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X‑ray Photoemission Spectroscopy Study of Cationic and Anionic Redox Processes in High-Capacity Li-Ion Battery Layered-Oxide Electrodes Dominique Foix,*,†,§,∥ Mariyappan Sathiya,‡,§,∥ Eric McCalla,‡,§,∥ Jean-Marie Tarascon,‡,§,∥ and Danielle Gonbeau*,†,§,∥ †
IPREM/ECP (UMR5254), Université de Pau, Hélioparc Pau Pyrénées, 2 Av. Pierre Angot, 64053 Pau Cedex 9, France Chimie du Solide-Energie, FRE 3677, Collège de France, 11 Place Marcelin Berthelot, 75231 Paris Cedex 05, France § ALISTORE-European Research Institute, FR CNRS 3104, 80039 Amiens Cedex, France ∥ Réseau sur le Stockage Electrochimique de l’Energie (RS2E), FR CNRS 3459, 80039 Amiens Cedex, France ‡
S Supporting Information *
ABSTRACT: Electrode materials based on Li-rich layered oxides are of growing interest for high-energy Li-ion battery applications because of their staggering capacities associated with the emergence of a novel, reversible anionic process. However, the fundamental science at work behind this new process needs to be well understood for further optimization. Here we report on the redox mechanisms in high-capacity Li-rich materials Li2Ru1−xMxO3 and Li2Ir1−xMxO3, by combining X-ray photoemission spectroscopy (XPS) core peaks and valence intensity analyses. We fully confirm that these materials electrochemically react with Li via cumulative reversible cationic/anionic redox processes, but more importantly we reveal that, depending on the nature of the metal (Ru or Ir), there is a delicate balance between metal and oxygen contributions. For instance, we show a greater implication of oxide ions for Irbased electrodes, consistent with the higher covalent character of Ir−O bonds compared to Ru−O bonds. We equally provide evidence that the oxygen redox process is responsible for the high capacity displayed by the Li-rich NMC Li1.2Ni0.13Co0.13Mn0.54O2 electrodes that are serious contenders for the next generation of Li-ion batteries. These combined results highlight the benefit of collecting both XPS core and valence spectra for a better understanding of anionic redox mechanisms in Li-rich layered oxides.
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INTRODUCTION
complementary techniques like X-ray photoemission spectroscopy (XPS) and electron paramagnetic resonance (EPR) on a model compound Li2Ru(1−x)SnxO3.3 The results in association with structural analysis and supported by density functional theory (DFT) calculations have revealed the existence of cumulative cationic and anionic reversible redox processes (Ru4+/Ru5+ and O2−/O2n−). This participation of anionic species O2− in the redox process is not specific to the Li2Ru(1−x)SnxO3 system, because with Li2Ru1−xTixO34 or a similar family of compounds (Li2Ir1−xSnxO3) for which we have recently reported, as deduced by combined microscopy and neutron studies, also observed has been the formation of O−O dimers in the oxygen sublattice during electrochemical oxidation.5 The involvement of O2− anion in the redox process was explained, based on DFT calculations, as the result of M(d)−O(sp) hybridization associated with a reductive coupling mechanism. However, the extent to which cationic
Layered lithium transition metal oxides LiMO2 (where M is composed of transition metals) with low molecular weight, high capacity, and reversibility are one of the best suitable electrode materials for Li-ion batteries. The first commercial Li-ion battery was assembled using LiCoO2 (∼140 mAh/g) as positive electrode material and graphite as negative. Since then, major advances in enhancing the capacity of layered transition metal oxides have been achieved via chemical substitution. First, the partial replacement of Co3+ with Ni2+ and Mn4+ has led to the Li(Ni1/3Mn1/3Co1/3)O2 layered oxides, coined as NMC, which display capacities as high as 200 mAh/g. More recently, further explorations of cationic substitutions in the layered oxide systems have led to materials made of Ni, Co, and Mn with some Li in the transition metal layers, in addition to the Li present in the van der Waals gap, termed as “Li-rich NMC” and showing capacities exceeding 250 mAh/g or greater.1,2 Such Lirich NMC electrodes are regarded as serious contenders for next-generation Li-ion batteries. The redox chemistry associated with the lithium insertion/deinsertion process of these Li-rich phases has been analyzed by our group recently using © 2015 American Chemical Society
Received: October 26, 2015 Revised: December 17, 2015 Published: December 18, 2015 862
DOI: 10.1021/acs.jpcc.5b10475 J. Phys. Chem. C 2016, 120, 862−874
Article
The Journal of Physical Chemistry C
tests were conducted at room temperature using a “Mac-Pile” or a VMP system (Biologic S.A., Claix, France). The cells were cycled with the current required to exchange 1 Li+ in 20 h from the formula Li2MO3. Ex situ samples were prepared in the glovebox where the cycled material was recovered, washed with anhydrous dimethyl carbonate (DMC) to remove any salt deposit formed during cycling, and then dried under vacuum. The completely vacuum-dried samples were carefully transferred to the XPS chamber without any air exposure. XPS. XPS measurements were carried out with a Kratos Axis Ultra spectrometer, using focused monochromatic Al Kα radiation (hν = 1486.6 eV). The XPS spectrometer was directly connected through a transfer chamber to an argon drybox, in order to avoid moisture/air exposure of the samples. The analyzed area of the samples was 300 × 700 μm2. Peaks were recorded with constant pass energy of 20 eV. For the Ag 3d5/2 line, the full width at half-maximum (fwhm) was 0.58 eV under the recording conditions. The pressure in the analysis chamber was ∼5 × 10−9 mbar. Short acquisition time spectra were recorded before and after each normal experiment to check that the samples did not suffer from degradation during the measurements. To check the reproducibility of the XPS results, analyses were done on several sets of samples. The binding energy scale was calibrated using the C 1s peak at 285.0 eV from the hydrocarbon contamination invariably present. To avoid any error on the calibration choice and for more precision, binding energy difference between the O 1s lattice and the metal core peaks was also examined. Core peaks were analyzed using a nonlinear Shirley-type background.12 The peak positions and areas were optimized by a weighted least-squares fitting method using 70% Gaussian and 30% Lorentzian line shapes. Quantification was performed on the basis of Scofield’s relative sensitivity factors.13 The curves fit for core peaks were obtained using a minimum number of components in order to fit the experimental curves. Moreover, for the O 1s spectra, the fwhm (full width at half maximum) were free but blocked to a maximum value to keep a consistency with the pristine electrode. For estimation of the uncertainty in the peroxo fraction, the O 1s spectrum of each pristine sample was fitted in two ways: with three components as reported and with four components like in the fully charged samples where peroxo-like species are present in significant quantities. Since no peroxo is present in the pristine material, this gives us an estimate of XPS detection limit for peroxo species (2−3% for Ru based electrodes, 2% for Ir-based electrodes, and 1.5% for Li1.2Ni0.13Co0.13Mn0.54O2) and allows the determination of the uncertainty on the peroxo fractions in the scans after electrochemical cycling (error bars in Figures 3, 10, and 12). The Ru 3d and Ir 4f spin-orbit doublets for pristine and cycled samples were fitted based on the values (intensity ratios, energy splitting between the two components (5/3−3/2 or 7/ 2−5/2), obtained for the active material of each set of compounds). For the valence band intensity analyses, a baseline (Shirley-type background) was chosen for each spectrum between 0 and ∼12 eV, constrained by the experimental profile of the valence spectrum in this energy range. Considering the whole valence spectrum (structures A, B, and B′), the area of the main band below Fermi level (band A) was estimated to obtain a satisfactory fitting of the global experimental profile of structure A (Figures 6, 7, 8, and 11). The area of band A was normalized to the area of Ru 4p or Ir 5p3/2, taking into account the differential photoionization cross
and anionic electronic transfers take place upon Li delithiation/ relithiation has not been thoroughly explored, particularly in relation to the complex cycling-driven changes in the M−O interactions. To get further insights into the redox processes in these Lirich oxides where electronic changes affect both oxygen and the transition metal, we have embarked on a systematic X-ray photoemission spectroscopy (XPS) study combining core and valence spectra at different stages of the charge/discharge process. To our knowledge, it is the first time that valence intensity analysis were done in conjunction with the study of core level spectra for better accessing the interplay between lithiation and electronic structure evolution in these Li-rich oxides. We fully confirm and specify the cumulative reversible cationic/anionic redox processes in these different materials with, therefore, the more important implication of oxide ions for Ir-based electrodes that will be discussed in terms of covalent bonding. For sake of clarity, the paper is organized as follows. First, we present electron transfers occurring in Li2Ru0.75Sn0.25O3, Li2Ru0.75Ti0.25O3, and Li2RuO3, for which a structural transition at ∼540 K and the tendency to Ru−Ru dimerization has been recently reported.6−11 Next, we extend the approach to Li2IrO3 and Li2Ir0.75Sn0.25O3 to examine the effect of introducing a 5d metal (Ir) so as to increase the covalence, prior to concluding by extending our analysis to Li-rich NMC (Li1+xNiyMnzCo1−x−y−zO2) for generalization of our findings.
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EXPERIMENTAL SECTION Synthesis. For all Li2MO3 compounds, stoichiometric amounts of Li2CO3 (with 10% excess to account for loss during synthesis) and the necessary combination of metal oxides RuO2 (Sigma-Aldrich 99.9%), SnO2 (Sigma-Aldrich, 99.8%), TiO2 (Sigma-Aldrich 99.9%), and IrO2 (Alfa Aesar, 99%) were ball-milled for 1 h. The resultant mixtures were pelletized and heated in a box furnace. For Li2RuO3, the sample was heated at 900 °C for 12 h followed by regrinding and heating at 950 °C for 12 h. For Li2Ru0.75Sn0.25O3, the mixture was heated at 800 °C for 6 h followed by heating at 900 °C for 12 h and at 1100 °C for 12 h with intermediate ball-milling and repelletizing. The Li2Ru0.75Ti0.25O3 samples were heated at 800 °C for 24 h. For Li2Ir1−xSnxO3 samples, the pellets were heated to 1000 °C for 15 h, ground, and pelletized again before heating to 900 °C for 36 h. The phase purities of all compounds were verified with X-ray diffraction as detailed elsewhere.3−5 The lithium-rich NMC compounds Li1.2Ni0.13Co0.13Mn0.54O2 have been synthesized by hydroxide coprecipitation of the metal hydroxide followed by mixing and ball-milling of metal hydroxide with a stoichiometric amount of lithium hydroxides and sintering at 900 °C. Electrochemical Studies. Electrochemical tests were done in Swagelok-type cells that could cycle up to 70 mg of the powder material. The cells were assembled in an argon-filled glovebox, using a Li metal disc as the negative electrode, a Whatman GF/D borosilicate glass fiber separator, and 1 M LiPF6 solution in a mixture of ethylene carbonate, propylene carbonate, and dimethyl carbonate in 1:1:3 ratio by weight (LP100) as electrolyte. The working electrodes were loose powder, typically made by ball-milling the active material powders with 10% in mass of Super P carbon black (except for Li2RuO3, for which no carbon black has been added) for 20 min in a Spex-800 ball mill. Usually, 40−70 mg of the mixed powders were used per cell. Galvanostatic charge−discharge 863
DOI: 10.1021/acs.jpcc.5b10475 J. Phys. Chem. C 2016, 120, 862−874
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The Journal of Physical Chemistry C
component, corresponding to a satellite structure as reported earlier,14 was necessary for the spectral decomposition. After charging to 3.6 V, the binding energy of the Ru 3d spectrum shifts positively by 0.7 eV, with the same being observed after the charge at 3.75 and 4 V (Figure 2b). These results are consistent with a change in the oxidation state of ruthenium from Ru4+ to Ru5+. In addition, in each case, an increase in the full width at half-maximum (fwhm) of Ru 3d peak is observed (for Ru 3d5/2: from 1.2 eV for the pristine to ∼1.5 eV for the charged states); this can be associated with some differences in ruthenium environments or can be due to a biphasic process as previously reported.17 With further removal of Li when charging at 4.15 and 4.6 V, no further positive shift is observed for Ru 3d core peak but rather a reverse evolution occurs (negative shift by ∼0.7 eV at 4.6 V compared to 4 V). On the basis of common XPS interpretation, this negative shift in binding energy of the transition metal ion, during charging to high voltage, is unexpected and will be discussed later. At 4.6 V, the Ru 3d peak is also narrowed as compared to the other charged samples (3.6, 3.75, and 4 V), suggesting a well-defined environment for ruthenium ions. Figure 2c shows the evolution of the O 1s core spectra upon removal of Li+ in Li2RuO3. The O 1s spectrum of the pristine electrode displays two main components: the one at ∼529.5 eV is characteristic of the O2− anions of the crystalline network, while the other at ∼531.7 eV is assigned to weakly adsorbed species at the surface. A weak component at ∼533.0 eV is also ascribed to adsorbed species. These signatures are preserved after charging to 3.6 and 4 V. A clear change is observed when pursuing the charge beyond 4 V. Indeed, the spectral analysis and the fitting procedure of the experimental curve imply the presence of an additional component at ∼530.5 eV, which can be explained by the existence of formal (O2)n− species or undercoordinated oxygen atoms, as previously discussed.3 A small decrease (∼15%) of (Olattice529.5eV/Ru) atomic ratio is observed on samples charged to 4.15 and 4.6 V, with the initial value being restored when considering the 530.5 eV component as a part of the O lattice. This new component could thus be assigned to the lattice oxygen of Li2−xRuO3 material. Similar analyses have been carried out for Li2Ru0.75Sn0.25O3 and Li2Ru0 0.75Ti0.25O3 at different stages of the charge process, included in Figures 3 and S1. The Ru 3d and O 1s spectra (Figure 3a) of the pristine electrodes exhibit the same characteristics as for the Li2RuO3 compound. A positive shift is noted up to 4 V for the Ru 3d spectrum, whereas further charging leads to a progressive shift back. For the O 1s core peak, the change of the experimental shape only occurs during charging through the 4.25 V plateau and again implies the introduction of an additional component at ∼530.5 eV. This O 1s component that disappears/reappears on subsequent discharges/charges was previously characterized for Li2Ru0.5Sn0.5O3 and attributed in line with EPR and theoretical results to peroxo-like species (O2)n−, demonstrating the redox activity of oxide ions.3 The data here obtained highlight, for all Ru-based electrodes, the appearance and increase of this O 1s signature between 4 V and charge at 4.6 V. This is illustrated by a peroxo fraction (probably higher than in the bulk) that shows a similar evolution upon charge for all compounds (Figure 3b). A fine analysis of the O 1s peak was possible for these Li2RuMO3 electrodes because of the fact that there were only small changes in the peak due to deposited species at the electrode/electrolyte interface at high potentials.
sections. Replicate measurements on several sets of compounds proved the reproducibility of the results.
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RESULTS AND DISCUSSION Li2Ru1−yMyO3 Series. Electrochemistry. The electrochemical performances of Li2Ru0.75M0.25O3 (M = Ru/Sn/Ti) compounds were tested versus Li in Swagelok cells cycled between 2 and 4.6 V at a C/10 rate (Figures 1a,b and S1). The
Figure 1. Voltage profiles over the first cycles at C/10 between 2.0 and 4.6 V for (a) Li2RuO3 and (b) Li2Ru0.75Sn0.25O3.
results are similar as reported earlier.3,4 All the samples show the staircase charge voltage on first charge followed by an S-like curve on first discharge and all subsequent cycling. Herein we have concentrated on studying the changes in the electronic arrangement of the metals and oxygens triggered by reversible uptake and release of Li in these materials. To this end, various samples at different states of charge/discharge (shown on the voltage curves in Figures 1 and S1c) have been studied by XPS combining core peaks and valence spectra (more representative of the bulk) analyses. XPS Core PeaksFirst Charge. The Ru 3d and O 1s spectra of the Li2RuO3 electrodes are shown in Figure 2. The lack of carbon black additive in this electrode composition allows more accurately the analysis of Ru 3d peak in spite of its partial overlap with the C 1s line. The Ru 3d spectrum of the pristine electrode (Figure 2a) corresponds to a simple spin-orbit doublet with the 3d5/2 and 3d3/2 components located at 282.2 and 284.6 eV, respectively. Similar binding energy (BE) values have been reported for other Ru(IV) oxides14,15 and ascribed, according to the theory developed by Kotani,16 to an “unscreened” final state. To fit the asymmetry on the higher binding energy side of the Ru 3d peak, an additional 864
DOI: 10.1021/acs.jpcc.5b10475 J. Phys. Chem. C 2016, 120, 862−874
Article
The Journal of Physical Chemistry C
Figure 2. Li2RuO3 first chargeXPS core peaks: (a) Ru 3d5/2−3/2−C 1s, (b) Ru 3d5/2, and (c) O 1s.
environment in conjunction with some (Ru5+ → O2n−) electronic redistribution could thus be the reason for the systematic Ru 3d reverse shift in XPS spectra. XPS Core PeaksFirst Discharge. For Li2RuO3, XPS analyses were performed on samples charged to 4.6 V and discharged to either 2.8 or 2 V. In both cases, the Ru 3d peak evolves compared to the fully charged state with still further negative shift of the maximum without recovering the shape of the pristine material (Figure 4a). However, a detailed analysis (Figure 4b) shows that a good fitting of the experimental curve is obtained when the Ru 3d peak is broken down into two components: one at the same binding energy as the pristine and the other associated with a negative shift compared to the pristine. Note that, for a sample discharged either to 2.8 or 2 V after charging only up to 4 V (not reported), only one component at the same binding energy as for the pristine is necessary to fit the experimental spectrum. The presence of two different Ru environments for the samples discharged to 2.8/2 V after charging to 4.6 V suggests a two-phase system probably due to the change in Ru environment after the formation of peroxo-like species as discussed earlier. We have checked that this result is intrinsic to the material and not due to poor electrode kinetics as was observed for electrodes cycled at different rates.
Universally, for all Li2Ru1−xMxO3 compounds investigated, the BE shifts of Ru 3d change from >0 with respect to the pristine material up to 4 V to
