The galvanostatic cyclability of the LCO electrodes was first evaluated using two representative electrolyte salts, namely, 0.5 m Li2SO4 and 1 m LiTFSI, which both provided 1 m Li+. The LCO electrode with 0.5 m Li2SO4 exhibited a gradual increase in capacity at 0.5C over 30 cycles as the wetting of the electrode increased,5 and the coulombic efficiency (CE) approached 97% (Figure 1a and 1c). By comparison, 1 m LiTFSI reduced the capacity from 110 to 30 mAh g–1 (Figure 1b) and afforded an inferior CE of 94% (Figure 1c). Rate-capability tests with 1 m LiTFSI displayed an ill-defined capacity at various current rates (Figure 1d). In contrast, complete recovery of the capacity was found for the LCO electrode with 0.5 m Li2SO4 at the terminating 0.2C after 10C operation. These results reveal that the cyclability and rate capability of the LCO electrode are both anion-dependent.
To determine the origin of the capacity fading, the short-range order of the LCO structure was analyzed using soft X-ray absorption fine structure (XAFS) spectroscopy.7 In surface-sensitive partial electron yield (PEY) mode (< 10 nm depth), the Co3+ 3d-O 2p hybridization peak in the oxygen (O) K edge spectra (529.8 eV) was attenuated for both LCO electrodes after 30 cycles. In addition, a new signal emerged at 531.4 eV, attributed to partial oxidation of the lattice oxide of LCO; this signal is denoted as O1 (Figure 1e‒f).25-27 The O1 signal was particularly intense for LCO with 1 m LiTFSI, indicating more extensive surface oxidation than in the case with 0.5 m Li2SO4. In bulk-sensitive partial fluorescence yield (PFY) mode (> 100 nm depth), the Co4+ 3d (t2g5eg0)-O 2p hybridization band was observed at 528 eV;25-27 this band is denoted as O2 (Figure 1f). The signal was more pronounced with 1 m LiTFSI, consistent with a subtle shift of the main Co3+ signals to higher energy than others in the Co L2 and L3 edge spectra (Figure S1). Because the O2 signal should disappear after Li+ intercalation during the discharging process, its existence after 30 cycles indicates incomplete Li+ intercalation into the LCO electrode. The remaining Li vacancy and Co4+ cause local distortion of the octahedral CoO6 unit.7,27 Long-range structural transformation was also observed. Powder X-ray diffraction analysis of the LCO electrode after operation with 1 m LiTFSI displayed split peaks related to the 003 reflection (Figure S2). The new peak appearing at 18.7° indicates widening of the interlayer distance due to the deficiency of Li+.4,28 In contrast, a less profound O2 signal was observed with 0.5 m Li2SO4 in the XAFS profile, along with a single 003 reflection, indicating complete Li+ intercalation into LCO after the same number of cycles.
Such inadequate Li+ intercalation has been ascribed to the insertion of H+ into the LCO electrode in previous studies.17-20 H+ can be produced through the dissociation of water at the electrode surface in neutral or weakly alkaline solutions. Electrochemical impedance spectroscopy (EIS) can be used to distinguish Li+ intercalation and H+ insertion in the LCO electrode.19 Successive Nyquist plots were acquired at open-circuit potential (OCP, Figure S3), during the charging process (from 0.66 V to 0.8 V, red to yellow curves), and the discharging process (from 0.8 V to 0.63 V, light green to dark green curves) for three galvanostatic cycles (Figure 2). The emerging semicircle at the beginning of the 1st charge is attributed to the charge-transfer resistance of Li+ (Rct,Li+).19,20,29 The increased electronic conductivity of LCO decreased the semicircle size during the charging process, as a result of transformation from the semiconductor to the semi-metal by Li+ extraction.30 During the discharging process, the semicircle gradually increased again with 0.5 m Li2SO4. Enlargement of the semicircle at the end of discharge (0.63 V) reflects Li+ filling at the topmost surface of the LCO electrode (Figure 2a–b).31 This behavior was repeated for all three cycles, and Rct,Li+ was estimated as 27.8 Ω after 30 cycles (Figure S4a and Table S1). In comparison, the Nyquist plot of the LCO electrode with 1 m LiTFSI showed a larger semicircle corresponding to Rct,Li+ at the beginning of the 1st charge (Figure 2c–d). More importantly, a new semicircle was observed in the low-frequency region at the terminal stage of the discharge process. Although this new semicircle vanished at the beginning of the following charging process over three cycles, it eventually grew irreversibly and accounted for a larger portion of the resistance (~147 Ω) after 30 cycles (Figure S4b and Table S1). This semicircle, which is apparently distinguished from that associated with Rct,Li+, is attributed to the resistance from H+ insertion (Rct,H+).19,20 The result of H+ insertion appeared at the end of the discharging process, that is, at the terminal stage of Li+ intercalation. This behavior was not observed with 0.5 m Li2SO4.
The intriguing question is which factor related to the electrolyte salt leads to the distinct degree of H+ insertion. Three possible factors were postulated: the pH, surface protective layer, and anion effects. The 1 m LiTFSI solution had a pH of 8, which is lower than that (pH 9.6) of the 0.5 m Li2SO4 solution (Figure S5). When the pH of the LiTFSI solution was adjusted to ~10 by adding LiOH, the capacity retention of LCO was still limited to ~34%, which is far inferior to the value of ~100% achieved with 0.5 m Li2SO4. The Nyquist plot for the LCO electrode with 1 m TFSI– (pH 10) revealed suppression of H+ insertion in the 1st cycle (Figure S6). Nonetheless, the Rct,H+ signal appeared at the end of the 3rd cycle (Figure S6), and both semicircles designated to Rct,Li+ and Rct,H+ grew over the course of 10 cycles (Figure S7), which is in line with the capacity fading.32 In stark contrast, the Rct,H+ signal did not appear for LCO with 0.5 m Li2SO4 at the 10th cycle, elucidating negligible capacity decay. Therefore, the pH effect was not critical under mildly alkaline conditions.
The formation of a solid-state protective layer on the LCO electrode was then considered. In non-aqueous media, decomposition of the electrolyte solution is accompanied by the construction of an interphase layer that serves to mitigate undesired surface reactions.33,34 However, both SO42– and TFSI– are inert in the given potential range. The X-ray photoelectron spectra (XPS) did not reveal any SO42–, TFSI–, or fragmental species of the electrolytes after 30 cycles (Figure S8). In the O 1s binding energy region, the signal of lattice Co–O at 529.7 eV was less intense, while the peak at 530.8 eV indicating both the oxide defects (O-2+δ) and the hydroxide became significant. This surface deformation was inevitably and chemically induced by the contact of the electrode with water, especially after Li+ extraction (Figure S9).4,6,16 Therefore, the above results indicate the absence of a solid-state protective layer over the LCO electrode. After 30 cycles, a thicker amorphous surface layer was formed on the LCO with 1 m LiTFSI, as observed in the transmission electron microscope images (Figure S10), demonstrating more acute surface degradation than that with 0.5 m Li2SO4.
Therefore, we focused on the last factor, the effect of the anion type. SO42– and TFSI– are well-known as strong kosmotropic and chaotropic ions, respectively, according to the Hofmeister series.35-38 Kosmotropic anions undergo intimate interactions with water and form a rigid solvation structure.39,40, In contrast, chaotropic anions, such as NO3–, ClO4–, and TFSI–, undergo weak interactions with water that induce disorder of the surrounding water structure. By investigating NO3– and ClO4– having less chaotropic character than TFSI–, the effect of the anions on the cell cyclability was elucidated (Figure 3a‒b), where the molal concentration of all electrolyte salts was similar to the molar concentration (M, mol L‒1, Table S2). The LCO electrode with 0.5 m Li2SO4 showed ~73% capacity retention after 100 cycles. In comparison, 1 m LiNO3 and 1 m LiClO4 exhibited lower retentions of 33% and 13%, respectively. With the use of 1 m LiTFSI, the capacity decreased immediately after 30 cycles, and the CE of ~94% was lower than that (>99%) achieved with all other electrolytes. Consequently, the cell performance was very distinct and followed the order: SO42– > NO3– > ClO4– > TFSI–, which corresponds to the reinforcing kosmotropic nature of the anion. Consistently, it appears to larger Rct,Li+ of LCO with NO3– than that with SO42– (Figure S11a). The presence of ClO4– promotes the formation of Rct,H+ (Figure S11b). All these evidences support the decisive correlation between the surface reaction and cyclability.
The anion interaction perturbs the O–H vibrations of the water molecules, which are sensitive to the strength of the hydrogen bond. Infrared (IR) spectroscopy allows distinction of the vibrational bands of the O–H stretching at ~3200 cm–1 for the strongest and ordered water structure versus that at ~3400 cm–1 for the weaker hydrogen bond.41 When Li2SO4 was added to the water in concentrations from 0.1 to 3 m, the O–H band at ~3200 cm‒1 gradually intensified as the hydrogen-bond became stronger (Figure 3c). This result may be attributed to the Li+-water interaction, because the hydrogen bond of the water network is negligibly perturbed by SO42–.42 In other words, the SO42––water interaction is as strong as the water network, thus providing a featureless signal. In sharp contrast, the IR band of LiNO3, LiClO4, and LiTFSI at ~3200 cm‒1 became less intense in the order: NO3– < ClO4– < TFSI–.41,43-45 This trend underpins the disordering of the solvation structure with increasing chaotropic nature of the anion. Additionally, the signals at 3500 ~ 3650 cm‒1 are related to both the free water and the chaotropic anion-water interaction.42,46-48 More pronounced absorption bands appeared at higher electrolyte concentrations, supporting the latter phenomenon. Figure S12 shows different O–H vibrations depending on the anion type in the presence of 1 m Li+. Note that the range of influence of ions on the structure of water has been debated, from the first solvation shell to the mid/long-range hydrogen-bonding network.39,40,49 Nonetheless, it is generally agreed that the anions have profound impact on the closest water molecules, and the strength of their interactions in the first solvation shell can determine the macroscopic properties of the electrolyte solution.40
More important is the effect of the anions on water structuring at the electrode surface, in accounting for the electrode deformation and the corresponding cell performance. The water structure in the bulk solution can be altered at the interfacial region by the local concentration of ions, surface charge of the electrode, material property, and applied electric field.50,51 The staircase-potential EIS in the non-faradaic potential region was quantified to electric double layer (EDL) capacitance (Cdl) on the LCO electrode (Figure 4a and Figure S13). In addition, the surface charge density (s) was calculated using Eq. 1 (Figure 4b):
where U is the electrode potential, and UPZC is the potential at the point of zero charge (PZC) where Cdl approaches the minimum value. Figure 4a shows that the Cdl of 1 m LiTFSI is 1.5 ~ 3 times higher than the Cdl of the other electrolytes in the given potential range of 0.3 ~ –0.3 V vs. Ag/AgCl. It is attributed to large van der Waals volume, high polarizability, and mild hydrophilicity of TFSI–,52-54 affording the excellent adsorption for both LCO material and carbon additive part in the EDL region (Figure S14). However, a higher Cdl does not necessarily mean a higher concentration of ion present at the “interface”, due to the possible formation of ion pairs. As the ion pairs do not contribute to the Cdl associated with the net surface charge, it is suggested to the necessity of full elucidation on the interfacial structure.
The detailed atomic arrangement of the interfacial structure formed at the LCO electrode was explicated using our recently developed mean-field QM/MM simulation.55 To investigate the electrolyte response to the cathodic charging, two different surface charge densities of s = 0 and –11.5 mC cm‒2 were compared. First, the local ion concentration along the surface normal direction (which was chosen as the z-direction) revealed the formation of two layers within ca. 6 Å from the topmost atoms of the LCO electrode (Figure S15a–b). Additionally, the local water concentration also showed a layering tendency near the electrode surface, and then converged into the bulk value at z > 6 Å (Figure S15c). Thus, it is reasonably assumed that the region at z < 6 Å is an interfacial region, that is, a part of the EDL (nanometer scale) in direct contact with the solid electrode, at which the ion concentration and free water density were analyzed.
Figure 4c shows the number of anions located in the interfacial region when the aqueous electrolyte was interfaced with the LCO electrode. Notably, the local anion population near the electrode increased at s = 0 mC cm–2 as the ion became more kosmotropic; SO42– showed the largest population. A water adlayer was formed on the LCO surface (Figure S16), which can attract ions to the interfacial region via ion-dipole interactions. Because the ion-dipole interaction can be maximized for ions with a high charge density, kosmotropic ions such as SO42– can accumulate in large quantities near the electrode surface.
Surprisingly, such a trend still holds even when the LCO electrode is cathodically polarized (s = –11.5 mC cm–2), where the anions experience electrostatic repulsion from the electrode (compare triangles and squares in Figure 4c). The LCO surface attracts Li+, which is further stabilized by the water adlayer (Figure 4d). In parallel, Li+ electrostatically attracts anions to form ion pairs. Kosmotropic anions having high charge density are now preferentially associated with small-size Li+ ions (Figure 4e), according to the “hard and soft acid and base (HSAB)” concept. The present simulation shows that contact ion pairs (CIPs) are more favorably formed when the anions have more kosmotropic character (Figure S17).
Along with the large increase in the local ion concentration (including Li+) with kosmotropic anions, the increased ionic strength near the electrode surface can minimize the number of uncoordinated water molecules, termed ‘free water’. This demonstrates the inverse relationship between the density of free water and the kosmotropic propensity of the anions (Figure 4f). Considering that the electrolyte pH is in the mildly alkaline range where water is regarded as the main H+ source, the lower free water density induced by the kosmotropic anions conceivably helps to suppress H+ insertion and maintain the stability of the electrode.
When the potential rose above 0.2 V, that is, positive s, the Cdl increased substantially for all electrolytes (Figure 4a) because many anions approach the EDL through electrostatic interactions. At 0.3 V, the s values for Li2SO4 and LiTFSI became comparable and were slightly higher than those of LiNO3 and LiClO4 (Figure 4b). The increased anion concentration can enhance the stability of the electrode against water, regardless of the anion type. It is demonstrated by the shift of the cutoff potential from –0.2 V to 0.3 V; the capacity retention was significantly increased from <35% to 56% with 1 m LiNO3, 65% with 1 m LiClO4, and 21% with 1 m LiTFSI over 200 cycles (Figures S18–19 and Table S3). The cell with 0.5 m Li2SO4 still exhibited superior performance (67%) despite utilizing approximately half the concentration of the other anions. From the increasing Cdl at positive potential, we anticipate that the concentration of anions at the electrode surface would be similar for all electrolytes in the Faradaic potential region where the Li+ intercalation/extraction process occurs (0.5 ~ 0.8 V vs. Ag/AgCl). If Li+ intercalation competes with H+ insertion, a high Li+ concentration near the LCO electrode is beneficial. Therefore, the ion pair formed by SO42– is more likely to reduce the detrimental effect of water.
Based on the insights gained from the above studies, we attempted to improve the cell performance by increasing the concentration of SO42–. The electrochemical potential window in the anodic region was slightly expanded with 3 m Li2SO4 (~130 mV, Figure 5a), implying that more SO42– was localized on the electrode and suppressed the water reaction.8,56 The O–H vibration of 3 m Li2SO4 was analogous to that with 0.5 m Li2SO4 in the Raman and ATR-IR spectra (Figure 3c and Figure S20), while the symmetric SO42– stretching mode in the Raman spectrum was blue-shifted at higher concentrations due to the significant formation of ion pairs (Figure 5b).57,58 Consistently, the appearance of an IR-inactive stretching band and the red-shift of the anti-symmetric SO42– vibration evidenced the growing number of ion associations (Figure S21).44
Full-cells composed of Li9/7–xNb2/7Mo3/7O2 (LNMO) as the negative electrode, LCO as the positive electrode,59,60 and aqueous electrolyte solutions with different salt concentrations (0.5, 3 m Li2SO4, 1 and 3 m LiTFSI, and 9 m LiNO3) were assembled. With the low-concentration electrolyte, LNMO limited the cell capacity to 40 ~ 90 mAh g–1LCO (Figure S22). The 0.5 m Li2SO4 and 1 m LiTFSI-based full-cells both experienced rapid capacity fading over 50 cycles (Figure 5c) owing to the deformation of LNMO and the hydrogen evolution reaction in the salt-in-water system. In comparison, 3 m Li2SO4 delivered 74% capacity retention over 500 cycles, which is superior to that achieved with 3 m LiTFSI (38%) and 9 m LiNO3 (34%) (Figures 5c–d, Figure S23, and Table S4). The outstanding cyclability attained using 3 m Li2SO4 is attributed to the extensive SO42–-based liquid protective layer that harnesses the water molecules. Although the cells with 1 ~ 3 m LiTFSI exhibited higher initial capacity than those with the corresponding Li2SO4, prompt capacity fading was observed during cycling. In addition, the poor cyclability of 9 m LiNO3 confirmed the indispensable role of the kosmotropic anions. Although the cell stability with the salt-in-water electrolyte is still inferior to that of WiSE (92% capacity retention with 21 m LiTFSI, Figure S24), this finding provides promise for developing aqueous LIBs with far lower electrolyte concentrations by tailoring the anion species.
In summary, we demonstrated the imperative role of anions in the salt-in-water electrolyte of aqueous LIBs. In the presence of chaotropic anions, severe capacity fading appeared due to deformation of the LCO surface by water and H+ insertion. In contrast, sulfate, as a kosmotrope, mitigated these detrimental effects by establishing stronger sulfate-water and sulfate-Li+ ion interactions. Atomic-scale multiscale simulations illuminated the higher local concentration of sulfate at the interfacial region of the LCO electrode. The stiff water structures and kosmotropic anion-Li+ pairs generated a liquid-phase protective layer at the interface and alleviated the detrimental effect of free water molecules in deforming the electrode structure. This fundamental understanding sheds light on the mechanistics of the improved cell cyclability with 0.5 m Li2SO4, providing 0.20% capacity-fading rate per cycle for 200 cycles in half-cells and 0.06% fading rate per cycle with 3 m Li2SO4 for 500 cycles in full-cells. The insight into the interfacial structure paves the way for the design of stable, inexpensive, and safe aqueous LIBs by tailoring the anion species and interfacial environments.