Effects of Li Salt Anions and O₂ Gas on Li Dissolution/Deposition Behavior at Li Metal Negative Electrode for Non-Aqueous Li-Air Batteries

To prepare 1.0 M glyme-based electrolytes for LABs, LiOTf (LiSO₃CF₃, 99.0%, Kishida), LiTFSI (LiN(SO₂CF₃)₂, 99.9%, Kishida), and LiFSI (LiN(SO₂F)₂, 99.0%, Kishida) were used as supporting salts, and dissolved in G4 (Japan Advanced Chemicals, H₂O content: < 30 ppm) as a solvent in an Ar-filled dry box (GBJF100E805, Glovebox Japan Inc.).

The σ values of the electrolytes were measured with a conductivity meter (S230 SevenCompact, Mettler Toledo).

The D values of Li⁺ (⁷Li), the anion (¹⁹F), and solvent G4 (¹H) in the electrolytes were measured by PGSE–NMR with a JEOL tunable pulsed-field gradient (PFG) probe (¹H resonance: 400 MHz).^19–21 Each sample was prepared in a NMR microtube (BMS-005J, Shigemi, Tokyo) to a height of less than 5 mm to prevent convection effects. The PFG was calibrated with H₂O (¹H resonance). Measurements were performed by setting the same PFG strength for each nuclear species with a proper irradiation time δ. The accuracy of the D values was confirmed by obtaining the same values at different δ. Measurements of η and ρ were performed with a falling ball-type viscometer (Lovis2000ME, Anton Paar). All the electrolyte properties were measured at 30°C for the following electrochemical cell tests.

The Li | Li symmetric cell was constructed with Li metal foil (thickness: 0.5 mm, Honjo Metal Co., Ltd.) used for both electrodes, a separator (Celgard 2400), and the G4-based electrolytes in an argon-filled glove box (Miwa, MDB-1BK-NT1). The polarization curves of the Li | Li symmetric cells were measured to evaluate the overpotentials for Li dissolution/deposition reactions at both Li metal electrodes and the changes occurring during cycling by using a VSP potentiostat (BioLogic Science Instruments). The applied current density and maximum discharge/charge capacity were 0.20 mA cm⁻² and 0.50 mAh cm⁻², respectively, in the cut off voltage range of −2.0 to 2.0 V. EIS analysis was also conducted during the Li dissolution/decomposition tests in Li | Li symmetric cells. The applied voltage and measured frequency range were 5 mV and from 100 kHz to 0.10 Hz, respectively.

For fabrication of the LAB (Li | O₂) cell, Ketjen Black (KB, Lion Corp. Ltd., EC600JD)-loaded carbon paper (TGP-H-060, Chemix. Co. Ltd.) was used as an air electrode, i.e., the positive electrode (PE). The loading weight of KB was ca. 1 mg cm⁻². The same separator and Li metal NE as that used in the Li | Li symmetric cell were used in this cell. The same applied current density (0.20 mA cm⁻²) and maximum discharge/charge capacity (0.50 mAh cm⁻²) were applied to the LAB cell in the cut off voltage range of 2.0 to 4.5 V.

To investigate the effects of O₂ gas on the Li deposition/dissolution behavior, Li | Li symmetric cells containing O₂ gas were fabricated based on the LAB-type cell with an inlet for O₂ gas. The polarization curves of these cells were measured.

For all the electrochemical tests, the cycle number was set to 15 cycles and the operating temperature was 30°C.

After 15 cycles of polarization and discharge/charge tests, the Li metal electrodes were removed from the Li | Li symmetric and LAB cells and rinsed with fresh G4 solvent without Li salts. After drying under vacuum, the Li metal electrodes were set in transfer vessels and the surface morphology and chemical composition of the SEI films were evaluated by SEM-energy dispersive X-ray spectroscopy (SEM-EDS, JSM-7800F, JEOL) and X-ray photoelectron microscopy (XPS, VersaProbe II, ULVAC-PHI), respectively.

As mentioned above, the Li dissolution/deposition behavior at the Li metal NE was influenced by the Li⁺ supply rate (Li⁺ diffusion rate and number per specific volume) and the surface conditions (chemical composition of the SEI film) of the Li metal NEs. Here, we first focus on the former factor of the Li⁺ supply rate to the surface of the Li metal NE.

Figure 2 shows the D values of the Li⁺, anion, and G4 solvent for 1.0 M of G4-based electrolytes containing LiOTf, LiTFSI, and LiFSI. For the LiTFSI and LiFSI electrolytes, the D values of Li⁺ and the G4 solvent were smaller than those of LiOTf. This result indicates that a larger amount of G4 strongly solvated Li⁺ leading to dissociation of the Li salts, and an increased Li⁺ diffusion radius. This finding was in good agreement with the difference between the D values of the anion and Li⁺. For the LiOTf electrolyte, the D values of TfO⁻ and Li⁺ were almost the same as each other, implying the existence of a large number of ionic pairs consisting of Li⁺ and TfO⁻. On the contrary, the D value of TFSI⁻ and FSI⁻ was clearly higher than that of Li⁺ owing to the high dissociation of the Li salts.

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In fact, the LiTFSI and LiFSI electrolytes exhibited higher viscosity η and density ρ owing to the stronger interaction between the G4 solvent and Li⁺ (Table I). As a result, higher σ values of 3.02 and 3.33 mS cm⁻¹ were obtained for the LiTFSI and LiFSI electrolytes, respectively. To confirm the degree of dissociation of the Li salts, the molar ionic conductivity Λ_NMR and apparent dissociation degree α_app were calculated from the Nernst–Einstein equation as follows:^19–22

where N is the number of isolated ions per specific volume, e is elementary charge, k is Boltzman's constant, T is temperature (K), Λ_imp is molar ionic conductivity calculated from σ and ρ values for the electrolytes by AC impedance method, respectively. The obtained Λ_NMR and α_app values are summarized in Table I. The LiTFSI and LiFSI electrolytes exhibited higher α_app values (0.52 and 0.53) than that of LiOTf (0.14).

Therefore, the LiTFSI and LiFSI electrolytes have advantages in terms of fast supply of Li⁺ to the Li Metal NE, owing in particular to the high dissociation degree of these Li salts.

Figure 3 shows the polarization curves measured from Li | Li symmetric cells containing the different G4-based electrolytes. For the LiOTf electrolyte, the overpotential for the Li deposition/dissolution reaction gradually increased upon cycling. Conversely, the LiTFSI and LiFSI electrolytes exhibited relatively stable overpotentials; for the LiFSI electrolyte the value gradually decreased upon cycling. Thus, the LiFSI electrolyte, which had the higher supply rate σ owing to its higher value of α_app, exhibited a notable advantage in terms of the Li deposition/dissolution reaction.

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The Li charge transfer resistance R_ct for the Li dissolution/deposition reaction was evaluated by EIS analysis of the Li | Li symmetric cells together with the SEI film resistance R_SEI. The obtained data were analyzed with the equivalent circuit, shown in Fig. 4. The Nyquist plots are shown in Figs. 5a–5c and the estimated cell resistance values are summarized in Fig. 5d. The R_ct for the cell with the LiFSI electrolyte was smaller than that for the cell with LiTFSI, and decreased with increasing cycle number. This result was in good agreement with the overpotential results. In addition, the LiFSI electrolyte exhibited a lower SEI film resistance R_SEI (< 5 Ω cm²) compared with that of LiTFSI (∼10 Ω cm²) during cycling. Therefore, LiFSI also stabilizes the Li metal NE at the SEI film. Conversely, the LiOTf electrolyte cell did not work regardless of the similar R_SEI values before the Li | Li symmetric cell test. This result is likely caused by the formation of an unstable SEI film derived from the LiOTf salt and/or a slow Li⁺ supply rate. However, the main cause is considered to be the low dissociation degree of LiOTf. In general, stronger interaction between Li⁺ and anion increases the reduction potential of anion.^23,24 Li⁺ adjacent to anion shares some negative charge and makes either the reduction of anion easier, resulting in reductive decomposition.

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Figure 6 shows SEM images of the Li metal electrodes after the Li | Li symmetric cell tests, i.e. after 15 cycling of polarization, with different G4 electrolytes. For the LiOTf electrolyte, almost no Li deposition and no Li dendrite formation were observed on the Li metal electrode. Conversely, many Li dendrites were found for the LiTFSI and LiFSI electrolytes. This result indicates that the core of Li dendrites was easily generated and grew in the G4 electrolytes having higher σ and α_app. Furthermore, as these values increased the average diameter of dendrites decreased. A decrease of R_ct was thought to enhance the formation of the Li dendrite core and increase the surface area of the Li metal electrode through the growth of fine Li dendrites. This finding concurs with the reduction of R_ct for the LiFSI electrolyte upon cycling.

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Figures 7a–7c and 7d–7f show the discharge/charge curves and cycling performance for LAB cells with different G4 electrolytes. Although the 1^st discharge curves showed quite similar behavior regardless of the type of electrolyte, the overpotential at the 1^st charge cycle were relatively low compared with those at the following cycles, especially for the electrolyte containing LiTFSI. After the 1^st cycle, the discharge curves for the LiOTf and LiFSI electrolyte cells deteriorated with further cycling. The discharge/charge curves are influenced by the PE reaction at the air electrode; however, we confirmed that the selection of the Li salt also affects the cell performance. Namely, the LiOTf electrolyte exhibited relatively good LAB operation regardless of the low σ value. For the LiFSI electrolyte, oxidative decomposition was confirmed at the end of 1^st charging over 4.05 V, which rose the overpotential and decreased the charge capacities at the following cycles. This degradation is clearly shown in the cycle performance results of the LAB cells (Fig. 7f). Therefore, O₂ gas introduced from the air electrode markedly influenced the LAB electrode reactions including the Li dissolution/deposition reaction at the Li metal NEs. The LiOTf electrolyte was particularly strongly affected.

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The LiTFSI and LiOTf electrolyte cells maintained lower cell voltages, particularly at the end of charging, and were relatively stable during cycling (Fig. 8). Conversely, for the LiFSI electrolyte the higher cell voltage due to the oxidative decomposition of FSI⁻ anion at charging led to deteriorate the cell performance, particularly the cycling performance, despite this electrolyte having the highest σ value.

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To confirm and investigate the effects of O₂ gas introduction into the LAB cells, SEM observations and the EDS analysis were conducted for the Li metal NEs after 15^th charging (Fig. 9). For all electrolytes, the surface of the Li metal NEs was clearly oxidized by O₂ and chemical compounds including carbon were also deposited. The Li dendrites almost disappeared, even for the LiFSI electrolyte. Only a few ribbon-like Li dendrites were observed for the LiTFSI electrolyte.

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Figure 10 shows XPS data for Li metal NEs after 15 discharge/charge cycles. For all electrolytes, a relatively large amount of carbon species, including Li₂CO₃, HCO₂Li, CH₃CO₂Li and plolyethers/ethers were detected in the SEI films on the Li metal NE. These carbon species were derived from electrolyte decomposition.^25,26

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The LiTFSI electrolyte produced a relative large amount of LiF. A LiF-rich SEI film covered the surface of the Li metal NE, which suppressed surface-oxidation and reduced the overpotential during discharge/charge cycles. This observation is supported by the fact that ribbon-like Li dendrites were only observed for the LiTFSI electrolyte. Although a LiF peak at 685.3 eV^27,28 for F 1s was also confirmed for the LiOTf electrolyte, the strength of the XPS peak was low, and a C-F peak derived from CF₃ (687.7 eV) became larger, indicating a different type of reductive decomposition of the anions. In the case of the LiFSI salt, an LiF peak was not detected because of the fewer F atoms present in the FSI⁻ anion than in the anions of other electrolytes. A relative small LiF peak confirmed for the Li | Li symmetric cell disappeared, indicating the covering by a large amount of Li₂O and Li₂CO₃ components as shown in SEM images. All XPS data implied that the surface of the Li metal NE was oxidized, and Li₂CO₃ and polymer-like SEI components formed a coating as a by-product of electrolyte decomposition in the operation of the LAB (Li | O₂) cell. These facts support the similar LAB cell performance compared with data of Li | Li symmetric cells.

The cell performance deteriorates owing to contributions from the Li metal NE and the air PE, although it is not clear which of these factors affect the deterioration the most. We used Li | Li symmetric cells, containing O₂ to investigate the effects of O₂ gas on the Li metal NE performance and the differences among the anions.

As mentioned above, the Li dissolution/deposition reaction is strongly influenced by the atmosphere in the cell (i.e., Ar or O₂). These effects were clearly confirmed for the LiOTf and LiFSI electrolytes. Although the Li dissolution/deposition reaction did not occur in the Li | Li symmetric cell with the LiOTf electrolyte, in the LAB cell this electrolyte exhibited relatively good performance, which was notably different from the behavior of the LiTFSI electrolyte. However, the LiFSI electrolyte deteriorated in the LAB cell. From SEM, EDS and XPS analyses, the surface of the Li metal NEs in the LAB cells was oxidized by O₂ introduced from the air electrode, and this influenced the cell performance. To confirm and elucidate the effects of O₂ gas on the Li metal NEs, the Li dissolution/deposition reaction was also investigated in Li | Li symmetric cells containing O₂ gas.

Figure 11 shows the polarization curves for O₂ gas containing Li | Li symmetric cells. For all electrolytes similar Li dissolution/deposition behaviors were confirmed and the overpotentials were reduced by the introduction of O₂ gas. The LiOTf electrolyte cell also operated stably and the overpotential decreased with cycling. This result indicates that proper surface-oxidation effectively stabilized the Li metal NEs and promoted the smooth Li dissolution/deposition reaction. Recently, similar effects of surface-oxidation of Li metal NE have been reported for LAB and LMB systems using N,N-dimethylacetamide (DMA)- and glyme-based electrolytes containing a LiNO₃ salt, respectively, where NO₃⁻ anions oxidize and stabilize the surface of the Li metal NE by formation of Li₂O passivating layer, and enable to improve the cycle performances.^29–31 Namely, the Li oxide layer containing Li₂CO₃ also stabilizes the surface of the Li metal NE to accelerate the Li dissolution/deposition and lower the overpotential of the air-electrode. As a result, the LAB cell containing LiOTf electrolyte also exhibited good cell operation. The Li | Li symmetric cell based on the LiFSI electrolyte exhibited a similar stable Li dissolution/deposition reaction. This finding indicates that the surface of the Li metal NE was oxidized by O₂ and the cell performance depends on the Li oxide layer. Furthermore, the FSI⁻ anion was stable at the Li metal NE in the LAB cell and was likely oxidized at the air electrode.

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Figure 12 shows the results of EIS analysis for the Li | Li symmetric cells containing O₂ gas. The Nyquist plots were analyzed with the equivalent circuit shown in Fig. 4, and compared with the results shown in Fig. 5. For all the electrolytes, both R_ct and R_SEI values were successfully reduced by the introduction of O₂ gas and the R_ct values clearly decreased with increased cycling. The lowest cell resistances were confirmed for the LiTFSI electrolyte. These results explain the LAB cells exhibiting similar cell performance, including those based on the LiOTf electrolyte. The best cell performance came from the cell containing the LiTFSI electrolyte.

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