Lithium Ion Coupled Electron-Transfer Rates in Superconcentrated

Jun 21, 2017 - The charge-transfer kinetics of lithium ion intercalation into LixMn2O4 cathode materials was examined in dilute and concentrated aqueo...
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Lithium Ion Coupled Electron-Transfer Rates in Superconcentrated Electrolytes: Exploring the Bottlenecks for Fast Charge-Transfer Rates with LiMn2O4 Cathode Materials Victoria A. Nikitina,*,†,‡ Maxim V. Zakharkin,†,‡ Sergey Yu. Vassiliev,† Lada V. Yashina,† Evgeny V. Antipov,† and Keith J. Stevenson‡ †

M. V. Lomonosov Moscow State University, Leninskie Gory 1/3, Moscow 119991, Russia Skolkovo Institute of Science and Technology, Moscow 143026, Russia

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‡

S Supporting Information *

ABSTRACT: The charge-transfer kinetics of lithium ion intercalation into LixMn2O4 cathode materials was examined in dilute and concentrated aqueous and carbonate LiTFSI solutions using electrochemical methods. Distinctive trends in ion intercalation rates were observed between water-based and ethylene carbonate/diethyl carbonate solutions. The influence of the solution concentration on the rate of lithium ion transfer in aqueous media can be tentatively attributed to the process associated with Mn dissolution, whereas in carbonate solutions the rate is influenced by the formation of a concentrationdependent solid electrolyte interface (SEI). Some indications of SEI layer formation at electrode surfaces in carbonate solutions after cycling are detected by X-ray photoelectron spectroscopy. The general consequences related to the application of superconcentrated electrolytes for use in advanced energy storage cathodes are outlined and discussed.

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INTRODUCTION The energy demands of modern society dictate the necessity for the practical implementation of higher energy density and higher power metal ion and metal−air batteries. Increasing the operation potential of the cathode material can simultaneously increase both capacity and power characteristics. Unfortunately, the design and study of 5 V lithium-intercalating cathode materials for energy storage have been hindered by the insufficiently wide potential windows of conventional electrolytes, which are mainly based on the mixtures of ethylene carbonate (EC) with linear alkyl carbonates and LiPF6 salt.1 This electrolyte composition provides a sufficient anodic (up to 4.5 V vs Li+/Li) as well as cathodic potential window of stability due to the formation of protective electrode/electrolyte interfaces, which kinetically hinder the reduction and oxidation of the electrolyte. However, there is a significant impetus for the discovery of more stable electrolyte systems to enable the operation of 5 V electrode materials in advanced energy storage applications. It is now generally understood that the thermodynamic stability of a majority of solvents in contact with a high-voltage cathode, which catalyzes the solvent decomposition reaction, cannot practically exceed 5.0−5.5 V. This range is true even for those solvents that have been reported to be stable toward oxidizing potentials, such as dinitriles, sulfones, and fluorinated carbonates. However, they still suffer from decreased reductive © 2017 American Chemical Society

stability because stable solid electrolyte interface (SEI) layers cannot be formed in these electrolytes.1,2 Under these conditions, the only option rests on decreasing the rate of the decomposition reaction at both oxidative and reductive potentials by creating a kinetic barrier for ion-coupled electrontransfer reactions at the electrode/electrolyte interface. Recently, it has been reported that the increase in the electrolyte salt concentration can be used to tune both the reductive and oxidative stability of electrode/electrolyte interfaces in a large variety of solvents3 (linear carbonates, acetonitrile, dimethyl sulfoxide, dimethoxyethane, tetrahydrofuran, glymes, etc.) as well as to suppress the corrosion of aluminum current collectors without using highly hygroscopic and thermally unstable LiPF6 salt.4−6 Organic EC-free concentrated electrolyte solutions (sometimes referred to as superconcentrated electrolytes) are compatible with graphite anodes because the initial reduction of salt anion (bis(trifluoromethane)sulfonamide (TFSI) or bis(fluorosulfonyl)amide (FSA) forms a protective film on the material surface, Special Issue: Fundamental Interfacial Science for Energy Applications Received: March 24, 2017 Revised: June 15, 2017 Published: June 21, 2017 9378

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decomposition. The observed trends allow for a more rational choice of salt concentration and solvent nature for the highervoltage battery systems. The in-depth electrochemical study of the intercalation kinetics in concentrated electrolytes is supplemented by the XPS analysis of the electrodes, which is aimed at finding correlations between the electrochemical behavior and chemical composition of the electrode/electrolyte interfaces.

which prevents graphite exfoliation and solvent co-intercalation.7−10 It was shown that using higher concentrations of conventional LiPF6 salt in carbonate solvents results in a decrease in charge-transfer resistance in Li[Ni0.4Mn0.4Co0.2O2]/ graphite cells11 and in the suppression of graphite exfoliation in propylene carbonate solutions12 due to the changes in the electrode/electrolyte interface structure in more concentrated solutions. Still, the application of organic superconcentrated electrolytes to high-voltage systems is questionable because only one example of the stable operation of a high-voltage LiNi0.5Mn1.5O4 spinel in LiFSA/DMC (dimethyl carbonate) electrolyte (1.3:1 solvent/salt molar ratio) has been reported.8 Other studies consider the application of concentrated solutions to sodium−air13 and lithium−sulfur14−16 battery systems, which are beyond comparison because many other parameters can effect the operational parameters. Superconcentrated aqueous systems also provide attractive possibilities to expand the potential window of stability for water-based solutions up to 2−2.5 V, greatly enhancing the operating voltage of an aqueous battery, although very high salt concentrations (up to 21 m LiTFSI) are required for the suppression of solvent decomposition.6,17 In some cases, the increased concentration of lithium salt is primarily used to minimize the dissolution of the transition metal from the electrode material in aqueous lithium ion18,19 and even calcium ion batteries.20 Attractive properties of superconcentrated electrolytes are dictated by both kinetic and thermodynamic factors. Kinetic suppression of solvent decomposition reactions, aluminum corrosion, and graphite exfoliation is primarily caused by the decrease in the number of free solvent molecules, which results in the preferential reduction/oxidation of the salt anion (TFSI or FSA) instead of the aprotic solvent, forming a compact layer at the electrode/electrolyte interface, which is enriched in fluoride and sulfur, as demonstrated in XPS studies.4,17 Thermodynamic factors result in the suppression of the undesired dissolution of system components, such as transition-metal ions in electrode materials,2,21 polysulfides in lithium−sulfur batteries,14,16 and aluminum current collectors.4,5 Although the thermodynamic factors can be easily rationalized for any particular electrode material/electrolyte solution combination, assessing the kinetic factors is much more complex because of diverse interfacial effects that are determined by the type and concentration of the salt as well as by the structure and composition of the material and the nature of the solvent. It is particularly important to understand the influence of the chemical composition and thickness of the surface layers on the rate of the intercalation reaction and to separate the contributions from the SEI and other surface layers’ resistance, charge-transfer resistance, and concentration effects into the intercalation rate. A deep understanding of the superconcentrated electrolyte effect on the kinetics of intercalation processes is required for the successful implementation of this strategy for high-voltage batteries. Therefore, in this study we focus on a well-studied 4 V cathode material LiMn2O4 in carbonate- and water-based concentrated electrolytes. This choice gives us the possibility to examine intrinsic kinetic effects, induced by the concentrationdependent electrode/electrolyte interface restructuring in aqueous and organic concentrated LiTFSI-based solutions, avoiding the complications induced by the electrolyte

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EXPERIMENTAL SECTION

Material Synthesis and Electrode Preparation. LiMn2O4 powder was prepared via a solid-state reaction from MnCO3 and Li2CO3 as described in ref 22. The particle mean weight diameter, as determined from SEM imaging, was 1.6 μm. Composite electrodes were fabricated by mixing 80% LiMn2O4, 10% Super P carbon black (Timcal), and 10% poly(vinylidenefluoride) (PVDF) as a binder. The electrode mass was suspended in an appropriate amount of N-methylpyrrolidone and spread over a platinum foil support. The resulting electrodes were dried at 120 °C in vacuum (p(O2) < 10−2 atm) for 4 h. The typical loading of the cathode material was 0.3−0.5 mg·cm−2. Electrolyte Preparation. An appropriate amount of ethylene carbonate, EC, (99%, anhydrous, Aldrich) was dissolved in diethyl carbonate, DEC, (99%, anhydrous, Aldrich) to achieve a 1:1 volume ratio. After the complete dissolution of EC, the mixture was dried with activated 3 Å molecular sieves overnight. The residual amount of water as determined by Karl Fisher titration was less than 30 ppm. The weighed amount of LiTFSI salt (>99.95%, Aldrich) was added to the EC/DEC solution to achieve molar ratios of EC/DEC to LiTFSI equal to 11:1 (ca. 1 mLiTFSI solution), 4:1 (ca. 2.5 m LiTFSI solution), and 2:1 (ca. 5 m LiTFSI solution). Aqueous solutions were prepared by dissolving an appropriate amount of LiTFSI in Millipore water (Milli-Q 18 MΩ·cm). Water to LiTFSI molar ratios were chosen to be 55:1 (ca. 1 m LiTFSI solution), 11:1 (ca. 5 m LiTFSI solution), and 6:1 (ca. 9 m LiTFSI solution). Electrochemical Measurements. Electrochemical measurements in both aqueous and carbonate electrolytes were performed in a threeelectrode glass cell. For measurements in EC/DEC solutions, a glassy carbon electrode was used as a counter electrode, and lithium foil served as a reference electrode. In aqueous solutions, 3 M KCl AgCl/ Cl− reference and glassy carbon counter electrodes were applied. Electrochemical data were registered with a Biologic VMP3 potentiostat/galvanostat. At least two electrochemical experiments (cyclic voltammetry and EIS) were performed for each electrolyte concentration to ensure the reproducibility of the results. The potentiometric intermittent titration technique (PITT) was employed to determine the variation of the lithium diffusion coefficient with the degree of intercalation. The potential was scanned in 10 or 20 mV steps. In aqueous electrolyte, the potential was restricted to 0.9−1.0 V vs AgCl/Cl− as a result of the start of the water decomposition reaction at higher potentials. At the end of each potential step, the residual current did not exceed the background value (typically 0.5−2 μA). Apparent diffusion coefficients were extracted from PITT data by accounting for the uncompensated ohmic resistance and slow interfacial kinetics following Montella’s model.23 The details of the calculations were reported previously.24,25 Impedance spectroscopy measurements (EIS) were performed to determine the kinetic parameters of the intercalation reaction in aqueous and carbonate concentrated electrolytes. Impedance spectra were registered in steps of 10 or 20 mV, after equilibration was reached for each potential value (usually 1−1.5 h). The frequency range was 10 kHz−10 mHz, with a 5 mV peak-to-peak alternating voltage. The MEISP program package was used to model experimental impedance spectra.26 To increase the accuracy of the approximation, constant phase elements (CPE) were used instead of capacitances. For all of the examined spectra, the CPE power was very close to unity. Because the deviations from ideal capacitive behavior are very low, the CPE elements were treated as capacitances in the data analysis and discussion. Impedance spectroscopy data were normalized per the true 9379

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Figure 1. (a) CVs of LiMn2O4 electrodes in 0.5 m Li2SO4 and 1 m LiTFSI solutions at 50 and 200 μV·s−1 scan rates. The potential scale is shifted to match the scale in 1 m LiTFSI solution (normalized potential scale). (b) CVs of LiMn2O4 electrodes in TFSI-based solutions: 55:1, 11:1, and 6:1 water/LiTFSI molar ratios. The scan rate is 100 μV·s−1. surface area of the material particles, which was computed from the coulometric mass assuming a monodisperse sample with a particle size of 1.6 μm. XPS Measurements. The electrodes for XPS experiments were cycled (10 cycles) in aqueous (55:1, 11:1, and 6:1 water/LiTFSI, 0.2− 1.1 V vs AgCl/Cl−) and carbonate (11:1, 4:1, and 2:1 EC-DEC/ LiTFSI, 3.5−4.5 V vs Li+/Li) solutions. After the final discharge, the electrodes cycled in carbonate solution were removed from the cell in a glovebox and washed with anhydrous DEC (soaking in the solvent for ca. 3 min) to remove the adsorbed LiTFSI. Washing with Millipore water under room conditions took place for electrodes cycled in aqueous solutions. The samples after cycling in carbonate solutions were loaded into the spectrometer without contact with the atmosphere using the Aldrich AtmosBag attached to the load lock. The as-synthesized LiMn2O4 composite electrode, carbon black, PVDF, and LiTFSI powders were also investigated as reference samples. XPS measurements were carried out using a Kratos Axis Ultra DLD spectrometer equipped with a monochromatic Al Kα X-ray source. XP spectra for Li 1s, Mn 3p, S 2p, C 1s, N 1s, O 1s, Mn 2p, and F 1s were acquired with an electron pass energy of 10−40 eV. The spectra were approximated with a Gaussian/Lorentzian convolution function with a Shirley/Tougaard background shape using the Unifit 2014 Software. Quantification of the composition was performed by taking into account the analyzer transmission function.

LiTFSI, which can be indicative of slower kinetics or lower diffusion coefficient values (vide infra). Figure 1b shows CVs in LiTFSI solutions with water/LiTFSI ratios of 55:1 (ca. 1 m), 11:1 (ca. 5 m), and 6:1 (ca. 9 m). Further increases in the electrolyte concentration were found to be incompatible with the three-electrode measurement setup with a standard aqueous 3 M KCl AgCl/Cl− reference electrode, so the highest concentration of the salt in our study corresponds to a 6:1 ratio. Given the lithium solvation number in aqueous solutions of 4, at a 6:1 water/LiTFSI ratio all of the solvent molecules are expected to belong to Li+ and TFSI− solvation shells. The formal potentials of LiMn2O4 redox processes exhibit a progressively positive shift, which amounts to 75 and 163 mV for 11:1 and 6:1 ratios relative to the potentials in 55:1 solutions. According to the Nernst equation for a Li+ intercalation reaction, the equilibrium potential of the reaction E depends on the lithium ion activity in solution (aLi+) and its activity in the crystal lattice (a[Li+]): a(Li+) RT E = E0 + ln F a([Li+]) (1) Therefore, the activity in the concentrations series increases accordingly so that the Li+ activity in the 6:1 solution is 3 orders of magnitude higher than that in the 1 m solution, which is typical behavior for highly concentrated solutions as a result of the strongly repulsive interactions between similar species. However, the increase in the concentration and the activity of LiTFSI does not seem to affect the oxidative stability of the electrolyte because the background current values do not decrease with the increase in salt concentration. Correspondingly, because the redox potentials are shifted positively, the CVs of LiMn2O4 electrodes are significantly distorted in the region of the second redox process, and no quantitative analysis of the curves is possible for 5 and 10 m concentrations. The absence of the suppression of the solvent oxidation in LiTFSI superconcentrated electrolytes is in agreement with the results reported by Suo et al.,17 where a considerable increase in the potential window width originated primarily from the suppression of solvent reduction and not oxidation. Figure S1a,c,e shows CVs in aqueous LITFSI solutions at scan rates of 50−200 μV·s−1. It can be noted that the peak-topeak separations in the 55:1 and 11:1 solutions exhibit a moderate increase from 20 to 55 mV with the increase in the scan rate, which is indicative of the transition from thermodynamic to diffusional reaction rate control.22 Peak-to-

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RESULTS AND DISCUSSION Concentrated Aqueous Solutions. Stable charge/discharge characteristics were repeatedly demonstrated for LiMn2O4 composite electrodes in a variety of aqueous electrolytes (LiNO3, LiCl, Li2SO4, and LiTFSI).17,19,27−30 However, no comparative data on the diffusion coefficient and charge-transfer kinetics or resistance values have been reported. To analyze possible anionic effects on the lithium intercalation reaction, we first compare the intercalation kinetics in 0.5 and 1 m aqueous solutions of Li2SO4 and LiTFSI, respectively. Figure 1a shows cyclic voltammograms (CVs) of LiMn2O4 electrodes in aqueous solutions involving two electrolytes, Li2SO4 and LiTFSI. The formal potential values for the two pairs of peaks associated with Li+ intercalation in LiMn2O4 in LiTFSI solution are shifted toward less-positive potential values by ca. 26 mV as compared to the formal potential values in Li2SO4, which agrees with the higher ionic association tendency of LiTFSI that apparently causes the decrease in lithium ion activity.31 The shape of CV is very similar in the two solutions, although the peak-to-peak separations are slightly higher in 9380

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Langmuir peak separations reach 80 mV in 6:1 solution, and given the low value of uncompensated ohmic resistance, this difference could be attributed to slower reaction kinetics in the most concentrated aqueous solution. Apparent diffusion coefficient values (Dapp) for Li+ solid-state diffusion in LiMn2O4 material in Li2SO4 and LiTFSI solutions were determined from PITT experiments under a spherical finite diffusion approximation (Figure 2). The positive potential

Figure 3. Nyquist plots for LiMn2O4 electrodes in aqueous solutions at 0.7 V (AgCl/Cl−, normalized potential scale). Z′ and Z″ values are normalized per the true surface area of the material particles.

at 0.7 V (normalized potential scale) in LiTFSI and Li2SO4 solutions, and Figures S2 and S3 in the Supporting Information (SI) present data taken at different potentials and the corresponding fits (Table S1). Nyquist plots in aqueous electrolytes show a single distorted semicircle, which contrasts with the LiMn2O4 EIS data in carbonate solutions, where a distinctive high-frequency semicircle appears.32,33 However, in all LiTFSI solutions the EIS spectra differ greatly from perfect semicircles, which can point to a more complex intercalation mechanism in LiTFSI, whereas the semicircle recorded in Li2SO4 solution is not distorted. This result is supported by the measured charge-transfer resistance (Rct) values, evaluated as the diameters of the corresponding semicircles using a simple Randles circuit (Figure 4a) and normalized per the true surface area of the material particles, which show only a minor potential dependence in the interval of 0.6−0.85 V (Figure 5a). Because this behavior clearly deviates from the Butler−Volmer model, which predicts a sharp rise in Rct values in regions of very low and very high intercalation degrees,24 it is reasonable to assume that the intercalation mechanism includes not only

Figure 2. Diffusion coefficients in concentrated aqueous electrolytes as determined from PITT data. Intercalation capacity in 55:1 (1 m) LiTFSI solution. Normalized potential scale.

limits were restricted to 0.85 V for the 11:1 LiTFSI solution and 0.75 V for the 6:1 solution (on the 55:1 LiTFSI potential scale) to avoid the influence of the start of solvent oxidative decomposition. The Dapp values in all of the electrolytes follow a general trend, which was observed for LiMn2O4 in organic carbonate solutions, with a shallow minimum in the vicinity of the formal potential of the first pair of LiMn2O4 peaks.32,33 The minimum in the Dapp value corresponds to the maximum in the intercalation capacity, Cint (Figure 2). The absolute values of Dapp in aqueous solutions are close to the values determined in organic carbonate solutions (on the order of 10−11 cm2·s−1). The Dapp values for all three LiTFSI solutions do not show a strong dependence on the LiTFSI concentration. However, diffusion coefficients in the Li2SO4 aqueous solution are still higher than those in LiTFSI (by a factor of 2), which points to the existence of some sort of specific interactions between the salt and the electrode material. Because the model used in the calculation of Dapp from PITT data disregards the short-time region, which may be distorted by some barrier layers at the material surface with a different diffusion coefficient, the physical reasons for the decrease in the Dapp in LiTFSI solutions are unclear. The variation of Dapp with the change in the solution composition implies changes in the material properties rather than the changes in the layers at the material surface. Because XRPD characterization of the electrodes after cycling does not point to any significant changes in the material bulk, the decrease in Dapp in the LiTFSI solutions may be tentatively attributed to the influence of the LiMn2O4 particle size polydispersity, which is not considered in our calculations. Further investigation of this effect will be the subject of a separate study. Intercalation Kinetics in Aqueous Electrolytes. Chargetransfer rates for the lithium intercalation reaction in aqueous electrolytes were determined from electrochemical impedance spectroscopy (EIS). Figure 3 shows the EIS response measured

Figure 4. Equivalent circuits used for fitting the experimental impedance plots. (a) Randles circuit−solution resistance, Rsol; charge-transfer resistance, Rct; double-layer capacity, Cdl; and reflective finite Warburg impedance, W. (b) Circuit for classical SEI formation− resistance of an SEI layer, RSEI; capacity of the SEI layer, CSEI. (c) Circuit for anion and cation conducting SEI−transmissive finite Warburg impedance, Wsh. 9381

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Figure 5. (a) Rct values as determined from EIS spectral fits using a Randles equivalent circuit on the normalized potential scale. (b) Rct values normalized per the lithium molar concentration.

LixMn2O4 electrodes, the XPS analyses were carried out for the as-received electrodes, samples kept in electrolyte, and those cycled in concentrated LiTFSI electrolytes. Prior to XPS measurements, the cycled samples were washed with the same solvent to remove LiTFSI residues (Figure S4). The composition quantification is presented in Figure 6. Electrode surfaces are mainly composed of carbon and fluorine originating from carbon black and PVDF and oxygen as well.

stages of simple charge-transfer kinetics but also stages associated with desolvation steps, or mass transport effects, such as diffusion through SEI layers. Another observation is related to the minimal Rct values observed in 0.5 m Li2SO4 solution (35 Ω·cm2 in 0.5 m Li2SO4 vs 300 Ω·cm2 in 1 m LiTFSI at 0.7 V). These values should reflect either a higher barrier for desolvation in the more associated, though dilute, electrolyte, or, more likely, kinetic effects related to the specific surface structure of LiTFSI at the electrode surface, or changes in the material surface layers after contact with the solution. The Rct values in 11:1 LiTFSI are on the order of 50 Ω·cm2, whereas for the 6:1 solution the chargetransfer resistance rises to ca. 120 Ω·cm2 (at 0.7 V). Importantly, the separation of the concentration effects matters even given the complex shape of the semicircles in LiTFSI solutions. Figure 5b shows Rct values, normalized per molar concentrations of LiTFSI and Li2SO4 solutions. The charge-transfer resistance value is given by R ct =

RT RT = 2 2 α 1−α nFi0 n F ksc R cO

(2)

where n is the number of transferred electrons, i0 is the exchange current density, ks is the heterogeneous rate constant of charge transfer, α is the transfer coefficient (α = 0.5), and cR and cO are the concentrations of lithium ion involved in the reaction and vacancy concentrations, respectively. Rate constant ks in this representation contains the molar concentration factor, i.e., ks = ks0cs(Li+)1−α. Thus, using higher concentrations of lithium salt should naturally result in a decrease in Rct, although the intrinsic concentration-independent rate constant is not varied if no additional factors contribute to its value. In the series of normalized Rct values, the lowest resistance is again observed in Li2SO4 solution. The Rct values in 55:1, 11:1, and 6:1 solutions amount to ca. 300, 90, and 210 Ω·cm2·mol1/2· L−1/2. The normalized Rct values in 11:1 LiTFSI electrolyte are considerably lower than those in 55:1 solution. This could point to the differences in surface layer structures, implying that a less-resistive interface is formed in 11:1 LiTFSI. Rct resistance in the 6:1 solution is much larger, which points to severe kinetic hindrances. Because the observed trends clearly demonstrate the complex nature of the intercalation process in LiTFSI solution, modeling the reaction with a simple charge transfer does not seem to reflect the physics of the process. Surface Composition in Aqueous Electrolytes. To examine the possible formation of surface layers at the

Figure 6. Surface composition for the pristine LiMn2O4 composite electrode and electrodes cycled in aqueous solutions.

For the pristine electrode, the most intense feature in the O 1s spectrum shown in Figure 7 is positioned at about 529.6 eV and can be assigned to oxide species such as O2− in the lattice of LiMn2O4. Other features at higher binding energies can be attributed to the surface OH groups (associated with residual oxygen contamination) and probably a small amount of water. In the literature, the presence of these species on uncycled electrodes is associated with residual oxygen contamination in electrode components such as PVDF and carbon black,34 weakly adsorbed species,35 and so forth. These observations fit our data for as-prepared LiMn2O4 and PVDF (not shown here), which were also found to possess the same photoemission intensity above 530 eV in O 1s spectra, which probably originated from oxygen-containing contamination during storage. In the C 1s spectrum shown in Figure 7, the major features at 284.3 eV (C−C, C−H), 285.9 eV (−CH2−), and 290.2 (−CF2−) indicate a high portion of carbon black and PVDF on the surface. F 1s shows a single broad component at 687.7 eV 9382

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low-energy components of Mn 2p and 3p peaks, which is likely to indicate the decreased amount of Mn3+ on the surface of the electrode upon cycling.36 Considering possible occasional variation of the carbon black and PVDF arrangement over the electrode surface, no pronounced effects of cycling in differently concentrated aqueous electrolytes on surface elemental composition over a probing depth of about 5−7 nm could be deduced from the available data. Aqueous SEI Modeling. XPS analysis does not provide definitive information regarding the concentration dependence of the surface composition of the electrodes, which can imply that either specific adsorbed species are dissolved after the electrodes are washed or that no adsorbate layers are formed in aqueous solutions. In this case, the difference in electrochemical behavior of the electrodes can originate from different thicknesses of defect layers (due to, for instance, Mn dissolution) at the electrode surface. On the basis of these considerations, the electrochemical data were modeled with an account of complex interface structure. Several equivalent circuits can be used to model the intercalation reaction. The simplest model is based on the Randles circuit, which does not provide good fits to our experimental data because the semicircles in the impedance spectra appear to be distorted. In the case when a single ion conducting surface layer is formed (i.e., a classical SEI layer), the transport of cations in the bulk of the solid electrolyte results in the appearance of the electric field gradient along with the concentration gradient. Under these conditions, the flux of the ions is affected by both diffusion and migration processes, which results in ohmic behavior with specific conductivity σ37

Figure 7. XPS spectra of LiMn2O4 composite electrodes cycled in aqueous concentrated electrolytes.

corresponding to C−F bonds. The Mn 2p peak shape is known to be practically impossible to analyze because of multiplet splitting and a shake-up satellite individual not only for each possible charge state but also for different local structure and the overlay of different states. Upon cycling in aqueous electrolytes, the C 1s spectrum practically does not change and the carbon atomic concentration also does not change much. At the same time, the O 1s spectral shape becomes essentially modified against a very small increase in oxygen atomic concentration. Because the O 1s relative intensity of lattice oxygen decreases after cycling, one can conclude that after cycling some surface layer forms that is not thicker than about 4 nm. This layer can be formed by LiOH or Li2CO3. Both are soluble in water but can form adsorbed layers. A possible formation of carbonate results in the increase of a peak at 290.2 eV by only 1 atom % as compared to the uncycled electrode (Figure S5), and this difference does not make it possible to trace both in the C 1s and Li 1s spectra. It should be noted that the scattering of the data over different spots on the samples (spot size 300 × 700 μm2) is high, and the effect of cycling as compared to the sample kept in electrolyte is not very pronounced. The fluorine spectrum is not modified except for a small energy shift (the same as for PVDF-related features in the C 1s spectrum) and for minor peak broadening of about 0.2 eV (Figure S6). This broadening was previously related to the distinct bonding between C and F in the PDVF-C composites as well as to the presence of (CxF)n chains with various C/F.34 The absence of features at 685 eV suggests that no inorganic fluoride forms upon cycling, which was also previously reported for LiMn2O4 cycled in concentrated aqueous LiTFSI solution.17 Sulfur and nitrogen are detected at the surface in very small amounts; their chemical state corresponds to that of the TFSI− anion. The cycling results in the vanishing of the low-energy component in Mn 2p and Mn 3p spectra (Figure S7). In summary, one can assume that cycling in aqueous solutions gives rise to a thin (adsorbed) layer of Li2CO3, LiOH, water, or other oxygen-containing species. We cannot exclude the formation of Mn hydroxides or carbonates, which are not soluble. The concentration of the solution has the most pronounced effects on the O 1s spectra above 530 eV and the

σ=

F 2Dc RT

(3)

where D is the diffusion coefficient of the mobile ion in the solid electrolyte and c is the concentration of the ion. Because double layers are formed both at solid electrolyte/material and solution/solid electrolyte interfaces, the system can be modeled by the equivalent circuit in Figure 4b. Two semicircles should appear in the Nyquist plots for a system with an SEI, but the resolution of the two semicircles will depend on the differences in the time constants of the two circuits, which can be determined, for instance, by the ratio of the double-layer capacity values. In our case, the overlap of the SEI and chargetransfer semicircles could result in the observed shape of the spectra. In cases in which both negative and positive ions are mobile in the surface layer, the electric field gradient does not develop and diffusion becomes the only ion transfer rate controlling process. The complex resistance of this layer can be modeled as finite-length diffusion with a transmissive boundary.38 In this case, a sloping line is observed in the region of high frequencies, which transforms into a semicircle in the region of lower frequencies. This kind of distortion is typical for aqueous LiTFSI solutions, so the model of a cation- and anionconducting SEI can be applicable to the experimental data. The double layer in this case is formed only at the solid electrolyte/ material interface. This case can be represented by the equivalent circuit in Figure 4c. The impedance spectra of all LiTFSI solutions were fitted to the equivalent circuits in Figure 4b,c (Figure S3), which imply the formation of either a classical lithium ion conducting solid electrolyte (SEI) or an interface, which conducts both cations 9383

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Figure 8. Rct (a) and RSEI(Wsh) (b) values on the normalized potential scale, as determined from the fitting of the experimental impedance spectra using circuits in Figure 4b,c.

and anions (Wsh). The RSEI (or RWsh) and Rct values cannot be unambiguously separated for all of the potential values as a result of substantial overlapping of the two semicircles, which results in significant scattering of the extracted values. However, the calculated dependencies allow us to distinguish some general trends (Figure 8). First, the two models produce rather similar potential dependencies of the Rct parameter. From the data available, it is impossible to draw conclusions about the more probable ion transfer mechanism. The Rct values (at 0.7 V, normalized per lithium molar concentration) for 55:1, 11:1, and 6:1 solutions amount to ca. 100, 60, and 190 Ω·cm2 (Figure 8a). This points to the same nonlinear concentration dependence with the minimum Rct value for the 11:1 solution. The Rct vs E plots for all of the solutions show a moderate rise at low intercalation degree values, which resembles Butler−Volmer behavior. However, this rise is too shallow to be compatible with the one-step ion-transfer mechanism. At higher potential values, the data for 11:1 and 55:1 solutions overlap, which points to rather close charge-transfer rates in these solutions. The chargetransfer rate in the 6:1 solution is apparently slower. Completely different trends are observed in RSEI (or RWsh) vs E dependencies (Figure 8b). RSEI values in 11:1 and 6:1 solutions scatter significantly; however, all of the values are in the range of 5−30 Ω·cm2. For the 55:1 solution, the range of values is considerably higher, 50−250 Ω·cm2. Contrasting data in Figure 8a,b suggests that for the most dilute solution low intercalation rates originate from the high resistance of the interfacial layers. For 11:1 and 6:1 solutions, the RSEI values are rather low. However, for 6:1 solution the Rct resistance is the highest in the series. Optimal values of Rct and RSEI for the 11:1 solution ensure the highest intercalation rates. Presumably, the enhanced resistance of the interfacial layers in the 55:1 solution originates from the dissolution of the surface Mn atoms due to Mn3+ disproportionation, which appears to be more pronounced in LiTFSI solutions due to complexation reactions.39 The defectiveness of the resulting surface layer should hinder the ion transfer. For the 6:1 solution, high Rct values dictate the low intercalation rate that can result from the formation of a dense resistive interface (i.e., the surface layer can be more soluble in less-concentrated solutions), which is caused by the aggregation of the LiTFSI salt because no signs of the LiTFSI oxidative decomposition products at the electrode surface were detected by XPS. The electrochemical analysis suggests that very different concepts can underlie the appearance of high-frequency

semicircles in the impedance spectra of oxide cathode materials. The RSEI or Wsh notations can take on the physical meaning of resistive layers, which are formed by ether (i) the dissolution of the material surface with the formation of defective structures or (ii) the oxidative decomposition of solution components on top of the material particles. For LiTFSI solutions in water, XPS does not allow for the determination of precisely the composition of the blocking layer. Still, our results allow us to suggest that in the case of aqueous solutions the SEI changes from a dissolution-induced defective layer at the material surface to a layer comprising adsorbed/precipitated LiTFSI salt or other solution components when going from dilute to more concentrated solutions. Concentrated EC/DEC Solutions. Dilute LiTFSI electrolytes cannot be applied in practical lithium-ion batteries because of severe corrosion of aluminum and stainless steel current collectors. However, the behavior of LiMn2O4 in LiTFSI solutions in EC/DEC is similar to the behavior in LiPF6-based carbonate electrolytes when platinum supports are used.22 In this study, 11:1, 4:1, and 2:1 carbonate/LiTFSI solutions were analyzed. Increasing LiTFSI concentration up to a 4:1 ratio has only a minor impact on the shape of CVs of LiMn2O4 electrodes at both lower and higher scan rates (Figure 9, Figure S1b,d,f). The peak-to-peak separations reveal the same trend in the transition from thermodynamic to diffusional control as was found for aqueous systems. It should also be noted that no shift of the redox potential value is observed as a result of the application of the Li+/Li reference electrode.

Figure 9. CVs of LiMn2O4 electrodes in 11:1, 4:1, and 2:1 carbonate/ LiTFSI solutions at a 100 μV·s−1 scan rate. 9384

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Langmuir Because the scale of the reference electrode exhibits the same potential dependence as that of the formal potential of the LiMn2O4 redox processes, the change in measured potential values cannot be observed. However, it should be kept in mind that the redox potentials are likely to shift in the direction of more positive and therefore more oxidizing potentials. In the most concentrated 2:1 electrolyte, the shape of the CV becomes distorted, with the peak-to-peak separations demonstrating a very pronounced scan rate dependence (Figure 9, Figure S1b,d,f). The peak-to-peak separations increase up to 500 mV at a 200 μV·s−1 scan rate. The increase in the peak-topeak separations points to a substantial decrease in the chargetransfer rate or the diffusion coefficient. The shape of the 2:1 CV is also affected by a high ohmic drop in solution, which amounts to ca. 200 Ω (for a separation between the working and reference electrodes of ca. 0.2−0.3 cm). The potential dependence of the apparent diffusion coefficients was determined from PITT data using the same model as that applied to aqueous systems. Figure 10 shows Dapp

Figure 11. Nyquist plots for LiMn2O4 electrodes in carbonate LiTFSI solutions at 3.980 V (Li+/Li). The ohmic resistance was subtracted from the plots for clarity.

Both equivalent circuits provide similar accuracy of the fits and reveal the same trends in the resistance potential dependencies (Table S2). Because it is impossible to distinguish between different types of ionic conductivity in the interfacial layers based on the electrochemical data alone, here we report only the results of the fitting to the classical SEI circuit (Figure 4b). Figure 12 shows Rct and RSEI values extracted from the experimental impedance spectra in the potential interval of 3.8−4.3 V (Li+/Li). The Rct values demonstrate a pronounced increase in the regions of low and high states of charge, which is compatible with the Butler−Volmer charge-transfer model (Figure 12a). In the region of intermediate intercalation degrees (3.9−4.1 V), the Rct values in 11:1, 4:1, and 2:1 solutions amount to 130, 450, and 4500 Ω·cm2·mol1/2·L−1/2 (normalized per molar concentration of LiTFSI salt). The charge-transfer rate decreases drastically with the increase in salt concentration, which points to the differences in the structure of the SEI/material interface in concentrated solutions. RSEI values show a limited potential dependence (Figure 12b). Increases in the salt concentration result in a monotonous increase in RSEI values: ca. 20 (11:1), 150 (4:1), and 900 (2:1) Ω·cm2. The increase in RSEI values clearly points to the change in SEI structure: a thicker, denser SEI is formed in concentrated carbonate LiTFSI solutions. Surface Composition in Carbonate Electrolytes. The atomic composition of electrode surfaces cycled in carbonate electrolytes (Figure 13) is similar to that for the cycled in aqueous solutions, as measured by XPS. Its results indicate the analogy among the general trends found in aqueous electrolytes mentioned for F 1s, Mn 2p, N 1s, and S 2p spectra. The lateral band of O 1s spectra above 530 eV in contrast to aqueous electrolytes is much more thoroughly discussed in the literature34,40−43 and is commonly assigned to C−O, CO, and more complex bonds of ethers, esters, alkoxides, and carboxylates originating from organic solvents. Nevertheless, an analysis of the band components revealed no correlations with the peak of carbonates at 290.5 eV in C 1s spectra. However, an increase in the hydrocarbon component at 285.2 eV in the C 1s spectra is typical for samples cycled in carbonate electrolytes, which suggests that residues of organic solvents or/and its decomposition products are the most plausible components of the surface coverage in the case of carbonate electrolytes. Here, however, again the data scattering over different spots on the samples is high, and the effect of cycling as compared to the sample kept in electrolyte is detectable but cannot be exactly quantified.

Figure 10. Diffusion coefficients in concentrated EC/DEC electrolytes as determined from PITT data. Intercalation capacity in 11:1 (1 m) LiTFSI solution.

values for Li+ in LiMn2O4 electrodes in carbonate LiTFSI solutions. In contrast to the aqueous systems, in carbonate solutions the diffusion coefficients vary systematically in the three electrolytes, although the overall shape of the Dapp vs E dependence remains unchanged, with the minima in Dapp values corresponding to the maxima in Cint (Figure 10). For instance, at 4.0 V, the D values decrease rather monotonously, following the sequence 1.5 × 10−11 (11:1), 7.0 × 10−12 (4:1), and 1.2 × 10−12 cm2·s−1 (2:1). As was discussed with respect to aqueous systems, the origin of the diffusion coefficient concentration dependence is unclear. In carbonate electrolytes, the drop in Dapp values is much more pronounced (ca. 1 order of magnitude). The account of the real particle size distribution is unlikely to compensate for this effect. Nyquist plots in carbonate LiTFSI solutions point to a significant increase in the overall resistance values with the increase in salt concentration (Figure 11). In contrast to aqueous solutions, a high-frequency semicircle can be observed even at the lowest solvent/salt ratio (11:1) (Figures S8 and S9). For 4:1 and 2:1 ratios, two partially overlapping semicircles can be observed, which are much better resolved compared to aqueous systems. The separation of high-frequency and low-frequency semicircles in carbonate solutions can be performed by fitting the experimental spectra to the circuits in Figure 4b,c (Figure S8). 9385

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Figure 12. Rct (a) and RSEI (b) potential dependencies, as extracted from the fitting of the experimental impedance spectra to the equivalent circuit in Figure 4b.

The F 1s spectra also shift and broaden upon cycling, and a tiny feature at 684−685 eV was observed (Figure 14, inset). It suggests the presence in our samples of fluoride species, which are different from PVDF, yet no interrelations between regions of F 1s and Li 1s spectra corresponding to LiF were observed (Figures S10 and S11) possibly because of the lower detection limit of Li 1s. The greatest influence of the solution concentration was again found for the shapes of the O 1s and Mn 2p and 3p peaks and for the width of the F 1s peaks. The appearance of a component at 534 eV of O 1s spectra for 11:1 and 4:1 samples was not present for the electrolyte cycled in 2:1 carbonate. This component was found to be particularly prominent for discharged electrodes repeatedly cycled over 30 times in 1 M LiPF6 carbonate solutions.34,41,44 It is noteworthy that this feature was not observed for any of the electrodes cycled in aqueous solutions or for the pristine electrode. Another feature of the 2:1 concentration case is that the low-energy component of Mn 2p and Mn 3p spectra does not change after cycling, in contrast to the 4:1 and 11:1 concentration cases. Therefore, we

Figure 13. Surface composition for the pristine LiMn2O4 composite electrode and electrodes cycled in carbonate solutions.

Figure 14. XPS spectra of LiMn2O4 composite electrodes cycled in carbonate electrolytes. 9386

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dance spectra. Clearly, the elucidation of the ion-transfer mechanism requires more experimental and theoretical studies to be performed. At this point, we call the attention of the electrochemical and material science community to the observed trends and correlations in the intercalation reaction kinetics.

might conclude that cycling in the most concentrated carbonate electrolyte resulted in the smallest difference in the formation of a surface layer with respect to that of the pristine electrode surface. XPS analysis of the electrodes cycled in carbonate electrolytes pinpoints the probable adsorption of solvent decomposition products at the electrode surfaces, and no implications for LiTFSI oxidative decomposition were found. An analysis of the electrochemical data suggests that the increase in salt concentration in organic solutions primarily results in the kinetic effects related to the changes in surface layer structures. At moderate salt concentrations (2 and 3 m), this can be beneficial because more thick and stable SEI should result in better cycling stability of the material. Figure S12 collects CVs in 1 and 3 m LiTFSI carbonate solutions. The effect of the salt concentration on the stabilization of the electrochemical response is obvious. However, it should be kept in mind that this stabilization comes at a price of a decreased charge-transfer rate.

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CONCLUSIONS In this work, we studied the intercalation kinetics in superconcentrated aqueous and carbonate electrolytes. It was found that the variation in the intercalation rate is nonlinear in aqueous solutions, with the maximal rate being observed for a 6:1 water/TFSI solution. In carbonate solutions, the increase in salt concentration results in a monotonous decrease in the intercalation rate. The rate of lithium ion intercalation was found to be highly dependent on the material surface structure with two different mechanisms of SEI formation in LiTFSI electrolytes. In dilute aqueous solutions (55:1 water/LiTFSI), the resistance of the interface is suggested to be determined by the formation of a defective layer due to the dissolution of Mn from the material particles’ surface. In the highly concentrated aqueous electrolytes, (6:1 water/LiTFSI) the results imply the formation of a very resistive interface, which cannot originate from the dissolution because this process is already suppressed in the 6:1 electrolyte. It can be assumed that some clustering or precipitation of LiTFSI at the particles’ surface takes place, which cannot be traced with XPS because these layers are very soluble and should disappear after washing the sample before the XPS measurements. In carbonate-based electrolytes, the charge transfer rate decreases monotonously with the increase in salt concentration as a result of the formation of a resistive SEI interface. In this case, SEI formation is more likely to be related to the oxidative decomposition of the solution components as supported indirectly by XPS analysis. In contrast to anodic SEI at lithium or graphite in superconcentrated electrolytes, where the reduction of the LiTFSI salt stabilizes the SEI, the cathodic SEI composition does not seem to include any salt decomposition products. The thickening of the SEI can be related to the higher intercalation/deintercalation potentials, which originate from the increase in LiTFSI activity. For the intermediate concentration, the formation of thicker SEI creates a barrier for Mn dissolution, while the rate of the reaction is still lower than that in more dilute solutions. The results of our study allow us to suggest that the enhancement of charge transfer rate in superconcentrated electrolytes is primarily related to the suppression of Mn dissolution. These results should be generally applicable to other solvents, which do not form stable SEI layers (acetonitrile, sulfones, etc.). If this suppression is accompanied by the formation of surface films from solvent decomposition products, then the reaction rate will be substantially decreased, limiting the performance of the intercalation system. Blocking effects are expected to manifest themselves in electrolytes based on cyclic carbonates with or without polymerizing electrolyte additives. A compromise between the rate of the intercalation process and cycling stability can be reached for some intermediate concentrations. We hope our findings would allow for a more precise design of salt/solvent systems, aimed at neutralizing the limitations of superconcentrated solutions.

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DISCUSSION The major differences in the electrolyte concentration dependence of the intercalation rate in aqueous and carbonate solutions stem from different surface chemistries formed in organic EC/DEC solvent and water. In aqueous media, the pronounced enhancement of the intercalation rate for the intermediate salt concentration can be tentatively attributed to the inhibition of Mn dissolution, which is apparently accompanied by defective surface-layer formation. This tendency is not observed in the EC/DEC LiTFSI electrolytes because the increase in the salt concentration with a simultaneous shift of the redox potentials toward more positive potential values results in an increase in the interfacial resistance, which is demonstrated by the increase in RSEI values. Higher electrolyte concentration in this case ensures the cycling stability of the electrode at the price of a reduced charge transfer rate, even though the dissolution of the material is naturally inhibited. At the highest concentrations, both in water and EC/DEC the reaction becomes essentially blocked by the surface layers, which in this case are likely to be formed by the adsorbed/aggregated LiTFSI in water and EC/DEC decomposition products in carbonate electrolytes. Unfortunately, the extreme complexity of the intercalation systems does not allow us to find immediate definitive answers to all of the intriguing questions regarding the specifics of the ion-transfer mechanism through the solution/SEI/material interface. The results of our study indicate that there is a linear correlation between the RSEI and Rct values in carbonate electrolytes. The same linear correlation is observed for the RSEI and diffusion resistance Rd ≈ 1/Dapp for EC-DEC/LiTFSI solutions. The situation is complicated by the fact that the Rct values extracted from the impedance spectra do not demonstrate sharp increases at the lowest and highest states of charge, associated with Butler−Volmer charge-transfer kinetics. This result could point to the complex nature of the observed medium frequency semicircle, which does not necessarily relate to the pure charge-transfer resistance. The extracted values can contain a resistive component associated with another slow step, whereas the charge-transfer step can be rather fast. For instance, our preliminary calculations suggest that diffusional transport through SEI with a slow Li ion transfer step at the solution/SEI interface results in two overlapping semicircles with correlated diameters on impe9387

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(10) Yamada, Y.; Usui, K.; Chiang, C. H.; Kikuchi, K.; Furukawa, K.; Yamada, A. General Observation of Lithium Intercalation into Graphite in Ethylene-Carbonate-Free Superconcentrated Electrolytes. ACS Appl. Mater. Interfaces 2014, 6, 10892−10899. (11) Petibon, R.; Madec, L.; Abarbanel, D. W.; Dahn, J. R. Effect of LiPF6 concentration in Li[Ni0.4Mn0.4Co0.2]O2/graphite pouch cells operated at 4.5 V. J. Power Sources 2015, 300, 419−429. (12) Nie, M.; Abraham, D. P.; Seo, D. M.; Chen, Y.; Bose, A.; Lucht, B. L. Role of Solution Structure in Solid Electrolyte Interphase Formation on Graphite with LiPF6 in Propylene Carbonate. J. Phys. Chem. C 2013, 117 (48), 25381−25389. (13) He, M.; Lau, K. C.; Ren, X.; Xiao, N.; McCulloch, W. D.; Curtiss, L. A.; Wu, Y. Concentrated Electrolyte for the Sodium− Oxygen Battery: Solvation Structure and Improved Cycle Life. Angew. Chem., Int. Ed. 2016, 55, 15310−15314. (14) Shin, E. S.; Kim, K.; Oh, S. H.; Cho, W. I. Polysulfide dissolution control: the common ion effect. Chem. Commun. 2013, 49, 2004−2006. (15) Urbonaite, S.; Novák, P. Importance of ‘unimportant’ experimental parameters in Li−S battery development. J. Power Sources 2014, 249, 497−502. (16) Cuisinier, M.; Cabelguen, P. E.; Adams, B. D.; Garsuch, A.; Balasubramanian, M.; Nazar, L. F. Unique behaviour of nonsolvents for polysulphides in lithium-sulphur batteries. Energy Environ. Sci. 2014, 7, 2697−2705. (17) Suo, L.; Borodin, O.; Gao, T.; Olguin, M.; Ho, J.; Fan, X.; Luo, C.; Wang, C.; Xu, K. Water-in-salt” electrolyte enables high-voltage aqueous lithium-ion chemistries. Science 2015, 350, 938. (18) Ruffo, R.; Wessells, C.; Huggins, R. A.; Cui, Y. Electrochemical behavior of LiCoO2 as aqueous lithium-ion battery electrodes. Electrochem. Commun. 2009, 11, 247−249. (19) Wang, G. J.; Zhang, H. P.; Fu, L. J.; Wang, B.; Wu, Y. P. Aqueous rechargeable lithium battery (ARLB) based on LiV3O8 and LiMn2O4 with good cycling performance. Electrochem. Commun. 2007, 9, 1873−1876. (20) Lee, C.; Jeong, S.-K. A Novel Superconcentrated Aqueous Electrolyte to Improve the Electrochemical Performance of Calciumion Batteries. Chem. Lett. 2016, 45, 1447−1449. (21) Manthiram, A.; Chemelewski, K.; Lee, E.-S. A perspective on the high-voltage LiMn1.5Ni0.5O4 spinel cathode for lithium-ion batteries. Energy Environ. Sci. 2014, 7, 1339−1350. (22) Vassiliev, S. Y.; Levin, E. E.; Nikitina, V. A. Kinetic analysis of lithium intercalating systems: cyclic voltammetry. Electrochim. Acta 2016, 190, 1087−1099. (23) Montella, C. Discussion of the potential step method for the determination of the diffusion coefficients of guest species in host materials: Part I. Influence of charge transfer kinetics and ohmic potential drop. J. Electroanal. Chem. 2002, 518, 61−83. (24) Levin, E. E.; Vassiliev, S. Y.; Nikitina, V. A. Solvent effect on the kinetics of lithium ion intercalation into LiCoO2. Electrochim. Acta 2017, 228, 114−124. (25) Nikitina, V. A.; Fedotov, S. S.; Vassiliev, S. Y.; Samarin, A. S.; Khasanova, N. R.; Antipov, E. V. Transport and Kinetic Aspects of Alkali Metal Ions Intercalation into AVPO4F Framework. J. Electrochem. Soc. 2017, 164, A6373−A6380. (26) MEISP. Kumho Chemical Laboratories, Korea, March 2002. (27) Tang, W.; Tian, S.; Liu, L. L.; Li, L.; Zhang, H. P.; Yue, Y. B.; Bai, Y.; Wu, Y. P.; Zhu, K. Nanochain LiMn2O4 as ultra-fast cathode material for aqueous rechargeable lithium batteries. Electrochem. Commun. 2011, 13, 205−208. (28) Luo, J. Y.; Xia, Y. Y. Aqueous Lithium-ion Battery LiTi2(PO4)3/ LiMn2O4 with High Power and Energy Densities as well as Superior Cycling Stability. Adv. Funct. Mater. 2007, 17, 3877−3884. (29) Liu, S.; Ye, S. H.; Li, C. Z.; Pan, G. L.; Gao, X. P. Rechargeable Aqueous Lithium-Ion Battery of TiO2/LiMn2O4 with a High Voltage. J. Electrochem. Soc. 2011, 158, A1490−A1497. (30) Jayalakshmi, M.; Mohan Rao, M.; Scholz, F. Electrochemical Behavior of Solid Lithium Manganate (LiMn2O4) in Aqueous Neutral Electrolyte Solutions. Langmuir 2003, 19, 8403−8408.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.langmuir.7b01016. CVs and Nyquist plots of LiMn2O4 electrodes. Fits of the experimental impedance spectra in aqueous LiTFSI solutions to the equivalent circuits. Parameters evaluated from the fitting of impedance spectra in aqueous and carbonate LiTFSI solutions. Comparisons of spectra of pristine electrodes and electrodes cycled in solution. Unwashed samples demonstrating the salt spectrum. F 1s peak of a pristine electrode and cycled electrodes. XPS spectra of Mn 2p peaks of the pristine electrode and cycled electrodes. (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Victoria A. Nikitina: 0000-0002-0491-3371 Lada V. Yashina: 0000-0002-8370-9140 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS V.N. acknowledges the support of RFBR, according to research project no. 16-33-60036 mol_a_dk. This work was also supported by the Skoltech Center for Electrochemical Energy Storage and Lomonosov Moscow State University Program of Development up to 2020.

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REFERENCES

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