Environmental In Situ X-ray Absorption Spectroscopy Evaluation of

Article pubs.acs.org/JPCC

Environmental In Situ X‑ray Absorption Spectroscopy Evaluation of Electrode Materials for Rechargeable Lithium−Oxygen Batteries Gregory S. Hutchings,† Jonathan Rosen,† Danielle Smiley,‡ Gillian R. Goward,‡ Peter G. Bruce,§ and Feng Jiao*,† †

Department of Chemical and Biomolecular Engineering, University of Delaware, Newark, Delaware 19716, United States Department of Chemistry and Chemical Biology, McMaster University, Hamilton, Ontario L8S 4L8, Canada § School of Chemistry, University of St. Andrews, St. Andrews, Fife KY16 9ST, United Kingdom ‡

S Supporting Information *

ABSTRACT: Lithium−oxygen batteries have attracted much recent attention due their high theoretical capacities, which exceeds that of Li-ion batteries. Among all the metal oxides that have been investigated in oxygen cathodes, α-MnO2 materials have shown unique electrochemical properties in rechargeable lithium oxygen batteries. Although extensive research has been performed to investigate the structure of αMnO2 upon lithium intercalation, its behavior upon reacting with lithium under an oxygen environment remains to be fully explored. Here, we performed a systematic study on the behavior of two forms of α-MnO2 nanowires (i.e., potassium and ammonia versions) together with bulk α-MnO2 in oxygen cathodes through environmental in situ X-ray absorption spectroscopy. The results show that the α-MnO2 materials undergo lithium insertion/removal and lithium peroxide formation/decomposition simultaneously. The former causes a self-switching of the oxidation state of Mn during cycling. Additionally, we found that potassium-containing α-MnO2 nanowires exhibit a suppression of Mn reduction until late in cell discharge under oxygen, retaining a higher degree of Mn4+ character for enhanced oxygen reduction activity than other, similar α-MnO2 materials. During cell recharge along with oxygen evolution, the materials were found to return to their initial states at low overpotential. can accommodate not only intercalated Li+ but also Li2O and other LixOy species within its structure.16,17 As shown in electrochemical studies of α-MnO2, the Li+ intercalation occurs within a potential range of 2.5−3 V vs the Li/Li+ couple.18 This potential is very close to the net Li−O2 electrochemical reaction, which theoretically occurs at 2.96 V vs Li/Li+: 2Li+ + O2 + 2e− ↔ Li2O2.19 In practice, overpotentials lower the potential of Li2O2 formation to between 2.5 and 2.8 V vs Li/ Li+, while the cell potential often exceeds 4 V vs Li/Li+ during charging. Since the potential range for Li+ insertion and removal lies within the discharge/charge potential window of the cell, it is likely that an insertion mechanism occurs simultaneously with Li2O2 formation. However, no direct evidence has been observed while cycling α-MnO2 in an oxygen environment, mainly due to the fact that it is difficult to monitor the electrochemical process at solid/liquid/gas triplephase boundaries. Many efforts have been devoted to the identification of the discharge products and elucidation of the reaction mechanism inside lithium−oxygen cells. Ex-situ structural characterization studies, such as Fourier transform infrared spectroscopy (FTIR), Raman spectroscopy, powder X-

1. INTRODUCTION Lithium-ion batteries (LIBs) dominate the current energy storage market for mobile applications. Traditional electrode materials for LIBs are usually evaluated under inert environments to avoid undesired side reactions, but advances in rechargeable lithium−oxygen (Li−O2) batteries have focused on electrode behaviors in O2 environments. These Li−O2 batteries operating with nonaqueous electrolytes offer the potential for a step-change in specific energy over traditional lithium-ion chemistries,1−3 and have attracted much recent attention. In a typical Li−O2 battery, O2 is reacted directly with Li+ at the cathode side to form Li2O2 in a net two-electron reduction reaction, 4−6 rather than storing energy via intercalation of Li+ in electrode materials, which are restricted in capacity due to limited interstitial volume. The introduction of metal oxides (e.g., Co3O4 and MnO2) into the oxygen cathode has been found to be an effective way to enhance discharge capacities;7−13 however, they often promote side reactions simultaneously.14,15 Therefore, it is critical to fully understand the behavior of metal oxides in the oxygen cathode so that we can design better oxygen cathodes with enhanced stabilities and performance. Unlike most other cathode materials characterized in the O2 environment of Li−O2 cells, such as Co3O4, Pt/C, and MnOx materials, α-MnO2 has a large 2 × 2 tunnel arrangement and © 2014 American Chemical Society

Received: February 18, 2014 Revised: May 27, 2014 Published: May 28, 2014 12617

dx.doi.org/10.1021/jp5017399 | J. Phys. Chem. C 2014, 118, 12617−12624

The Journal of Physical Chemistry C

Article

2.3. Li−O2 XAS Electrochemical Cell Assembly. The cathodes consisted of 24 wt % Super P carbon, 42 wt % manganese oxide, and 34 wt % polyvinylidene difluoride (PVdF) binder, which is more typical for Li−O2 cathode studies than purely intercalation-based compositions. The powders were mixed into a slurry with N-methyl-2-pyrrolidone (NMP) as a solvent. Dibutyl phthalate (DBP) was added to the mixture as a plasticizer. The slurry was cast on glass, allowed to dry in air, then peeled and punched into freestanding disks. The DBP plasticizer was then extracted in methanol. All materials were dried under vacuum and stored under Ar prior to environmental electrochemical cell construction. The environmental electrochemical cells were constructed from commercial CR2032 stainless steel coin cell casings, with a modified cathode backing incorporating a meshed stainless steel screen (MTI Corporation). For the in situ XAS tests, vacuum-dried 1 M lithium bis(trifluoromethanesulfonyl)imide (LiTFSI) in tetraethylene glycol dimethyl ether (TEGDME) was used as the electrolyte. The purified TEGDME solvent was further dried over 4 Å molecular sieves before use. A glass fiber separator (Whatman GF/D) was saturated with the electrolyte and placed between the cathode and the lithium metal negative electrode. Assembly of the electrochemical cells was conducted in an Ar-filled glovebox (MBraun). The completed cells were heat sealed inside foil packaging (Shield Pack, class PPD), which had been modified to include a standard coin cell holder and external electrode leads. The bags were either kept sealed under Ar or purged in a pure O2 stream before being resealed under ambient pressure. The internal O2 or Ar environment was kept static during cycling. A schematic of the completed cell is shown in Figure 1.

ray diffraction, solid state NMR, X-ray photoelectron spectroscopy (XPS), and electron microscopy, have been performed.20−26 However, environmental in situ diagnostic tools that allow us to probe the dynamics of structural evolution with the presence of oxygen gas during the electrochemical reaction should offer a better insight into the changes that occur. In this paper, we report the first environmental in situ X-ray absorption spectroscopy (XAS) investigation of real-time structural changes to α-MnO2 cycling under an oxygen environment. Synchrotron XAS allows us to monitor oxidation state and coordination environment information on Mn without any undue restrictions on pressure or atmospheric conditions. Using this structural diagnostic tool, we follow the lithium insertion into α-MnO2 that occurs during discharge and extraction on charge with or without the presence of oxygen gas molecules. The detailed electrochemical behaviors of a variety of α-MnO2 materials under various environments have been investigated systematically.

2. EXPERIMENTAL DETAILS 2.1. Preparation of Manganese Oxides. Crystalline αMnO2 bulk material was synthesized by reflux of Mn2O3 (Sigma-Aldrich) in H2SO4, as described in the literature.16 The final product was then dehydrated at 300 °C, and stored under Ar. For the synthesis of xLi2O·MnO2, the α-MnO2 material was mixed in methanol with a proportional amount of LiOH·H2O, then calcined in air at 275 °C for 3 h. The actual molar ratio of xLi2O·MnO2 was determined by ICP-OES (Li2O/MnO2 = 0.116 for 0.15Li2O·MnO2, and Li2O/MnO2 = 0.199 for 0.25Li2O·MnO2). Nanowires of crystalline α-MnO2 were prepared using two separate literature methods, resulting in K+ or NH4+ remaining in the structure. For the K+ version,27 1 mmol KMnO4 was combined with 1 mmol NH4Cl in 50 mL of distilled water under vigorous stirring. The resulting mixture was then transferred to Teflon-lined autoclaves and kept at 140 °C for 24 h. For the NH4+ version (only metal ions are Mn),28 (NH4)2SO4, MnSO4·H2O, and (NH4)2S2O8 were combined in a 1:0.4:0.4 molar ratio in distilled water (typically, 1.35 g MnSO4·H2O in 30 mL of water), then transferred to Teflonlined autoclaves. This mixture was then kept at 140 °C for 12 h. For both materials, the final product was obtained after washing with distilled water and ethanol, then drying at 60 °C overnight. 2.2. Structural Characterization. Phases were confirmed by X-ray diffraction, using a PANalytical X’Pert MRD with Cu Kα radiation (45 kV, 40 mA). Morphology was assessed by scanning electron microscopy (SEM), using a JEOL JSM7400F, as well as with transmission electron microscopy (TEM), using a JEOL JEM-2010F (images of the morphology are shown in Supporting Figure S1). The surface areas was measured by N2 adsorption, using a Micromeritics ASAP 2020 operated at 77.36 K, and are 24 m2 g−1 for α-MnO2 bulk, 57 m2 g−1 for α-MnO2 nanowires synthesized with K+, and 45 m2 g−1 for α-MnO2 nanowires synthesized with NH4+. 7 Li magic angle spinning nuclear magnetic resonance (MAS NMR) spectra were acquired at a Larmor frequency of 116.6 MHz on a Bruker AV-300 spectrometer with a custom-built double-resonance probe supporting rotors with 1.8 mm diameter. The sample size was ∼2−5 mg, extracted from the current collector mesh. A Hahn-echo pulse sequence was used with a 90° pulse of 2.5 μs and 10 s recycle delay. 30−35 kHz MAS was used. Signals were accumulated over 1k scans.

Figure 1. Schematic of the environmental electrochemical cell used for in situ XAS measurements.

Galvanostatic testing of these cells was performed at a current rate of 85 mA g−1oxide. All cycling during the in situ XAS analysis was conducted on a MTI battery analyzer. A MACCOR 4000 series system was used for all additional electrochemical characterization. For cycling α-MnO2 under O2, one critical issue is the stability of electrolyte, carbon, and binder during the electrochemical cycling. Several groups have reported that Li−O2 cells operated with propylene carbonate electrolyte produce Li2CO3 and other carbonate species as the main discharge products instead of Li2O2 or Li2O.29,30 Thorough characterization of evolved gases and cycled electrodes have shown at least some degree of degradation in most organic electrolytes (including ether-based electrolytes, as used in this study).13,14,31−34 To date, fully reversible operation with Li2O2 as the sole discharge product has only been confirmed for a limited number of materials, including freestanding Au cathodes in a DMSO-based electrolyte.20 In this paper, we used a TEGDME-based 12618



Article

Figure 2. First cycle load curves for α-MnO2 K+ nanowires (black line) and bulk (red dashed-dotted line) cells, cycled at a rate of 85 mA g−1oxide under (a) O2 and (b) Ar. Points on the plot designated with “●” marks correspond to the state of discharge or charge during the in situ XAS data collection.

Figure 2a. The large 2 × 2 channels of α-MnO2 are capable of reversibly intercalating Li+,16 which leads to the reduction of Mn4+ to Mn3+. This intercalation results in a natural selfswitching mechanism during electrochemical cycling, with an influence on the formation of Li2O2 when the cell is cycled under O2. Additionally, it is anticipated that Li2O insertion can occur in the presence of O2, acting as a stabilizer of the α-MnO2 structure.17 The insertion of either Li+ or Li oxide species can occur during discharge as an insertion happening simultaneously with ORR, as the plateau voltages of both the pure Li+ intercalation mechanism and the Li2O insertion mechanism occur in the same range as ORR.36,37 Determining the precise timing of structural changes due to insertion can be critical for designing an optimal electrode for cycling under O2. In situ XANES (X-ray absorption near edge structure) spectra were collected for each catalyst material at key points during the first electrochemical cycle, as indicated in Figure 2a,b. For reference, examples of the raw Mn K-edge XANES spectra are shown in Supporting Figure S3 for the in situ αMnO2 bulk environmental electrochemical cell under O2. A primary use for XANES is to identify shifts in oxidation state, which is usually observed at half the height of the largest peak that marks the edge. By exploiting the known linear relationship between Mn oxidation state and edge energy,38 the oxidation states at each position in the electrochemical cycle have been calculated based on commercial bulk ß-MnO2 and Mn2O3 as Mn4+ and Mn3+ standards, respectively, and are shown in Figure 3a. For comparison with a purely Li+ intercalation-based system, in situ XAS data of α-MnO2 bulk and K+-α-MnO2 nanowires cycled in an Ar (O2-free) environment were recorded at cycle end-points, and the calculated oxidation states are also shown in Figure 3b. The displayed error bars reflect beamline resolution, which restricts the certainty of the obtained oxidation states to within ±0.1. Point (1) corresponds to the as-made state of the α-MnO2 materials, before any electrochemistry has taken place. Deviation from the expected oxidation state for each material can be attributed to reaction with the electrolyte at open circuit. Some small differences at point (1) between the two cycling conditions for the K+ nanowires is likely due to an additional surface reaction occurring in the presence of O2, which is mostly removed by the discharge to point (2). The surface reaction would only noticeably affect the oxidation state for the high surface area material, which is why the same behavior is

electrolyte because of its enhanced stability in O2 compared to carbonate-based electrolytes, even though Li+ intercalation performance is reduced. Additionally, only the first cycle is studied here in order to minimize side reaction issues. 2.4. Mn K-Edge XAS Measurements. XAS measurement of the in situ environmental electrochemical cells and powder standards was performed on beamline X10C of the National Synchrotron Light Source at Brookhaven National Laboratory (BNL). Due to geometric restrictions of the electrochemical cells, fluorescence data collected with a 7 element Si drift detector were used this analysis. The free Demeter software was used for XAS alignment and data processing.35 Energy shift parameters for data alignment were determined by the first peak of the pre-edge region. As the Fe K-edge is relatively close to the end of useful data for EXAFS (extended X-ray absorption fine structure) transformation, the data for processing is restricted to a measurement end of 11.3 Å−1 in k-spacing (last data point at 7046 eV). A k-range of 1 to 10.5 Å−1 is included in the Fourier transform, with a background spline determined from a start point of 6552 eV (E0). For simplicity, the EXAFS spectra are presented without phase correction, meaning that the apparent distances described here are approximately 0.3−0.4 Å less than the real distances. In order to qualitatively assess the structural changes to αMnO2 during electrochemical cycle in O2, single-scattering paths were calculated from the 0.15Li2O·MnO2 structure, which is proposed to be the lithia-stabilized α-MnO 2 structure.22 Standard α-MnO2 has the tetragonal space group I4/m, with a = 9.815 and c = 2.847 (JCPDS 44−141). The lithia-stabilized 0.15Li2O·MnO2 version has an expanded lattice with a = 9.942 and c = 2.851, and new generators for Li at (0.204, 0.078, 0) and O at (0, 0, 0.40), as proposed by previous neutron diffraction study.16 As Mn−Li scattering cannot be directly detected by XAS, distortions to the α-MnO2 structure resulting from the addition of Li2O or Li+ are considered instead.

3. RESULTS AND DISCUSSION In early lithium−oxygen screening tests, α-MnO2 nanowires synthesized with K+ showed greatly increased performance over other nanostructured and bulk MnOx catalysts.7 This improved performance is also seen over K+-free nanowires (referred to as the NH4+ version, Supporting Figure S2a) in the TEGDME system under O2, as shown for the first electrochemical cycle in 12619



Article

rays during XAS measurement, the XANES result will be weighted toward the oxidation state near the surface if the surface and bulk properties are not the same). In order to investigate this insertion behavior further, 7Li MAS NMR experiments were conducted ex situ for K+ nanowire cathodes and chemically lithiated 0.15Li2O·MnO2 bulk powder (Supporting Figure S4). The chemical shift of the lithium species formed on electrochemical reaction of lithium and oxygen with the cathode material is consistent with that of Li2O or Li2O2, as expected for cycling in an O2 environment,21 and the narrow peak at 0 ppm is in good agreement with the expected chemical shift for Li2O and Li2O2 in all three spectra. Nevertheless, the 7Li NMR spectra indicate Li2O or Li2O2 species in close contact with the α-MnO2 materials, as evidenced by the broader spinning sideband manifold (Supporting Figure S5a). Since Li2O2 and Li2O are virtually indistinguishable in 7Li NMR, the spectrum for the 0.15Li2O· MnO2 bulk powder is very similar to the discharged K+ nanowire cathode, even though only Li2O is expected in the chemically lithiated powder while the discharged cathode should contain a large portion of Li2O2. The increasing breadth of the sideband manifold is attributed to a difference in intimate mixing of the particles, where the Li2O or Li2O2 species in the fully discharged cathode appears to be in closest contact with the α-MnO2 materials. Typically, Li+ ions intercalated within a paramagnetic host exhibit a large chemical shift (several hundred ppm), as well as a short spin−lattice (on the order of milliseconds) relaxation compared to a diamagnetic lithium site (on the order of seconds).39 There is no evidence of a paramagnetically shifted lithium species in any of the spectra, as only a single site with a diamagnetic chemical shift (∼0 ppm) and long spin−lattice relaxation time (>5 s for all spectra) is observed. Chemical shift regions between +2500 to −2000 were carefully examined in an attempt to reveal a paramagnetically shifted Li peak. The lack of such a peak indicates that the electrochemical reaction involves the formation of Li2O or Li2O2 as a separate diamagnetic phase, rather than an intercalation of discrete lithium ions within α-MnO2 channels. The apparent lack of any Li+ in the discharged K+ nanowire material in the 7Li NMR results, compared to the expectation of some Li+ insertion from in situ XANES experiments, may be due to large Li2O2 signal from discharge under O2 compared to the relatively small amount of Li+ in the α-MnO2 structure. Additional insights into the insertion behavior are provided by considering the full EXAFS sequences, which are presented in Figure 4a,b (q-space data for the K+ nanowires and bulk catalyst are presented in Supporting Figure S5). It should first be restated that, unless otherwise noted, all the distances presented in this paper are apparent distances without phase corrections. To convert the apparent distances to phasecorrected real distances, approximately 0.3−0.4 Å should be added to the apparent values. For all three materials, the large peak at approximately 1.4 Å in k3-weighted R-space can be primarily attributed to Mn−O interaction as part of the octahedral coordination in the α-MnO2 structure (1.892 Å phase-corrected scattering length). The second large peak near 2.5 Å is a combination of Mn−Mn contributions (N = 2 for each), corresponding to distance between Mn atoms in edgeshared octahedra (2.847 and 2.889 Å phase-corrected scattering lengths). The third peak, at 2.9 Å, shows a strong connection to the insertion processes, and becomes much more distinct in the discharged in situ cells. During the course of lithiation, this third peak grows relative to neighboring peaks as seen in cell

Figure 3. Calculated Mn oxidation state for (a) K+ nanowire and (b) bulk environmental electrochemical cells, as estimated from XANES data. The black lines correspond to the O2 cells, while red lines correspond to the Ar (O2-free) systems. Error bars refer to machine resolution limitations, which effectively add ±0.1 to oxidation state measurements.

not observed for the bulk material. During discharge (between points (1) and (5)), the bulk α-MnO2 clearly shows a downward trend in oxidation state to point (5), while the K+ nanowires do not show any drop in oxidation state at all until point (5). As there is relatively little capacity between points (4) and (5), the lack of an oxidation state change for K+ until after point (4) means that the material remains at an oxidation state of around +3.6 for the majority of the discharge under O2. The final oxidation state change on discharge for the bulk αMnO2 material, from +3.37 to +3.09, is quite large, and represents a significant switch in the catalyst state. Moreover, this final state is consistent with the predicted oxidation state based on charge passed (+3.15), based on the first discharge of the O2-free materials (Figure 2b and Supporting Figure S4b). Note that the low intercalation capacities in the present study are likely due to poor compatibility with the TEGDME electrolyte. The NH4+ nanowires also exhibit a Mn oxidation state drop from +3.53 to +3.21 that occurs over the entire discharge range (not shown), which explains the overall similarity in the electrochemistry of these two materials. The final change in the K+ nanowires (from +3.60 to +3.45) is within the error bars of the measurement, and is far from the final oxidation state of +3.27 expected for Li+ intercalation based on the charge passed. In the case of the O2 system, a possible cause is greater degree of insertion of Li2O during discharge rather than Li+, which does not result in a change to the Mn oxidation state. For the Ar system, the final oxidation state indicated by XANES is higher than expected for Li+ insertion, which can be explained by a surface reaction with the TEGDME electrolyte (due to reabsorption of fluorescence X12620



Article

discharged to around 70 mAh g−1oxide at this stage. Indeed, there are no significant changes to the third EXAFS peak around 2.9 Å, or any significant change in oxidation state, which would both be associated with lithiation. Thus, the major contribution to the peak changes is new Mn−O coordination with typical Li−O2 discharge products at the surface of the materials. Such a mechanism is be more prevalent for the nanowires, which have much higher surface areas than the bulk α-MnO2. EXAFS spectra taken at points (3), (4), and (5) show the rest of the cell discharge in an O2 environment, during which Li+ intercalation, Li2O insertion, and formation of typical Li− O2 discharge products can all occur. There is clearly no major change to peaks at 2.9 and 3.2 Å for the K+-α-MnO2 nanowires, even by point (4), while changes in this region are visible for both the NH4+-α-MnO2 nanowire and bulk materials. In the case of point (3), this phenomenon is partially explained by the higher discharge voltage at that state of discharge for K+-αMnO2 nanowires than for the other materials, which means that less Li+ would have intercalated by that point. This is confirmed by the oxidation state analysis, which also shows no change to the K+ material at point (3). By point (4), however, the cell potential is 2.4 V for all materials, and a significant amount of Li+ would have intercalated into the K+-α-MnO2 nanowires if no other mechanism was dominating. However, only a relatively small increase in the peak at 2.9 Å is observed. This suggests an alternate mechanism may be dominant in this material, such as Li2O insertion rather than Li+ intercalation. The 0.15Li2O·MnO2 and 0.25Li2O·MnO2 ex situ materials exhibit similar peak shape in this region, and the relative intensity of this peak also grows only slightly between these materials despite an increase of Li2O content by 0.1 formula units. Such an insertion of Li2O is consistent with the 7Li MAS NMR result. A possible explanation of this behavior is that Lix(Li2O)yMnO2 is formed in situ, which contains an unbalanced ratio of Li to O, but significantly more O2 in the α-MnO2 channels than LixMnO2. Such materials exhibit significant distortions from the normal tetragonal symmetry of α-MnO2, with a/b ratios as high as 1.060.37 The net result for these α-MnO2 materials when cycled under O2 is that the K+-α-MnO2 nanowires should still retain an oxidation state closer to Mn4+ late in the discharge cycle, compared to the other α-MnO2 materials. By point (5), even the K+-α-MnO2 nanowires show a distinct peak around 2.9 Å, but still to a lower extent than either NH4+-α-MnO2 nanowires or the bulk. For comparison with the O2 system, in situ EXAFS spectra of α-MnO2 bulk and K+-α-MnO2 nanowires cycled in an Ar (O2free) environment are shown in Figure 5a,b. The spectra for the discharged materials closely match the shape of their counterparts cycled in O2, which suggests similar α-MnO2 states upon discharge and agrees with the oxidation state trends calculated in Figure 3a,b. In all three materials, charging to point (6) results in a near-complete return to the original asmade state, which is exhibited by immediate returns to the initial Mn oxidation state and qualitative EXAFS behavior. This means that delithiation occurs early in the charge cycle (∼3.2 V vs Li/Li+) for α-MnO2, between points (5) and (6). Further charging to 4.5 V vs Li/Li+ (7) results in near-complete reversibility for the NH4+-α-MnO2 nanowires and bulk materials for peaks below 3.2 Å. In the Ar system, similar reversibility is exhibited in both XANES and EXAFS spectra. For the NH4+-α-MnO2 nanowires and bulk materials, there is also a reduction in the Mn−O to Mn−Mn peak ratio back

Figure 4. Sequence of k3-weighted phase-uncorrected EXAFS spectra for cells with (a) K+ nanowires and (b) bulk α-MnO2, cycled in O2. Labels indicate the state of discharge/charge for each spectrum, as indicated in Figure 2a.

discharge in O2 between points (1) and (5). In this region, two potential contributions could occur: Mn−Li (3.327 Å phasecorrected scattering length, N up to 2) and Mn−O (3.437 Å phase-corrected scattering length, N = 6). Despite previous studies which included metal to Li single-scattering interaction as part of EXAFS fitting,40 the very weak X-ray scattering of the Li atom compared to Mn makes it less likely for this interaction to contribute significantly, particularly considering the low expected coordination number. Instead, the Mn−O single scattering provides the largest contribution. This path corresponds to scattering between the central Mn atom of an edge-shared octahedron and the O atom on the opposite side of the adjacent octahedron, and the growth in this peak is likely due to bending of the 2 × 2 framework at the edge-shared octahedral position to accommodate Li+ or Li2O insertion. In ex situ tests, chemically lithiated xLi2O·MnO2 materials show distinct peaks in this area compared to the as-made α-MnO2, which further supports attribution of this peak to the insertion behavior (Supporting Figure S6). The growth of this peak at 2.9 Å is fairly dramatic in the case of α-MnO2 bulk, but is much more minor for the nanowire materials, indicating that insertion during cycling in O2 does not follow precisely the same mechanism in both particle size scales. The next peak, located at approximately 3.2 Å, contains contributions from Mn−Mn single-scattering from corner-shared octahedral (3.45−3.49 Å phase-corrected scattering length) and Mn−O corresponding to interaction with inserted O from Li2O in the structure. As the interplay of these peak contributions during cycling is complex, and this peak and those farther out in R-space also contains many multiple-scattering interactions, only the first three peaks are considered during the bulk of the analysis. Upon cell discharge to point (2), which reflects the material state after polarization down to the plateau potential, there is a distinct change in the relative intensity of the first two peaks (Mn−O/Mn-Mn peak ratio increases by about 0.3 for the bulk and NH4+-α-MnO2 nanowires (Supporting Figure S7) and by 0.45 for the K+-α-MnO2 nanowires, see Supporting Table S1). In this range, it is not expected that significant lithiation would have occurred in any of the structures, and the cell is only 12621



Article

nanowires is below 30 mAh g−1oxide for the first cycle. As this Li+ insertion capacity exhibited under Ar accounts for less than half of the plateau around 3.5 V vs Li/Li+ for all materials, the electrochemistry results suggest that OER is also occurring in this region, as would be expected for decomposition of Li2O2 close to the surface. This explains why the α-MnO2 bulk does not show any enhanced performance over the K+ material in cell recharge, despite having an oxidation state closer to +3 during OER. The likely reason is that Li2O2 close to the surface is consumed in the lower plateau, which occurs alongside delithiation of the αMnO2 materials, leaving only the products farther away from the α-MnO2 surface. For the three α-MnO2 materials, there is not a significant enough difference in activity compared to the large overpotential imposed by diffusion limitations for the upper plateau. Ultimately, the enhanced behavior of the K+ nanowires during ORR is the primary improvement over the other studied α-MnO2 materials, which is important for influencing O2 redox behavior going forward. The major discovery of this analysis is that the oxidation state change for K+ nanowires does not occur until the very end of ORR, meaning that the manganese oxidation state of +3.60 is present for almost the entire cell cycle. In contrast, the α-MnO2 bulk and NH4+ nanowires both exhibit continuous changes in the oxidation state during that range. This behavior runs counter to the expected intercalation, and suggests that, in this case, the oxidation state self-switching behavior may actually harm the performance of α-MnO2 during cell cycling under O2. Through a combination of Li2O insertion and suppression of oxidation state change, the K+ version out-performs other synthesized α-MnO2 materials.

Figure 5. k3-weighted phase-uncorrected EXAFS data for O2-free cells under Ar with (a) K+ nanowires and (b) bulk α-MnO2 as the cathode active material.

toward the original values, which is consistent with surfaceadsorbed discharge products decomposing during charge. For the K+-α-MnO2 nanowires, however, there is an increase in this peak ratio, and the peaks at 2.9 and 3.2 Å actually show an increase in magnitude. No clear change of Mn−O or Mn−Mn distance is observed. From the electrochemistry in Figure 2a, the K+-α-MnO2 nanowires show a plateau above 4.4 V that occurs after the discharge capacity has been consumed, which is a feature that is not evident for the other materials (which show a higher overpotential during the majority of cell charge). This can be associated with catalyzed electrolyte decomposition in the presence of oxygen, which could return the α-MnO2 to a state similar to point (2) instead of point (1), except with some residual Li+ species remaining due to partial nonreversibility. In the first-cycle recharge shown in Figure 2a, the electrochemical behavior of these materials consists of two regions: a small low-voltage plateau between the points marked (5) and (6), and a large high-voltage plateau between (6) and (7). This double-plateau phenomenon during oxygen evolution reaction (OER) when cycling in O2 has been recently explored for nanoscale Co3O4,13 as well as with analysis of evolved gases for Li−O2 cells.33 The lower plateau is attributed to fast decomposition in the initial plateau region of Li2O2 and other side reaction products close to the catalyst surface. This decomposition is then followed by a slower, diffusion-limited decomposition of the remaining discharge products in the highvoltage portion. The width of the lower plateau region varies for each material, but is approximately 110 mAh g−1oxide, which is well above the amount of charge passed via lithiation for the Ar cells, as shown in Figure 2b. Both the K+-α-MnO2 nanowires and bulk α-MnO2 show Li+ insertion ability (∼65 mAh g−1oxide), while the insertion capability of NH4+-α-MnO2

4. CONCLUSION In this work, a novel environmental in situ XAS technique was employed to detect changes to α-MnO2 in operation under O2 and Ar environments. Through analysis of the XANES and EXAFS spectra recorded at key points in the electrochemical cycle, we discovered an oxidation state self-switching behavior occurring due to a combined Li+ and lithium oxide insertion mechanism when cycled in O2 . In particular, α-MnO 2 nanowires synthesized with K+ were found to suppress a Mn oxidation state change toward Mn3+ during discharge, maintaining a high degree of Mn4+ character that may favorably influence O2 redox behavior. In contrast, nanowires synthesized with only NH4+ and bulk α-MnO2 both showed more dramatic changes in oxidation state and new coordination due to insertion of Li+ species throughout the course of cell discharge. Upon cell recharge, the α-MnO2 materials return to their initial states at low overpotential, with partial Mn3+ character for OER. With knowledge of this self-switching mechanism, new bifunctional materials can be designed to improve the O2 redox behavior in the O2 system without the need for separate materials optimizing each reaction.

■

ASSOCIATED CONTENT

S Supporting Information *

Electron microscopy of the α-MnO2 materials; electrochemistry and XAS data for the NH4+-α-MnO2 nanowires; XANES spectra; solid-state NMR spectra; k3-weighted q-space EXAFS data for the cells cycled under O2; EXAFS data for standard αMnO2-based materials; Mn−O/Mn−Mn peak ratio analysis. This material is available free of charge via the Internet at http://pubs.acs.org. 12622



■

Article

(14) Freunberger, S. A.; Chen, Y.; Drewett, N. E.; Hardwick, L. J.; Bardé, F.; Bruce, P. G. The Lithium-Oxygen Battery with Ether-Based Electrolytes. Angew. Chem., Int. Ed. 2011, 50, 8609−8613. (15) McCloskey, B. D.; Scheffler, R.; Speidel, A.; Bethune, D. S.; Shelby, R. M.; Luntz, A. C. On the Efficacy of Electrocatalysis in NonAqueous Li−O2 Batteries. J. Am. Chem. Soc. 2011, 133, 18038−18041. (16) Johnson, C. S.; Dees, D. W.; Mansuetto, M. F.; Thackeray, M. M.; Vissers, D. R.; Argyriou, D.; Loong, C.-K.; Christensen, L. Structural and Electrochemical Studies of α-Manganese Dioxide (αMnO2). J. Power Sources 1997, 68, 570−577. (17) Tompsett, D. A.; Islam, M. S. Electrochemistry of Hollandite αMnO2: Li-Ion and Na-Ion Insertion and Li2O Incorporation. Chem. Mater. 2013, 25, 2515−2526. (18) Ohzuku, T. Topotactic Reduction of Alpha-Manganese (Di)Oxide in Nonaqueous Lithium Cells. J. Electrochem. Soc. 1991, 138, 360−365. (19) Lu, Y.-C.; Gasteiger, H. A.; Parent, M. C.; Chiloyan, V.; ShaoHorn, Y. The Influence of Catalysts on Discharge and Charge Voltages of Rechargeable Li−Oxygen Batteries. Electrochem. Solid-State Lett. 2010, 13, A69−A72. (20) Peng, Z.; Freunberger, S. A.; Chen, Y.; Bruce, P. G. A Reversible and Higher-Rate Li−O2 Battery. Science 2012, 337, 563−566. (21) Leskes, M.; Drewett, N. E.; Hardwick, L. J.; Bruce, P. G.; Goward, G. R.; Grey, C. P. Direct Detection of Discharge Products in Lithium−Oxygen Batteries by Solid-State NMR Spectroscopy. Angew. Chem., Int. Ed. 2012, 51, 8560−8563. (22) Trahey, L.; Karan, N. K.; Chan, M. K. Y.; Lu, J.; Ren, Y.; Greeley, J.; Balasubramanian, M.; Burrell, A. K.; Curtiss, L. A.; Thackeray, M. M. Synthesis, Characterization, and Structural Modeling of High-Capacity, Dual Functioning MnO2 Electrode/ Electrocatalysts for Li−O2 Cells. Adv. Energy Mater. 2013, 3, 75−84. (23) Adams, B.; Radtke, C.; Black, R.; Trudeau, M.; Zaghib, K.; Nazar, L. Current Density Dependence of Peroxide Formation in the Li−O2 Battery and Its Effect on Charge. Energy Environ. Sci. 2013, 6, 1772−1778. (24) Du, P.; Lu, J.; Lau, K. C.; Luo, X.; Bareño, J.; Zhang, X.; Ren, Y.; Zhang, Z.; Curtiss, L. A.; Sun, Y.-K.; et al. Compatibility of Lithium Salts with Solvent of the Non-Aqueous Electrolyte in Li−O2 Batteries. Phys. Chem. Chem. Phys. 2013, 15, 5572−5581. (25) Harding, J. R.; Lu, Y.-C.; Tsukada, Y.; Shao-Horn, Y. Evidence of Catalyzed Oxidation of Li2O2 for Rechargeable Li−Air Battery Applications. Phys. Chem. Chem. Phys. 2012, 14, 10540−10546. (26) Yang, J.; Zhai, D.; Wang, H.-H.; Lau, K. C.; Schlueter, J. A.; Du, P.; Myers, D. J.; Sun, Y.-K.; Curtiss, L. A.; Amine, K. Evidence for Lithium Superoxide-like Species in the Discharge Product of a Li−O2 Battery. Phys. Chem. Chem. Phys. 2013, 15, 3764−3771. (27) Gao, Y.; Wang, Z.; Wan, J.; Zou, G.; Qian, Y. A Facile Route to Synthesize Uniform Single-Crystalline α-MnO2 Nanowires. J. Cryst. Growth 2005, 279, 415−419. (28) Wang, X.; Li, Y. Synthesis and Formation Mechanism of Manganese Dioxide Nanowires/Nanorods. Chem.Eur. J. 2003, 9, 300−306. (29) Mizuno, F.; Nakanishi, S.; Kotani, Y.; Yokoishi, S.; Iba, H. Rechargeable Li−Air Batteries with Carbonate-Based Liquid Electrolytes. Electrochemistry 2010, 78, 403−405. (30) Freunberger, S. A.; Chen, Y.; Peng, Z.; Griffin, J. M.; Hardwick, L. J.; Bardé, F.; Novák, P.; Bruce, P. G. Reactions in the Rechargeable Lithium−O2 Battery with Alkyl Carbonate Electrolytes. J. Am. Chem. Soc. 2011, 133, 8040−8047. (31) McCloskey, B. D.; Bethune, D. S.; Shelby, R. M.; Girishkumar, G.; Luntz, A. C. Solvents’ Critical Role in Nonaqueous Lithium− Oxygen Battery Electrochemistry. J. Phys. Chem. Lett. 2011, 2, 1161− 1166. (32) Chen, Y.; Freunberger, S. A.; Peng, Z.; Bardé, F.; Bruce, P. G. Li−O2 Battery with a Dimethylformamide Electrolyte. J. Am. Chem. Soc. 2012, 134, 7952−7957. (33) Ottakam Thotiyl, M. M.; Freunberger, S. A.; Peng, Z.; Bruce, P. G. The Carbon Electrode in Nonaqueous Li−O2 Cells. J. Am. Chem. Soc. 2013, 135, 494−500.

AUTHOR INFORMATION

Corresponding Author

*Mailing address: 150 Academy St., Newark, DE 19716 (USA). Tel: +13028313679. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS

■

REFERENCES

The work was sponsored by the University of Delaware Research Foundation (UDRF) and the startup support from the University of Delaware. G.S.H. would like to acknowledge the Collins Fellowship from the Department of Chemical and Biomolecular Engineering at the University of Delaware, and also C. Ni, F. Deng, and F. Kriss of the W. M. Keck Electron Microscopy Facility at the University of Delaware for electron microscopy time and assistance. G.S.H. and F.J. would like to acknowledge Dr. Jeffrey A. Read for helpful discussions during experimental design. 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 No. DE-AC02-98CH10886. The authors would also like to acknowledge the X10C beamline support staff at BNL.

(1) Bruce, P. G.; Freunberger, S. A.; Hardwick, L. J.; Tarascon, J.-M. Li−O2 and Li−S Batteries with High Energy Storage. Nat. Mater. 2011, 11, 19−29. (2) Christensen, J.; Albertus, P.; Sanchez-Carrera, R. S.; Lohmann, T.; Kozinsky, B.; Liedtke, R.; Ahmed, J.; Kojic, A. A Critical Review of Li/Air Batteries. J. Electrochem. Soc. 2012, 159, R1−R30. (3) Armand, M.; Tarascon, J.-M. Building Better Batteries. Nature 2008, 451, 652−657. (4) Abraham, K. M.; Jiang, Z. A Polymer Electrolyte-Based Rechargeable Lithium/Oxygen Battery. J. Electrochem. Soc. 1996, 143, 1−5. (5) Laoire, C. O.; Mukerjee, S.; Abraham, K. M.; Plichta, E. J.; Hendrickson, M. A. Elucidating the Mechanism of Oxygen Reduction for Lithium−Air Battery Applications. J. Phys. Chem. C 2009, 113, 20127−20134. (6) Allen, C. J.; Mukerjee, S.; Plichta, E. J.; Hendrickson, M. A.; Abraham, K. M. Oxygen Electrode Rechargeability in an Ionic Liquid for the Li−Air Battery. J. Phys. Chem. Lett. 2011, 2, 2420−2424. (7) Débart, A.; Paterson, A.; Bao, J.; Bruce, P. α-MnO2 Nanowires: A Catalyst for the O2 Electrode in Rechargeable Lithium Batteries. Angew. Chem., Int. Ed. 2008, 47, 4521−4524. (8) Ida, S.; Thapa, A. K.; Hidaka, Y.; Okamoto, Y.; Matsuka, M.; Hagiwara, H.; Ishihara, T. Manganese Oxide with a Card-House-Like Structure Reassembled from Nanosheets for Rechargeable Li−Air Battery. J. Power Sources 2012, 203, 159−164. (9) Xiao, J.; Xu, W.; Wang, D.; Zhang, J.-G. Hybrid Air-Electrode for Li/Air Batteries. J. Electrochem. Soc. 2010, 157, A294−A297. (10) Cheng, H.; Scott, K. Carbon-Supported Manganese Oxide Nanocatalysts for Rechargeable Lithium−Air Batteries. J. Power Sources 2010, 195, 1370−1374. (11) Cui, Y.; Wen, Z.; Sun, S.; Lu, Y.; Jin, J. Mesoporous Co3O4 with Different Porosities as Catalysts for the Lithium−Oxygen Cell. Solid State Ion. 2012, 225, 598−603. (12) Riaz, A.; Jung, K.-N.; Chang, W.; Lee, S.-B.; Lim, T.-H.; Park, S.J.; Song, R.-H.; Yoon, S.; Shin, K.-H.; Lee, J.-W. Carbon-Free Cobalt Oxide Cathodes with Tunable Nanoarchitectures for Rechargeable Lithium−Oxygen Batteries. Chem. Commun. 2013, 49, 5984−5986. (13) Black, R.; Lee, J.-H.; Adams, B.; Mims, C. A.; Nazar, L. F. The Role of Catalysts and Peroxide Oxidation in Lithium−Oxygen Batteries. Angew. Chem., Int. Ed. 2013, 52, 392−396. 12623



Article

(34) McCloskey, B. D.; Speidel, A.; Scheffler, R.; Miller, D. C.; Viswanathan, V.; Hummelshøj, J. S.; Nørskov, J. K.; Luntz, A. C. Twin Problems of Interfacial Carbonate Formation in Nonaqueous Li−O2 Batteries. J. Phys. Chem. Lett. 2012, 3, 997−1001. (35) Ravel, B.; Newville, M. ATHENA, ARTEMIS, HEPHAESTUS: Data Analysis for X-ray Absorption Spectroscopy Using IFEFFIT. J. Synchrotron Radiat. 2005, 12, 537−541. (36) Rossouw, M. H.; Liles, D. C.; Thackeray, M. M.; David, W. I. F.; Hull, S. Alpha Manganese Dioxide for Lithium Batteries: A Structural and Electrochemical Study. Mater. Res. Bull. 1992, 27, 221−230. (37) Ling, C.; Mizuno, F. Capture Lithium in αMnO2: Insights from First Principles. Chem. Mater. 2012, 24, 3943−3951. (38) Jiao, F.; Frei, H. Nanostructured Manganese Oxide Clusters Supported on Mesoporous Silica as Efficient Oxygen-Evolving Catalysts. Chem. Commun. 2010, 46, 2920−2922. (39) Grey, C. P.; Dupré, N. NMR Studies of Cathode Materials for Lithium-Ion Rechargeable Batteries. Chem. Rev. 2004, 104, 4493− 4512. (40) Pickering, I. J.; George, G. N.; Lewandowski, J. T.; Jacobson, A. J. Nickel K-Edge X-ray Absorption Fine Structure of Lithium Nickel Oxides. J. Am. Chem. Soc. 1993, 115, 4137−4144.

12624


Environmental In Situ X-ray Absorption Spectroscopy Evaluation of

Recommend Documents