Enhanced Electrochemical Lithium Ion Charge Storage of Iron Oxide

Article pubs.acs.org/cm

Enhanced Electrochemical Lithium-Ion Charge Storage of Iron Oxide Nanosheets Sibo Niu,† Ryan McFeron,† Fernando Godínez-Salomón,† Brian S. Chapman,‡ Craig A. Damin,† Joseph B. Tracy,‡ Veronica Augustyn,‡ and Christopher P. Rhodes*,† †

Department of Chemistry and Biochemistry, Texas State University, San Marcos, Texas 78666, United States Department of Materials Science and Engineering, North Carolina State University, Raleigh, North Carolina 27695, United States

‡

S Supporting Information *

ABSTRACT: Iron oxides are appealing cathode materials for low-cost electrochemical energy storage, but iron oxide nanoparticles (NPs) exhibit very low capacities, particularly at fast charging and discharging times, which are increasingly important for numerous applications. We report that synthesis and stabilization of iron oxide in nanosheets results in significantly improved lithium-ion charge storage capacities compared to those of iron oxide NPs at both slow and fast charging/discharging times. The iron oxide nanosheets have lateral dimensions of ∼50 nm and thicknesses of ∼1 nm and are composed of smaller crystallites. The structure of the nanosheets is consistent with maghemite, γ-Fe2O3, which contains cation defects. The γ-Fe2O3 phase is not typically observed within a nanosheet form, and γ-Fe2O3 nanosheets transform to NPs at a relatively low temperature of 200 °C. The transformation of γ-Fe2O3 from a nanosheet to an NP occurs in conjunction with removal of structural H2O. The γ-Fe2O3 nanosheets exhibited lithium-ion charge storage capacities of up to 148 mA h g−1, which is significantly greater than that of commercial γ-Fe2O3 NPs (32 mA h g−1). γ-Fe2O3 nanosheets showed the ability to be rapidly charged and discharged (93.2 mA h g−1 at a 9 min discharge time) with significantly higher capacities than γ-Fe2O3 NPs. The electronic conductivity of the nanosheets was 3 times higher than that of NPs, which is attributed to facilitated electron conduction within the nanosheets. Kinetic analysis of the charge storage mechanism suggests the nanosheets store charge predominantly via a capacitive charge storage process rather than conventional intercalation. The understanding of how to synthesize and stabilize iron oxide nanosheets, their unique electrochemical properties, and their distinct charge storage mechanism furthers the design of charge storage materials with improved capacities, enhanced rate capabilities, and lower cost.

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low cost, natural abundance, and low toxicity.6−8 Several different phases of iron oxide, including magnetite (Fe3O4),9−11 maghemite (γ-Fe2O3),9,12−14 and hematite (α-Fe2O3),10,12 have been investigated as Li-ion battery cathodes. Within the potential range of ∼1.5−4.2 vs Li, iron oxides electrochemically store lithium cations and electrons with a concomitant change in the iron oxidation state (Fe3+/2+).14 Electrochemical charge storage using Fe2O3 can be described by the following equation:

INTRODUCTION Electrochemical energy storage materials with high capacities and the ability to be rapidly charged and discharged and that are extremely low cost are needed for numerous applications, including electric vehicles, consumer devices, and grid-level energy storage.1,2 For current lithium-ion (Li-ion) batteries, the cathode material constitutes a significant proportion of battery material costs, and low-cost, high-performance cathodes are needed for next-generation batteries.3 In addition to high costs, the capacities and voltages of conventional cathodes drop significantly at fast discharge times,4 and cathodes that can be rapidly charged and discharged with high capacities are especially needed for electric vehicles.5 Among candidate materials for cathodes, iron oxides (general formula FeOx) have attracted significant interest due to their © 2017 American Chemical Society

Fe2O3 + x Li+ + x e− ⇌ LixFe2O3

(1)

Received: June 5, 2017 Revised: August 15, 2017 Published: August 16, 2017 7794

DOI: 10.1021/acs.chemmater.7b02315 Chem. Mater. 2017, 29, 7794−7807

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Chemistry of Materials

evaluated as arsenic sorbents.42 A previous study reported growth of γ-Fe2O3 nanosheets on graphene and explored their absorption of electromagnetic radiation.43 γ-Fe2O3 nanosheets for cathode materials have not been previously reported. The γFe2O3 phase is of particular interest and importance as a cathode material since the γ-Fe2O3 phase contains cation vacancies within the structure that provide the potential to increase charge storage sites and facilitate cation diffusion.9,14 In this work, we report the synthesis and unique electrochemical Li-ion charge storage properties of γ-Fe2O3 nanosheets and have evaluated their performance as cathode materials for Li-ion batteries. The γ-Fe2O3 nanosheets are transformed to NPs at relatively low temperatures (200 °C). Analysis of the structure and composition of the material supports that ethylene glycol is present within the nanosheets, and the transformation of γ-Fe2O3 from a nanosheet to an NP occurs with removal of structural H2O. Our results show that γFe2O3 nanosheets exhibit significantly higher capacities and improved rate capabilities compared with γ-Fe2O3 NPs. Kinetic analysis shows that the significantly improved charge storage properties, particularly at high rates, result from a surface-based charge storage mechanism.

Iron oxides show reasonable cathodic voltages (∼2.5 V vs Li), and the theoretical capacities for a one-electron reaction per iron center (x = 2 in eq 1) are 336 mA h g−1 for Fe2O3 and 232 mA h g−1 for Fe3O4, which are significantly higher than the capacities of current commercial cathodes (e.g., ∼140−170 mA h g−1 for LiCoO2 and LiFePO4).15,16 Despite their high theoretical capacities, iron oxide cathodes typically exhibit very low Li-ion capacities, which is attributed to their low solid-state Li-ion diffusion and low electronic conductivity.13 Prior work reported that the specific capacities of γ-Fe2O3 and Fe3O4 nanoparticle (NP) cathodes were between 1 and 15 mA h g−1.9,11 Several approaches have been explored to increase the capacity of FeOx cathodes. Cation defects have been shown to significantly influence the electrochemical properties of iron oxides.17,18 Substitution of a fraction of the Fe3+ sites with highly oxidized Mo6+ to increase the concentration of cation vacancies was used as a route to increase the capacity of defect spinel γ-Fe2O3 NPs.9 Similarly, vanadium (V5+) substitution into FeOx aerogels has been demonstrated to result in increased cathode capacities.18 Hollow FeOx NPs with a significant concentration of cation defects have shown improved capacities and stability.14 In addition, prior work has shown that decreasing the size of αFe2O3 NPs12,19 and γ-Fe2O3 NPs20 resulted in increased capacity. Two-dimensional (2D) materials consisting of a single layer or a few layers of atoms exhibit remarkable and in many cases vastly superior properties compared with their bulk counterparts and other types of nanomaterials, such as NPs and nanotubes.21−25 For example, the 2D structure of metal chalcogenides enhances the conductivity by a factor of ∼102−103 compared with that of their bulk forms.26 2D materials can provide improved electrochemical energy storage on the basis of their unique features, including quantum confinement,27 exposure of the entire surface to the electrolyte,28 reduced or no solid-state ion diffusion,29 anisotropic growth directions,30 and their ability to accommodate structural strain.31 Different exposed facets can give different redox potentials, charge transfer rates, and diffusion coefficients.32 Slow diffusion of ions in the solid state has been identified as the rate-limiting step in conventional cathode materials.4 Surface-based charge storage in 2D materials circumvents the need for solid-state ion diffusion, which enables rapid charging and discharging.29,33 Single-crystalline nanosheets of LiFePO4 with a high proportion of (001) facets were demonstrated to provide excellent rate capabilities, which was attributed to the reduced diffusion distance for Li ions.34 Single- and few-layer forms of TiO2 and TiS2 have been shown to exhibit significantly improved Li ion storage capacity at high rates enabled by capacitive charge storage.28,29 Although nanosheet architectures have low volumetric energy densities, the understanding of charge storage in 2D materials can also contribute to the design of layered cathodes with controlled interlayer structures and dynamics.35,36 Iron oxide nanosheets with α-Fe2O337 and Fe3O438 phases have been reported. Although α-Fe2O3 nanosheets have been studied as anode materials for Li ion batteries in the low electrochemical potential range,37 prior studies have not evaluated iron oxide nanosheets as cathode materials. Nanosheets of iron oxyhydroxides δ-FeOOH,39,40 and γ-FeO(OH)41 and characterization of their magnetic properties has been reported. The magnetic properties of γ-Fe2O3 nanosheets have been previously studied.41 γ-Fe2O3 nanosheets have also been

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RESULTS AND DISCUSSION Iron Oxide Nanosheets: Synthesis and Structural Analysis Using Microscopy. FeOx nanosheets were synthesized by adapting an approach used to synthesize δ-FeOOH nanosheets.39,40 The two-step synthesis process consists of (i) formation of Fe(OH)2 nanosheets within an ethylene glycol/ H2O solution followed by (ii) treatment in a H2O2 solution to oxidize Fe2+ to Fe3+ and convert Fe(OH)2 to FeOx (additional synthesis details are provided in the Supporting Information, Figures S1 and S2 and supporting text). Representative transmission electron microscopy (TEM) and scanning electron microscopy (SEM) images of the as-prepared FeOx nanomaterial are presented in Figure 1 and show the nanosheet structure. The nanosheets have approximate lateral dimensions of 50 nm and a thickness of ∼1 nm (Supporting Information, Figure S3). High-resolution TEM (Figure 1C) shows that the nanosheets consist of multiple crystallites. Observed lattice spacings of 2.5 and 3.0 Å are consistent with the (311) and (220) planes of γ-Fe 2 O3 (JCPDS card no. 39-1346), respectively. As discussed below, the assignment of the γFe2O3 phase is further supported by Raman spectroscopy. From the high-resolution TEM images, higher numbers of crystallites with lattice spacings of 2.5 Å were observed, which suggests that the nanosheet growth in the lateral dimension occurs preferentially along the (311) planes of the crystal. Pores were also observed within the nanosheets (Supporting Information, Figure S4). On the basis of the unit cell dimensions of γ-Fe2O3 (a = b = c = 8.35 Å) and thickness from TEM images (Figure S3), the nanosheets were ∼1 unit cell thick, indicating very few layers of atoms exist within the sheet. Effect of Inclusion of Ethylene Glycol on the Iron Oxide Structure and Morphology. Layered materials such as graphene, MoS2, and numerous other materials have weak van der Waals interactions between the layers that allow these materials to be relatively easily formed into single- or few-layer sheets by exfoliation.21,27 Unlike these layered materials, iron oxide phases, including γ-Fe2O3, Fe3O4, and α-Fe2O3, do not possess layered crystalline structures, and therefore, additional compounds are needed to facilitate the formation of an iron 7795


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performance with a specific capacity of ∼15 mA h g−1 (Supporting Information, Figures S6 and S7). The effect of ethylene glycol on facilitating the formation of iron oxide nanosheets may be related to the complexation of ethylene glycol with aqueous iron species during the hydrolysis and condensation reactions; however, further work is needed to understand the specific role of ethylene glycol in the synthesis mechanism. Prior work suggested that ethylene glycol chelates Fe2+ in the reactive FeO5 group at {010} facets of LiFePO4, which suppresses crystal growth along the {010} surfaces.44 The complexation of ethylenediaminetetraacetic acid (EDTA) with aqueous Fe3+ species was found to increase the kinetics of condensation reactions,45 which suggests that ethylene glycol could similarly alter the reaction kinetics. Complexation of ethylene glycol with aqueous Ti species was also previously suggested as an important factor in the synthesis of TiO2 nanosheets.46 Effect of Thermal Treatments. Following the formation of the γ-Fe2O3 nanosheets, we investigated the effect of thermal treatments on the phase and morphology. Temperature treatments have been shown to significantly alter the structure and properties of iron oxides.47 Thermogravimetric analysis (TGA) of the as-prepared γ-Fe2O3 nanosheets (Supporting Information, Figure S9) showed a gradual mass loss up to ∼400 °C followed by a more significant mass loss from ∼400 to ∼480 °C. The mass loss observed from 25 to 400 °C in the TGA is generally consistent with a prior TGA study of iron oxides48 and is attributed to removal of loosely bound physisorbed H2O at lower temperatures followed by removal of strongly bound chemisorbed or “structural” H2O. Prior work clearly supports that metal oxides can have tightly bound structural H2O, for example, within hydrated V2O5 (V2O5·nH2O),49 that can significantly affect the material structure and properties. In addition to mass loss below 400 °C, the TGA of the asprepared γ-Fe2O3 nanosheets also showed a substantial mass loss from ∼400 to ∼480 °C. Given the important role of ethylene glycol within the synthesis of the nanosheets, we considered that, in addition to the presence of H2O within the structure, ethylene glycol may also be present within the samples at low temperatures. Attenuated total (internal) reflectance-Fourier transform infrared (ATR-FT-IR) spectra were obtained for the sample heated to 200 °C to determine if ethylene glycol was present and for the sample heated to 450 °C to evaluate if ethylene glycol was removed by treatment at a higher temperature. The infrared spectrum of the 200 °Ctreated sample shows the presence of absorptions which are attributed to vibrational modes of ethylene glycol, suggesting an interaction with the surface of the iron oxide (Supporting Information, Figure S10 and supporting text). The two vibrational bands observed between 1300 and 1000 cm−1 in the infrared spectra of the 200 °C-treated samples are attributed to the asymmetric C−O stretching mode, νas(CO), and symmetric C−O stretching mode, νs(CO), of ethylene glycol50 that interacts with the iron oxide surface. Shifting of the ethylene glycol absorptions relative to their frequencies within the liquid phase as a result of interaction with the surface of the iron oxides is consistent with the results of a prior study by Jansen et al. that reported shifting of the νs(CO), as well as other vibrational modes of ethylene glycol, to higher frequencies upon interaction with the surfaces of a Rh(100) single crystal.51 The hydroxyl groups of ethylene glycol may potentially interact with the surface via either a monodentate or a bidentate configuration.51 The strong interaction of ethylene

Figure 1. Iron oxide nanosheets prepared using ethylene glycol and H2O: (A) transmission electron microscopy (TEM) image; (B) scanning electron microscopy (SEM) image; (C) high-resolution TEM image (inset: calculated electron diffraction pattern from fast Fourier transform (FFT) of the high-resolution TEM image); (D) unit cell of γ-Fe2O3 showing Fe and O coordination and cation vacancies.

oxide nanosheet rather than a nanoparticle form. Ethylene glycol has been reported to play an important role in the synthesis of LiFePO434 and δ-FeOOH39 nanosheets. To evaluate the effect of ethylene glycol on the synthesis of γFe2O3 nanosheets, for comparison, samples were synthesized using the same conditions except using H2O in place of ethylene glycol. When ethylene glycol was omitted during the synthesis, a mixture of nanorods and NPs (particle sizes of ∼50−100 nm) consistent with α-FeOOH (goethite) and Fe3O4 phases formed, and the material exhibited poor electrochemical 7796


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Figure 2. Microscopy of thermally treated samples: TEM images of (A) EG/H2O-200, (B) EG/H2O-350, and (C) EG/H2O-450; SEM images of (D) EG/H2O-200, (E) EG/H2O-350, and (F) EG/H2O-450.

% was also within experimental error of the amount of ethylene glycol determined from elemental analysis (8.7 ± 1.4 wt %). Using the calculated mass loss from the TGA data, molecular formulas using the general formula FeOx(C2H6O2)y(OH2)w were determined (Table S1). The as-prepared material was determined to have a formula of Fe2O3(C2H6O2)0.16(OH2)1.50, and the samples heated to 120 °C under vacuum and to 200 °C were determined to have formulas of Fe2O3(C2H6O2)0.17(OH2)1.08 and Fe2O3(C2H6O2)0.15(OH2)0.71, respectively. Notably, the relative amounts of ethylene glycol within these samples were determined to be very similar; however, the H2O content within the samples decreased upon heating to higher temperatures. The strong interaction of ethylene glycol with the iron oxide surface, as supported by FT-IR data described above, is consistent with ethylene glycol being retained within the structure. The TGA data of the 120 °C vacuum-treated and 200 °Ctreated samples also showed the presence of a mass loss region between ∼200 and 325 °C. The differences between the 120 °C vacuum-treated and 200 °C-treated samples in this region are likely due to loss of different amounts of structural H2O present within the samples. Different coordination environments of H2O to Fe3O4 surfaces have been previously reported.53 We also consider that H2O within the structure is likely in the form of both OH groups and H2O, consistent with prior analysis of hydrated RuO2,54 and therefore, the formula may be more accurately represented as Fe2O3(C2H6O2)y(OH)z(OH2)w. Further investigation is needed to determine the specific nature of OH, H2O, and ethylene glycol within the iron oxide nanomaterials. On the basis of the transformation temperatures observed from TGA, we investigated relatively low temperature treatments (200 °C) along with higher temperatures (350 and 450 °C) to determine the effect of thermal treatments on the structure and electrochemical properties of the material

glycol with oxide surfaces is further supported by calculations that reported a strong, negative binding energy of −1.0 eV mol−1 for binding of ethylene glycol with the (010) facet of LiFePO4.34 In contrast with the data for the 200 °C sample, the ATR-FT-IR spectrum of the 450 °C-treated sample shows no absorptions corresponding to ethylene glycol bands with respect to the normalized relative intensities of iron−oxygen stretching modes, ν(Fe−O) (Figure S10). The absence of ethylene glycol absorptions in the 450 °C-treated sample supports the conclusion that the major mass loss observed from ∼400 to ∼480 °C in the TGA data results from the loss of ethylene glycol from the structure. On the basis of our X-ray diffraction (XRD) and Raman analysis described below, the iron oxide at 450 °C is within the α-Fe2O3 phase. We also note that a previous study that investigated iron oxide nanoparticles synthesized in the presence of glycols also showed mass loss above 400 °C.52 Based on the presence of both H2O and ethylene glycol within the structure and the removal of ethylene glycol in the 400−480 °C region, we considered the structure of the asprepared, 120 °C vacuum-treated, and 200 °C-treated samples to be generally represented as composed of Fe2O3, ethylene glycol, and H2O. The general mechanisms of mass loss are presented in the Supporting Information. The mass loss from the TGA within the 25−400 °C region (region I) was used to calculate the amount of H2O within the structure, and mass loss within the 400−600 °C region of the TGA (region II) was used to calculate the amount of ethylene glycol within the structure (Supporting Information, Table S1). Using this approach, the amount of ethylene glycol within the 200 °C-treated sample was determined to be 5.2 ± 2.7 wt %, which is in good agreement with elemental analysis of the percentage of carbon which corresponds to 5.4 ± 1.5 wt % ethylene glycol within the sample. For the nanosheets heated to 120 °C, the amount of ethylene glycol determined from TGA analysis of 5.6 ± 2.5 wt 7797


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Chemistry of Materials compared with the as-prepared nanosheet sample. Our motivation for investigating low temperature treatments was that in a number of oxide systems, namely, TiO2 and RuO2, relatively low temperature treatments that result in structures with significant disorder have been shown to result in improved electrochemical properties compared with the electrochemical properties of highly crystalline materials obtained at higher temperatures.4,5 The TEM and SEM images for the samples heated to 200, 350, and 450 °C (Figure 2) show that the as-prepared γ-Fe2O3 nanosheets (Figure 1) transform into NPs upon thermal treatment. The 200 °C sample showed primarily NPs; however, a small number of nanosheets were observed within the TEM images, indicating a partial transformation to NPs at 200 °C. At higher temperatures of 350 and 450 °C, only nanoparticles were observed. From the TEM images of the thermally treated samples (Figure 2), the NPs are significantly smaller than the lateral dimensions of the nanosheets (Figure 1). The TEM images of the thermally treated samples shown in Figure 1 also indicate that the sizes of the NPs for the 450 °C-treated sample are larger than the sizes of the NPs of the 350 °C-treated sample, indicating that growth of the NPs occurs at higher temperatures. On the basis of the transformation of the γ-Fe2O3 nanosheets to NPs at 200 °C, we further investigated the stability of the nanosheets at lower temperatures. After treatment at 120 °C under vacuum, which is the treatment used for our electrochemical testing, the structure and phase of the γ-Fe2O3 nanosheets remained similar to those of the as-prepared material, as shown from SEM images, XRD, and Raman spectroscopy (Supporting Information, Figures S11−S13). The observation that the conversion from a nanosheet to an NP occurs at a relatively low temperature (200 °C) suggests a low transformation energy. Because the crystal structure of γ-Fe2O3 is not intrinsically layered, we considered that the surface of the nanosheet must be stabilized to prevent surface energy-driven transformation to an NP. From our analysis of the TGA data, the sample at 200 °C has a lower amount of H2O than the 120 °C sample, and the amounts of ethylene glycol within both the 120 and 200 °C samples are relatively similar. We consider that structural H2O, likely surface hydroxyl groups, and ethylene glycol play an important role in stabilizing the nanosheet form of γ-Fe2O3; however, further investigation is needed to determine the specific nature of H2O and ethylene glycol and their effect on stabilizing γ-Fe2O3 nanosheets. Previous studies of FeOOH phases identified that adsorbed water altered the effective surface enthalpy.55,56 The importance of the interaction of coordinated species on a nanosheet structure is supported by previous studies that reported that rolling up of titanate nanosheets was related to the interaction of coordinated Na+ or H2O with the TiO6 octahedron.57 Numerous prior studies have reported the importance of chelating groups, including ethylene glycol and other ligands in solution, to control the iron oxide crystal growth direction, phase, and morphology,34,39,40,44,45 and surface interactions were determined to be important in the stabilization of Au nanosheets.58 X-ray Diffraction Analysis. Powder XRD was used to characterize the γ-Fe2O3 nanosheets and heated samples (200, 350, and 450 °C) as shown in Figure 3. The XRD peak positions for the as-prepared FeOx nanosheets are close to the peak positions of both Fe3O4, magnetite (JCPDS card no. 750033), and γ-Fe2O3, maghemite (JCPDS card no. 39-1346).

Figure 3. Comparison of powder X-ray diffraction (XRD) data of iron oxide samples prepared using ethylene glycol (EG) and H2O and treated at different temperatures. Shown also are reference data for αFe2O3 (hematite; JCPDS card no. 33-0664), Fe3O4 (magnetite; JCPDS card no.75-0033), and γ-Fe2O3 (maghemite; JCPDS card no. 39-1346).

The XRD (311) peak of the as-prepared FeOx nanosheets is slightly closer to that of the Fe3O4 phase. Mild heating of the as-prepared nanosheets at 120 °C vacuum results in a small shift of the (311) peak at 2θ = 35.67° to 2θ = 35.63°, and the XRD peak position for the 120 °C vacuum-treated sample is consistent with that of γ-Fe2O3 (Figure S12). The small shift of the (311) peak between the as-prepared 120 °C vacuumtreated samples was confirmed by using graphite as an internal standard. On the basis of the TGA analysis discussed above, which showed that mass loss occurred under vacuum treatment at 120 °C, we attribute the small (311) peak shift to removal of bound H2O within the structure that results in local strain that alters the spacing of the iron and oxygen planes (Δd‑spacing = 0.003 Å). Prior work showed that the iron and oxygen planes were shifted by 0.3−0.5 Å in thin layers that interacted with a Pt substrate.59 The challenge of clearly distinguishing between Fe3O4 and γFe2O3 phases using XRD is related to the similarity of the Fe3O4 and γ-Fe2O3 crystal structures, which have very small differences between the unit cells (“a” spacings of Δ = 0.01 Å between the two phases).60 The broadening of the XRD peaks of nanocrystalline Fe3O4 and γ-Fe2O3 makes assignment of these phases more challenging.61,62 High-resolution TEM and Raman spectroscopy, discussed below, support that the local structure of the as-prepared iron oxide nanosheets is more consistent with γ-Fe2O3 rather than Fe3O4. The crystallite size of the as-prepared nanosheets of 10.9 nm from Scherrer analysis of the XRD peak width (Table 1) and the crystallites observed in the high-resolution TEM image (Figure 1) confirm that the nanosheets are polycrystalline. XRD measurements of thermally treated FeOx nanosheets are shown in Figure 3. The reflections of the sample heated at 200 °C are more consistent with γ-Fe2O3 rather than Fe3O4. 7798


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Table 1. Average Crystallite Sizes, Brunauer−Emmett−Teller (BET) Surface Areas, Cumulative Pore Volumes, and Electronic Conductivities (σelec) of Iron Oxide Nanomaterials Synthesized Using Ethylene Glycol/H2O and Heated to Different Temperatures As Described in the Texta material ID

crystallite sizeb (nm)

BET surface area (m2 g−1)

cumulative pore volume (cm3 g−1)

mean pore diameter (nm)

FeOx-EG-as-prep FeOx-EG-200 FeOx-EG-350 FeOx-EG-450 γ-Fe2O3-np

10.9 12.3 12.9 23.7

138.9 98.5 70.1 42 52

0.46 0.44 0.53 0.42 0.19

12.8 17.2 29.9 39.6 14.5

σelec (S/cm) 2.2 7.2 7.2 2.1

× × × ×

10−8 10−9 10−9 10−8

a “γ-Fe2O3-np” denotes commercial γ-Fe2O3 nanoparticles. bThe crystallite size was determined via Scherrer analysis of X-ray diffraction as described in the Materials and Methods.

Upon heating, the crystallite size increases slightly from 10.9 to 12.3 nm during treatment at 200 °C (Table 1) as the nanosheets transform to NPs. The peaks for the sample treated at 350 °C are similar to those at 200 °C, with the emergence of additional low-intensity peaks at 2θ = 24.0°, 33.1°, and 40.9° that are consistent with α-Fe2O3. From XRD, the 350 °Ctreated sample is therefore predominately γ-Fe2O3 with a small amount of α-Fe2O3. The reflections in the sample treated at 450 °C are consistent with α-Fe2O3 (JCPDS card no. 33-0664), thus indicating conversion of γ-Fe2O3 into α-Fe2O3. Respective average crystallite sizes of 12.9 and 23.7 nm measured at 350 and 450 °C are in general agreement with the TEM images (Figure 1). Analysis Using Raman Spectroscopy. Raman spectroscopy was used to characterize the frequencies of the vibrational modes of the iron oxide nanomaterials. Since the vibrational mode frequencies are highly sensitive to the local structure, Raman spectroscopy provides information on the local structure and has been shown to be more useful than diffraction techniques to probe the structural differences between nanocrystalline iron oxides, particularly γ-Fe2O3 and Fe3O4.63 The structures of Fe3O4 and γ-Fe2O3 are closely related; however, local structure differences provide the basis for differentiating the two phases using vibrational spectroscopy. At room temperature, Fe3O4 belongs to the inverse spinel cubic structure under the Oh7 (P4332) point group and consists of two nonequivalent Fe positions in the unit cell: “A” positions (occupied by Fe3+) involve tetrahedral Fe coordination, and “B” positions (equally populated by Fe3+ and Fe2+ ions) involve octahedral Fe coordination.64,65 Within the Fe3O4 structure, the presence of Fe2+ within octahedral sites is more stable as explained by crystal field theory.66 The space group of γ-Fe2O3 can be identical to that of Fe3O4 or symmetry-reduced from cubic (P4332) to tetragonal (P412121) with vacancy ordering within the structure.67,68 The γ-Fe2O3 structure has cation vacancies either randomly distributed throughout the tetrahedral and octahedral sites69 or preferentially located on the octahedral sites (as shown in Figure 1D).67,70 The structure of γ-Fe2O3 can be represented as (Fe3+)[Fe5/33+□1/3]O4, where “(Fe3+)” refers to tetrahedral coordination, “[Fe5/3]” designates octahedral coordination, and “□” denotes an octahedrally coordinated cation vacancy.71 The different local structures of Fe3O4 and γ-Fe2O3 affect the frequencies of their vibrational modes, and previous studies have shown that Raman spectroscopy can be used to differentiate γ-Fe2O3 and Fe3O4.63,72 Magnetite has five Raman active modes (A1g + Eg + 3T2g)73 for the smallest unit cell. The band centered at ∼665 cm−1 (Figure 4) observed in the Raman spectrum of Fe3O4 exhibits A1g symmetry and is

Figure 4. Comparison of Raman spectra in the 100−900 cm−1 spectral region of iron oxide samples prepared using ethylene glycol and H2O and treated at different temperatures; shown also are spectra for reference phases, as labeled.

attributed to vibrational modes consisting of symmetric stretching of oxygen atoms along Fe−O bonds in the A positions.73,74 In contrast, the Raman spectrum of γ-Fe2O3 exhibits a broad band which can be deconvoluted into two bands centered at 665 and 721 cm−1 (Figure S5, Supporting Information). Prior work supports that the Raman spectrum of γ-Fe2O3 within the 600−800 cm−1 region can be deconvoluted into a band at ∼665 cm−1, which is attributed to in-phase vibrations of FeO4 tetrahedra in the absence of cation vacancies,73,74 and a band at ∼721 cm−1, which is attributed to vibrational modes of local Fe−O structures in the vicinity of cation vacancies.63 The relative intensities of these two bands may be related to the degree of cation vacancies within the structure.63 In addition, the band at 538 cm−1 (T2g symmetry) within Fe3O4 is not present within γ-Fe2O3 (Figure S5), which further supports the differentiation of Fe3O4 and γ-Fe2O3 using Raman spectroscopy. The Raman spectra for the synthesized iron oxide nanomaterials in the 100−900 cm−1 spectral region are shown in Figure 4 along with the spectra of commercially available γ-Fe2O3, Fe3O4, and α-Fe2O3 for comparison. The frequencies of the Raman bands of the as-prepared nanosheets are consistent with γ-Fe2O3 rather than Fe3O4 as supported by the comparison 7799


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Figure 5. Galvanostatic charge and discharge voltage profiles (second cycle) of iron oxide samples: (A) comparison of EG/H2O-as-prep and commercial γ-Fe2O3 at a mass-normalized current of 3 mA g−1; (B) comparison of EG/H2O-as-prep with samples heated to different temperatures (200, 300, and 400 °C) at a mass-normalized current of 3 mA g−1; (C) comparison of EG/H2O-as-prep and commercial γ-Fe2O3 at a massnormalized current of 300 mA g−1; (D) comparison of EG/H2O-as-prep with samples heated to different temperatures (200, 300, and 400 °C) at a mass-normalized current of 300 mA g−1 (electrolyte 1 M LiPF6 in EC/DEC, 1:1, v/v; counter/reference metallic Li).

consistent with α-Fe2O3, which is also consistent with data obtained from XRD. Electrochemical Properties of Iron Oxide Nanosheets. Electrochemical tests were performed on the γ-Fe2O3 nanosheets and heat-treated samples (200, 350, and 450 °C). The electrodes were dried under vacuum at 120 °C prior to electrochemical testing to prevent H2O contamination, which can react with the PF6 anion and lead to degradation within the cell. As discussed above, the nanosheet structure was maintained and the local structure remained γ-Fe2O3 after the 120 °C vacuum treatment (Figures S11−S13). Shown in Figure S14 (Supporting Information) is the comparison of the voltage profiles for the first, second, and fifth cycles. The as-prepared γFe2O3 nanosheets showed an initial discharge capacity of 190.1 mA h g−1, which decreased to 147.6 mA h g−1 for the second cycle. The capacity decrease after the first cycle, as is often observed, is attributed to formation of a solid−electrolyte interface (SEI) layer and/or structural changes occurring during the first few cycles. The thermally treated samples showed more significant changes in the voltage profiles and capacities from the first to second cycle, suggesting more substantial structural changes occur for these materials over the first charge/ discharge cycles. Shown in Figure 5A are voltage profiles (second cycle) of the as-prepared γ-Fe2O3 nanosheets compared with commercial γFe2O3 nanoparticles from galvanostatic charging/discharging at a mass-normalized current of 3 mA g−1 between 1.5 and 4.2 V vs Li. The γ-Fe2O3 nanosheets exhibited a discharge capacity of 147.6 mA h g−1 (second cycle), which was significantly (4.5

with the commercial γ-Fe2O3 and Fe3O4 samples and previously reported spectra for these phases.63 The similarity of the band structure of the as-prepared nanosheets to that of γ-Fe2O3 within the 600−800 cm−1 region is further supported by the deconvolution of the spectrum into two peaks at 665 and 721 cm−1 (Figure S5), with the higher frequency 721 cm−1 band attributed to vibrations of FeO4 tetrahedra in the local vicinity of cation vacancies as described above. The local γ-Fe2O3 structure of the as-prepared iron oxide nanosheets can be notated as (Fe3+)[Fe5/33+□1/3]O4, showing that the nanosheets contain cation vacancies within the structure. We note that the assignment of the phase of the as-prepared FeOx nanosheets to γ-Fe2O3 rather than the related feroxyhyte, δ-FeOOH, structure is supported by the Raman spectra, where the low-frequency spectral region does not show bands at 222 and 292 cm−1 which have been reported for feroxyhyte.75 Raman spectroscopy was also used to probe the effect of temperature treatments on the local structure. As shown in Figure 4, the Raman spectrum of the sample heated to 200 °C (EG/H2O-200) is consistent with the local structure of γFe2O3. Heating to 350 °C (EG/H2O-350) results in a local structure similar to that present within the 200 °C-treated sample; however, a small additional peak at 407 cm−1 is observed, indicating a local structure of α-Fe2O3 is present to a small degree. The presence of predominately γ-Fe2O3 and a small amount of α-Fe2O3 within the 350 °C-treated sample is consistent with the XRD results (Figure 3). Further heating to 450 °C (EG/H2O-450) results in Raman bands that are 7800


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Table 2. Electrochemical Properties of Iron Oxide Nanomaterials: Average Discharge Capacity from the Second Cycle at Different Mass-Normalized Currents of 3, 30, 60, 150, and 300 mA g−1 average discharge capacity (mA h g−1) at various mass-normalized currents

a b

material ID

3 mA g−1

30 mA g−1

60 mA g−1

150 mA g−1

300 mA g−1

capacity retentiona (%)

Coulombic efficiencyb (%)

EG/H2O-as-prep EG/H2O-200 EG/H2O-350 EG/H2O-450 γ-Fe2O3-comm

147.6 100.2 83.1 66.2 34.2

136.7 90.6 80.9 53.3 30.4

120.9 80.3 70.4 46.3 27.8

102.0 73.6 52.9 43.0 24.0

82.2 64.8 46.2 40.4 19.2

86.0 86.7 90.9 82.6 86.6

98.2 99.2 99.2 98.6 99.4

Average capacity retention as a percentage of the discharge capacity of the 2nd and 40th cycles at a mass-normalized current of 10 mA g−1. Average Coulombic efficiency from the 2nd to 40th cycles at a mass-normalized current of 10 mA g−1.

times) higher than the capacities of commercial γ-Fe2O3 NPs (31.8 mA h g−1) and iron oxide NPs from prior work (∼5−15 mA h/g).9,18 Shown in Figure 5B is the comparison of the capacities of γ-Fe2O3 nanosheets at a mass-normalized current of 3 mA g−1 to those of the FeOx samples heated to 200, 350, and 450 °C. All samples showed average discharge voltages in the range of ∼2.4 V vs Li. Notably, the γ-Fe2O3 nanosheets exhibited ∼50% higher capacities than the γ-Fe2O3 NPs obtained from thermal treatment at 200 °C. The nanosheet architecture of the material enabled the storage of 0.9 mol of Li/mol of Fe2O3 (Li0.9Fe2O3), which was significantly higher than 0.6 and 0.2 mol for the 200 °C NPs and commercial NPs, respectively. In addition to exhibiting high capacities at low charge and discharge rates, the iron oxide nanosheets were tested at a 100fold higher current to evaluate the materials’ rate capabilities at fast charging and discharging times. The discharge capacities at different mass-normalized currents are presented in Table 2. The discharge time of the γ-Fe2O3 nanosheets ranged from 49 h (C/49) at 3 mA g−1 to 17 min (3.6C) at 300 mA g−1. As is typically observed, the discharge capacity decreases at faster charge/discharge times. However, the γ-Fe2O3 nanosheets showed remarkably high discharge capacities at high rates. The specific discharge capacities of the γ-Fe2O3 nanosheets compared with commercial γ-Fe2O3 NPs and the thermally treated samples at a mass-normalized current of 300 mA g−1 are shown in parts C and D, respectively, of Figure 5. Even at high rates corresponding to an average discharge time of 17 min (3.6C rate), the iron oxide nanosheets showed a capacity of 82.2 mA h g−1, which was 4.3 times higher than the capacity of commercial γ-Fe2O3 NPs and 27% higher than that of the γFe2O3 NPs obtained from thermal treatment at 200 °C. To further probe the capabilities of the γ-Fe2O3 nanosheets at rapid charge and discharge times, we tested the capacities of the γFe2O3 nanosheets using a filter paper separator rather than the commercial separator used for other tests (Supporting Information, Figure S16, Table S2). Using this configuration, the γ-Fe2O3 nanosheets showed capacities of 111.9 mA h g−1 (2.8C rate, 21 min discharge time) and 93.2 mA h g−1 (6.7C rate, 9 min discharge time) at mass-normalized currents of 300 and 600 mA g−1, respectively, which further demonstrates that the γ-Fe2O3 nanosheets exhibit high capacities at very fast charge and discharge times. The γ-Fe2O3 nanosheets show significantly higher capacities of up to 147.6 mA h g−1 compared with γ-Fe2O3 NPs from prior studies, which showed capacities of ∼15 mA h g−1.9 We note that the capacities obtained for the γ-Fe2O3 nanosheets were achieved within electrodes prepared using a typical slurrycast electrode with 10 wt % carbon black. Hollow γ-Fe2O3 nanoparticles with a significant proportion of cation vacancies

that were within a conductive matrix of carbon nanotubes with significantly higher carbon content (20 wt % carbon) were shown to exhibit higher capacities.14 Higher amounts of cation vacancies9,14 in addition to using a more conductive matrix may further increase the capacities of the γ-Fe2O3 nanosheets. In addition to higher specific discharge capacities and improved rate capabilities compared with NPs, the γ-Fe2O3 nanosheets exhibited high capacity retention of 86.0% of the initial capacity after 40 cycles (2nd cycle to 40th cycle) as shown in Figure 6 and Table 2. Ex situ SEM images of the γ-

Figure 6. Capacity retention upon cycling for iron oxide samples prepared with ethylene glycol with different thermal treatments (electrolyte 1 M LiPF6 in EC/DEC, 1:1, v/v; counter/reference metallic Li; voltage range of 1.5−4.2 V vs Li; mass-normalized current of 30 mA g−1).

Fe2O3 nanosheet electrodes after 10 charge/discharge cycles (Supporting Information, Figure S15) support that the initial nanosheet architecture (Figure 1) is maintained after electrochemical cycling. The 350 °C-treated sample had a higher capacity retention of 90.9%; however, this is attributed to its lower initial capacity. The iron oxide nanomaterials also exhibited high Coulombic efficiencies of >98%, indicating high discharge/charge efficiencies and good interfacial stability with the electrolyte. Analysis of Factors That Contribute to Improved Electrochemical Performance of Iron Oxide Nanosheets. On the basis of the higher electrochemical performance of iron oxide nanosheet cathodes, we were particularly interested in the factors that resulted in the improved electrochemical performance with the aim to understand these factors and provide a route to design materials with improved properties. The phase 7801


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Figure 7. (A) Cyclic voltammograms (CVs) of as-prepared γ-Fe2O3 nanosheets, γ-Fe2O3 NPs (heated at 200 °C), and commercial γ-Fe2O3 nanoparticles (scan rate 0.5 mV/s; electrolyte 1 M LiPF6 in EC/DEC, 1:1, v/v; counter/reference metallic Li). (B) Comparison of CVs of asprepared γ-Fe2O3 nanosheets at different scan rates of 0.1, 0.5, 1.0, 5.0, and 10 mV s−1. (C) Capacity vs (scan rate)−1/2 for as-prepared γ-Fe2O3 nanosheets. (D) log(peak cathodic current) vs log(scan rate) for iron oxide nanosheets compared with commercial γ-Fe2O3 NPs and 200 °C γ-Fe2O3 NPs.

To probe the effect of the surface area, we determined the Brunauer−Emmett−Teller (BET) surface areas as well as pore size distributions and pore diameters for the nanomaterials (Table 1). The BET surface areas, cumulative pore volumes, and mean pore diameters for the as-prepared γ-Fe 2 O 3 nanosheets are significantly higher than those of the 200 °C γ-Fe2O3 NPs and commercial γ-Fe2O3 NPs. As the temperature treatment increases to 350 and 450 °C, the BET surface area decreases while the mean pore diameter increases. Plotting the specific discharge capacity vs the BET surface area for the different materials resulted in a generally linear trend (Supporting Information, Figure S17) suggesting that the charge storage capacity is predominately directly correlated with the surface area. To provide additional insight into the charge storage mechanism, we analyzed the kinetics of charge storage using the scan rate dependence of the cyclic voltammetric response that has been shown to allow differentiation of diffusioncontrolled and capacitive charge storage processes.28,29 Shown in Figure 7A are cyclic voltammograms (CVs) of γ-Fe2O3 nanosheets compared with commercial γ-Fe2O3 NPs and 200 °C γ-Fe2O3 NPs at a scan rate of 0.5 mV s−1. For both materials, the CVs exhibit predominant anodic peaks at ∼2.3 V vs Li. For the cathodic region, the γ-Fe2O3 nanosheets and 200 °C γ-Fe2O3 NPs have peaks at ∼1.7 V vs Li, and the commercial γ-Fe2O3 NPs exhibit a peak at a slightly lower voltage of ∼1.5 V vs Li. The potentials observed in the CV of γFe2O3 nanosheets are consistent with the voltage profiles from

of the material plays an important role, particularly in light of prior work which showed that γ-Fe2O3 containing cation vacancies exhibited higher capacities than the closely related Fe3O4 structure that did not contain cation vacancies.9,14 The observation that the nanosheets have a γ-Fe2O3 phase which contains cation vacancies may result in both improved charge storage sites within the material and improved solid-state Li ion diffusion.9 Beyond the phase, the comparison of the structure and electrochemical properties of the as-prepared nanosheets and 200 °C-treated samples shows that the improved properties are directly related to the nanosheet architecture. Both the asprepared nanosheets and 200 °C-treated sample have similar γFe2O3 phases as supported by XRD (Figure 4) and Raman spectroscopy (Figure 5). However, the comparison of the TEM and SEM images for the two materials (Figures 1 and 2) clearly shows the different morphologies: the as-prepared γ-Fe2O3 exists as a nanosheet, whereas the 200 °C-treated γ-Fe2O3 material is a NP. Similar crystallite sizes are also observed (Table 1). However, the γ-Fe2O3 nanosheet has significantly improved capacities compared with the NP form at both low and fast rates. For the NPs obtained from treating the samples to 450 °C, the phase is transformed from γ-Fe2O3 to α-Fe2O3, and lower electrochemical performance is obtained. Small changes in the particle size are also observed, which can also influence the electrochemical performance as supported by prior work.19 7802


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Chemistry of Materials the galvanostatic tests (Figure 5). In the CV, the γ-Fe2O3 nanosheets exhibit dramatically higher mass-normalized currents compared to the γ-Fe2O3 NPs, indicating that more charge can be stored within the nanosheet architecture for the same mass of material and at the same time scale. In Figure 7B, the CVs at different scan rates for the γ-Fe2O3 nanosheets are shown (CVs vs scan rate for commercial γFe2O3 NPs and 200 °C γ-Fe2O3 NPs are included in the Supporting Information, Figures S18 and S19). At higher scan rates, the peaks for the γ-Fe2O3 nanosheets are broader and shifted to positive potentials for the anodic scan, which is consistent with ohmic losses in the electrochemical cell including the electrode. Intercalation of ions within a material involves solid-state ion diffusion that can be described by a semi-infinite diffusion process where the peak current is proportional to the square root of the scan rate.76 The time dependence of the current can be useful in understanding the energy storage mechanism. In the analysis proposed by Trasatti et al.,77 the sweep rate dependence of the capacity (Q(v)) can be described by adding the time-independent capacity (Qcapacitive) to a time-dependent capacity controlled by semiinfinite diffusion ((constant)v−1/2): Q (v) = Q capacitive + (constant)v−1/2

surface region, as supported by prior studies of hydrous RuO2 that reported charge storage within outer and inner surface regions.77 Electronic conductivity measurements (Table 1) were obtained to determine if this influenced the electrochemical performance, particularly at high rates. The electronic conductivity of the γ-Fe2O3 nanosheets, σelec of 2.2 × 10−8 S cm−1, was low but was 3 times higher than that of the 200 °C nanoparticle sample, σelec = 7.2 × 10−9 S cm−1. The higher electronic conductivity for the nanosheets may contribute to the improved rate capabilities. The origin of the higher electronic conductivity in γ-Fe2O3 nanosheets requires further investigation. Electronic conduction in iron oxides occurs via a small polaron hopping mechanism,78 and numerous factors contribute to the nanomaterial’s overall electronic conductivity, including the phase, defect density, surface structure, particle size, interparticle junctions, and morphology.79 In addition, quantum confinement effects and surface interactions can contribute to the electronic conduction within 2D materials.21,22

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CONCLUSIONS Iron oxide, γ-Fe2O3, nanosheets were determined to exhibit substantially improved electrochemical charge storage properties compared with NPs. The synthesis of γ-Fe2O3 nanosheets was enabled by inclusion of ethylene glycol as a structuredirecting agent during the synthesis. The nanosheets had lateral dimensions of ∼50 nm and thicknesses of ∼1 nm and therefore consisted of very few layers of atoms within the sheet. X-ray diffraction, Raman spectroscopy, and high resolution microscopy supported that the phase of the nanosheets was γ-Fe2O3, which contains cation defects. The nanosheets were found to be stable only under mild heating conditions (120 °C under vacuum). Temperature treatments of 200 °C resulted in transformation from nanosheets to an NP. The γ-Fe2O3 phase does not possess a layered crystal structure and is typically not observed as a nanosheet. Ethylene glycol was determined to be present within the γ-Fe2O3 nanosheets. The transformation from a nanosheet to a NP involved removal of H2O from the structure, and therefore, structural H2O and ethylene glycol may play an important role in stabilizing γ-Fe2O3 within a nanosheet form. We consider that this work supports the important role of surface interactions in the stabilization of nanosheet forms of phases that are not inherently twodimensional (i.e., that do not have weak van der Waals interactions between layers). The γ-Fe2O3 nanosheets demonstrated remarkably higher Li ion electrochemical charge storage at low rates: 4.6 times higher capacities than those of commercial γ-Fe2O3 NPs. Of particular significance is the remarkable rate capabilities of the nanosheets, which showed the ability to be rapidly charged and discharged at rates up to 6.7C (9 min discharge times) with high capacities. The capacity was directly correlated with the BET surface area. Kinetic analysis supported the conclusion that the nanosheets exhibit a primarily capacitive charge storage process. In addition, the electronic conductivity of the nanosheets was determined to be higher than that of the nanoparticles, which is attributed to the nanosheet structure. The significantly improved electrochemical properties of the γFe2O3 nanosheets therefore arise from the combination of the nanosheet architecture structure, which allows surface-based charge storage, reduced diffusion distances, unique electrical properties, and the cation-defective γ-Fe2O3 phase. The finding

(2)

Shown in Figure 7C is a plot of the capacity as a function of (scan rate)1/2 for the γ-Fe2O3 nanosheets. At low scan rates (v = 0.1−0.5 mV s−1), linear behavior is observed (labeled as “region 1”), consistent with a semi-infinite diffusion process. Extrapolation of this linear dependence to the y-intercept (v−1/2 = 0 or v = ∞) gives Qcapacitive. At fast scan rates (v > 0.5 mV s−1), significant deviation from this linear behavior occurs (labeled as “region 2”) due to increased ohmic losses and/or increased diffusion limitations stemming from both the electrode and electrolyte. Another method to evaluate the kinetics of an energy storage process is to determine the dependence of the current at a particular potential with the sweep rate. Prior work has established that a power law can be used to describe the relationship of the current, i, with the sweep rate, v: i = avb

(3)

where a is a constant and b can range from 0.5 (semi-infinite diffusion) to 1.0 (capacitive processes).28,29 The plot of log(i) vs log(v) provides the b-value from the slope. From the slope of the plot (Figure 7D), the γ-Fe2O3 nanosheets exhibit a b-value of 0.8 over the scan rate of 0.1−10 mV s−1, indicating that a majority of the current at the peak potential results from a capacitive-type process rather than semi-infinite diffusion. For comparison, the b-values of 200 °C γ-Fe2O3 NPs and commercial γ-Fe2O3 NPs are 0.7 and 0.6, respectively, which are lower than the b-value of the nanosheets but still primarily capacitive. The higher b-value of the γ-Fe2O3 nanosheets supports the conclusion that the primary charge storage mechanism is capacitive and that the synthesized nanosheets exhibit a higher degree of capacitive charge storage compared to γ-Fe2O3 NPs. The predominance of capacitive charge storage is consistent with the higher surface area of the γ-Fe2O3 nanosheets (Table 1) compared with Fe2O3 NPs, which can both increase the number of surface storage sites and decrease the solid-state diffusion distance required for the intercalation of Li+ in the structure. Within the γ-Fe2O3 nanosheets, there may exist a distribution of Li+ charge storage sites within the 7803


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by depositing a solution of the dried powder suspended in 2-propanol. SEM images were obtained using a Helios NanoLab 400 DualBeam field emission scanning electron microscope. TGA (TA Instruments Q50) was performed at a constant heating rate of 10 °C min−1 in air. Analysis of the mass loss from TGA was performed by using an αFe2O3 phase above 500 °C, which is consistent with prior work80 and our X-ray diffraction data at 450 °C. Nitrogen physisorption adsorption/desorption isotherms were measured using a Micromeritics ASAP 2020 surface area and porosimetry analyzer. Samples were degassed at 100 °C for 16 h prior to characterization. Cumulative pore volumes and mean pore diameters were calculated from the adsorption isotherm representing the volume in pores between 1.7 and 300 nm using the Barrett−Joyner−Halenda (BJH) method (Micromeritics Microactive software, version 4.02). Raman spectra were obtained with a Horiba LabRam HR Evolution confocal Raman spectrometer using a backscattering geometry and an 1800 grooves/mm grating. The 514 nm line of an argon ion laser was focused through an Olympus microscope with a 50× lens. Raman spectra are averages of 50 individual spectra that were obtained using a 24 s acquisition time. Laser-induced thermal effects were observed in prior Raman studies of iron oxides, and low laser powers are necessary to minimize spectral changes due to local heating.72,81 To avoid sample degradation, the laser power was reduced to below 1 mW using neutral density filters. Visual inspection of the samples, pre- and postanalysis, using white light illumination did not reveal any laser-induced changes. To further verify that the laser did not induce spectral changes, additional spectra were obtained at laser powers below 0.1 mW using the same acquisition and sampling conditions as above; no differences were observed compared to the spectra taken at 1 mW. The spectra of commercially available γ-Fe2O3 (details provided above), Fe3O4 (Alfa Aesar, 99.99%), and α-Fe2O3 (Sigma-Aldrich, >99%) were obtained for comparison to the synthesized nanomaterials. The Raman spectra were analyzed using Grams/AI spectral analysis software, version 9.2 (Thermo Scientific). Attenuated total (internal) reflectance-Fourier transform infrared (ATR-FT-IR) spectra were collected over the range of 4000−400 cm−1 using a Harrick Scientific (Pleasantville, NY) SplitPea attenuated total (internal) reflectance (ATR) microsampling accessory coupled to a Bruker (Billerica, MA) Tensor II FT-IR spectrometer that utilized a liquid nitrogen-cooled mercury cadmium telluride (MCT) detector and that was controlled using Bruker OPUS 7.5 software (version 7.5.18). The ATR accessory contained a diamond hemispherical internal reflection element (IRE), and the samples were brought into contact with the IRE using a loading of 0.5 kg. The infrared spectra represent the average of 64 individual scans collected at a spectral resolution of 4 cm−1. Analysis of the resulting spectra was performed using Grams/AI spectral analysis software (details provided above). To obtain electronic conductivities of the FeOx nanomaterials, 200−400 mg of powder was placed in a cell (Pred Materials, HS flat cell) with a 5 kg force spring. Two-point probe measurements were obtained using a constant voltage (±0.1 V) applied to the cell using an Arbin Instruments BT-2043 potentiostat/galvanostat. The current was monitored until quasi steady state was reached (∼3 min), and the resistance was determined using Ohm’s law, R = V/I The thickness and diameter of the pellet were measured using a micrometer (Mitutoyo, United States). The electrical conductivity, σelec (S cm−1), was calculated from the experimentally measured values using the equation σ = l/(RA), where l is the thickness of the sample, R is the measured resistance (Ω), and A is the cross-sectional area. Electrochemical Measurements. Electrodes were fabricated from a slurry composed of 80 wt % FeOx material (active material), 10 wt % conductive carbon (Timcal, Super C65), and 10 wt % binder (Aremka, Kynar HSV900) in 1-methyl-2-pyrrolidinone (NMP; Alfa Aesar, anhydrous). The slurry was stirred overnight and then cast onto a cleaned aluminum foil current collector. The resulting electrode sheet was dried overnight in a fume hood and then transferred to a 60 °C oven and allowed to dry overnight. Disks (0.5 in. diameter) of the dried electrode sheets were then pressed out of the aluminum foil and dried in a vacuum oven at 120 °C for 16 h. The electrodes had loadings of 1.5 ± 0.2 mg cm−2 of active material. For electrochemical

of unique electrochemical properties, the charge storage mechanism, and factors that result in stabilization of iron oxide nanosheets furthers the design of improved charge storage materials with high rate performance and lower cost through using nanosheet architectures to overcome limitations of rate-limiting solid-state ion diffusion. Further work is needed to understand the specific coordination environment of surface H2O and ethylene glycol and their role in stabilizing γ-Fe2O3 nanosheets and determine the nature of structural changes in the material over extended electrochemical cycles.

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MATERIALS AND METHODS

Materials. Sodium hydroxide (NaOH; 97%, EM Science) and ethylene glycol (C2H6O2; 99%, BDH) were obtained from VWR. Iron(II) sulfate heptahydrate (FeSO4·7H2O; 99%), sulfuric acid (H2SO4; 95%−98%), ethanol (C2H6O; 99.5%), and hydrogen peroxide (H2O2; 30%) were obtained from Sigma-Aldrich. γ-Fe2O3 NPs (Nanoarc) were obtained from Alfa Aesar. Iron Oxide Nanomaterial Synthesis. Iron oxide (FeOx) nanosheets were synthesized by modifying methods reported for synthesizing δ-FeOOH nanosheets. 39,40 For synthesizing FeOx nanosheets, 0.5 g of sodium hydroxide (NaOH) was dissolved in 5 mL of deionized (DI) water (≥13 MΩ cm), and then 50 mL of ethylene glycol (EG) was added to the solution. In a separate solution, 0.04 g of iron(II) sulfate heptahydrate (FeSO4·7H2O) was dissolved in 6 mL of 0.01 M sulfuric acid (H2SO4). The NaOH/EG and FeSO4/ H2SO4 aqueous solutions were then degassed with bubbling argon for at least 1 h. After this step, the FeSO4/H2SO4 solution was slowly added to the NaOH/EG solution using a pressure-equalizing funnel, and the reaction was allowed to proceed at room temperature for 3 h. The resulting material was collected by centrifugation at 5000 rpm (Sorvall LYNX 600, Thermo Scientific, 4304 rcf) for 15 min and then rinsed three times with a 50 vol % water/ethanol solution. The product was then suspended in 10 mL of water/ethanol, and 20 mL of 3 wt % H2O2 was added to the solution at a rate of 0.04 mL min−1 using a syringe pump. The as-prepared FeOx nanosheets were collected by centrifuge and then dried at room temperature. To determine the effect of ethylene glycol on the synthesis, the material was prepared by the same method, except 50 mL of DI water was used as the solvent instead of EG. The dried FeOx nanomaterials were used either as-prepared or after being heated within a muffle furnace (Thermo, Thermolyne) to 200, 350, or 450 °C for 24 h in air using a ramp rate of 5 °C min−1 from room temperature. The amount of carbon (wt %) within the as-prepared FeOx nanosheets heated to 120 °C for 16 h and the 200 °C-treated samples was determined from elemental analysis carried out by Galbraith Laboratories, Inc. (Knoxville, TN). Structural, Thermal, Physical, and Electrical Characterization. Powder XRD was performed with a Bruker D8 Focus powder X-ray diffractometer using Cu Kα radiation (λ = 1.54060 Å). XRD patterns were recorded for 2θ between 20° and 70° using a step size of 0.0002° and an integration time of 5 s per step. The crystallite size was calculated using the Scherrer equation, L = Kλ/(β cos θ), where L is the crystallite size, K is the Scherrer constant, λ is the X-ray wavelength, β is the line broadening at half the maximum intensity (fwhm), and θ is the Bragg angle (rad). A Scherrer constant of 0.9 was used for the analysis reported here. The FWHM was determined from the experimental XRD pattern for either the (311) peak for γ-Fe2O3 or the (110) peak for α-Fe2O3 (additional details are provided in the text). Graphite was used as an internal standard to calibrate 2θ values for the as-prepared FeOx nanosheets and the samples treated at 120 °C under vacuum. TEM images were obtained using a JEOL JEM 1200EXII microscope with an accelerating voltage of 120 kV. A JEOL 2010F microscope operated at 200 kV was used for high-resolution TEM. TEM images were analyzed using ImageJ software (version 1.50g). Electron diffraction patterns were calculated from FFT of highresolution TEM images using Digital Micrograph software, version 3.7.0 (Gatan, Inc.). TEM samples were prepared on lacy carbon grids 7804


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(2) Goodenough, J. B.; Kim, Y. Challenges for Rechargeable Li Batteries. Chem. Mater. 2010, 22, 587−603. (3) Xu, J. T.; Dou, S. X.; Liu, H. K.; Dai, L. M. Cathode Materials for Next Generation Lithium Ion Batteries. Nano Energy 2013, 2, 439− 442. (4) Zheng, H.; Li, J.; Song, X.; Liu, G.; Battaglia, V. S. A Comprehensive Understanding of Electrode Thickness Effects on the Electrochemical Performances of Li-ion Battery Cathodes. Electrochim. Acta 2012, 71, 258−265. (5) Fergus, J. W. Recent Developments in Cathode Materials for Lithium Ion Batteries. J. Power Sources 2010, 195, 939−954. (6) Xie, K. Y.; Li, J.; Lai, Y. Q.; Lu, W.; Zhang, Z. A.; Liu, Y. X.; Zhou, L. M.; Huang, H. T. Highly Ordered Iron Oxide Nanotube Arrays as Electrodes for Electrochemical Energy Storage. Electrochem. Commun. 2011, 13, 657−660. (7) Xavier, C. S.; Paskocimas, C. A.; da Motta, F. V.; Araujo, V. D.; Aragon, M. J.; Tirado, J. L.; Lavela, P.; Longo, E.; Delmonte, M. R. B. Microwave-assisted Hydrothermal Synthesis of Magnetite Nanoparticles with Potential Use as Anode in Lithium Ion Batteries. Mater. Res. 2014, 17, 1065−1070. (8) Yun, S.; Lee, Y. C.; Park, H. S. Phase-Controlled Iron Oxide Nanobox Deposited on Hierarchically Structured Graphene Networks for Lithium Ion Storage and Photocatalysis. Sci. Rep. 2016, 6, 19959. (9) Hahn, B. P.; Long, J. W.; Mansour, A. N.; Pettigrew, K. A.; Osofsky, M. S.; Rolison, D. R. Electrochemical Li-ion Storage in Defect Spinel Iron Oxides: the Critical Role of Cation Vacancies. Energy Environ. Sci. 2011, 4, 1495−1502. (10) Komaba, S.; Mikumo, T.; Yabuuchi, N.; Ogata, A.; Yoshida, H.; Yamada, Y. Electrochemical Insertion of Li and Na Ions into Nanocrystalline Fe3O4 and α-Fe2O3 for Rechargeable Batteries. J. Electrochem. Soc. 2010, 157, A60−A65. (11) Komaba, S.; Mikumo, T.; Ogata, A. Electrochemical Activity of Nanocrystalline Fe3O4 in Aprotic Li and Na Salt Electrolytes. Electrochem. Commun. 2008, 10, 1276−1279. (12) Larcher, D.; Masquelier, C.; Bonnin, D.; Chabre, Y.; Masson, V.; Leriche, J. B.; Tarascon, J. M. Effect of Particle Size on Lithium Intercalation into α-Fe2O3. J. Electrochem. Soc. 2003, 150, A133−A139. (13) Hibino, M.; Terashima, J.; Yao, T. Reversible and rapid discharge-charge performance of γ-Fe2O3 prepared by aqueous solution method as the cathode for lithium-ion battery. J. Electrochem. Soc. 2007, 154, A1107−A1111. (14) Koo, B.; Xiong, H.; Slater, M. D.; Prakapenka, V. B.; Balasubramanian, M.; Podsiadlo, P.; Johnson, C. S.; Rajh, T.; Shevchenko, E. V. Hollow Iron Oxide Nanoparticles for Application in Lithium Ion Batteries. Nano Lett. 2012, 12, 2429−2435. (15) Chen, Z. H.; Dahn, J. R. Methods to Obtain Excellent Capacity Retention in LiCoO2 Cycled to 4.5 V. Electrochim. Acta 2004, 49, 1079−1090. (16) Wu, X. L.; Jiang, L. Y.; Cao, F. F.; Guo, Y. G.; Wan, L. J. LiFePO4 Nanoparticles Embedded in a Nanoporous Carbon Matrix: Superior Cathode Material for Electrochemical Energy-Storage Devices. Adv. Mater. 2009, 21, 2710−2714. (17) Gillot, B.; Nivoix, V. New Cation-deficient Vanadium-Iron Spinels with a High Vacancy Content. Mater. Res. Bull. 1999, 34, 1735−1747. (18) Chervin, C. N.; Ko, J. S.; Miller, B. W.; Dudek, L.; Mansour, A. N.; Donakowski, M. D.; Brintlinger, T.; Gogotsi, P.; Chattopadhyay, S.; Shibata, T.; Parker, J. F.; Hahn, B. P.; Rolison, D. R.; Long, J. W. Defective by Design: Vanadium-substituted Iron Oxide Nanoarchitectures as Cation-insertion Hosts for Electrochemical Charge Storage. J. Mater. Chem. A 2015, 3, 12059−12068. (19) Kitaura, H.; Takahashi, K.; Mizuno, F.; Hayashi, A.; Tadanaga, K.; Tatsumisago, M. Preparation of α-Fe2O3 Electrode Materials via Solution Process and their Electrochemical Properties in All-SolidState Lithium Batteries. J. Electrochem. Soc. 2007, 154, A725−A729. (20) Abbasi, A.; Mirhabibi, A.; Arabi, H.; Golmohammad, M.; Brydson, R. Synthesis, Characterization and Electrochemical Performances of γ-Fe2O3 Cathode Material for Li-ion Batteries. J. Mater. Sci.: Mater. Electron. 2016, 27, 7953−7961.

testing, coin cells (2032, Pred Materials) were fabricated in an argon atmosphere (≤1 ppm H2O) glovebox using the electrode disk, a separator (MTI Corp. or glass microfiber filter, GF/C, Whatman), a metallic lithium counter/reference electrode, and the electrolyte (1 M LiPF6 in ethylene carbonate (EC)/diethyl carbonate (DEC), 1:1, v/v, Sigma-Aldrich). Galvanostatic charge/discharge measurements were performed over a voltage range of 1.5−4.2 V vs Li on an Arbin Instruments BT2043 test station using mass-normalized currents of 3.0−20 mA g−1 based on the active material mass. The discharge time at different mass-normalized currents was determined from the experimental data and reported as the C-rate (C = discharge full capacity in 1 h). Cyclic voltammograms were measured using the same configuration over a voltage range of 1.5−4.2 V vs Li at scan rates of 0.1−10 mV s−1. To characterize the morphology of the FeOx material (120 °C vacuum-treated) after electrochemical cycling, electrodes were tested within a cell (HS test cell, Pred Materials) which was galvanostatically cycled over a voltage window of 1.5−4.2 V vs Li at 30 mA g−1 using the configuration and instrumental setup describe above. After 10 charge/discharge cycles, the electrodes (with the discharged state) were removed from the cell within the glovebox and then rinsed in a mixture of EC and DEC (EC/DEC, 1:1, v/v) to remove the salt before imaging. The electrodes were then allowed to dry inside the glovebox and then placed inside the vacuum chamber of the glovebox for 1 h. SEM images of the electrodes were then obtained as described above.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.chemmater.7b02315. Synthesis details and TEM, SEM, XRD, TGA, ATR-FTIR, Raman spectroscopy, galvanostatic cycling, and cyclic voltammetry data of iron oxide nanomaterials (PDF)

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AUTHOR INFORMATION

Corresponding Author

*Phone: 512-245-6721. Fax: 512-245-2374. E-mail: cprhodes@ txstate.edu. ORCID

Joseph B. Tracy: 0000-0002-3358-3703 Veronica Augustyn: 0000-0001-9885-2882 Christopher P. Rhodes: 0000-0003-4886-9875 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS C.P.R. gratefully acknowledges funding from the National Science Foundation (PREM Center, Grant DMR-1205670) for support of this research. J.B.T. acknowledges support from the National Science Foundation (Research Triangle MRSEC, DMR-1121107). This work was performed in part at the Analytical Instrumentation Facility (AIF) at North Carolina State University, which is supported by the State of North Carolina and the National Science Foundation (ECCS1542015). The AIF is a member of the North Carolina Research Triangle Nanotechnology Network (RTNN), a site in the National Nanotechnology Coordinated Infrastructure (NNCI).

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REFERENCES

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