Facile Fabrication of Hierarchical α-Fe 2 O 3 ...

Published: August 09, 2011

r 2011 American Chemical Society 18164 dx.doi.org/10.1021/jp205834m | J. Phys. Chem. C 2011, 115, 18164–18173

ARTICLE

pubs.acs.org/JPCC

Facile Fabrication of Hierarchical α-Fe2O3: Self-Assembly andIts Magnetic and Electrochemical PropertiesManickavachagam Muruganandham,*,†,‡ Ramakrishnan Amutha,‡ Marappan Sathish,§ TonySarvinder Singh,† Rominder P. S. Suri,† and Mika Sillanp€a€a‡

†Water & Environmental Technology (WET) Center, Department of Civil and Environmental Engineering, Temple University,Philadelphia, Pennsylvania 19122, United States‡Laboratory of Green Chemistry, Faculty of Technology, Lappeenranta University of Technology, Patteristonkatu 1,FI-50100 Mikkeli, Finland§Institute of Multidisciplinary Research for Advanced Materials, Tohoku University, Sendai 980-8577, Japan

bS Supporting Information

1. INTRODUCTION

Nano/micro structured functional material synthesis is cur-rently a topic of interest, because of their special propertiescompared to their bulk counterparts. Designs of well-definedmorphology and the synthesis of colloidal nanocrystals con-trolled by surface properties have been intensively studied in thepast decade for their size- and shape-dependent properties.1

Hematite (α-Fe2O3), an n-type semiconductor (Eg = 2.1 eV), isan attractive multifunctional material due to its low cost, goodstability, high resistance to corrosion, nontoxicity, and environ-mentally friendly properties.2 Its potential applications includecatalysis, gas sensors, magnetic devices, photoelectrodes, pig-ments, and adsorbents.2�6 Therefore, nano/micro-structured α-Fe2O3 fabrication is important from both fundamental andapplication points of views. Various morphological hematiteshave been successfully prepared under different experimentalconditions by many synthetic methodologies.7�10 These meth-ods tend to be complicated, however, with the obvious drawbackthat the removal process often compromises the structuralintegrity of the final products and limits the practical applications.Scientists are continuously exploring new and simple syntheticmethodologies for the preparation of various materials.

An attractive hematite preparation method is solid-statethermal decomposition of the iron oxalate complex.5,6,11 Moststudies were, however, conducted with commercially availableiron(II) and iron(III) oxalates, but except for our recent report,6

none reported fabrication of nano/micro-structured hematite.To the best of our knowledge, no reports are available for organicsolvent-assisted ferric oxalate preparation and its thermal decom-position studies. We earlier confirmed that ferric oxalate pre-paration is very crucial for hematite morphology. Therefore,investigation of the above-mentioned topic may result in newinteresting findings.

Recently, we successfully prepared self-assembled mesopor-ous amorphous iron oxide and wormlike hierarchical poroushematite by thermal decomposition of ferric oxalate prepared inwater.5,6 We further found that the solvent used for ferric oxalatepreparation played a major role in hematite morphology. Here,we report surface-structure-controlled hematites of various mor-phologies on a large scale without the use of any templates or

Received: June 21, 2011Revised: August 5, 2011

ABSTRACT: We report novel fabrication of various surface-structure-controlled morphologies of α-Fe2O3 in a simplethermal decomposition process. Ferric nitrate�oxalic acidcomplex formed in ethanol solvent facilitates chain- and rodlikemorphological hematites after thermal evaporation of solventfollowed by decomposition at 400 �C. This simple methodology offers industrially applicable nanostructured hematite preparation.The oxalic acid concentration and method of solvent evaporation played an important role in the final morphology. Increasing theoxalic acid concentration from 0.125 to 0.4 M facilitated hematite chainlike morphology of porous aggregates. Thermaldecomposition temperature played an important role on the surface structure of hematite. The synthesized hematite wascharacterized by X-ray diffraction (XRD), field emission scanning electron microscopy (FE-SEM), transmission electronmicroscopy (TEM), Raman spectra, Fourier transform infrared spectra (FTIR), and nitrogen adsorption analysis. The influenceof various solvents such as ethanol, methanol, propanol, t-butanol, and various iron precursors on the final morphology wasinvestigated, and the results indicated that both solvents and iron precursors are crucial for the final hematite morphology. Themagnetic properties of the synthesized chainlike hematite were studied, and the remnant magnetization and coercivity of thechainlike α-Fe2O3 at 300 K are found to be 1.36 emu/g and 147 Oe, respectively. The electrochemical Li ion battery performance ofsynthesized hematite was investigated. A plausible mechanism of hematite self-assembly was proposed.

18165 dx.doi.org/10.1021/jp205834m |J. Phys. Chem. C 2011, 115, 18164–18173

The Journal of Physical Chemistry C ARTICLE

surfactants by a self-assembly process. We studied the influenceof various experimental parameters on the ferric oxalate prepara-tion and its decomposition product (α-Fe2O3) morphology. Wealso investigated the magnetic and electrochemical properties ofthe synthesized hematite.

2. EXPERIMENTAL SECTION

Ferric nitrate nonahydrate, anhydrous ferric chloride, iron-(III) oxalate hexahydrate, oxalic acid, ethanol, methanol, n-propanol, and t-butanol were purchased from Aldrich ChemicalCo. Ltd. (Helsinki, Finland). All chemicals were of analyticalgrade and used without further purification. For all experimentalwork Milli Q-Plus water (resistance = 18.2 MΩ) was used.

The synthesis of α-Fe2O3 involved two steps. First, the ferricnitrate�oxalic acid complex was prepared in ethanol by mixingan equal volume (100 mL) of 0.1 M ferric nitrate nonahydratewith the required concentration of oxalic acid in ethanol. Whileusing the same volume of solvent (200 mL), we prepared fourtypes of ferric oxalate solids with 0.125, 0.175, 0.225, and 0.4 Moxalic acid. During magnetic stirring, which was conducted in thedark, oxalic acid was added drop by drop (30�45 min) into aferric nitrate solution. Solvent evaporation was avoided bycovering the solution with parafilm while it was stirred overnightand then evaporated on a hot plate until becoming a greenish gel.In the second step, prepared ferric oxalate complex is decom-posed at a desired temperature and time under open atmosphericconditions. After decomposition, the oven was cooled to roomtemperature and the samples were washed with plenty of waterand ethanol and dried at 120 �C for 2 h.2.1. Electrochemical Measurements. For electrochemical

measurements, a three-electrode cell was constructed, in whichthe working electrode was fabricated by pressing the mixtureof 85% hematite, 10% acetylene black, and 5% poly(tetra-fluoroethylene) (PTFE) on a nickel mesh. Li metal pressed onnickel mesh was used as the counter electrode and referenceelectrode. The galvanostatic charge�discharge was conducted ina potential cutoff range of 0.001 to 3.5 V (vs Li/Li+) at thecurrent rate of 0.1 C. There was a 3 h rest period before chargeand discharge measurement. For cyclic voltammetric measure-ment, the mass of the active material in the working electrodeis ∼10 mg and the scan rate was 0.1 mV/s. The electrolytesolution was 1 M LiClO4 dissolved in a mixture of ethylenecarbonate (EC) and diethyl carbonate (DEC), with an EC/DECvolume ratio of 1:1 (v/v). Electrochemical performance wasinvestigated by use of the Solartron cell test analyzer.X-ray diffraction (XRD) patterns were recorded on an X0 Pert

PRO PAN analytical diffractometer, with scanning angles of10��100�. Transmission electron microscope (TEM) imageswere recorded on a field emission (FE) TEM [Philips CM-200FEG - (S) TEM - Super Twin]. Samples for TEM were preparedby ultrasonically dispersing the hematite into ethanol and thenplacing a drop of this suspension onto a carbon-coated coppergrid and air-drying. The working voltage of TEMwas 80 kV. Themorphology of the hematite was examined on a Hitachi S-4100scanning electron microscope (SEM). Prior to SEM measure-ments, the samples were mounted on a carbon platform, whichwas then coated by platinum via a magnetron sputter for 10 min.The plate containing the sample was placed in the SEM foranalysis at desired magnifications. The X-ray photoelectronspectra were collected on an ESCA-1000 X-ray photoelectronspectrometer (XPS), with Mg Kα X-ray as the excitation source.

Fourier transform infrared (FTIR) spectra were recorded by useof a Bruker Vertex 70 with a platinum ATR-QL, diamondaccessory. Raman spectra were acquired by use of a RenishawinVia microscope with a 514.5 nmArgon laser at 50% power witha 50� aperture. Five acquisitions were taken for each samplefrom 100 to 1000 cm�1 for 10 s. Thermal analysis was done on aNetzsch TG 209 F1 thermomicrobalance under air/oxygenatmosphere at a flow rate of 10 �C/min.

3. RESULTS AND DISCUSSION

We used a new synthetic methodology for the preparation ofvarious morphological hematites. Many groups have studiedpreparation of ferric oxalates in water but not synthesis by useof organic solvents.5,6,12 Controlled experiments indicated thatthe method of preparation of ferric oxalate is crucial for the finalmorphology of hematite. Therefore, investigating how the ferricoxalate preparation conditions influence the hematite morphol-ogy is important. The ferric oxalate complex appears as emeraldgreen; however,mixing both iron nitrate and oxalic acid in ethanolresults in an orange-red solution as shown in Figure 1a. Theseresults indicate that mixing of ferric nitrate and oxalic acid inethanol solvent does not facilitate the ferric oxalate complex.Moreover, room temperature (25 �C) evaporation did not changethe solution color. Solvent evaporation on a hot plate, however,induced color change fromorange-red to emerald green, implyingformation of ferric oxalate complex as shown in Figure 1b.Therefore, we prepared all ferric oxalate complexes via the hot-plate evaporation process. Decomposition of the final ferricoxalate solid formed during both room-temperature and hot-plate evaporation indicates that the method of solvent evapora-tion is crucial to the final hematite morphology.

Unlike dry ferric oxalate salt formed in water, the ferric oxalateformed in ethanol solvent appeared gel-like (Figure 1b) anddiffered from that formed in water. To the best of our knowledge,no studies on ethanol solvent-assisted ferric oxalate preparationexist, so comparison is difficult. Therefore, we attempted tocompare the structure of ferric oxalate formed in both water andethanol by FTIR and Raman spectra. We have previouslyanalyzed the structure of ferric oxalates formed in water, andwe compared these results with commercial ferric oxalatehexahydrate.5,6 The peak intensities and position of the synthe-sized iron oxalates in ethanol are different from the commercialiron(III) oxalate hexahydrates as shown in Figure S1 (SupportingInformation). The FTIR spectra of commercial ferric oxalatehexahydrate showed strong absorbance at 1598, 1643, and1730 cm�1, which are characteristic of υ(CO) bands, and at1345, 1240, 758, and 816 cm�1, which are assigned to υ(C�C)

Figure 1. Ferric oxalate complex formed in ethanol: (a) before and (b)after solvent evaporation on a hot plate.



vibration mode. Ferric oxalate prepared in ethanol, however,showed strong absorbance peaks at 1608 and 1672 cm�1 and avery weak absorbance peak at 1763 cm�1, which are assigned toυ(CO) mode. The peaks at 1349, 1304, and 801 cm�1 areassigned to υ(C�C) mode. Some new additional medium-intensity peaks were also noted from the 600�1500 cm�1

region; such peaks were not noted in the commercial ferricoxalate hexahydrate.

We earlier confirmed that the ferric oxalate synthesized inwater solvent is similar to the commercial ferric oxalate hexahy-drate. Therefore, the structural difference of ferric oxalate ismainly caused by the solvent used for preparation. Interestingly,however, a weak υ(OH) mode observed at 3524 cm�1 reflectsthe hydrated form of ferric oxalates. Although we did not usewater at any preparation stage, ferric oxalate occurred as thehydrated form. Since we have used ferric nitrate nonahydratesalts for synthesis, water molecules could come from the above-mentioned iron precursor. Therefore, the synthesized ferricoxalate in ethanol may possess the hydrated form. We cannotcompletely rule out the possibility of ethanol being coordinatedwith ferric oxalates, since we observed many additional FTIRpeaks between 1000 and 1500 cm�1. Generally, alcohol CHstretching vibrations will appear around 3000 cm�1, and a ferricoxalate absorbance band developed at 2994 cm�1. Thus, a peakat 1076 cm�1 is assigned to υ(CO) mode, which is in the normalregion of alcoholic CO absorbance. These results clearly indicatethat the ferric oxalate structure formed in ethanol solvent isdifferent from the commercial ferric oxalate hexahydrate.

Raman spectra of both the above-mentioned ferric oxalates areshown in Figure 2. The results showed a peak shift toward thelower wavenumber in the synthesized ferric oxalate, which isexpected due to differences in ferric oxalate complex structures.Strong peaks appeared at 555 and 539 cm�1 in commercial andsynthesized ferric oxalate. These results confirm ferric oxalateformation. One notable difference in the FTIR spectra of thesesamples appears at 1730 cm�1, characteristic of terminal carbo-nyl group presence. Edwards and Russell13 studied the possiblestructure of iron(III) oxalate complex at various oxalate concen-trations by FTIR. We also elucidated the possible complexstructure of ferric oxalate formed in water in our earlierpublication.6 In the absence of a possible terminal uncoordinatedpeak, the structure of ferric oxalate formed in ethanol is proposedas presented in Figure S2a (Supporting Information). The

Raman spectra of ferric oxalate formed in ethanol solvent isquite different from that formed in water solvent.6 Further ferricoxalate complex structure details were not investigated, how-ever, since it is used as a precursor for hematite synthesis and isexpected to decompose completely during the thermal decom-position process.

We studied the thermal decomposition of ferric oxalateformed in ethanol under air atmospheric conditions usingstandard and derivative thermogravimetry (TG and DTG) asshown in Figure S2b (Supporting Information). The TG analysisshowed approximately 8.4% decomposition at the initial stages(below 120 �C). This might be due to elimination of eithercoordinated water or ethanol molecules present in ferric oxalatesalts. Moreover, 47.8% of decomposition was observed at 150�260 �C, confirming iron oxide formation. DTG results showedone sharp peak at 254 �C, confirming the formation of iron oxidefrom ferric oxalate. All ferric oxalate preparation was initiallyconducted in 0.1 M ferric nitrate with 0.17 M oxalic acid inethanol, and thermal decomposition was studied at 400 �C for2 h. After solvent evaporation on a hot plate, the gel-like ferricoxalate was dried completely and further used for decomposition.We confirmed α-Fe2O3 formation by XRD and XPS andexamined the Raman spectra, and we studied morphology byFE-SEM. The XRD pattern in Figure 3 confirmed formation ofα-Fe2O3. All XRD peaks can be indexed to the orthorhombic-centered α-Fe2O3 (JCPDS 80-2377) with lattice parameters ofa = 0.5035 nm and c = 1.374 nm, indicating high purity and goodcrystallinity of the synthesized hematite.

Figure 4 shows the XPS spectra of hematite prepared bythermal decomposition of ferric oxalates at 400 �C for 2 h; panelsa and b depict the XPS spectra of Fe 2p and O 1s. The Fe 2pspectrum indicates the existence of doublet Fe 2p3/2 and Fe 2p1/2,with binding energies of 711.5 and 724.5 eV. The Fe 2p3/2 peakis also associated with a satellite peak located approximately7.5�8 eV higher than the main peak, which is characteristic ofα-Fe2O3. The corresponding satellite peak noted at 720 eV, aresult of charge transfer screening, can be attributed solely tothe presence of Fe3+ ions of α-Fe2O3. A high-resolution O 1sspectrum is shown in Figure 4b. The peak at 532.2 eV is mainlyattributed to Fe�O in α-Fe2O3. The XPS results also confirmedthe formation of α-Fe2O3, which is in good agreement withearlier reports.14

Figure 2. Raman spectra of ferric oxalates: (A) commercial ferricoxalate hexahydrate and (B) ferric oxalate synthesized in ethanol solvent.

Figure 3. XRDpattern of hematite formed at 400 �C for 2 h in a thermaldecomposition process of ferric oxalate formed in ethanol solvent.



The FE-SEM picture in Figure 5 shows the morphology of α-Fe2O3 formed in the thermal decomposition process. As shownin Figure 5a, these results clearly indicate formation of anaggregate-free, porous rodlike hematite. The average width andlength of the starlike hematite are in the range of 300�500 nmand 1�4 μm. The rodlike morphology consists of wormlikesegments, and such segmented self-assembly induces poroussurface structures. To the best of our knowledge, we are thefirst to report a unique surface structured rodlike hematite.P�erez-Maqueda et al.15 prepared similar porous needlelikeBaTiO3 by thermal decomposition of a citrate precursor. Undersimilar experimental conditions, decomposition of ferric oxalateprepared in water results in a different hematite morphology.6

Figure 5b also shows the results from the thermal decompositioncarried out in gel-like ferric oxalates formed after solventevaporation. The morphology of the formed hematite is quitesimilar to that in Figure 5a. These results showed that gel-likeferric oxalate solid drying does not affect the morphology.Moreover, since samples spread in the oven during the geldecomposition process, care should be taken; however, we didnot note this in ferric oxalate prepared in water or in commercialsamples. This may be caused by fast solvent or byproduct (nitricacid) evaporation during the decomposition process. We alsoperformed experiments to understand the influence of ferricoxalate preparation conditions on the final hematite morphol-ogy. After ferric nitrate and oxalic acid were mixed, the ethanolsolvent was partially evaporated overnight. Figure 5c shows thedecomposition of this residue. Thus, the hematite formedunder these conditions showed quite different morphologycompared to the FE-SEM in Figure 5a,b. These results implythat ferric oxalate preparation conditions directly influence thehematite morphology.

We examined different conditions to understand the influenceof various experimental parameters on the morphology ofhematite. We investigated the influence of decompositiontemperature on morphology from 300 to 400 �C. FE-SEMresults clearly indicate that hematite morphology formed at300 �C is different from that formed at 400 �C (Figure 5b) asshown in Figure S3 (Supporting Information). Hematiteformed at lower temperatures does not have a porous surfacestructure and star shape. Decomposition at lower tempera-tures (175 �C), however, results in only aggregates, as shownin Figure S3b ( Supporting Information ). These results differfrom the aforementioned discussion, and therefore decom-position temperatures play an important role in the finalmorphology.

We investigated the influence of oxalic acid concentration onhematite morphology. While ferric nitrate was kept at 0.1 M, theoxalic acid concentration was changed from 0.125 to 0.4 M. TheFE-SEM results in Figure 6 clearly show that the oxalic acidconcentration influenced the hematite morphology. At low initialoxalic acid concentrations (0.125M), hematite chainlike bundleswere noted, as shown in Figure 6a. Further increase of the oxalicacid concentration to 0.225M, however, results in well-dispersedchainlike rods as shown in Figure 6b. Moreover, all three oxalicacid concentrations (0.125, 0.17, and 0.225 M) used for ferricoxalate preparation showed a similar hematite self-assembly unitbut different morphologies as shown in Figures 5 and 6a,b.Interestingly, at higher concentrations (0.4 M), irregular porousaggregates occurred (Figure 6c), which is quite different from theabove-mentioned morphology. These results confirmed theinfluence of oxalic acid concentration on hematite morphology.Therefore, it is necessary to study how varying oxalic acidconcentrations influence the hematite morphology. We havealready confirmed that, at higher oxalic acid concentrations(0.4 M), the ferric oxalate complex structure differs from thatformed at low concentrations (0.17 M).6 Moreover, at higheroxalic acid concentrations, uncoordinated oxalic acid also con-tains ferric oxalate salt.6 On the basis of these observations, twopossible ways of addressing these issues exist. One is changingoxalic acid concentrations, which may change the ferric oxalatecomplex structure, and the other is an oxalic acid-induced etchingprocess. We have previously suggested an oxalic acid-inducedetching mechanism for the formation of hematite microsheetsfrom microbundles at higher oxalic acid concentration (0.4 M).6

Moreover, in this study no obvious changes, except their size,emerge in morphology at various oxalic acid concentrations. Inthis study, however, hematite morphology changes occurredeven at low initial oxalic acid concentrations (0.125 and 0.17 M),where there is no uncoordinated oxalic acid present alongwith ferric oxalate. These results suggest that the possiblechanges in ferric oxalate structure with oxalic acid may beresponsible for final hematite morphology. Our experimentalobservations also support these speculations. As discussed ear-lier, the ferric oxalate structure formed in water and ethanoldiffers, which could yield different results. On the other hand, theferric oxalate complex is expected to completely decomposebefore being converted into hematite during the thermal decom-position process. Therefore, it is not possible to directly convertthe ferric oxalate structure into hematite. We propose that thedecomposition of differently structured ferric oxalate couldinfluence the initial growth of nuclei by a dissolution/etching

Figure 4. XPS spectra of hematite prepared at 400 �C for 2 h in a thermal decomposition process of ferric oxalate formed in ethanol solvent: (a) Fe 2p;(b) O 1s.



process, which further self-assembles into different morpholo-gies. At high oxalic acid concentrations (0.4 M), however, wecannot rule out the involvement of oxalic acid-induced etchingalong with the above proposed mechanism.

Another important variable possibly influencing hematitemorphology formation is the iron precursor used for ferric oxalatepreparation. Under similar conditions, we exchanged ferric nitratewith ferric chloride as the iron precursor and carried out decom-position at 400 �C for 2 h. The FE-SEM results, in Figure S4(Supporting Information), show hematite porous nanoparticleformation. These results differ from the decomposition results offerric oxalate prepared from ferric nitrate as a precursor, presentedin Figure 5. Therefore, the hematite morphology is also influ-enced by the nature of iron precursors, and we address thereason for the morphology changes in the later part of thismanuscript. We have also examined how solvents used for ferricoxalate preparation influence hematite morphology. While other

experimental parameters were kept constant, ethanol was replacedby an equal volume of methanol, n-propanol, or t-butanol. Aftersolvent evaporation, decomposition was carried out at 400 �C for2 h. The XRD results in (not shown) imply formation of hematitein all these solvents, and the results showed no influence onhematite phase formation. The FE-SEM results are in Figure S5(Supporting Information). Ferric oxalate prepared in methanol orn-propanol yields hematite aggregates and hierarchical poroussurface structured aggregates, shown in Figure S5a,b (SupportingInformation). Interestingly, Figure S5c (Supporting Information)shows that ferric oxalates prepared in t-butanol produced hier-archical porous rod/plate-shaped hematite. These hematite mor-phologies are quite different from those prepared with ethanolor water as a solvent. Therefore, solvents used for ferric oxalatepreparation played a crucial role in hematite morphology.

It is interesting to study the influence solvents may have onferric oxalate structure, which eventually affects final hematite

Figure 5. FESEM of hematite formed at 400 �C for 2 h in a thermal decomposition process of ferric oxalate formed in ethanol: (a) dry soliddecomposition, (b) gel decomposition, and (c) partial evaporation of solvent. The corresponding higher magnification panel is presented in a1,b1,c1.



morphology. After complete solvent evaporation on a hot plate,the final ferric oxalate solid showed respective solvent odor,which could indicate coordination of solvent with ferric oxalatesolid. Our speculation is also supported by TG results, which

showed almost 8.4% weight loss at the initial stages (below120 �C), and IR spectra showed additional absorbance peaks from1500 to 1000 cm�1, which are not observed in ferric oxalate pre-pared in water as shown in Figure S1 (Supporting Information) .

Figure 6. FE-SEM of hematite formed at 400 �C for 2 h in a thermal decomposition process of various ferric oxalates prepared at different oxalic acidconcentrations in ethanol solvent: (a) 0.125 M; (b) 0.225 M; (c) 0.4 M; (d) 0.4 M at 300 �C for 2 h. The corresponding higher magnification panel ispresented in a1,b1,c1.



This solvent coordination may change the ferric oxalate structure,or during the decomposition process the solvent evaporation couldinfluence the initial crystal growth of hematite, resulting in variousmorphological hematites. These observations are also supported byour experiments. Thus, on the basis of Raman and FTIR spectra,the ferric oxalate prepared in water and ethanol solvents showeddifferent ferric oxalate structures resulting in morphologicallydifferent hematite. Similarly, decomposition of synthesized andcommercial ferric oxalate yields differentmorphologies of hematite.A detailed understanding of the relationship of ferric oxalatestructures formed in different solvents and their decompositionproduct (hematite) morphology needs multifold microscopicanalysis. These issues will be addressed in a future publication.

We have not used surfactants or templates or any matrices foreither ferric oxalate or hematite synthesis (in decompositionprocess). Therefore, various morphological hematites could haveformed by self-assembly mechanisms.We studiedmechanisms ofvarious morphological hematite self-assembly via experimentalobservations. To understand these mechanisms, it is necessary toanswer two important questions. First, how do varying oxalic acidconcentrations change the hematite morphology? Second, whywere rod- and chainlike morphologies not formed under otherexperimental conditions? The possible answer for the firstquestion was discussed earlier.

We examined temperature-dependent (200�400 �C) mor-phology and degree of crystallization of iron oxide formed fromferric oxalate by FE-SEM and XRD analysis. XRD results indicate

that as the decomposition temperature increased from 200 to400 �C, the degree of crystallization of iron oxide increased, asevident from the conversion of amorphous iron oxide to crystal-lized hematite. Similarly, at the same temperatures the FE-SEMresults show conversion of amorphous aggregated iron oxide intoporous rodlike hematite, shown in Figures 5a and 6b (at lowoxalic acid concentration) and in Figure 6c,d (at higher oxalicacid concentration). Most importantly, the crystallized hematiteat 400 �C possesses a different morphology compared to that atlower temperatures. Obviously, these results imply that rodlikehematite formed during the crystallization process. Therefore,we propose a crystallization-induced self-assembly mechanismfor the formation of various morphologies of porous hematite.Recently we reported novel crystallization-induced wormlikehierarchical formation of porous hematite microbundles.6 Wealso experimentally proved that decomposition of both commer-cial ferric oxalate hexahydrate and ferric oxalate formed fromferric chloride�oxalic acid yields crystallized hematites at200 �C. Therefore, the porous rodlike hematite was not formedat 400 �C if we used ferric chloride for ferric oxalate preparationand from commercial ferric oxalate hexahydrate. On the basisof these conclusions, it is understood that the formation ofamorphous-phase iron oxide at low temperatures is a prerequisitefor formation of both rod- and chainlike hematite at 400 �C.The schematic mechanism is shown in Scheme 1. TEM analysisindicates the formation of porous surface structured rodlikehematite, and the selected area electron diffraction (SAED)pattern revealed a diffused ring pattern confirming the polycrys-talline nature of the hematite, shown in Figure 7.3.1. Magnetic Properties. We studied the magnetic proper-

ties of the synthesized chainlike hematite by using a quantum-designed physical propertymeasurement system (PPMS).More-over, we measured the temperature-dependent magnetization(M�T) curves with an applied field of 100 Oe between 10 and300 K via field-cooling (FC) and zero-field-cooling (ZFC)procedures. Initially, we performed the ZFC measurements bycooling samples to 5 K in a zero field and then raising the field to100 Oe. The magnetic moment was recorded by the warmingcycle and the field-cooled curve (FC) measurement was donedirectly after cooling in the same field (H = 100 Oe). Bulkhematite has a Morin transition from the low-temperatureantiferromagnetic phase to a weakly ferromagnetic phase at ca.260 K. Morin temperature reportedly decreases as particle sizedecreases, tending to disappear for particles smaller than ca. 20 nm.16

Scheme 1. Formation of Porous Rodlike Hematite by Self-Assembly

Figure 7. (a) TEM of rodlike hematite image; (b) SAED pattern.



The left top inset of Figure 8 shows the temperature-dependentmagnetization curve of the chainlike hematite. The character-istics of the ZFC and FC magnetization curves indicate noapparent Morin transition, and the FC and ZFC curves of thechainlike Fe2O3 split significantly. Zhao et al.

17 previously reportedthe absence of Morin transistion in hematite nanorods. The ZFCcurve shows thatmagnetization of chainlike hematite increases withtemperature up to 300 K, representing a ferromagnetic property,whereas the FC magnetization was almost constant up to 300 K.18

The magnetic features of hematite, including the temperaturevalues of magnetic transitions, can be influenced by numerousfactors, such as crystal defects, strains, surface phenomena, andparticle size.19 The absence of a Morin transition in our study maybe due to the porous nature of the chainlike hematite surface.We measured the magnetic hysteresis (M�H) curves of

α-Fe2O3 at different temperatures to further confirm the magneticproperties. Figure 8 shows the field-dependentmagnetization curveof the chainlike hematite measured at 5 and 300 K. The magnetichysteresis (M�H) measurements were carried out at 5 and 300 Kwith the field sweeping from�20 to +20 kOe. The hysteresis loopsof chainlike Fe2O3 display the ferromagnetic behavior of thesample. The remnant magnetization and coercivity of the chainlikeα-Fe2O3 at 300 K are 1.36 emu/g and 147 Oe, respectively, and at5 K the values are 2.3 emu/g and 294 Oe, respectively. Themagnetic hysteresis curves measured at different temperatures aresaturated within the applied field, and the saturation magnetizationat 5 and 300 K are 6.7 and 6.2 emu/g, respectively. The coercivityand remnant magnetization at 5 K is greater than at 300 K,indicating that the magnetic moment may get further freezingand the magnetic region may change.20

3.2. Electrochemical Measurements. α-Fe2O3 has beenintensively studied as a promising anode material for next-generation lithium-ion batteries.16,21,22 Earlier studies confirmedthat lithium intercalation performance could be influenced byvarious properties such as pore size, surface area, and hematiteparticle size.23 Generally, increasing the surface area or porosityof the hematite increases lithium storage performance. Previousreports also demonstrate that morphology plays a significant rolein the discharge performance.6,24,25 Therefore, it is interesting tostudy the electrochemical performance for lithium storage ofsynthesized hematites at various oxalic acid concentrations.

Figure 9 shows the charge�discharge cycling profiles of as-prepared hematites by using 0.0125, 0.175, and 0.225 M initialoxalic acid concentrations. Corresponding hematite morpholo-gies are presented in Figures 5a and 6a,b, hereafter denotedas H-1, H-2, and H-3. We conducted all experiments in thevoltage window of 0.001�3.5 V (vs Li) at the current rate of0.1 C. During the first discharge, we observed three successiveplateaus for all the three synthesized hematites electrodes.In the first cycle, the voltage decreased steeply to 1.62 Vfrom the open-circuit voltage of about 3.2 V, and a capacity ofabout 40�50 mA 3 h 3 g

�1 was conquered. The second plateauwas observed at about 1.10 V up to a capacity of approximately

Figure 8. M�H curves of chainlike hematite measured at 5 and 300 K.(Insets) Left top, ZFC�FC curve; right bottom, center-expandedM�H curve.

Figure 9. Charge�discharge profile of hematite electrodes in potentialrange 0.001�3.5 V at current rate of 0.1 V. Hematite was prepared at400 �C for 2 h at various oxalic acid concentrations: (A) 0.125,(B) 0.175, and (C) 0.225 M.



230�250 mA 3 h 3 g�1. Another voltage plateau occurred at 0.8 V,

whereupon a plateau sets in remains until a discharge capacity of1300� 1400 mA 3 h 3 g

�1, followed by a gradual drop in voltageuntil the end of discharge. In the second cycle, however, thedischarge profile only shows one small voltage plateau around1.1 V and the discharge capacity decreases gradually. Thedischarge capacities of the electrode in the first and 10th cyclesare 1400 and 80 mA 3 h 3 g

�1 for H-3, 990 and 75 mA 3 h 3 g�1 for

H-2, and 1400 and 65mA 3 h 3 g�1 for H-1. The higher irreversible

capacity obtained in the first discharge could be attributed toLi2O formation, fromwhich Li extraction is almost impossible. Inaddition, formation of metallic Fe and the subsequent reversibleLi�Fe alloy formation enhance hematite capacity. A fast capacityfading rate was, however, observed for the prepared hematite.

When the particle size decreases to nanosize, the thickness ofthe Li2O formation can be minimized and reversibility of theLi�Fe alloy can be considerably improved. Thus, particle sizeplays a key role in the initial capacity and the reversibility.These results clearly indicate that Li-ion intercalation inFe2O3 was significantly influenced by the hematite size. Thehigher capacity observed for the H-3 in the first cycle iscomparable or slightly higher than the reported hematitenanorods (1151 mA 3 h 3 g

�1).26 This may result from differ-ences in morphology as well as the various surface propertiesof these hematites. Thus, the simple preparation route andpromising Li-ion storage properties makes the as-prepared α-Fe2O3 a potential candidate for Li-ion batteries.Cyclic voltammetry (CV) was recorded for the as-prepared

samples in the potential range of 0.001�3 V at a slow scan rate of0.1 mV/s with Li metal as the counterelectrode and referenceelectrode. The electrolyte solution was 1MLiClO4 dissolved in amixture of ethylene carbonate (EC) and diethyl carbonate(DEC), with a volume ratio of EC/DEC of 1:1(v/v). Figure 10shows the CV profile of the as-prepared samples. In the firstcycle, the cathode sweep (Li insertion) shows two well-resolvedpeaks at 1.55 and 0.55 V, corresponding to Li intercalation andFe3+ to Fe0 reduction processes. The observed cathodic peakfrom 0.1 to 0.001 V might be associated with the irreversiblereaction of Fe2O3 with metallic lithium into Li2O and metallic Feand the reversible alloying reaction of Fe with Li. The anodicsweep (Li removal) shows a hump at 1.75 V owing to theoxidation of Fe0 to Fe3+. The characteristic peaks in the secondand subsequent cycles disappear slowly, indicating fading capa-city. The capacity fading rate observed for these samples,however, is lower than our previous reported samples and isdue to the nanosize and rod- and chainlike morphology of thehematite.6

4. CONCLUSIONS

We successfully prepared various morphologies of α-Fe2O3

by a simple template-free method in large scale. Ferric oxalateprepared in ethanol followed by thermal decomposition at400 �C yields porous rod- and chainlike morphologies atdifferent oxalic acid concentrations. All the above-mentionedhematite morphologies were formed by self-assembly mechan-isms. The oxalic acid concentration and solvent used for ferricoxalate preparation and its decomposition temperature stronglyinfluenced the hematite morphology. Magnetic measurementsindicate that synthesized hematite possesses weak ferromagneticproperties, and the magnetization value was found to be 6.2 emuat room temperature. Electrochemical measurements indicatethat the hematite morphology strongly influences the electro-chemical performance.

’ASSOCIATED CONTENT

bS Supporting Information. Five figures showing FTIR,thermal analysis, structure of ferric oxalate complex, and FE-SEM of hematite. This material is available free of charge via theInternet at http://pubs.acs.org.

’AUTHOR INFORMATION

Corresponding Author*E-mail [email protected] or [email protected]; phone +1 215 204 7802; fax +1 215 204 0622.

Figure 10. Cyclic voltammogram profile of hematite electrode inpotential range 0.001 to 3.0 V at scan rate 0.1 mV/s. Hematite wasprepared at 400 �C for 2 h at various oxalic acid concentrations: (A)0.125, (B) 0.175, and (C) 0.225 M.



’ACKNOWLEDGMENT

We appreciate the financial support of the EU Transfer ofKnowledge fellowshipMarieCurieGrantMKTD-CT-2006-042637.

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