Electrolyte for Aluminum Rechargeable Batteries Using Dimethyl Sulfone and Aluminum Chloride

Abstract

The electrochemical deposition and dissolution of Al with an electrolyte solution containing dimethyl sulfone and aluminum chloride was investigated at 30 °C to determine the reversibility of the process. Although aluminum deposition was observed in the electrolyte solution containing dimethyl sulfone and aluminum chloride, multiple oxidation peaks were observed in the dissolution reaction. The morphology of the deposited aluminum was observed using a scanning electron microscope, and it was found that the shape of the aluminum deposits varied greatly depending on the deposition potential. Electrochemical measurements under various conditions revealed that the oxidative decomposition of the electrolyte at the Al counter electrode during the aluminum deposition on the working electrode resulted in a large amount of aluminum oxide contamination in the deposited aluminum. To prevent such a reaction, we designed a cell in which the potential of the counter electrode did not exceed the electrochemical potential window and observed a very reversible aluminum deposition dissolution reaction. Finally, we successfully charged/discharged the Al negative electrode using a dimethyl sulfone-based electrolyte. These findings are important because the electrolytes with dimethyl sulfone and aluminum chloride are compatible as electrolytes for rechargeable aluminum batteries and because they highlight the precautions that need to be taken in their application.

1. Introduction

Electrochemical energy storage is essential for the efficient use of renewable energy to realize a sustainable society. In particular, the development of electric vehicles requires high performance rechargeable batteries. Lithium-ion rechargeable batteries (LIBs) have high energy density and long cycle life; therefore, LIBs are used as energy storage for various applications. Nevertheless, the use of rare metals such as lithium, cobalt, and nickel, which are employed as electrode and electrolyte materials, is expensive and unreliable in terms of sustainability.^1–3

Therefore, various post-LIB technologies have been studied in recent years. One of the post-LIBs is a rechargeable battery that uses a metal that becomes a multivalent cation as the negative electrode.^4–6 In particular, aluminum becomes a trivalent cation and is known to have four times larger capacity per volume (8042 mAh cm⁻³) than that of lithium (2060 mAh cm⁻³). In addition, aluminum has a higher melting point than lithium, is stable in the air, and is abundant on the earth.^7–9 However, compared to other multivalent cation rechargeable batteries such as magnesium and calcium, the major problem of the aluminum rechargeable battery is the limited number of electrolytes capable of reversibly depositing and dissolving aluminum metal, which is essential for the aluminum negative electrode of batteries. The well-known ionic liquid mixture of aluminum chloride (AlCl₃) and 1-ethyl-3methylimidazolium chloride is the most widely used electrolyte for rechargeable aluminum batteries, but it is very sensitive to moisture. In this regard, Abbott has successfully prepared a new deep eutectic solvent that is relatively insensitive to water by mixing AlCl₃ with acetamide or urea at room temperature.¹⁰ Dai et al. have also successfully constructed an aluminum rechargeable battery using the AlCl₃/urea deep eutectic solvent.¹¹ In such electrolytes, urea functions as a Lewis base to induce disproportionation of AlCl₃, producing electrochemically active Al³⁺ contained cations. Meanwhile, we have reported aluminum rechargeable batteries using a mixed electrolyte solution of AlCl₃, dipropylsulfone (DPSO₂), and toluene.^12–14 This electrolyte has the advantages of low cost and high safety because of its low concentration of AlCl₃, which corrosive and toxic. Furthermore, DPSO₂ has a melting point of 30 °C and can be mixed with toluene to maintain liquid at room temperature.¹⁵ Similar to DPSO₂, dimethylsulfone (DMSO₂) was studied as an electrolyte for aluminum plating by Legrand.^15–17

Legrand et al. analyzed the aluminum chemical species contained in the electrolyte with DMSO₂ using NMR and Raman spectroscopy, and concluded that the main hexacoordinated aluminum species was Al(DMSO₂)₃³⁺. Dokko et al. reported the effects of the solvent structure on the hopping conduction of lithium ions in highly concentrated LiBF₄/sulfone electrolyte.^18,19 The oxygen atoms bonding the S atom in sulfones are coordinated to lithium ions to realize lithium ions hopping conduction. These suggest that the sulfone functions as a Lewis base, causing a cleavage in AlCl₃ and lowering the melting point, as in the electrolyte mixture of AlCl₃ and urea reported by Abbott. Moreover, Okamoto et al. reported that electrolytes with AlCl₃/DMSO₂ mole ratio from 0.38 to 0.42 was liquid even at room temperature.^20–22 However, although it has been shown that it is a liquid at room temperature, there are no studies that have investigated the electrochemical properties of electrolytes using aluminum chloride and dimethyl sulfone at room temperature. Moreover, research on electrolytes using aluminum chloride and dimethyl sulfone has so far been almost entirely aimed at aluminum metal plating. Therefore, the cell configuration is significantly different from that of a rechargeable battery. For example, in a battery, the electrodes are in close contact with each other via a separator, whereas in plating, the distance between the electrodes is larger. In addition, the amount of electrolyte used in a battery is small, whereas in plating, a large amount is used.

In this study, we investigated for detailed analysis of the deposited aluminum in a liquid electrolyte with an AlCl₃/DMSO₂ ratio of 0.38 at room temperature under various electrochemical conditions that affect the deposited material. In order to explore the potential of aluminum chloride and DPSO₂ as an electrolyte for rechargeable batteries, we conducted charge/discharge measurements using an aluminum negative electrode.

2. Experimental

2.1 Preparation of DMSO₂/AlCl₃ electrolyte

DMSO₂/AlCl₃ electrolyte was prepared according to Ref. 15. 3.00 g (31.9 mmol) of DMSO₂ (Tokyo Kasei, 99.0 %) and 1.616 g (12.1 mmol) of AlCl₃ (99.9 %) were put in a glass cup and heated at 110 °C to melt. AlCl₃ was supplied by Nippon Light Metal Company, Ltd. The resultant DMSO₂/AlCl₃ electrolyte was liquid even at room temperature. All cell production and reagent preparation was carried out inside a glove box filled with argon gas. (UNICO LTD., YSD-800L, dew point was below −60 °C and oxygen concentration was not controlled.)

2.2 Electrochemical measurements

A home-made glass cell was fabricated using a molybdenum plate (Nilaco Corporation, 1 × 1 cm or 5 × 5 mm, 0.1 mm thick) as the working electrode, an Al plate (Nilaco Corporation, 1 × 1 cm, 0.1 mm thick, 99 % purity) as the counter and reference electrodes, glass fiber filter paper (ADVANTEC, GA-100) as the separator, and DMSO₂/AlCl₃ electrolyte as the electrolyte solution. The volume of the electrolyte was 0.75 ml, and the distance between the working and counter electrodes was controlled by the thickness of the separator, which was approximately 0.2 mm. The aluminum electrode was polished using # 400 emery paper in a glove box. The molybdenum electrode and current collector were soaked in dilute hydrochloric acid for several minutes before use. The cell was placed in a glass pot with electrical continuity to the outside and was closed to the outside air. The potentiostat (VPS, BioLogic) was used for the electrochemical measurements. VPS has a potential measurement probe for the counter electrode and realizes the working and counter potential measurement simultaneously. The scanning speed and temperature were standardized at 10 mV s⁻¹ and 60 °C for all CV measurements. The time for aluminum deposition using the constant potential method was adjusted to 10 hours. For charging and discharging measurements, HJ1001SD8 (MEIDEN HOKUTO CORPORATION) was used. Charging was controlled by charging time for 10 minutes, and discharge was controlled by the cut-off voltage at −0.5 V and discharging time for 10 minutes.

2.3 Characterization

After potentiostatic Al deposition, the Mo plate working electrode was taken out of the cell, washed with methanol, and vacuum-dried at 60 °C. Scanning electron microscope (SEM) images were observed using VE-9800 (KEYENCE). Elemental analysis was performed using an Energy-dispersive X-ray spectroscopy (EDS) analyzer (AMETEK) equipped with VE9800. X-ray diffraction (XRD) spectra were observed with XRD-6100 (Shimadzu, 50 kV, 30 mA, equipped with a CuKα source (λ = 0.1541 nm), and scanning rate 2° min⁻¹). X-ray photoelectron spectroscopy (XPS) was performed with ESCA-3400 (SHIMADZU). The tube current for the XPS measurement was 20 mA, the tube voltage was 10 kV, and the number of integration times was 8. The Ar⁺ etching time was 120 s for −0.3 V constant potential deposits and 240 s for −0.8 V constant potential deposits. The X-ray irradiation area during XPS measurement was a circle with a diameter of 6 mm. Before XPS measurement, the sample was washed once with ethanol, dried at room temperature under an argon atmosphere, and then transferred to XPS by passing through air.

3. Results and Discussions

Figure 1 shows the cyclic voltammograms (3rd cycle) of a cell with a Mo plate working electrode, an Al plate counter electrode, and an Al wire as the quasi-reference electrode. The lower potential limit was fixed at −0.5 V vs. Al/Al³⁺, and the upper potential limit was varied from 2.6 V to 3.0 V vs. Al/Al³⁺ (Fig. 1a). The decomposition of the electrolyte was observed as the upper potential elevated. The oxidative decomposition reaction was assumed to be the generation of chlorine by oxidation of AlCl₃. A clear reduction peak was observed at 2.2 V vs. Al/Al³⁺ when the upper potential reached at 2.9 V vs. Al/Al³⁺, indicating the reduction of chlorine. From the above, the upper limit of the potential window of DMSO₂/AlCl₃ electrolyte is approximately 2.5 V vs. Al/Al³⁺.

Figure 1.

(a) CV of the molybdenum electrode in the electrolyte with AlCl₃ and DMSO₂. The lower potential limit was fixed at −0.5 V vs. Al/Al³⁺, and the upper potential limit varied from 2.6 V to 3.0 V vs. Al/Al³⁺. (b) CV of the molybdenum electrode in the electrolyte with AlCl₃ and DMSO₂. The upper potential limit was fixed at 2.5 V vs. Al/Al³⁺, and the lower potential limit varied from −0.3 V vs. Al/Al³⁺ to −0.8 V vs. Al/Al³⁺.

Next, the upper potential limit was fixed at 2.5 V vs. Al/Al³⁺, and the lower potential limit varied from −0.3 V vs. Al/Al³⁺ to −0.8 V vs. Al/Al³⁺ (Fig. 1b). At the lower potential limit of −0.3 V vs. Al/Al³⁺, only one oxidation peak (E1) was observed at 0.8 V vs. Al/Al³⁺, while another oxidation peak (E2) was observed at 1.4 V vs. Al/Al³⁺ when the lower potential limit was set below −0.6 V vs. Al/Al³⁺. The potential of E2 shifted positively as the lower potential limit was shifted negatively. The reduction decomposition of DMSO₂ was not observed in the range of the applied potential down to −0.3 V vs. Al/Al³⁺. On the other hand, a slight hysteresis was observed in the CV in the range of the applied potential down to −0.6 V vs. Al/Al³⁺, and it is thought that the reduction decomposition of DMSO₂ has started. When the applied potential was expanded to a negative potential, a reduction peak of DMSO₂ was observed at around −0.5 V vs. Al/Al³⁺.²³

In order to clarify the reason for the difference of the oxidation reaction for different deposition potential, aluminum was potentiostatically deposited at −0.3 V and −0.8 V vs. Al/Al³⁺, and the deposits on the Mo electrode were analyzed by SEM. Aluminum deposited for 10 hours at −0.3 V and −0.8 V vs. Al/Al³⁺ had a coulombic amount of 35.5 C and 51.2 C. The coulombic efficiency was 62 % and 40 %, respectively, assuming the deposited material was metallic aluminum. The coulomb efficiency of Al deposition on the molybdenum electrodes was estimated to be low because aluminum peeled off when deposited for long periods of time and adhered to the separator. Figures 2a and 2b show SEM images of the deposits at −0.3 V and −0.8 V vs. Al/Al³⁺, respectively. The deposits at −0.3 V vs. Al/Al³⁺ showed submicron-sized grains, while those at −0.8 V vs. Al/Al³⁺ had a very smooth surface. Elemental analysis of these deposits was performed by EDS (Figs. 2c and 2d). For the Al deposited at −0.3 V vs. Al/Al³⁺, a small intensity of O and Mo as substrate were observed together with a strong intensity of Al. It is known that the surface of metallic aluminum is easily oxidized in the air, forming a dense and very thin oxide layer, which was detected as a small amount of oxygen observed in the Al deposited at −0.3 V vs. Al/Al³⁺ since the sample was exposed to air for the SEM measurement. On the other hand, the deposited Al at −0.8 V vs. Al/Al³⁺ showed a significant increase in the intensity of O. The intensity of Mo also increased due to numerous cracks in the deposits.

Figure 2.

(a, b) SEM, (c, d) EDS, (e, f) XRD, and (g, h) XPS results of the deposited aluminum on molybdenum electrode with (a, c, e, g) −0.3 V vs. Al/Al³⁺, and (b, d, g, h) −0.8 V vs. Al/Al³⁺, respectively.

To further investigate the significantly increased O content at −0.8 V vs. Al/Al³⁺, XRD and XPS measurements were performed. Figures 2e and 2f show the XRD patterns of the deposits at −0.3 V and −0.8 V vs. Al/Al³⁺, respectively. In both cases, diffraction peaks identified as metallic aluminum, metallic molybdenum as the substrate, and a small amount of MoO₃ were observed. Figures 2g and 2h show Al2p core-level spectra after Ar⁺-sputtering for different periods of time for the deposits at −0.3 V and −0.8 V vs. Al/Al³⁺, respectively. Before Ar⁺-sputtering, peaks assigned to Al₂O₃ and metallic aluminum were observed for the deposits at −0.3 V vs. Al/Al³⁺, whereas for those at −0.8 V vs. Al/Al³⁺, the only peak due to Al₂O₃ were observed. With increasing the Ar⁺-sputtering period of time, the intensity of metallic aluminum increased and that of Al₂O₃ decreased for the deposits at −0.3 V vs. Al/Al³⁺, while for the deposits at −0.8 V vs. Al/Al³⁺, the peak due to Al₂O₃ was dominant even after the Ar⁺-sputtering for 20 min. The rate of argon etching was 2 nm min⁻¹, and XPS analysis was performed on a depth of approximately 40 nm after 20 minutes of etching. These results suggested that the surface of the deposits at −0.8 V vs. Al/Al³⁺ were a mixture of metallic aluminum and a large amount of Al₂O₃. Since Al₂O₃ was not detected in the XRD measurement, there is a possibility that the precipitate contained amorphous Al₂O₃. From the SEM images, for the deposits at −0.8 V vs. Al/Al³⁺, many cracks were observed due to the presence of a large amount of Al₂O₃, which is less ductile than metallic aluminum. The absence of diffraction peaks due to Al₂O₃ in the XRD pattern suggests that the majority of the deposits consisted of metallic aluminum. In the CV measurements (Fig. 2b), the positive shift of the oxidation peak E2 with lowering the lower potential limit is due to the overpotential caused by Al₂O₃ in the deposits, which interfered with the dissolution of deposited Al metal.

To investigate the reason for the difference in deposits, the potential of the counter electrode was simultaneously recorded during CV measurements. Figures 3a and 3b show CVs measured from −0.3 V to 2.5 V vs. Al/Al³⁺ using the Mo working and Al counter electrodes with the same surface area (1 cm × 1 cm). In the first cycle (Fig. 3a), a reduction current due to metallic aluminum deposition and no oxidation peak was observed, but after 20 cycles, the oxidation peaks due to Al dissolution were observed. In addition, the potential of the counter electrode did not exceed +2.2 V vs. Al/Al³⁺ in both the first and the 20th cycle. Since an oxidation reaction occurs at the counter electrode when a reduction reaction occurs at the working electrode, the counter electrode potential was +2.2 V vs. Al/Al³⁺ when the working electrode potential was −0.3 V vs. Al/Al³⁺. This result suggests that the overpotential required for oxidation and reduction was significantly different for the commercial Al metal and deposited Al metal. The aluminum that precipitated at the beginning of the cycle reacted with impurities on the molybdenum surface and was deposited as an irreversible substance such as an oxide, therefore no oxidation current was observed in the first cycle. Figures 3c and 3d shows the CV and counter electrode potentials measured from −0.8 V to 2.5 V vs. Al/Al³⁺. In the first cycle, the counter electrode potential was kept below 2.0 V vs. Al/Al³⁺, but increased to 3.5 V vs. Al/Al³⁺ at the 20th cycle. The potential of 3.5 V vs. Al/Al³⁺ is beyond the stable potential window of the DMSO₂/AlCl₃ electrolyte, and the electrolyte solution can decompose at the Al counter electrode during Al deposition at the Mo electrode. These results suggest that the large amount of Al₂O₃ in the deposits observed by XPS when the potential of the working electrode was swept to −0.8 V vs. Al/Al³⁺ is due to the oxidative decomposition products, including oxygen, of the electrolyte reacting with the deposited aluminum. For further clarification, the CV measured after holding the Mo working electrode at 3.0 V vs. Al/Al³⁺ for 10 h (Figs. 3e and 3f). Reduction peaks were observed around 1.3 V and −0.3 V vs. Al/Al³⁺ for the first cycle. After 5 cycles, the reduction peaks became smaller, and aluminum deposition and dissolution currents were observed. The onset potential of the chlorine reduction current was 2.5 V vs. Al/Al³⁺, and no reduction current was observed above 2.2 V vs. Al/Al³⁺, meaning that the reduction current with peaks at 1.8 V and 1.3 V vs. Al/Al³⁺ was not chlorine reduction, and they will be attributed to the reduction of the oxidation products of the DMSO₂.

Figure 3.

CVs measured from (a, b) −0.3 V to 2.5 V vs. Al/Al³⁺ and (c, d) −0.8 V to 2.5 V vs. Al/Al³⁺ indicating with the counter electrode potential using the working and counter electrodes with the same surface area (10 mm square). (e, f) The CV measured after holding the working electrode at 3.0 V vs. Al/Al³⁺ for 10 hours.

Reduction decomposition of dimethyl sulfone occurred at around −0.5 V, but oxidation decomposition progressed significantly when the potential of the counter electrode rose up to +3 V or more. In the cell used in this study, the glass fiber filter was used as a separator to attach the working electrode and the counter electrode, and the oxygen produced at the counter electrode easily forms an oxide with the highly reactive aluminum immediately after electrochemical deposition. Generally, the counter electrode is placed separately in the case of electroplating, but considering the structure of the battery, the positive electrode and negative electrode are placed in very close locations; therefore, it is quite possible for the aluminum negative electrode to be passivated.

Depending on the deposition potential, two types of precipitates were observed in deposited aluminum. In the case of the aluminum deposited at −0.3 V, EDS analysis showed almost no elements other than aluminum, and XPS analysis also suggested the existence of metallic aluminum. These results suggested the deposition of high-purity metallic aluminum, and the following two points are the main differences from commercially available aluminum.

The first is the difference in surface area caused by the difference in surface shape. The electrochemically deposited aluminum showed a rough surface shape and commercially available aluminum sheet had smooth surface. Therefore, the electrochemical surface area differs from that of commercially available aluminum sheets, and the overpotential of the deposited aluminum for a constant current density per geometric area should be lower than commercially available aluminum.

Secondly, differences in the oxygen content that affected the electrochemical activity of aluminum metal. Commercially available aluminum has a purity of 99 %, and it contains less than 1 % oxygen. During electrochemical deposition and dissolution, the corrosion resistance changes significantly depending on the trace amounts of oxygen. Therefore, oxygen content in the aluminum metal caused the difference of the electrochemical activity.

Furthermore, in the aluminum precipitated at −0.8 V, the results of EDS elemental analysis showed large amounts of oxygen and sulfur. The presence of these elements causes an increase in overpotential.

In research for rechargeable batteries, the working electrode and the counter electrode are often the same sizes in coin cell configuration. On the other hand, the counter electrode should be larger than the working electrode for general electrochemical experiments. Therefore, we also performed CV measurements using a 10 mm square molybdenum plate for the working electrode (Fig. 4a) and a 5 mm square molybdenum plate for the working electrode (Fig. 4b), and the counter electrode potential was recorded simultaneously. As a result, even when the working electrode potential was swept from −0.8 V to 2.5 V, the potential of the counter electrode remained below +1.0 V. The Figs. 4c and 4d also show the results of measurements under more extreme conditions, where the area of the counter electrode was 10 times larger than that of the working electrode. Under these conditions, a very linear reduction current and sharp oxidation peak were observed, and an extra redox current or peak was not observed.

Figure 4.

CVs measured from (a) −0.3 V to 2.5 V vs. Al/Al³⁺ and (b) −0.8 V to 2.5 V vs. Al/Al³⁺ indicating with the counter electrode potential using a 5 mm square molybdenum plate for the working electrode and a 10 mm square aluminum plate for the counter electrode. (c, d) CVs indicating with the counter electrode potential using the counter electrode with 10 times larger surface area than that of the working electrode.

These results suggest that the metallic aluminum in the counter electrode is resistant to dissolution compared to deposited metallic aluminum, which is electrochemically deposited from the electrolyte using dimethyl sulfone. The most significant difference between aluminum plate and deposited aluminum is the surface morphology. As shown in the Fig. 2a, electrochemically deposited metallic aluminum has a very bumpy shape, resulting in a high surface area. On the other hand, the surfaces of commercially available aluminum sheets are smooth, suggesting that a large overpotential was required for electrochemical dissolution. Although the counter electrode used in this study was polished with No. 400 emery paper to increase the surface roughness, it was insufficient.

In order to evaluate the possibility of using DMSO₂ as an electrolyte for aluminum rechargeable batteries, we conducted an aluminum deposition and dissolution experiment on a molybdenum metal plate current collector. A 10 mm square aluminum plate was used as the positive electrode, and a 5 mm square molybdenum metal plate was used as the negative electrode. As shown in Fig. 5a, the overpotential for the charging reaction, i.e. the deposition of aluminum, was approximately 0.2 V, and the overpotential for the discharging reaction (dissolution of aluminum) was almost the same to the charging reaction. Figure 5b shows the charge/discharge curve and coulombic efficiency for the reaction repeated up to 450 cycles. The coulombic efficiency was 100 % up to 30 cycles, but after 30 cycles it decreased to around 90 %. The charge/discharge curve showed that after the 50th cycle, the voltage drops rapidly at the end of discharging, and the cell voltage reached to the cutoff voltage. In addition, the overvoltage associated with charging and discharging increased slightly after 300 cycles. The linear decrease in coulombic efficiency suggests that some of the aluminum deposited by charging changed into inactive compounds such as aluminum oxide and increased resistance for the charge/discharge reaction. The XPS results for deposited aluminum suggested the deposited aluminum was slightly oxidized (Fig. 2g). However, even after 450 cycles, the coulombic efficiency remains above 80 %, and it can be said that, excluding an extremely long-life positive electrode such as graphite, the electrolyte using DMSO₂ and aluminum chloride can be used adequately in the evaluation of positive electrodes for aluminum rechargeable batteries.

Figure 5.

(a) Charge/discharge curves of Al negative electrode deposited on Mo plate current collector. (b) Charge/discharge curves and Coulombic efficiency of Al negative electrode deposited on Mo plate current collector for 450 cycles.

4. Conclusion

CV measurements and analysis of the deposited aluminum were performed for an electrolyte with aluminum chloride using dimethyl sulfone as the solvent. When the working and counter electrodes had the same size, the potential of the counter electrode was observed to exceed the stable electrochemical potential window of the electrolyte. We concluded that a large overvoltage is required for the oxidation reaction on the counter electrode with the metallic aluminum plate compared to electrochemically deposited aluminum. Because the rough morphology of the precipitates caused a larger surface area than the aluminum plate electrode. When aluminum is deposited with the potential of the counter electrode exceeding the stable potential window of the electrolyte, aluminum with high oxide content was deposited, which not only results in a different deposit morphology but also requires a large overvoltage when electrochemically oxidized. Charge/discharge measurement of the Al negative electrode indicated 100 % Coulombic efficiency for initial cycle, and maintained at high value for 450 cycles.

Acknowledgment

AlCl₃ was kindly supplied from Nippon Light Metal Company, Ltd.

CRediT Authorship Contribution Statement

Masanobu Chiku: Conceptualization (Lead), Writing – original draft (Lead)

Takumi Miyake: Investigation (Lead)

Hikaru Kamei: Investigation (Supporting)

Eiji Higuchi: Writing – review & editing (Equal)

Hiroshi Inoue: Supervision (Equal), Writing – review & editing (Equal)

Conflict of Interest

The authors declare no conflict of interest in the manuscript.

Footnotes

M. Chiku and E. Higuchi: ECSJ Active Members

H. Inoue: ECSJ Fellow

References

• 1)  R. Schmuch, R. Wagner, G. Hörpel, T. Placke, and M. Winter, Nat. Energy, 3, 267 (2018).
• 2)  M. Li, J. Lu, Z. Chen, and K. Amine, Adv. Mater., 30, 1800561 (2018).
• 3)  R. Xiong, Y. Pan, W. Shen, H. Li, and F. Sun, Renewable Sustainable Energy Rev., 131, 110048 (2020).
• 4)  F. F. Liu, T. T. Wang, X. B. Liu, and L.-Z. Fan, Adv. Energy Mater., 11, 2000787 (2021).
• 5)  Z. Li, J. Häcker, M. Fichtner, and Z. Zhao-Karger, Adv. Energy Mater., 13, 2300682 (2023).
• 6)  L. Stievano, I. de Meatza, J. Bitenc, C. Cavallo, S. Brutti, and M. A. Navarra, J. Power Sources, 482, 228875 (2021).
• 7)  J. Tu, W. Song, H. Lei, Z. Yu, L. Chen, M. Wang, and S. Jiao, Chem. Rev., 121, 4903 (2021).
• 8)  H. Yang, H. Li, J. Li, Z. Sun, K. He, H.-M. Cheng, and F. Li, Angew. Chem. Int. Ed., 58, 11978 (2019).
• 9)  K. L. Ng, B. Amrithraj, and G. Azimi, Joule, 6, 134 (2022).
• 10)  Y. Fang, X. Jiang, X.-G. Sun, and S. Dai, Chem. Commun., 51, 13286 (2015).
• 11)  M. Angell, C.-J. Pan, Y. Rong, C. Yuan, M.-C. Lin, B.-J. Hwang, and H. Dai, Proc. Natl. Acad. Sci. U.S.A., 114, 834 (2017).
• 12)  M. Chiku, H. Takeda, S. Matsumura, E. Higuchi, and H. Inoue, ACS Appl. Mater. Interfaces, 7, 24385 (2015).
• 13)  M. Chiku, T. Kunisawa, E. Higuchi, and H. Inoue, RSC Adv., 9, 41475 (2019).
• 14)  M. Chiku, T. Yihan, E. Higuchi, and H. Inoue, J. Alloys Compd., 969, 172430 (2023).
• 15)  L. Legrand, M. Heintz, A. Tranchant, and R. Messina, Electrochim. Acta, 40, 1711 (1995).
• 16)  L. Legrand, A. Tranchant, R. Messina, F. Romain, and A. Lautie, Inorg. Chem., 35, 1310 (1996).
• 17)  L. Legrand, A. Tranchant, and R. Messina, Electrochim. Acta, 41, 2715 (1996).
• 18)  K. Dokko, D. Watanabe, Y. Ugata, M. L. Thomas, S. Tsuzuki, W. Shinoda, K. Hashimoto, K. Ueno, Y. Umebayashi, and M. Watanabe, J. Phys. Chem. B, 122, 10736 (2018).
• 19)  Y. Ugata, S. Sasagawa, R. Tatara, K. Ueno, M. Watanabe, and K. Dokko, J. Phys. Chem. B, 125, 6600 (2021).
• 20)  A. Okamoto, M. Morita, and N. Yoshimoto, J. Surf. Finish. Soc. Jpn., 65, 47 (2014).
• 21)  A. Okamoto, M. Morita, and N. Yoshimoto, J. Surf. Finish. Soc. Jpn., 63, 641 (2012).
• 22)  A. Okamoto, N. Niwa, M. Egashira, M. Morita, and N. Yoshimoto, Electrochemistry, 81, 906 (2013).
• 23)  J. Matsuda, A. Okamoto, H. Akahoshi, S. Shironita, and M. Umeda, J. Electrochem. Soc., 169, 112516 (2022).