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The cycle stability of lithium negative electrodes for Li–air secondary batteries was studied under oxygen atmosphere using Li∣Li symmetric cells with three organic electrolyte solutions: 1.0 M LiCF3SO3/tetraglyme (G4), 1.0 M LiN(SO2CF3)2/G4, and 1.0 M LiNO3/G4. Of these, 1.0 M LiNO3/G4 showed excellent stability without dendrite deposition, even for increased dissolution/deposition capacity from 0.50 to 2.0 mAh cm−2. These results are considered to be due to the stable Li2O passivation layer that was formed, not only by the direct reaction with oxygen, but also by the action of NO3 − as an oxidant, which released NO2 − as a redox mediator. Li–O2 cells with 1.0 M LiNO3/G4 showed a clear charging voltage plateau at 3.7 V, which evidenced the redox mediator effect of NO2 −, and cell cycleability was enhanced to 25 cycles.

In recent years, rechargeable Li-air batteries (LAB) have attracted much attention because of their potential high energy density over 5 times larger than that of the conventional Li-ion batteries [1]. However, there are many problems still to be solved for all the cell components: the cathode, anode and electrolyte solution. As for the anode, Li dendrite growth during discharge/charge cycles should be suppressed to stabilize the cycle life of the LAB. In order to improve the reversibility of the Li electrode, we have been evaluating tetraglyme (G4)-based electrolyte solutions by using a Li|Li symmetric cell [2]. We confirmed that a 1.0 M LiNO3/G4 electrolyte solution containing O2 formed a Li2O protective layer on the Li surface effectively, which suppressed the formation of Li dendrite at the current density of 0.2 mA cm-2. In this study, high rate tests up to 2.0 mA cm-2 above our target of 0.4 mA cm-2 were applied to verify and analyze the reversibility of Li electrode reaction. LiN(SO2CF3)2 (LiTFSI, 99.9%, KISHIDA) and LiNO3 (99.9% purity, Wako) were vacuum- dried at 110 °C and dissolved in tetraglyme (G4, Japan Advanced Chemicals, H2O < 30 ppm) solvent to prepare 1.0 M glyme electrolyte solution of each electrolyte salt. Li|Li symmetric cells with 1.0 M LiNO3/G4 or 1.0 M LiTFSI/G4 electrolyte solution were assembled in an Ar-filled glove box to be used for Li dissolution/deposition tests in an isothermal chamber set at 30°C at a constant current from 0.2 to 2.0 mA cm-2. The surface of the Li electrodes after the dissolution/deposition test and its cross-section were observed by SEM-EDS and analyzed by elementary analysis. In addition, the outermost surface was analysed by using XPS. The flatness of the polarization curve for dissolution/deposition tests in the 1.0 M LiNO3/G4 electrolyte suggests that Li dendrite growth and electrolyte decomposition are suppressed at an applied current of 0.6 mA cm-2 or less. Fig. 1 shows the SEM image of the Li electrode surface after the dissolution/deposition test. Although no large dendrites were observed in the LiNO3/G4 electrolyte, the entire surface was covered with fine Li dendrites of about 100 nm in diameter and about 800 nm long in the LiTFSI one (Fig. 1a & c). In addition, the cross-sectional SEM image of the Li surface tested in the LiNO3/G4 electrolyte shows some pits, while that in the LiTFSI/G4 electrolyte solution exhibits bulky decomposition product layer of about 20 μm thick accompanying a larger crevices (Fig. 1b & d). C 1s spectra of XPS showed a dominant signal assigned to Li2CO3 on the electrodes tested at higher rates, which suggested the formation of the protective layer of Li2CO3 and Li2O. Moreover, weak signals such as C-C and C-O detected on the Li surface in LiNO3/G4 electrolyte solution suggested the suppressed degradation. The details of the mechanism will be discussed on our presentation. This study was supported by JST “Next Generation Batteries Area in Advanced Low Carbon Technology Research and Development Program (ALCA)” from MEXT, Japan. [1] P. G. Bruce, S. A. Freunber, L. J. Hardwick, J-M. Tarascon, Nature materials, 11, 19 (2012). [2] M. Saito, S. Yamada, T. Fujinami, S. Kosaka, Y. Tachikawa, K. Ito, Y. Kubo, ECS Trans., 75(22), 53 (2017). Figure 1

Introduction Rechargeable non-aqueous Li-air batteries (LABs) have attracted much attention as one of the next-generation batteries because of their potential for possessing large specific energy (theoretical value: 3.5 kWh kg-1) more than 5 times of those of Li-ion batteries [1]. For practical use, some technical problems have to be overcome; during discharging, the choking of air electrode by the Li2O2 generated; during charging, the requirement of applying high overpotential for the electrochemical oxidation of Li2O2 or the short circuit of LAB cells caused by Li dendrite growth on Li electrode. Li dissolution/deposition reactions using various Li salts, solvents, and additives have been investigated under oxygen free condition, which is far from practical use, to find that the repetition of the reaction provides Li deposition on Li electrode non-uniformly with clear boundary. This is probably caused by the chemical composition and uniformity of the solid electrolyte interphase (SEI) films formed on Li electrode. Such Li deposition can raise the problems mentioned above. These remaining problems, moreover, restrict the Li dissolution/deposition reactions only at low depth of charge/discharge though LABs are promising due to their large specific energy. We have reported that the presence of oxygen improved the Li dissolution/deposition reversibility, due to the formation of proper SEI films [2]. In this study, focusing on not only presence of oxygen gas but also LiNO3-based electrolyte [3], we evaluated the reversibility of Li | Li symmetric cell and Li-air (Li | O2) cell at relatively high state of charge/discharge. Experimental Li | Li symmetric cells was constructed with Li metal foils (thickness：0.5 mm) used for both electrodes, a separator (Celgard 2400), and 1.0 M glyme-based electrolytes containing LiSO3CF3 (LiOTf), LiN(SO2CF3)2 (LiTFSI) or LiNO3 in an Ar-filled dry box. Li-air (Li | O2) cells were constructed in the similar method but Ketjen Black-loaded carbon paper was used as an air electrode, the positive electrode. Li dissolution/deposition tests were conducted under Ar or O2 atmosphere. The applied current density and maximum discharge/charge capacity were 0.20 mA cm−2 and 2.0 mAh cm−2 (corresponding to 10 μm of Li dissolution/deposition), respectively, in the cut off voltage range of −2.0 to 2.0 V. The Li metal electrodes after 15 cycles of the tests were rinsed with tetraglyme to remove Li salts, and the obtained electrodes were analyzed using scanning electron micrograph (SEM) equipped with energy dispersive X-ray spectroscopy (EDS) and X-ray photoelectron spectroscopy (XPS) to characterize the resulting SEI films. Results and Discussion Fig. 1 showed the surface and cross-sectional SEM images of Li electrodes of Li | Li symmetric cells after 15 cycles of Li dissolution/deposition reactions at 2.0 mAh cm−2. In the case of LiTFSI-based electrolyte, oxygen suppressed crack and Li dendrite growth (thick deposition) compared to oxygen free condition. In addition, oxygen suppressed the decomposition of the electrolyte, LiOTf tetraglyme solution, and enhanced the deposition of Li2O, which works as a protective layer, though rather non-uniformly. These are indicated by chemical compositions determined by EDS (F and S contents are lower under oxygen atmosphere) and XPS spectra (O1s, Li1s, C1s, and F1s spectra). In the case of LiNO3-based electrolyte under oxygen, larger overpotential was observed than that in the case of LiOTf and LiTFSI but the suppression of dendrite growth and the uniform Li2O deposition were observed on the Li electrodes. These must contribute to improve the reversibility. The similar tendency was also observed for Li-air (Li | O2) cells. The cell using LiNO3 tetraglyme solution as an electrolyte showed the suppression of Li dendrite growth and the uniform deposition of Li2O, though the charge/discharge reversibility was slightly improved in this experimental condition. These results can be associated with the formation of uniform Li2O protective layer on Li negative electrode, probably promoted by NO3 -: NO3 -+ 2Li →NO2 -+ Li2O [4]. Of particular note is that these Li dissolution/deposition reactions using LiNO3-based electrolyte, which showed the suppression of Li dendrite growth and the uniform deposition of Li2O, were performed at high depth of charge/discharge, where the thickness of dissolved/deposited layer reached about 10 μm, comparable to practical use. References [1] P. G. Bruce,et al., Nature Mater., 11, 19 (2012). [2] M. Saito, et al., Electrochimica Acta, 252, 192 (2017). [3] J. Uddin, et al., J. Phys. Chem. Lett., 4, 3760 (2013). [4] D. Sharon, et al., ACS. Appl. Mater. Interfaces. 7, 16590, (2015). Figure 1

In recent years, development of new-generation batteries with high energy densities exceeding that of Li-ion batteries (LIBs) has been intensively studied. One of the candidates is a non-aqueous Li-air (O2) battery (LAB) and the energy density is more than 5 times larger than those of LIBs. Also, the cost is relatively low because of no containing of precious metals in the positive electrode (PE). However, for the practical use, the LAB still have some problems, e.g., Li metal dendrite growth at the negative electrode, electrochemical durability for electrolytes especially against O2 - radical at the PE, large overpotential to decompose the Li2O2 deposition as a discharge product, because it leads to deteriorate the air-electrode and electrolyte, and then lowers the energy efficiency of LABs. Therefore, many researchers have researched on air-electrode catalysts, e.g., noble metal nanoparticles (Pt, Au) and oxides (Co3O4, RuO2) loaded on nano-carbon materials. However, such kinds of solid electro-catalysts have not exhibited sufficient performance yet because the Li2O2 product generates and covers on the catalysts. To prevent this problem, we have synthesized a new composite catalyst composed of MnO2 nanosheets (Mn-NS) and Ketjen black (KB), i.e. a KB-comp. Mn-NS catalyst, and investigated the electro-catalytic activity against Li2O2 deposition/ decomposition reaction using a LAB cell, together with the durability [1]. In this study, we focused on the effect of interfacial cations for the MnO2 nanosheets on the ORR/OER reactions by comparing the performance of Li+, Na+ and K+-form KB-comp. MnNS catalysts. Mn-NS colloid solution was synthesized by a one-step solution method [1, 2], and then a KB ethanol dispersion was added to form a mixed solution. Afterward, 0.10 M LiCl, NaCl and KCl aqueous solutions were gradually dropped into the mixture to form the KB-comp. MnNS catalysts with each cation (Li+, Na+ and K+). Characterization of the obtained catalysts was carried out by a XRD, TG-DTA, BET measurement and TEM observation. The electro-catalytic activities were evaluated by using a LAB cell using the catalysts, Li metal NE and 0.20 M LiN(SO2CF3)2/diglyme (G2) electrolyte, and tested in constant current (CC) mode at 0.20 mA cm-2 at 30oC. Cyclic voltammetry (CV) and AC impedance measurements were also conducted to elcidate the reaction mechanism for ORR and OER at the PE. From XRD analysis, crystal phases of the synthesized KB-comp. Mn-NS catalysts clearly correspond to the birnessite-type MnO2 layer (JCPDS No. 80-1098) with different distance between the nanosheets, implying the existence of each cation. Fig. 1 shows charge/discharge curves and cycleability of the capacities for the LAB cells using the KB-comp. Mn-NS catalysts, respectively. As a result, the Li+ form clearly exhibited the lowest overpotential during both discharge and charge processes and the best cycle performance among them. This indicates the Li+ ions around the MnO2 nanosheets enabled to promote a Li2O2 generation near them, and keep good contact between the catalytic sites and Li2O2 product. From the CV measurement (Fig. 2), the order of magnitude for both ORR and OER currents was Li+ > Na+ > K+. This was equal to the trend of cell performance. The AC impedance suggested that the Li-form KB-comp. Mn-NS catalyst enhanced the Li2O2 decomposition reaction during charge process as compared with the other cation-form ones. The enhancement mechanism in more detail will be reported in the meeting. This work was supported by JST “A Tenure-track Program” and JSPS “KAKENHI” (25870899), Japan. [1] M. Saito et al., Electrochim. Acta, 252, 192 (2017). [2] K. Kai et al., J. Am. Chem. Soc., 130, 15938 (2008). Figure 1

In recent years, non-aqueous type rechargeable Li-air (O2) batteries (LABs) have attracted much attention as large-scale energy storage devices for electric vehicles because of the high energy density over 5 times larger than that of the conventional Li-ion batteries (LIBs) [1]. However, there are some problems to be solved for the practical use such as smooth deposition/decomposition reaction of Li2O2 at air electrode and suppression of Li dendrite growth at Li metal negative electrode (NE). To address these problems, we selected 1.0 M LiNO3/tetraglyme(G4) electrolyte and added acetonitrile (AN) or dimethyl sulfoxide (DMSO) with both high dielectric constant ε and low viscosity η to enhance the Li salt dissociation and lowering the electrolyte viscosity. Namely, the NO3 - anion was reported to work as a mediator to decompose the Li2O2 product at air electrode [2] and to form Li2O layer on the surface of Li metal NE to suppress the Li dendrite growth [3,4] and extra electrolyte decomposition. In this study, we investigated the bifunctional effects of NO3 - anion and its enhancement by using the dual solvent system for electrolyte. As a reference, 1.0 M LiOTf/G4 was also examined in the same way. Figure 1 shows the η values for 1.0 M LiNO3/G4+X and 1.0 M LiOTf/G4+X (X= DMSO, AN) electrolytes. The both η values decreased with an increase in the content of mixed solvents especially for AN because of one tenth lower η (0.37 mPas) than G4. As a result, the σ value was drastically improved (Fig. 2). For the addition of DMSO, the decrease in the η was not so big. However, the σ value effectively increased as well as those for the AN. This indicates that the relative high ε (47) of DMSO enhanced to dissociate the Li salts and improved the σ value by increasing the number of carrier ions. In fact, the effect was confirmed by Raman spectra and Walden plots. Figure 3 shows the discharge/charge curves for the LAB cells using the dual solvent electrolytes at the applied current of 0.20 mA cm-2. By mixing DMSO, the overpotential was drastically reduced especially during discharge process. This effect was enhanced at a higher rate operation. In addition, Li deposition/dissolution tests using Li foil | Cu mesh cell also exhibited improved performances for the 1.0 M LiNO3/G4 electrolytes. The effects for the dual solvent system will be reported in more detail at the meeting. This study was supported by JST Project “ALCA-SPRING”, Japan. [1] P. G. Bruce et al., Nature materials, 11, 19 (2012). [2] D. Sharon, et al., ACS App. Mater. & Int., 7, 16590 (2015). [3] J. Uddin, et al., J. Phy. Chem. Lett., 4, 3760 (2013). [4] M. Saito et al., J. Electrochem. Soc., 164(12), A2872 (2017). Figure 1

Lithium (Li)-air (O2) batteries (LABs) have been expected to be used for electric vehicles because of the high theoretical energy density of ca. 3500 Wh kg-1. However, there are some problems in the practical use, i.e. Li dendrite growth at Li metal negative electrode (NE), high overpotential during discharge process, etc. Recently, to solve these problems, LiNO3/ tetraglyme (G4)-based electrolytes were studied because of the bi-functional effects of surface oxidation for Li metal NE [1] and Li2O2 decomposition mediator at air electrode [2] by the NO3 -. However, the LiNO3 salt is quite low dissociation degree in G4 solvent, which causes to low ionic conductivity. In this study, we mixed dimethyl sulfoxide (DMSO) or acetonitrile (AN) with high dielectric constant e and low viscosity h as binary solvent to 1.0 M LiNO3/G4 electrolyte and investigated the effects on the electrolyte properties and LAB cell performance. G4 (< 30 ppm H2O) and AN or DMSO were mixed at the volume ratio of 9:1, 7:3 and 5:5, and 1.0 M LiNO3/G4+X (X =DMSO, AN) was prepared by dissolution of LiNO3 as supporting salt in an Ar filled dry box. The viscosity η and ionic conductivity σ were measured. The self-diffusion coefficient D of ions and solvents were evaluated by a PGSE-NMR [3]. The dissociation degree of Li salt was evaluated by using a Raman spectroscopy and Walden plot. Discharge/charge properties were also tested by using LAB cells using the electrolytes to discuss the effect of dual solvent. The η values of 1.0 M LiNO3/G4+X (DMSO, AN) decreased with increase in the content of dual solvents especially for AN because of the lower viscosity (0.37 mPa s) than that of DMSO (2.0 mPa s). However, DMSO-mixed electrolytes exhibited a similar increase in the σ value as well as AN mixed ones. Namely, the Li salt dissociation was assumed to be enhanced by the high e value (47) of DMSO. Fig. 1 shows the Raman spectra for the dual solvent electrolytes. In fact, the peak corresponding [Li+-(G4)] complex, i.e. Li+ solvation structure of G4, at 870 cm-1 disappeared for the DMSO-mixed electrolyte, indicating the strong interaction with DMSO solvent molecules compared with G4 ones. Walden plots exhibited and supported the enhancement of Li salt dissociation. Therefore, the DMSO and AN mixing as binary solvent was effective to increase in the number and mobility of carrier ions, respectively. Fig. 2 shows the discharge/charge properties of LAB cells using the 1.0 M LiNO3/G4+X (X = DMSO, AN, vol. ratio = 5:5). The both dual electrolytes successfully reduced the overpotential during discharge process. Especially for the DMSO-mixing a significant effect was achieved owing to the good supplying rate of NO3 - as mediator to the air-electrode. The effects for the dual solvent system in more detail, e.g. stability of the DMSO and AN in the dual electrolytes will be also reported at the meeting. This study was supported by JST Project “ALCA-SPRING”, Japan. [1] J. Uddin et al., J. Phys. Chem., 4, 3760 (2013). [2] D. Sharon et al., ACS Appl. Mater. Interfaces, 7, 16590, (2015). [3] M. Saito et al., RSC Adv., 7, 49031-49040 ,(2017). Figure 1

Li Deposition/dissolution Behavior at the Interphase between Li Metal Anode and Glyme-based Electrolytes for Li-air Batteries Taichi Fujinami a, Shinya Yamada a, Yusuke Tachikawa a, Kimihiko Ito b, Yoshimi Kubo b, and Morihiro Saito a a Department of Applied Chemistry , Tokyo University of Agriculture & Technology, 2-24-16 Koganei, Tokyo 184-8588, Japan b National Institute for Materials Science (NIMS) , 1-1 Namiki, Tukuba, Ibaraki 305-044, Japan E- mail: mosaito @ cc . tuat.ac.jp In recent years, rechargeable Li-air batteries (LABs) have attracted much attention as large scale energy storages for electric vehicles and stationary energy storage systems because of the high energy density over 5 times larger than that of the conventional Li-ion batteries (LIBs) [1]. However, there are some problems to be solved for the practical use such as the choking of air electrode by the generated L2O2 at discharging and the high overpotential for the electrochemical oxidation of Li2O2 at charging. Especially for the anode side, Li dendrite growth during discharge/charge cycles should be also suppressed to prevent the short circuit of LAB cells. In this study, we investigated the relationship between ion transport behavior and Li deposition/dissolution reaction at Li metal anode from two aspects; i.e. i) Li+supplying rate to the Li metal anode and ii) charge transfer rate through the surface layer containing SEI films at Li metal anode to understand the Li deposition/dissolution behavior. For the preparation of electrolytes for LABs, LiSO3CF3 (LiOTf, 99.0%, KISHIDA) and LiN(SO2CF3)2 (LiTFSI, 99.9%, KISHIDA) and LiN(SO2F3)2 (LiFSI, 99.0%, KISHIDA) were used as supporting salts, and dissolved in tetraglyme (G4, Japan Advanced Chemicals, H2O content: < 30 ppm) as a solvent in an Ar-filled dry box to prepare the 1.0 M glyme-based electrolytes. For the former aspect i), self-diffusion coefficient of Li+ (D Li +) in the electrolytes was measured by PGSE-NMR [2] along with the viscosity h, density r and ionic conductivity s. The degree of apparent dissociation a app of Li salts was also estimated from the D and s by using Nernst-Einstein equation [2]. The later aspect ii) was examined with [Li | Li] symmetric and [Li | KB-loaded (1 mg cm-2) carbon paper] LAB cells containing the same electrolytes. Fig. 1 shows the polarization curves of [Li|Li] symmetric cells with the three kinds of electrolytes at 30oC. The magnitude of overpotential h for the Li deposition/dissolution reaction was LiOTf >> LiTFSI > LiFSI. The order of h was in good agreement with that of ionic conductivity for the electrolytes. For the LiFSI-based electrolytes, the h value rather became smaller than those at the initial cycling. In fact, the SEM images of the surface of Li metal after the test exhibited many Li dendrites especially for the LiTFSI and LiFSI-based electrolytes (Fig. 2). Namely, Li+ was more easily reacted with Li metal under the faster Li supplying condition. In contrast, no Li dendrite was observed on the Li metal anode for the LiTfO-based one. Therefore, the LiTFSI and LiFSI electrolytes are considered to have an advantage for Li+ supplying process. On the other hand, the LAB test cell showed a different Li deposition behavior against the results for the [Li|Li] symmetric cell. The surface of Li metal anode was basically covered with Li oxides and oxygen-rich polymeric film layer. Only for the LiTFSI electrolyte Li dendrite-like compound was observed. Also, the LAB test cell using LiOTf one was successfully operated as well as the other electrolytes. These different behaviors were attributed to the existence of O2 gas, which is easily reacted with Li metal anode. The detail of the effects of SEI film by the different electrolytes and the oxidation by supplied O2gas into the LAB test cells will be reported in the meeting. This study was supported by JST “Next Generation Batteries Area in Advanced Low Carbon Technology Research and Development Program (ALCA)” from MEXT, Japan. [1] P. G. Bruce, S. A. Freunber, L. J. Hardwick, J-M. Tarascon, Nature materials, 11, 19 (2012). [2] M. Saito, S. Yamada, T. Fujinami, S. Kosaka, Y. Tachikawa, K. Ito, Y. Kubo, ECS Trans., 75(22), 53 (2017). Figure 1