ARTICLE pubs.acs.org/JPCC
Effects of Liquid Electrolytes on the Charge Discharge Performance of Rechargeable Lithium/Sulfur Batteries: Electrochemical and in-Situ X-ray Absorption Spectroscopic Studies Jie Gao,† Michael A. Lowe,† Yasuyuki Kiya,‡ and Hector D. Abru~na*,† † ‡
Department of Chemistry and Chemical Biology, Cornell University, Ithaca, New York 14853-1301, United States Battery & Capacitor Development, Subaru Technical Research Center, Fuji Heavy Industries, Ltd., Mitaka, Tokyo 181-8577, Japan ABSTRACT: A sulfur/carbon composite has been prepared to serve as a cathode for lithium/sulfur batteries. The effects of seven different liquid electrolytes on the electrochemical performance were investigated using galvanostatic discharge charge tests on coin cells. The electrolytes included ether, sulfone, and carbonate solvents with common lithium salts. It was found that the solvent plays a key role on the electrochemical performance of the lithium/sulfur battery cathode while the lithium salt has no significant effects. Additional characterization, using in situ sulfur K-edge X-ray absorption spectroscopy (XAS), provided insights into the soluble sulfur species in the discharged and charged batteries. We find that the use of lowviscosity ethereal solvents results in a more complete reduction of soluble polysulfides, while soluble polysulfides remained more oxidized in viscous ethereal solvents. Moreover, XAS revealed that reduced sulfur species chemically react with carbonate-based solvents, making this class of solvents inappropriate for elemental sulfur cathodes of lithium batteries.
1. INTRODUCTION Rechargeable lithium/sulfur batteries, which use sulfur as the cathode and lithium as the anode, have been the subject of intense research in the recent past due to their high theoretical specific capacity of 1672 mAh/g and energy density of 2600 Wh/kg, assuming the complete reaction of lithium and sulfur to the ultimate product Li2S.1,2 The capacity is about 10 times higher than those of traditional lithium transition metal oxide cathodes like LiCoO2 or LiFePO4 which attain typical capacities of 120 160 mAh/g.3 5 Also, elemental sulfur is nontoxic and naturally abundant and has a much lower cost and easier ability of manufacture than lithium transition metal oxides. Realization of the considerable advantages listed above has been precluded by insufficient control and understanding of the sulfur reduction reactions. However, despite the significant interest in developing sulfur as a cathode, very few studies have directly probed sulfur reduction intermediates and products under battery relevant conditions.6 As a result, there are very limited data on the chemical failure mechanisms of lithium sulfur batteries. This is due, in large part, to the complexity of sulfur and polysulfide (electro)chemistry and the difficulty in identifying specific polysulfide species (e.g., to study rates of equilibria or relative solubilities). An illustrative example is a recent study that electrochemically monitored soluble polysulfides, using different solvents during cycling of a lithium sulfur Swagelok cell.6 While the total concentration of polysulfides could be estimated at different points during the discharge, the identity of specific species and their relative concentrations remained uncertain. The end members of the reaction, elemental sulfur and Li2S, can be isolated and both are highly electronically and ionically insulating, which has been proposed as the explanation of the r 2011 American Chemical Society
rapid capacity fade upon cycling.5 To enhance the utilization of sulfur, researchers have been able to incorporate sulfur into conductive materials such as conducting polymers7 10 and various carbons to form sulfur/polymer and sulfur/carbon composites, respectively.11 18 However, the observed charge storage capacities, even for the first cycle of the best performing composites, are at least 10 20% below the theoretical capacity17 and, more often, the capacity is less than 50%. Additional characterization techniques are required to understand the shortfall. A likely reason is the high solubility of polysulfides in many battery solvents and the inability to recapture all of the dissolved species at a reasonable current density. Sulfur K-edge X-ray absorption spectroscopy (XAS) is a promising probe for elemental sulfur reactions and has been used to probe the local bonding and oxidation state of sulfurcontaining organic compounds,19 transition-metal complexes,20 and inorganic crystals,21 but to our knowledge, none have systematically investigated the lithium elemental sulfur system. Oxidized sulfur species do exhibit significant and systematic differences depending on their chemical state and local environment, and an enhanced signal due to the improved overlap of the 1s 3p orbitals.19 However, high resolution XAS of the near-edge region still yields valuable data for the elemental sulfur system.23 The characteristic energy of the sulfur K edge is at 2472 eV, which is high enough in energy to permit carefully designed in situ XAS studies at a synchrotron light source. For example,
Received: August 11, 2011 Revised: October 26, 2011 Published: October 31, 2011 25132
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The Journal of Physical Chemistry C previous in-situ studies have explained the sulfur-centered changes during lithiation of FeS2.22 As part of our efforts to rationally improve sulfur cathode performance, our research group has recently used this technique to perform a detailed study of sulfur reduction in a single electrolyte system.23 When combined with electrochemical data, in-situ XAS represents a powerful technique to probe soluble sulfur species in battery electrolytes. In previous work, various electrolytes have been used for different sulfur, sulfur/polymer, or sulfur/carbon composite cathode materials. It is well-known that the electrochemical behavior of an electroactive material can be dramatically different in different electrolytes, and this is expected to be especially true for the sulfur system due to the presumably different solubility of various polysulfides. Therefore, it is important to study how electrolytes will affect the electrochemical performance of lithium/sulfur batteries. Generally, battery electrolytes can be categorized into several groups by the specific solvents and salts employed. One group involves organic carbonates including ethylene carbonate (EC), propylene carbonate (PC), and diethyl carbonate (DEC). Another group includes ethers such as 1,3dioxolane (DOL), 1,2-dimethoxyethane (DME), and tetra(ethylene glycol) dimethyl ether (TEGDME). In addition, there are limited reports on the use of ethyl methyl sulfone (EMS) as a solvent.6,17 Common lithium salts are LiClO4, LiPF6, or LiCF3SO3. Several researchers have conducted comparative studies of specific solvent mixtures in lithium sulfur batteries, including TEGDME and DOL,24 DME and DOL,25 and DME, DOL, and di(ethylene glycol) dimethyl ether.26 However, there is a lack of systematic comparison of the effects of different electrolyte systems on the electrochemical performance of the same sulfur cathode. In the present work, we prepared a sulfur/carbon composite and investigated its electrochemical performance in seven different electrolytes which are composed of common battery solvents and lithium salts as the solutes. We found that the solvent plays a key role in the electrochemical performance of a lithium/sulfur cell while the lithium salt has no significant effects. In particular, we determined that carbonate-based solvents are not appropriate for elemental sulfur cathodes of lithium batteries. Using in-situ XAS at the sulfur K-edge, we probed the sulfur reduction intermediates and products in DOL/DME, TEGDME, and EC/DEC solvents and observed a reaction between reduced sulfur species and the carbonate solvent during the course of battery discharge. The XAS data provide otherwise inaccessible insights into the reactions of sulfur in these solvents.
2. EXPERIMENTAL SECTION Conductive carbon blacks (Super P Li) were provided by Timcal Graphite & Carbon. All other chemicals were purchased from Sigma-Aldrich. All chemicals were used without further purification. The sulfur/carbon composite for the electrochemical study was prepared via the following procedure. First, the required amount of elemental sulfur and carbon black (50:45 by mass) was mixed thoroughly by grinding it in a mortar. The mixture was transferred into a small vial and put into a Teflon-lined steel autoclave. Then the container was sealed and heated at 155 °C for 10 h and allowed to cool to room temperature. Nine electrolytes were prepared by dissolving the appropriate lithium salts into the desired solvents in an argon-filled glovebox.
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These included: 1.0 M LiCF3SO3 in each of the following TEGDME, DOL/DME (1:1), PC/EC/DEC (1:4:5), EMS/ DEC (8:1), and EMS/DOL/DME (4:1:1); 1.0 M LiPF6 in TEGDME; and 1.0 M LiClO4 in TEGDME, DOL/DME (1:1), and EC/DEC (1:1). All ratios were based on volume. Electrochemical measurements were conducted in CR2032 coin cells. The cathodes for the sulfur/carbon composite were prepared from a mixture of 95 wt % sulfur/carbon and 5 wt % polytetrafluoroethylene used as a binder. The mixture was rolled into a thin sheet, dried under vacuum at 70 °C overnight, and cut into a circular electrode with a typical area of 0.71 cm2. It was then pressed onto a nickel foam as a current collector. Coin cells were assembled in an argon-filled glovebox with lithium foil as the anode, Celgard 2320 as the separator, and one of the above solutions as the electrolyte. Galvanostatic discharge charge tests were carried out with an Arbin battery testing system. The cells were discharged to 1.5 V at 100 mA/g and charged to 3.0 V at 500 mA/g for each cycle. The capacity was normalized to the mass of sulfur. X-ray absorption spectra were collected at beamline X19A of the National Synchrotron Light Source at Brookhaven National Laboratory. CR2032-type coin cells were modified as described in a recent report23 but with 8 μm thick aluminized Kapton windows (Sheldahl Specialty Materials). Electrodes were prepared by grinding together in a mortar elemental sulfur, carbon black, and poly(vinylene difluoride) (50:45:5 by mass) and then casting on a carbon-coated aluminum mesh current collector. Samples were prepared in an argon-filled glovebox, stored in hermetically sealed vials for 1 3 days before use, and then rapidly installed in a He-purged chamber with electrochemical feedthroughs to minimize any possible air diffusion through the aluminized window. The X-ray beam was incident on the cathode casing and probed the layer of electrolyte solution between the sulfur cathode and the Kapton window. Spectra were collected in fluorescence mode using a large area Si detector (Canberra Industries, Inc.) and then normalized and backgroundsubtracted using the Athena software.27
3. RESULTS AND DISCUSSION 3.1. Electrochemical Characterization of Sulfur/Carbon Composite. We first investigated solvent effects on the charge/
discharge performance of sulfur/carbon cathodes. Figure 1a shows the first discharge profiles of sulfur/carbon composite cathodes in 1.0 M LiCF3SO3 with five different solvents: TEGDME; DOL/ DME (1:1); PC/EC/DEC (1:4:5); EMS/DEC (8:1); EMS/DOL/ DME (4:1:1). We chose these solvents since they are commonly used in most reported work on lithium/sulfur batteries.11 18 For the TEGDME and DOL/DME electrolytes, two distinct voltage plateaus were observed at around 2.4 and 2.1 V, respectively, which is consistent with previous results.14,24 While the mechanism is still the subject of debate, it is generally accepted that the plateau at about 2.4 V corresponds to the reduction of elemental sulfur to soluble, long-chain polysulfides. The lower plateau at 2.1 V reflects the reduction of short-chain polysulfides, which results in the formation of insoluble products, such as Li2S2 and Li2S. The first discharge plateau is 50 mV higher in TEGDME than in DOL/DME, but the second discharge plateau is 75 100 mV lower in TEGDME. The differences may be explained by the significant difference in viscosity between the solvents,28 which will influence the local concentration (and possibly equilibria) of the elemental sulfur and polysulfides. Note that ionic resistivity, which is also influenced by solvent viscosity, should shift both of the 25133
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Figure 2. Cycling performance of sulfur/carbon composites in 1.0 M LiCF3SO3 with different solvents.
Figure 1. First (a) discharge and (b) charge profiles of sulfur/carbon composites in 1.0 M LiCF3SO3 with different solvents: TEGDME, DOL/DME (1:1), PC/EC/DEC (1:4:5), EMS/DEC (8:1), EMS/ DOL/DME (4:1:1).
TEGDME discharge plateaus relative to DOL/DME. The sulfur/carbon cathode gave a similar discharge capacity of about 1000 mAh/g for the TEGDME and DOL/DME electrolytes. For the PC/EC/DEC electrolyte, only one voltage plateau is seen around 2.4 V, and no further discharge reaction was evident afterward. Similar results have been reported for this electrolyte when the electrode used had high loadings of sulfur in a highly porous carbon (HPC)16 matrix or the sulfur was ball-milled with a carbon black and not melt-impregnated as in the current work.15 Under lower loadings of sulfur in HPC16 and at higher sulfur impregnation temperatures,15 researchers have observed higher capacities and a well-defined plateau at around 2.0 V. The differences in response appear to be strongly related to how thoroughly the sulfur is incorporated into the carbon, leading previous researchers to propose that polysulfides are only chemically stable when physically adsorbed within the carbon mesostructure.16 More recently, a highly ordered nanostructured carbon sulfur cathode has shown very attractive performance in an EMS electrolyte.17 Therefore, it was of interest to look at this electrolyte. However, the EMS-based electrolyte did not wet the Celgard 2320 separator, so we added a small amount of DEC or DOL/DME to form an EMS/DEC (8:1) or EMS/DOL/DME(4:1:1) cosolvent system. For the EMS/DEC electrolyte, while the plateau around 2.4 V was similarly observed, only a voltage slope was seen during subsequent discharge. The overall discharge capacity was about 800 mAh/g. For the EMS/DOL/DME electrolyte, two voltage plateaus were observed at around 2.4 and 2.1 V. The plateau at 2.1 V became less flat from about the half-discharge point although the overall discharge capacity was similar to that of TEGDME or DOL/DME. This may result from the high viscosity of the EMS solvent since only a small amount of DOL/DME was added. It is worth noting that the more viscous solvents (TEGDME, EMS/DEC, and EMS/DOL/DME) have a higher first voltage plateau and a lower second voltage plateau, supporting the conclusion that the voltages of the discharge plateaus reflect concentration-dependent chemical equilibria.
Figure 1b shows the first charge profiles of the same sulfur/ carbon cathodes in the above five electrolytes. For the TEGDME and DOL/DME electrolytes, the charge capacities were both over 1000 mAh/g. The two voltage plateaus for each electrolyte reflect the reverse reactions observed during the discharge process of the sulfur/carbon cathode. The charge voltage for the DOL/DME electrolyte is 100 150 mV lower than that for the TEGDME electrolyte, indicating that the reversibility of the sulfur cathode is better in DOL/DME electrolyte than that in the TEGDME electrolyte. Note that during both discharging and charging processes the voltage difference between plateaus is larger for TEGDME than DOL/DME, supporting the argument for solvent-influenced changes in chemical equilibria. For the PC/EC/DEC electrolyte, almost no charge capacity was obtained. For the EMS/DEC electrolyte, only a very small capacity (about 100 mAh/g) with a high voltage slope was seen. For the EMS/DOL/DME electrolyte, a reasonable charge capacity of about 900 mAh/g was observed although this value is slightly less than that for TEGDME or DOL/DME electrolyte. The dramatic difference in the EMS/DEC and EMS/DOL/DME electrolytes clearly indicates that DEC is likely responsible for the failure of the EMS/DEC cell. Although the EMS solvent shows some potential capabilities, its high viscosity and melting point will likely preclude its application for lithium/sulfur battery electrolytes. Figure 2 presents the cycling performances of the sulfur/ carbon composites in 1.0 M LiCF3SO3 with different solvents. While the capacity remained relatively stable after the second cycle in DOL/DME electrolyte, it faded gradually in the TEGDME or EMS/DOL/DME electrolyte. However, almost no capacity was obtained after the first discharge in carbonatecontaining electrolytes, consistent with a previous report for high sulfur loading in HPC.16 This indicates that some irreversible reactions likely occur during the first discharge charge cycle in PC/EC/DEC and EMS/DEC electrolytes and that the carbonate-based solvent is responsible for battery failure. However, the electrochemical response cannot distinguish between chemical reactions with the solvent and formation of an insulating film (e.g., precipitation of polysulfides), and additional characterization was performed as described in the next section. We then investigated the influence of different solutes (i.e., lithium salts) on the charge/discharge performance of sulfur/ carbon cathodes. In this case, we chose TEGDME as a solvent because LiPF6 does not dissociate well in DOL or DME but does 25134
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Figure 3. First (a) discharge and (b) charge profiles of sulfur/carbon composites in TEGDME with different 1.0 M lithium salts.
Figure 5. In-situ sulfur K-edge XAS for 1.0 M LiClO4 in DOL/DME, TEGDME, and EC/DEC and S8 and Li2S standards.
Figure 4. Cycling performance of sulfur/carbon composites in TEGDME with different 1.0 M lithium salts.
dissociate well in TEGDME. Figure 3a shows the first discharge profiles of sulfur/carbon composites in TEGDME with 1.0 M of three different lithium salts: LiPF6, LiCF3SO3, and LiClO4. Two distinct voltage plateaus are observed for all three lithium salts, and the overall discharge capacities are virtually the same considering the experimental uncertainties. Figure 3b shows the first charge profiles of the same sulfur/carbon cathodes in the three electrolytes, which also show no significant differences. Figure 4 presents the cycling performances of the sulfur/carbon composites in TEGDME with 1.0 M of the three lithium salts. While the capacity fades gradually after the second cycle in LiPF6 and LiCF3SO3 electrolytes, it remains somewhat more stable in LiClO4 electrolyte. However, from the point of practical applications, LiClO4 is not a suitable lithium salt for lithium/sulfur batteries due to safety concerns.28 3.2. XAS Characterization of Sulfur Reduction Intermediates and Products. To better understand the influence of solvent on the sulfur reaction mechanism, sulfur K-edge XAS data were collected during galvanostatic cycling (at a 80 mA/g rate) for a subset of the solvents. Note that the spectra correspond to predominantly polysulfide species in solution, since the
cell design traps a comparatively thick (∼200 μm) layer of solution between the electrode and the X-ray window.23 At the “tender” energies near the sulfur K-edge, the solution and window will attenuate the incident and emitted X-rays and signal from species in solution will be up to 200 times greater than from sulfur on the electrode. This enhanced signal from species in solution is well suited for probing sulfur electrolyte interactions. Figure 5 shows the initial and discharged XAS for sulfur electrodes in DOL/DME, TEGDME, and EC/DEC, along with the spectra for the solid-state standards of elemental sulfur and lithium sulfide (Li2S). Elemental sulfur XAS exhibits a single sharp peak at 2472.7 eV due to the electronic transition from 1s to unoccupied 3p states, while lithium sulfide shows two peaks at 2473.7 and 2476.3 eV. As discussed previously,23 the main peak for solid-state lithium sulfide shifts to higher incident energies due to a strong Coulombic interaction between the S2‑ and eight neighboring Li+ ions.21 This is in contrast to the correlation between shifts to higher energy and an increased sulfur oxidation state which is commonly reported in the XAS literature.19 Initially, DOL/DME and TEGDME exhibited similar features in the near-edge spectra, with a main peak for elemental sulfur near 2472.7 eV and a pre-edge feature at 2471 eV due to reduced sulfur radical anions.23 The anions are a product of self-discharge, an issue for sulfur cathodes because of the non-negligible solubility of sulfur in many battery electrolytes. The more intense prepeak for DOL/DME is in agreement with a greater degree of self-discharge relative to TEGDME. In the discharged batteries, the main peak decreased in intensity and shifted to higher energies. In TEGDME, it retained a distinct prepeak at 2471 eV due to radical anions and a comparatively slight (+0.3 eV) shift in the peak position, but discharged sulfur in DOL/DME exhibited no prepeak and a much more distinct peak shift to near 2473.5 eV, close to the position of the main peak for the Li2S standard. In both solvents, 25135
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Scheme 1. Proposed Reactions between Polysulfides and Carbonates
the discharge products still had small shoulders at lower energies ( 0.4 eV from elemental sulfur), possibly due to intermediatelength polysulfides. Additional theoretical interpretation is necessary to explain the observed peak shifts, but it is likely that the magnitude of the shifts is related to how strongly the sulfur species is coordinated by lithium ions. For shorter chain polysulfides, the shift would be similar to solid-state lithium sulfide, while longer chain polysulfides would exhibit less of a shift. If this model is correct, then the data would indicate that in both solvents a significant quantity of sulfur in solution remains in higher oxidation states at the end of discharge, having become electronically isolated by diffusion away from the electrode. In DOL/DME, the soluble sulfur species are more highly reduced, resulting in a peak shifted almost all the way to 2473 eV. The lower viscosity of DOL/DME likely facilitates reduction of the polysulfides both by direct electrochemical reduction and via disproportionation reactions. Upon charging, the spectra for sulfur in TEGDME and DOL/ DME essentially regain their original features. The intensity of the peak due to neutral sulfur increases for DOL/DME, since the charged battery is at a higher state of charge than the battery during the initial spectrum. Even in the charged batteries, both solvents retain a pre-edge feature due to the presence of radical anions in solution. Sulfur in EC/DEC-based electrolyte exhibits XAS signals very different from sulfur in DOL/DME and TEGDME, indicating that the sulfur reduction mechanism is significantly different in carbonate-based solvents. As seen in Figure 5, the spectrum for the initial battery has no prepeak at 2471 eV, indicating that the battery has not formed radical anions during self-discharge. Upon discharge, the biggest spectral change is the growth of a peak at 2474.2 eV. A careful comparison of the initial spectrum for each solvent shows that this peak is present in the EC/DEC sample before discharge begins, suggesting that it may also be formed during self-discharge. The chemical species associated with the new peak has not yet been identified, but its shift from the main peak of elemental sulfur (+1.3 eV) is roughly consistent with previous reports of thioether (+0.6 eV) and sulfonium (+1.0 eV) functionalities,19 especially since these shifts may vary with the specific chemical environment.29 Two possible routes for reaction are shown in Scheme 1. In (a), nucleophilic sulfide anions may chemically react with the alkyl carbonate solvent, similarly to the methylation of thiolates by dimethylcarbonate.30 Reaction (b) is based on thioether formation from ethylene carbonate, which is catalyzed by alkali metals, hydroxides, and carbonates.31 While the spectroscopic data set cannot establish if the reaction proceeds with both cyclic and linear carbonates, the poor reversibility of the EMS/DEC cosolvent mixture demonstrates that, at a minimum, the linear carbonates are unstable. We expect this nucleophilic attack to be an issue in any of the carbonate
Figure 6. Correlation between cell voltage and peak height of the reaction product from reduced sulfur and EC/DEC solvent. Peak height was determined by subtracting a linear background between 2473.5 and 2474.9 eV.
solvents. Of particular interest is the recent report by Bruce et al. in which reduced oxygen products were shown to rapidly attack carbonate solvents. They proposed that the reaction proceeds through a nucleophilic attack, similar to Scheme 1b.32 Generally, previous reports in which carbonate solvents were successfully used for sulfur-based batteries involved sulfur that was trapped within a highly mesoporous architecture.15,16 Under these conditions, the different relative concentrations of solvent, polysulfide, and decomposition products may slow the reaction rate. The proposed reactions are further supported by the temporal variation in the peak, as displayed in Figure 6. The peak appears to increase throughout the discharge, but the peak grows most rapidly at the end of the discharge (>250 mAh/g). These capacities correspond to an average oxidation state near S0.5 or the polysulfide S42 . The nucleophilicity of polysulfides has been reported to increase with decreasing chain length as the charge is spread over fewer sulfur atoms, with S22 and S42 being the most nucleophilic polysulfides.33 Thus, the decomposition product appears when the most nucleophilic polysulfides are generated.
4. CONCLUSIONS The effects of liquid electrolytes on the electrochemical performance of lithium/sulfur batteries have been investigated by using galvanostatic discharge charge tests. It was found that the solvent in the electrolyte plays a key role in the electrochemical performance of a lithium/sulfur cell while the lithium salt has no significant effects. The sulfur/carbon cathode exhibits 25136
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The Journal of Physical Chemistry C a similar discharge capacity of about 1000 mAh/g for the TEGDME and DOL/DME electrolytes. Almost no capacity is obtained after the first cycle in PC/EC/DEC and EMS/DEC electrolyte. On the basis of the electrochemical study, DOL/ DME and TEGDME could be appropriate and promising solvents for the electrolytes of lithium/sulfur batteries. EMS may only be interesting if combined with less viscous solvents of low melting point. XAS studies revealed that the discharge products of sulfur in DOL/DME and TEGDME are similar, while the reduced sulfur species chemically reacted with the carbonate-based solvents. In DOL/DME, sulfur exhibited faster self-discharge and a more complete reduction of solution-phase sulfur species upon discharge. Reduction of sulfur in the ether-based solvents involved distinct radical anion intermediates that persisted in solution even in the discharged and charged states. In combination with the electrochemical data, XAS yields valuable chemical information about the effects of solvent on sulfur electrochemistry. Additional theoretical work is ongoing to improve the assignments of spectral features to specific polysulfides.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected].
’ ACKNOWLEDGMENT This work was supported by Fuji Heavy Industries, Ltd. M.A.L. is supported by a National Defense Science and Engineering graduate fellowship. Use of the National Synchrotron Light Source, Brookhaven National Laboratory, was supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, under Contract DE-AC02-98CH10886.
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