ARTICLE pubs.acs.org/JPCC
Crystallization-Induced Top-Down Wormlike Hierarchical Porous r-Fe2O3 Self-Assembly Ramakrishnan Amutha,†,‡ Manickavachagam Muruganandham,*,‡,§ Marappan Sathish,^ Sambandam Akilandeswari,† Rominder P. S. Suri,§ Eveliina Repo,‡ and Mika Sillanp€a€a‡ †
Department of Physics, Annamalai University, Annamalainagar 602 008, India Laboratory of Green Chemistry, Faculty of Technology, Lappeenranta University of Technology, Patteristonkatu 1, FI-50100 Mikkeli, Finland § Water & Environmental Technology (WET) Center, Department of Civil and Environmental Engineering, Temple University, Philadelphia, Pennsylvania 19122, United States ^ Institute of Multidisciplinary Research for Advanced Materials, Tohoku University, Sendai 980-8577, Japan ‡
bS Supporting Information ABSTRACT: In the first large-scale investigation we found a novel crystallization-induced fabrication of top-down wormlike hierarchical porous hematite self-assembly (WHHS). WHHS results in the crystallization of stable amorphous iron oxide microbundles at higher temperatures (400 C). Oxalic acid was successfully used as both a complexing and an etching agent for the preparation; etching induced a morphological transition from microbundles to microsheets. The preparation method of the ferric oxalate complex is a crucial step for the formation of WHHS. The synthesized ferric oxalate and WHHS was characterized using X-ray diffraction (XRD), field emission scanning electron microscopy (FE-SEM), high-resolution transmission electron microscopy (HR-TEM), Raman spectra, FTIR, and nitrogen adsorption analysis. We investigated the influence of decomposition temperatures and concentration of both oxalate and iron precursors on the formation of WHHS and its surface properties as well as the magnetic and electrochemical properties of the synthesized WHHS. The WHHS showed weak ferromagnetic properties. We also examined a plausible WHHS formation mechanism.
1. INTRODUCTION Hematite (R-Fe2O3), an n-type semiconductor (Eg = 2.1 eV), is an attractive multifunctional material due to its low cost, good stability, high resistance to corrosion, nontoxicity, and environmentally friendly properties.1 Its potential applications are various such as catalysis, gas sensors, magnetic devices, photoelectrodes, pigments, and adsorbents.24 Scientists are continuously exploring new methodologies to fabricate desired morphology and surface properties for these applications. The conventional simple metalorganic complex is a significant asset for material fabrication. Thus, oxalic acid has been successfully used as a coordination reagent (ligand) to prepare various nano/micromaterials.1,5,6 Du et al. reported oxalic acidassisted fabrication of Fe2O3 hollow urchins by hydrothermal treatment of a Fe(NO3)3oxalic acid coordination complex without using any other reagents.1 Most of these synthetic methods are based on hydrothermal methods, which limit largescale and industrial production applications. One of the attractive methods for iron oxide preparation is by solid-state thermal decomposition of the iron oxalate complex at suitable temperature. Most studies were, however, conducted using commercially available iron(II) and iron(III) oxalates under different experimental conditions, but none have prepared nano/microstructured iron oxide synthesis.710 r 2011 American Chemical Society
The wormlike morphological surfactants, micelles, and polymers were successfully used as templates or morphology directing agents to fabricate materials with wormlike morphology.1113 The wormlike R-Fe2O3 film has previously been successfully synthesized using pluronic copolymer as a structure directing agent.14 Such synthesis requires template removal, however, which is an additional step. On the other hand, oxalic acid is considered the most effective iron oxide contaminant dissolving agent in minerals.15,16 Chen et al. successfully prepared porous R-Fe2O3 nanodisks via oxalic acid etching.17 Therefore, oxalic acid could possibly be used as both a complexing and an etching agent for iron oxide fabrication. Development of simple, reliable synthetic methods for hierarchically selfassembled architectures with controlled morphologies is still a major challenge. Generally, crystallization of amorphous materials results in a more stable crystalline state. Such crystallization-induced nano/ microstructured fabrication is a novel approach, however, and has not been studied. We recently reported the synthesis of stable amorphous mesoporous iron oxides (AMIOs) using ferric oxalate complex under open atmospheric conditions at low Received: January 11, 2011 Revised: February 26, 2011 Published: March 21, 2011 6367
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The Journal of Physical Chemistry C temperature.18 We now report crystallization-induced oxalic acid-assisted (complexing and etching agent) fabrication of hematite microbundles and microsheets without using any templates and solvents. This hierarchical porous hematite microsheets and microbundles morphology has not been previously reported; we thereby introduce a new morphology in the iron oxide family. We also discuss the influence of various experimental parameters on the formation of WHHS and its mechanisms as well as the magnetic, electrochemical properties, and adsorption of the synthesized WHHS.
2. EXPERIMENTAL SECTION Ferric nitrate nonahydrate, anhydrous ferric chloride, iron(III) oxalate hexahydrate, and oxalic acid were purchased from Aldrich Chemical Co. Ltd. (Helsinki, Finland). All chemicals were of analytical grade and used without further purification. For all experimental work Milli Q-Plus water (resistance = 18.2 MΩ) was used. The synthesis of iron oxide involves two steps. First, the preparation of a ferric nitrateoxalic acid complex (ferric oxalate) by mixing an equal volume (100 mL) of 0.1 M ferric nitrate nonahydrate with the required concentration of oxalic acid. We prepared three types of ferric oxalate solids using 0.17, 0.24, and 0.4 M oxalic acid resulting in 2.5, 3.25, and 5 g of solids, hereafter denoted as Feox-1, Feox-2, and Feox-3. During magnetic stirring, conducted in the dark, oxalic acid was added drop by drop (3045 min) into a ferric nitrate solution. The solution was stirred for several hours and then evaporated on the hot plate until becoming a dry greenish solid. Care was taken at the final stages of solvent evaporation allowing performance of nitric acid evaporation in the fume cupboard. In the second step, prepared ferric oxalate complex is decomposed at a desired temperature and time under open atmospheric conditions. After decomposition, the oven was cooled to room temperature, and the samples were washed with plenty of water and ethanol and dried at 120 C for 2 h. The X-ray diffraction (XRD) patterns were recorded using an X’Pert PRO PAN analytical diffractometer, scanning angles of 10 to 100. High-resolution transmission electron microscope images were recorded using FE-TEM (Philips CM-200 FEG - (S) TEM Super Twin). Samples for HR-TEM were prepared by ultrasonically dispersing the catalyst into ethanol and then placing a drop of this suspension onto a carbon-coated copper grid and air-drying. The working voltage of TEM was 80 kV. The morphology of the catalyst was examined using a Hitachi S-4100 scanning electron microscope (SEM). Prior to SEM measurements, the samples were mounted on a carbon platform, which was then coated by platinum using a magnetron sputter for 10 min. The plate containing the sample was placed in the SEM for analysis at desired magnifications. The surface properties (surface area, pore size, and pore volume) of the iron oxides were measured with a Autosorb-1-C surface area and pore size analyzer (Quantachrome UK). All the samples were degassed overnight at 120 C and then analyzed for nitrogen adsorption. The metal concentrations in the filtrates were analyzed by an inductively coupled plasma optical atomic emission spectrometry (ICP-OES) model iCAP 6300 (Thermo Electron Corp.). 3. RESULTS AND DISCUSSION We used a new synthetic methodology for WHHS fabrication that utilized thermal decomposition of ferric oxalate salt. We also used a new synthetic method for the ferric oxalate preparation and utilized the excess oxalate ion as an etching agent during the
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hematite synthesis (decomposition process). Since we employed a new synthetic method, it is important to study structure and composition of synthesized ferric oxalate salt. The three ferric oxalate complexes were not precipitated in solvent, and we did not study the complex structure in solution phase. All studies were conducted in the solid state after solvent evaporation. Though many studies have focused on various aspects of iron oxalate complex, ours studies differ since we used it as a precursor for WHHS synthesis.710,18 In our earlier studies, we reported FTIR spectra of Feox-1 and compared them with commercial iron(III) oxalate hexahydrates.18 We used various oxalic acid concentrations for ferric oxalate preparation, since it is necessary to study how oxalic acid is coordinated with iron at various concentrations. The FTIR spectra of three ferric oxalate solids were analyzed and compared with commercial iron(III) oxalate hexahydrate. Figure 1 showed the FTIR spectra of all three ferric oxalate solids, which were similar to commerical iron(III) oxalate hexahydrate, and FTIR values are shown in Table S1. All FTIR spectra showed two strong absorbance bands at 1240 and 1600 cm1, characteristic of ν(CC) and ν(CO) stretching vibrations. The v(CC) mode has also been clearly noted at 758 and 816 cm1. The νOH stretching vibration is noted at 3540 cm1, confirming the hydrated form of ferric oxalates. The other characteristic peak presence at 1730 cm1 corresponds to the terminal carbonyl group. All the observed peak values resemble those of earlier reports assigned for iron oxalates.19,20 Edward et al. studied the possible structure of iron(III) oxalate complex and suggested that the iron(III) oxalate complex formed by bidentate oxalate coordination at low oxalic acid concentration (1:2) and hexacoordinated at 1:3 iron:oxalate ratio, as shown in the Supporting Information Figure S1. Therefore, a different type of iron:oxalate coordination complex may exist at various oxalic acid concentrations.20 These conclusions are based on the presence of terminal, noncoordinated CO bonds at 1730 cm1 in the FTIR spectra of all three ferric oxalate salts. Interestingly, the Raman spectrum of Feox-1 is quite different from the other two ferric oxalates (Feox-2 and Feox-3) as shown in Figure S2 of the Supporting Information. The Feox-1 showed strong bands at 1475 and 1705 cm1, which were expected to be the v(CO) mode. The v(CC) mode apeared at 840 and 920 cm1, and the v(CO2) mode appeared at 108, 180, 480, and 538 cm1. A weak OH mode observed at 3400 cm1 also reflects the hydrated form of ferric oxalates. The peak intensities of the
Figure 1. FTIR spectra of ferric oxalates prepared by using oxalic acid concentration of (A) 0.17 M, (B) 0.24 M, (C) 0.4 M, and (D) commercial ferric oxalate hexahydrates (Aldrich). 6368
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Figure 2. TGA analysis of ferric oxalate under air atmospheric conditions. (A) Feox-1, (B) Feox-2 and (C) Feox-3.
synthesized iron oxalates are different from the commercial iron(III) oxalate hexahydrates. The Raman spectra of ferric oxalate, however, correlate well with the FTIR spectra and earlier reports assigned for iron oxalates.19 The Feox-2 and Feox-3 showed very similar Raman spectra, indicating a similar complex structure. Thus, all the intense peaks are noted at a lower wavenumber, and no charecteristic peaks are noted at a longer wavenumber, except one weak mode at 1300 cm1. Generally, v(FeO) and δ(FeO2) bands appeared at lower wavenumbers, and peaks appeared at 222 cm1. The strong peak appearing at 292, 400, and 485 cm1 is assigned for the δ(CO2) mode. The aforementioned discussion clearly indicated that the structure of Feox-1 differs from the other two ferric oxalates. We earlier reported that in thermal analysis of ferric oxalate, prepared using 0.17 M, amorphous mesoporous iron oxide (Fe2O3) formed at 200 C.18 We used various oxalic acid concentrations for ferric oxalate preparation; however, therefore it is necessary to study and compare the thermal analysis of all three ferric oxalates. The thermal decomposition patterns of all three synthesized ferric oxalates, under air atmosphere, are shown in Figure 2. The TGA and DTG analysis indicated that the decomposition of all three ferric oxalates results in iron oxide as a final decomposition product. Moreover, all three decomposition patterns revealed that the decomposition proceeds without producing stable intermediates. A small difference in the decomposition pattern occurs between the three ferric oxalate salts and might be due to the use of various oxalic acid concentrations. The Feox-1 only showed a 2% weight loss at the early stages of decomposition process, and the total weight loss of 64.77% is in good agreement with the theoretical weight loss expected for the formation of Fe2O3. Feox-2 and Feox3 showed about 13% and 9.9% of weight loss at the early stages and 78.87% and 57.83% total weight loss. The higher percentage of weight loss at the early stages may be due to the presence of hydrated oxalic acid, and the difference in total weight loss is due to presence of FeO phase along with Fe2O3 phase at 200 C. FeO is converted into Fe2O3 during the thermal decomposition process, however, as evident from the appearance of the hematite phase at 400 C. The formation of hematite in the thermal decomposition of all three ferric oxalates at 400 C was confirmed by EDX, XRD, XPS, and Raman spectra analysis. The Raman spectra of all three RFe2O3 formed at 400 C were analyzed and compared with
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Figure 3. Raman spectra of WHHS synthesized at 400 C for 2 h by thermal decomposition of (A) Feox-1, (B) Feox-2, and (C) Feox-3, and (D) commercial hematite.
Figure 4. XRD pattern of hematite synthesized at 400 C by using ferric oxalates prepared: (A) Feox-1, (B) Feox-2, and (C) Feox-3.
commercial R-Fe2O3. The Raman spectra of all three hematites are consistent with commercial hematite, shown in Figure 3. Generally, hematite yields seven phonon lines (two A1g modes, 225 and 498 cm1, and five Eg modes, 247, 293, 299, 412, and 613 cm1). All the peak values in Table S1 are close to the abovementioned hematite peak values. These results confirmed formation of R-Fe2O3 in the decomposition process. Thus, compared with the bulk material, a small shift toward a lower wavenumber was observed, which might be related with the size of the particles. Earlier studies noted similar observations.21,22 The XRD pattern of hematite (R-Fe2O3) formed in the thermal decomposition of three ferric oxalates under open air atmosphere is shown in Figure 4. All XRD peaks can be indexed to the rhombo-centered R-Fe2O3 (JCPDS 80-2377) with lattice parameters of a = 0.5035 nm and c = 1.374 nm, indicating high purity and good crystallinity of the synthesized hematite. No other impurities were noted in XRD. We have examined the morphology of three ferric oxalates by FESEM, and the results clearly indicated that the ferric oxalates did not possess any definite morphology. A representative FESEM picture of Feox-3 is shown in Figure S3a. Thermal decomposition of the three ferric oxalates under open atmosphere at 400 C results in WHHS as shown in Figure 5. The 6369
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The Journal of Physical Chemistry C FESEM results clearly indicated formation of aggeragates of free, homogeneous hierarchical porous microbundles. The thickness and length of the microsheets were not uniform; the wormlike hierarchical porous surface structure was homogeneously distributed from top to bottom. One major difference among the three hematite microbundles is their size. Thus, the microbundles size decreased when the oxalic acid concentration is increased for ferric oxalate synthesis. The decrease in microbundle size is due to oxalic acid induced etching during the decomposition process, and the details are discussed in the mechanism section. It should be mentioned that as the oxalic acid concentration increases (0.5 and 0.8 M) for ferric oxalate preparation, it induces an adverse effect on WHHS formation, evident in the FESEM picture presented in Supporting Information Figure S5. Therefore, the oxalic acid concentration plays a major role on the formation of WHHS.
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In order to understand the formation of the wormlike structure, we performed HR-TEM analysis. The HR-TEM analysis clearly indicated that the formation of the wormlike morphology was assembled by oriented attachment with the assistance of thermal energy shown in Figure 6a,b. Moreover, Figure 6a displays the HRTEM lattice fringes of three different worm segments with a spacing of 0.25 nm, corresponding to the lattice spacing (d-value) of the (110) hematite planes. These results confirm the oriented attachment mechanism for the formation of wormlike hierarchical porous hematite. Thus, it is interesting to note that the wormlike hematite is self-assembled by a series of segments, quite similar to worm morphology. We have not used any templates or catalysts for the fabrication of either ferric oxalate or hematite synthesis (during the decomposition process). Therefore, as discussed before, the WHHS was formed by a self-assembled process. Various controlled
Figure 5. FESEM pictures of WHHS formed at 400 C by thermal decomposition of (A) Feox-1, (B) Feox-2, and (C) Feox-3. 6370
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Figure 6. HRTEM analysis of WHHS (A) images and (B) lattice fringes.
experiments were performed to understand the mechanism of WHHS formation . All controlled experiments used 0.17 M oxalic acid, while keeping ferric nitrate at 0.1 M, and decomposition was carried out at 400 C for 2 h. Interestingly, except for ferric nitrate, no other iron(III) salts, such as ferric chloride or ferric sulfate, facilitated WHHS as shown in the Supporting Information Figure S3b. Similarly, only oxalic acid, not sodium oxalate or potassium oxalates, facilitated WHHS, resulting in hematite aggregates as shown in Figure S3c,d. While keeping the ferric nitrate to oxalic acid molar ratio at 1:1.7, experiments were performed by changing the ferric nitrate concentration from 0.1 to 1 M. The decomposition results clearly indicated that increasing the ferric nitrate concentration increases the microbundles shape of aggregates, shown in Figure S4. A porous surface structure is noted at both concentrations but is not similar to one prepared using low initial concentration. Controlled experiments also indicated that the method of ferric oxalate salt preparation is a crucial step in WHHS formation. We also performed two methods of solvent evaporation: one at room temperature and another on a hot plate. Thermal decomposition of the aforementioned two ferric oxalate salts indicated that the latter is suitable for WHHS synthesis and the former yields aggregates (Figure S6a). Similarly, decomposition of commercial iron(III) oxalate hexahydrate did not yield WHHS (Figure S6b). Another important experimental parameter influencing the formation of WHHS is decomposition temperature. The decomposition temperature depended on morphology, and degree of crystallization was studied by FESEM and XRD. Thus, decomposition results indicated that crystallized WHHS is formed above 300 C. Moreover, the degree of wormlike hierarchical surface structure increases when the decomposition temperature increases from 300 to 400 C, depicted in Figure 5 and Figure S7a. At low temperatures (200 C) amorphous iron oxide forming microbundles occurred; however, they did not possess a hierarchical surface structure as in Figure S7b,c. These results clearly indicated that the crystallized WHHS should have formed from the amorphous iron oxide microbundles at early stages. Notably, as the decomposition temperature increases, the degree of crystallization of iron oxide increases, which eventually converts mesoporous amorphous iron oxide to macroporous WHHS surface structure as is evident from XRD and nitrogen adsorption analysis. We conducted controlled experiments using the iron oxide formed at 200 C for 2 h (Figure S7b,c) by using Feox-2 and Feox-3. About 0.5 g of this iron oxide was calcined at 400 C, shown in Figure S8. Thus, the FESEM results clearly indicated formation of WHHS, which imply that the WHHS formed from amorphous iron oxide. Therefore, we propose a new mechanism: crystallization-induced WHHS formation. During the crystallization process the microbundles are expected to shrink and the
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oxalic acid induced etching process the hematite bundles as a result of the WHHS formed at 400 C. As we earlier reported, the thermal decomposition of ferric oxalate prepared using ferric chlorideoxalic acid results in wellcrystallized hematite aggregates at 200 C. Similarly, the commercial iron(III) oxalates also yields crystallized hematite aggregates at 200 C.18 These two results clearly indicate why WHHS is not formed if ferric oxalate is prepared in other conditions/ methods. No other oxalate precursors and ferric oxalate solvent evaporation methods, except oxalic acid, yields hematite aggregates. Therefore, formation of amorphous iron oxide microbundles at early stages of decomposition is a prerequisite for WHHS formation. The second important mechanism is oxalic acid induced etching of microbundles. As the oxalic acid concentration increased, the sizes of the microbundles were changed into microsheets (Figure 5). The size of the microbundles could possibly be changed by either change in ferric oxalate complex structure using various oxalic acid concentration or oxalic acid induced etching process. The Raman spectra suggests a quite similar complex structure of Feox-2 and Feox-3. The size of the WHHS microbundles, however, decreased as the oxalic acid concentration increased for ferric oxalate preparation from Feox2 to Feox-3. Therefore, the decrease in microbundles size being changed by the ferric oxalate complex structure is completely ruled out. Moreover, during the thermal decomposition process the ferric oxalate is decomposed; hence, it is not possible to directly convert ferric oxalate complex structure into WHHS. Consequently, the decrease in bundles size is due to oxalic acid etching during the decomposition process. We performed additional experiments with a 10% silver nitrate solution to confirm the presence of excess/uncoordinated oxalic acid. Thus, silver ions form an insoluble silver oxalate complex with oxalic acid. The ferric oxalate solution formed at 0.17 M oxalic acid did not yield any precipitate with silver nitrate solution. The other two ferric oxalate solutions (0.24 and 0.4 M), however, immediately yield white precipitate. These experiments confirmed the presence of uncoordinated oxalate ion in ferric oxalate solution. Additional experiments with commercial iron(III) oxalate hexahydrate with additional oxalic acid confirmed the oxalic acid induced etching. Approximately 0.5 g of oxalic acid was mixed with 0.25 g of iron(III) oxalate hexahydrate in 200 mL of water, and the final solution was evaporated on the hot plate. Decomposition at 400 C results in hematite aggregates (Figure S9). The FESEM picture in Figure S9a shows that oxalic acid etched porous aggregates as marked by the red line. Such porous aggregates were not observed in the absence of oxalic acid (Figure S6b), however, clearly indicating oxalic acidinduced etching during the decomposition process. Chen et al. proposed a similar oxalic acid-induced etching of hematite.16 We examined the magnetic properties of porous microsheets (WHHS) formed from Feox-3. The temperature requirement of the magnetization M(T) was measured in zero-field-cooling (ZFC) and field-cooling (FC) conditions by applying the field H = 100 Oe. For zero field cooled magnetization measurements, the sample was first cooled to 5 K. After applying a 100 Oe field at 5 K, the magnetization was measured in the warming cycle. For field-cooled magnetization measurements, the sample was cooled in the same field (100 Oe) to 5 K, and the magnetization was measured in the warming cycle under the field. The rate of cooling is constant in both ZFC and FC conditions. Temperature-dependent ZFCFC curves of the WHHS are shown in 6371
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Figure 7. (a) MH loop of the synthesized WHHS by using Feox-3. Inset: magnetization between 700 and þ700 Oe. (b) Temperature dependence of ZFC and FC magnetization of R-Fe2O3 microsheets at the applied field of 100 Oe.
Figure 7b. The ZFCFC curves of the Feox-3 show a broad cusp corresponding to the blocking temperature TB. The irreversibility of ZFC and FC curves and the presence of ZFC maximum indicate that above TB the particles were characterized by superparamagnetic behavior, whereas below TB, the magnetic moment of each particle was frozen in the local field direction.23 The FC magnetization curve increases very slowly with decreasing temperature and tends to flatten out. There is no overlap of ZFC and FC branches at any measured temperature which bifurcate at high temperatures, Tirr ≈ 300 K (irreversibility temperature). The second sharp feature in the FC and ZFC curves is observed at 221 K—the so-called Morin transition— and is accordingly accompanied by the AF to WF transition.24 The value of TM is lower than the bulk value (TM ∼ 260 K), indicating a well-crystallized condition for the wormlike structure as already shown by the HRTEM in Figure 6. Meng et al. reported the TM value in R-Fe2O3 nanochains at 237 K, which is higher than our observation.25 The magnetic (MH) hysteresis measurements of the sample were carried out in an applied magnetic field at 5 and 300 K with the field sweeping from 10 to þ10 kOe. Figure 7a shows the hysteresis loops of the R-Fe2O3 microsheets at 5 and 300 K. The remanent magnetization and coercivity of the R-Fe2O3 microsheets at 300 K are 0.089 emu g1 and 51 Oe, respectively. The smaller remanent magnetization for as-obtained wormlike R-Fe2O3 compared to the commercial hematite (0.61 emu/g) is probably associated with wormlike assembly of hematite nanoparticles, since the remanent magnetization is strongly dependent upon the particle shape.26,27 At 5 K, the remnant magnetization and coercivity force for the sample are determined to be 0.325 emu g1 and 408 Oe, respectively. The remanent
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magnetization and coercivity force of R-Fe2O3 microsheets at 5 K are larger than under 300 K, indicating that the magnetic moment may get further freezing and the region may change.28 The smaller remanent magnetization and coercivity value of the WHHS in the present study suggest that it could be due to the superparamagnetic nature of WHHS. However, as the average particle size is larger compared to the single domain size of hematite required for superparamagnetism, we believe that the smaller coercivity value could account for the weak ferromagnetic behavior only.29 The R-Fe2O3 has been intensively studied as a promising anode material for next-generation lithium ion batteries,6,17,30 Earlier studies confirmed that the lithium intercalation performance could be influenced by various properties such as pore size, surface area, and particle size of the hematite.31 Generally, increasing the surface area or porosity of the hematite increases lithium storage performance. Previous reports also demonstrate that the morphology plays a significant role in the discharge performance.32,33 Therefore, it is interesting to study the electrochemical performance for lithium storage of synthesized WHHS. Using Li metal as the counter and reference electrodes, we recorded cyclic voltammetry (CV) for the WHHS electrode in the potential range of 0.0053 V at a slow scan rate of 0.5 mV/s. Figure 8a shows the CV profile of the WHHS prepared by using Feox-3. The cathode sweep (Li insertion) shows two wellresolved peaks at 1.46 and 0.38 V corresponding to the Liintercalation and Fe3þ to Fe0 reduction processes. The observed cathodic peak from 0.2 to 0 V might be associated with the irreversible reaction of Fe2O3 with metallic lithium into Li2O and metallic Fe and the reversible alloying reaction of Fe with Li. The anodic sweep (Li removal) only shows a small hump at 1.8 V, owing to the oxidation of Fe0 to Fe3þ. Figure 8b shows the chargedischarge cycling profiles of the three WHHS in the voltage window of 0.013.5 V (vs Li) at the current rate of 0.1 C. During the first discharge, three successive plateaus were observed for all the three WHHS electrodes. In the first cycle, the voltage decreased steeply to 1.65 V from the open-circuit voltage of about 3.2 V, and a capacity of about 6090 mA h g1 was conquered. The second plateau was observed at about 1.10 V up to a capacity of approximately 200250 mA h g1. Another voltage plateau is observed at 0.8 V, whereupon a plateau sets in and maintains until a discharge capacity of 800900 mAh/g, followed by a gradual drop in voltage until the end of discharge. After 10 cycles, however, the discharge profile only shows one small voltage plateau around 0.9 V. For 68 mA h g1, the discharge capacities of the electrode in the first and third cycles are 1216 and 247 mA h g1 for WHHS-1, 1198 and 162 mA h g1 for WHHS-2, and 1343 and 203 mA h g1 for WHHS-3. It could be attributed to the formation of Li2O, from which the extraction of Li is almost impossible. In addition, it is speculated that the formation of Li2O, metallic Fe, and LiFe alloy in the WHHS microbundles alters the bundles size and its porosity. Thus, a fast capacity fading rate was observed for the WHHS-1, which may be due to larger microbundles. The Li ion intercalation in the Fe2O3 was significantly influenced by the microbundles size and porosity of the WHHS. The higher capacity observed for the WHHS in the first cycle is comparable to the reported hematite nanorods (1151 mA h g1),34 which may result from the wormlike hierarchical surface structures of the hematite. The first discharge capacity of the three WHHS considerably exceeds the theoretical capacity of R-Fe2O3 (1007 mA h g1).35 Thus, the simple 6372
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Figure 8. (a) Cyclic voltammogram profile of WHHS electrode in the potential range of 0.0053.0 V at the scan rate of 0.5 mV/s and chargedischarge profile of WHHS electrodes in the potential range of 0.0053.0 V at the current rate of 0.1 C. Solid lines show the first cycle and dashed lines show the third cycle: (b) WHHS-1, (c) WHHS-2, and (d) WHHS-3.
preparation route and promising Li ion storage properties makes the WHHS as potential candidate for Li ion batteries. Because of its nontoxic nature and low cost, R-Fe2O3 is one of the ideal candidates for wastewater treatment process. The synthesized WHHS has the following advantages over conventional hematite nanoparticles: (1) due to the WHHS microsize, it can be separated after use in water treatment process; (2) due to the presence of a top-down wormlike hierarchical porous surface structure, the WHHS can be used as membrane for water treatment process. The nitrogen adsorption analysis revealed the presence of a macroporous surface structure, which is in good agreement with FESEM results. The surface area of the three WHHS prepared using Feox-1 to Feox-3 is found to be 30, 58, and 40 m2/g. Interestingly, the pore size was noted as 2542 nm in the above-mentioned WHHS. The arsenic(V) removal ability of the synthesized WHHS was tested at pH 3. The adsorption experiments indicated that the synthesized WHHS had an excellent As(V) adsorption capacity. The adsorption equilibrium is reached within 30 min as shown in Figure S10. Thus, about 7.3, 12.5, and 9.6 mg g1 of As(V) was noted in the above three WHHS prepared using Feox-1 to Feox-3. Similarly, under the same conditions commercial hematite (Merck) showed about 0.78 mg g1 of As(V) removal, which indicated that the
synthesized WHHS is an excellent candidate for the arsenic removal. The adsorption capacity of the three synthesized WHHS is much higher than earlier reports used for As(V) removal. The high adsorption capacity of the synthesized WHHS is due to its high surface area and hierarchical porous surface structure. Zhong et al. reported about 7.6 mg g1 As(V) removal by using the flowerlike hierarchical hematite.36 Thus, we allege that the synthesized WHHS is an excellent candidate for heavy metal removal due to its unique top-down porous surface structure. Thus, during the adsorption process the heavy metals may adsorb on the surface as well as on the pore walls inside the bundles. Because of top-down porous structures, wastewater containing heavy metals may penetrate through the macropores into inside bundles. Moreover, during the penetration, due to fast As(V) adsorption on WHHS (fast equilibrium), the heavy metals may adsorb on the catalysts while purified water can pass through hierarchical pores so we can finally collect the purified water at the bottom (Figure 9). Therefore, the synthesized WHHS can be used as a potential material for membrane application and will be addressed in future publications. We have, fabricated wormlike hierarchical porous hematite by using an industrially applicable synthetic method without using templates or catalysts. The method of ferric oxalate preparation is 6373
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Figure 9. Schematic representation of As(V) adsorption on WHHS.
very important for the formation of wormlike hierarchical porous microbundles and microsheets. The influence of various experimental parameters on porous hematite formation and the wormlike morphology was self-assembled by oriented attachment. Magnetic studies of the synthesized microsheets indicated a weak ferromagnetic behavior at room temperature. The results of the adsorption of arsenic(V) on the synthesized hematite indicated that the hierarchical porous hematite has excellent arsenic removal capacity.
’ ASSOCIATED CONTENT
bS
Supporting Information. Raman spectra and FESEM pictures of various iron oxides. This material is free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected], Tel þ358403553415, Fax þ35815336013.
’ ACKNOWLEDGMENT EU Transfer of Knowledge fellowship Marie Curie grant MKTD-CT-2006-042637 is thanked for financial support. ’ REFERENCES
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