organic Contaminants in Paper Pulp Treatment 10467


Article

The Journal of Physical Chemistry C Effluents under Fenton and Photo-Fenton Conditions. Appl. Catal. B: Environ. 2002, 36, 63−74. (66) Bandara, J.; Morrison, C.; Kiwi, J.; Pulgarin, C.; Peringer, P. Degradation/decoloration of Concentrated Solutions of Orange II. Kinetics and Quantum Yield for Sunlight Induced Reactions via Fenton Type Reagents. J. Photochem. Photobiol. A: Chem. 1996, 99, 57−66. (67) Gutowska, A.; Kałuzna-Czaplińska, J.; Józw ́ iak, W. K. Degradation Mechanism of Reactive Orange 113 Dye by H2O2/Fe2+ and Ozone in Aqueous Solution. Dyes Pigm. 2006, 74, 41−46. (68) Abbott, L. C.; Batchelor, S. N.; Lindsay Smith, J. R.; Moore, J. N. Reductive Reaction Mechanisms of the Azo Dye Orange II in Aqueous Solution and in Cellulose: From Radical Intermediates to Products. J. Phys. Chem. A 2009, 113, 6091−6103. (69) Rahmani, A.; Zarrabi, M. Degradation of Azo Dye Reactive Black 5 and Acid Orange 7 by Fenton-Like Mechanism. Iran. J. Chem. Eng. 2010, 7, 87−94. (70) Stylidi, M.; Kondarides, D. I.; Verykios, X. E. Visible LightInduced Photocatalytic Degradation of Acid Orange 7 in Aqueous TiO2 Suspensions. Appl. Catal. B: Environ. 2004, 47, 189−201. (71) Chen, X.; Chen, J.; Qiao, X.; Wang, D.; Cai, X. Performance of Nano-Co3O4/Peroxymonosulfate System: Kinetics and Mechanism Study Using Acid Orange 7 as a Model Compound. Appl. Catal. B: Environ. 2008, 80, 116−121. (72) Bilgi, S.; Demir, C. Identification of Photooxidation Degradation Products of C.I. Reactive Orange 16 Dye by Gas ChromatographyMass Spectrometry. Dyes Pigm. 2005, 66, 69−76.

10468


Synthesis of Mesoporous Iron Oxides by an Inverse Micelle Method

Recommend Documents