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Dissociation and Diffusion of Glyme-Sodium Bis(trifluoromethanesulfonyl)amide Complexes in Hydrofluoroether-Based Electrolytes for Sodium Batteries Shoshi Terada, Hiroko Susa, Seiji Tsuzuki, Toshihiko Mandai, Kazuhide Ueno, Yasuhiro Umebayashi, Kaoru Dokko, and Masayoshi Watanabe J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.6b06804 • Publication Date (Web): 27 Sep 2016 Downloaded from http://pubs.acs.org on September 30, 2016

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The Journal of Physical Chemistry

Dissociation and Diffusion of Glyme-Sodium Bis(trifluoromethanesulfonyl)amide Complexes in Hydrofluoroether-Based Electrolytes for Sodium Batteries

Shoshi Terada,† Hiroko Susa,† Seiji Tsuzuki,‡ Toshihiko Mandai,†,¶ Kazuhide Ueno,†,# Yasuhiro Umebayashi,§ Kaoru Dokko,†,∥,* Masayoshi Watanabe†

†

Department of Chemistry and Biotechnology, Yokohama National University, 79-5 Tokiwadai,

Hodogaya-ku, Yokohama 240-8501, Japan ‡

Research Center for Computational Design of Advanced Functional Materials (CD-FMat), National

Institute of Advanced Industrial Science and Technology (AIST), Tsukuba Central 2, 1-1-1 Umezono, Tsukuba, Ibaraki 305-8568, Japan §

Graduate School of Science and Technology, Niigata University, 8050 Ikarashi, 2-no-cho, Nishi-ku,

Niigata City 950-2181, Japan ∥

Unit of Elements Strategy Initiative for Catalysts & Batteries (ESICB), Kyoto University, Kyoto

615-8510, Japan

*CORRESPONDING AUTHOR FOOTNOTE: To whom correspondence should be addressed. Telephone/Fax: +81-45-339-3942. E-mail: [email protected]

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ABSTRACT. Physicochemical properties and battery performance of [Na(glyme)][TFSA] complexes (TFSA: bis(trifluoromethanesulfonyl)amide) dissolved in a hydrofluoroether (HFE) were investigated. Glyme (tetraglyme (G4) or pentaglyme (G5)) coordinates to Na+ to form a 1:1 complex [Na(G4 or G5)]+ cation. Raman spectroscopy revealed that the complex structure of [Na(glyme)]+ is maintained in the HFE solution, and free (uncoordinated) glymes are not liberated on adding HFE. HFE molecules are scarcely involved in the first solvation shell of Na+ because of their low electron-pair donating ability. Raman spectra of the [TFSA]− anion suggests that the attractive interaction between the complex [Na(glyme)]+ cation and [TFSA]− anion is enhanced on adding HFE. The population of contact ion-pair (CIP) and/or aggregate (AGG) is smaller for the G5 system than for the G4 one, and the [Na(G5)][TFSA]/HFE has higher ionic conductivity. The self-diffusion coefficients of the [Na(glyme)]+ complex and [TFSA]− were measured by pulsed field gradient (PFG) NMR and the dissociativity of [Na(glyme)][TFSA] was assessed. The dissociativity of the G5 system is greater than that of the G4 one, and the dissociativity can be correlated with the attractive interaction between [Na(glyme)]+ and [TFSA]− as evaluated by ab initio calculations. The dissociativity of the complexes gave significant effects on the battery performance.


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INTRODUCTION Glymes (Gn, CH3−O−(CH2−CH2−O)n−CH3) form complexes in certain molar ratios with metal ions such as Li+, Na+, K+, Mg2+, and Al3+ owing to the electron-pair donating ability of ether oxygen atoms.1−13 Ether oxygen atoms have lone pairs, and the electrostatic and induction interactions between the oxygen atoms and the metal ion result in complex formation.14 The crystal structures and physicochemical properties of [M(glyme)x]Xy (M: metal ion, X: counter anion) complexes have been previously established.1-13 The metal ion interacts with the ligand (glyme) as well as with the anion in the [M(glyme)x]Xy complex with the ligand and anion conferring significant effects on the physicochemical properties of the complex.15−19 Previously, we reported the electrochemical properties and battery applications of such complexes.20−29 In the case of the [Li(glyme)][TFSA] (TFSA: bis(trifluoromethanesulfonyl)amide) complex, the Li+ forms a 1:1 complex with the glyme molecule when the glyme chain length (n) is 3 or 4. [Li(G3 or G4)][TFSA] complexes are low melting and remain in the liquid state at room temperature.30,31 Pappenfus et al. pointed out that the [Li(G4)][TFSA] can be regarded as an ionic liquid and shows ionic conductivity at ambient temperature.30 Our recent studies demonstrated that the molten [Li(glyme)][TFSA] complexes are solvate ionic liquids consisting of [Li(glyme)]+ and [TFSA]−.16,17 The molten [Li(glyme)][TFSA] partially dissociates into complex cation and anion, has ionic conductivity, and can be used as electrolytes for Li batteries.22−27 In this study, we investigated the [Na(glyme)][TFSA] complexes as electrolytes for Na batteries. Na batteries are attracting attention because Na is earth-abundant compared to Li. In the case of the [Na(glyme)][TFSA] complex, Na+ forms a 1:1 complex with a glyme molecule when the chain length, n, is 4 or 5.9,10 The melting points of [Na(G4)][TFSA] and [Na(G5)][TFSA] are 72 and 32 °C, respectively. [Na(G5)][TFSA] remains in the liquid state as a supercooled liquid at room temperature ACS Paragon Plus Environment

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while [Na(G4)][TFSA] remains in the solid state. Molten [Na(G5)][TFSA] shows a low ionic conductivity of 0.6 mS cm−1 at 30 °C.29 Using a molten complex as an electrolyte for Na batteries at room temperature is challenging because the low ionic conductivity causes high internal resistance in the cells. To utilize [Na(glyme)][TFSA] complexes in Na batteries at room temperature, the addition of another solvent is a facile and effective way to lower viscosity and enhance ionic conductivity. The complexes of [M(glyme)][TFSA] can dissolve into various solvents.32 Interestingly, low polar solvents such as toluene are also miscible with [M(glyme)][TFSA] and can be used as electrolyte solutions. In this work, a hydrofluoroether (HFE), 1,1,2,2–tetrafluoroethyl 2,2,3,3–tetrafluoropropyl ether (CF2H−CF2−O−CH2−CF2−CF2H), was used as the solvent for electrolyte solutions of [Na(G4 or G5)][TFSA]. The HFE is a low polar solvent, and is chemically stable and non-flammable.25,33 Previously, we prepared electrolyte solutions of [Li(glyme)][TFSA]/HFE and used these as electrolytes for Li-S batteries.25 The addition of HFE increased ionic conductivity and was effective in improving battery performance. In this work, we elucidated the solvation structures of Na+ and ion-pair formations in the [Na(G4 or G5)][TFSA]/HFE solutions using Raman spectroscopy. In addition, the self-diffusion coefficients of ions in [Na(G4 or G5)][TFSA]/HFE were measured using pulsed field gradient (PFG) NMR, and the dissociativity, so called “ionicity”,34 of [Na(G4 or G5)][TFSA] was evaluated in solution. Ionicity is correlated with Na+ solvation structures and ion-pair formation. Finally, the [Na(G4 or G5)][TFSA]/HFE solutions were used as electrolytes in Na batteries. The relationships between electrolyte properties and battery performance are discussed.

EXPERIMENTAL Materials ACS Paragon Plus Environment

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Bis(trifluoromethanesulfonyl)amide (H[TFSA]) and Na2CO3 were purchased from Kanto Chemical and used as received. Na[TFSA] was synthesized by neutralization of H[TFSA] with Na2CO3 according to a previously published procedure.35 The obtained Na[TFSA] was dried under vacuum at 80 °C then at 100 °C for 12 h each and then stored in an Ar atmosphere glove box. Triglyme (G3), tetraglyme (G4), and pentaglyme (G5) were kindly supplied by Nippon Nyukazai. G3 and G4 were used as received and G5 was distilled under reduced pressure over sodium metal. Na[TFSA] and glyme (G4 or G5) were mixed in a 1:1 ratio to obtain a [Na(glyme)][TFSA] complex. HFE solvent, 1,1,2,2– tetrafluoroethyl 2,2,3,3–tetrafluoropropyl ether, was purchased from Daikin Industries and used without further treatment. [Na(glyme)][TFSA] was mixed with HFE in various ratios to prepare electrolyte solutions (Table S1). The water content of the solutions was measured by Karl-Fischer titration and was less than 50 ppm for all the electrolyte solutions. Na2S was purchased from Kojundo Chemical Laboratory. The acetylene black (AB, Denka Black) and Ketjen black (KB, EC600JD) were supplied by Denki Kagaku Kogyo and Lion Corporation, respectively. Poly(vinylidene fluoride) (PVDF) was purchased from Kishida Chemical. Elemental sulfur (S8), poly(vinyl alcohol) (PVA, saponification degree 86−90 mol%, average degree of polymerization 3100–3900), N-methyl-2-pyrrolidinone (NMP), and Na metal were purchased from Kanto Chemical. Na0.44MnO2 was synthesized by a solid-state reaction according to a reported procedure.36 Na2CO3 and MnCO3 were ground in the a molar ratio of 0.48:2, pressed into pellets, and heated for 8 hours at 300 °C under air followed by grounding, pelletizing, and heating for 9 hours at 800 °C under air. Measurements The ionic conductivities (σ) of [Na(glyme)][TFSA]/HFE solutions were determined by the complex impedance method using an impedance analyzer (VMP3, Biologic) in the frequency range of ACS Paragon Plus Environment

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500 kHz–1 Hz with a sinusoidal alternating voltage amplitude of 10 mV root-mean-square (rms). A cell equipped with two platinized platinum electrodes (CG–511B, TOA Electronics) was utilized for conductivity measurements, and the cell constant was determined using a 0.01 M KCl aqueous solution at 25 °C prior to the measurements. The cell was placed in a temperature-controlled chamber and conductivity was measured at 30 °C. The liquid density and viscosity were determined using a viscometer (SVM 3000, Anton Paar). Raman spectra of the samples were collected using a Raman spectrometer equipped with a 532 nm laser (RMP–330, JASCO). The instrument was calibrated using a polypropylene standard. The spectral resolution was 4.5 cm−1. Sample temperature was adjusted to 30 ± 0.1 °C using a Peltier microscope stage (TS62, INSTEC) with a temperature controller (mk1000, INSTEC). Pulsed field gradient (PFG) NMR measurements were carried out to determine the self-diffusion coefficients of glyme, TFSA anion, and HFE. A JEOL-ECX 400 NMR spectrometer with a 9.4 T narrow-bore super-conducting magnet equipped with a pulsed-field gradient probe and current amplifier was used for the measurement. Self-diffusion coefficients were calculated with the Hahn spin-echo sequence using the signals of 1H of terminal methyl group of glyme and 19F of [TFSA]− and HFE. The detailed experimental procedures have been reported elsewhere.37 The diffusion echo signal attenuation, E, is related to the experimental parameters by the Stejskal equation with a sinusoidal pulsed-field gradient: ln = ln =

− g 4 −

where S is the spin-echo signal intensity, δ is the duration of the field gradient with magnitude g, γ is the gyromagnetic ratio, and ∆ is the interval between the two gradient pulses. The value of ∆ and δ were set at 50 ms and 5 ms, whereas g was set at 0.01~1.9 T m−1 depending on the electrolyte. The sample was ACS Paragon Plus Environment

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inserted into an NMR microtube to a height of 3 mm to exclude convection. All measurements were conducted at 30 °C. Each sample was placed in a sample tube with an outer diameter of 4 mm, and that tube was inserted into a 5 mm standard NMR sample tube. Cyclic voltammetry (CV) and linear sweep voltammetry (LSV) were performed with a three-electrode cell using an electrochemical analyzer (VMP3, Biologic). Cu and Pt disk electrodes (3 mm in diameter) were used as the working electrodes for CV and LSV measurements, respectively. Pt wire was used as the counter electrode. The reference electrode was Li metal soaked in 1 mol dm−3 Li[TFSA]/G3 solution, confined in a glass tube with a liquid junction (Vycor glass). All electrochemical measurements were performed at 30 °C. [Na(glyme)][TFSA]/HFE was tested as an electrolyte for Na batteries. In this study, Na-Na0.44MnO2 and Na-S cells were prepared. The Na0.44MnO2 composite cathode was prepared according to a previously reported method by thoroughly mixing Na0.44MnO2, AB, and PVDF in a weight ratio of 80:10:10 and adding NMP to obtain a slurry. The slurry was pasted onto Al foil and dried overnight at 80 °C. The composite cathode was compressed at 50 MPa followed by drying in a vacuum at 120 °C for 12 h. The sulfur cathode was also prepared using a similar procedure. Sulfur, KB, and PVA were mixed in a weight ratio of 40:45:15 in NMP to obtain a slurry. The slurry was pasted onto Al foil and dried overnight at 80 °C. The S composite cathode was compressed at 50 MPa and used without further drying. Coin cells (2032-type) were fabricated using Na0.44MnO2 or S composite cathodes in an Ar-filled glovebox. Na metal was used as the anode and glass fiber filter (GA–55, ADVANTEC) was used as the separator. 200 µL of [Na(glyme)][TFSA]/HFE was used as the electrolyte. The charge– discharge tests were carried out using an automatic charge/discharge instrument (HJ1001SD, Hokuto Denko) at 30 °C. ACS Paragon Plus Environment

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Computational Methods The Gaussian 09 program was used for the ab initio molecular orbital calculations.38 The basis sets implemented in the program were used. Electron correlation was accounted for by the second-order Mϕller-Plesset perturbation (MP2) method.39,40 The geometries of the complexes were fully optimized at the HF/6-311G** level (Figures S1−S4). The coordinates of the optimized geometries were shown in Table S2. The geometries of complexes in crystals were used for initial geometries for the optimizations.9 The basis set superposition error (BSSE)41 was corrected for the calculations of interaction energies (∆Eint) using the counterpoise method.42 The accuracy of the interaction energies obtained using HF-level optimized geometries is described in Supporting Information (Table S3). The interaction energy (∆Eint) for [Na(glyme)]+ complex is the interaction between Na+ and glyme. The ∆Eint for [Na(glyme)][TFSA] complex is sum of the interactions among three species. The stabilization energy by the formation of complex from isolated species (∆Eform) was calculated by the equation ∆Eform = ∆Eint + ∆Edef. The ∆Edef for [Na(glyme)]+ complex is the energy difference between the optimized geometry of the glyme isolated in the gas phase and the distorted geometry which the glyme takes in the complex.43 The ∆Edef for [Na(glyme)][TFSA] complex is the sum of the deformation energies of glyme and [TFSA]− anion. The interaction energy (∆Eint) and deformations energy (∆Edef) were calculated at the MP2/6-311G** level. The binding energy of [Na(glyme)]+ and [TFSA]− (∆Ebind) was calculated by the equation ∆Ebind = ∆Eform ([Na(glyme)][TFSA]) − ∆Eform ([Na(glyme)]+), where the ∆Eform ([Na(glyme)][TFSA]) and ∆Eform ([Na(glyme)]+) are ∆Eform for [Na(glyme)][TFSA] and [Na(glyme)]+, respectively. The HOMO energy levels of isolated glymes and those of the glymes in the complexes were calculated at the HF/6- 311G** level using the optimized geometries. ACS Paragon Plus Environment

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The optimized geometries of the [Na(G4)]+, [Na(G5)]+, [Na(G4)][TFSA], and [Na(G5)][TFSA] at the HF/6-311G** level were further optimized and vibrational frequency analysis was carried out at the B3LYP/6-311+G** level. The vibrational analysis shows that the optimized geometries correspond to potential energy minimum structures. The zero point, thermal energy and entropy corrections were carried out using the calculated frequencies for obtaining the change of the Gibbs free energy by the formation of the complex from isolated species (∆Gform) (Table S4).

RESULTS AND DISCUSSION Solvation Structures of Na+ [Na(G5)][TFSA] is in the liquid state at room temperature and is fully miscible with HFE. [Na(G4)][TFSA] is a solid at room temperature and precipitation remains when the molar ratio of HFE/[Na(G4)][TFSA] is less than 2 in the mixture. Solvation structures of Na+ in electrolyte solutions were investigated using Raman spectroscopy. Glymes have a fingerprint mode at ~870 cm−1 when the glyme molecule coordinates to a metal cation in a crown-ether like structure.1-3 This Raman peak is called the breathing mode and is a combination of COC stretching vibration and CH2 rocking vibration. Previously, we reported the crystal structures of [Na(G4)][TFSA] and [Na(G5)][TFSA] complexes.9,10 In the crystals, glyme and Na+ form a 1:1 complex cation, [Na(glyme)]+, with a crown-ether like coordination structure resulting in an intense Raman peak at 870 cm−1 assignable to the breathing mode. The breathing mode is also observed for the molten complexes, suggesting that the 1:1 complex structures of [Na(G4)]+ and [Na(G5)]+ are kept in the molten state. Detailed analysis of the molten [Na(G4)][TFSA] and [Na(G5)][TFSA] complexes suggested that almost all the glyme molecules coordinate to Na+ and presence of free glymes in the liquids is negligible.9,10 In the case of the ACS Paragon Plus Environment

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[Na(glyme)][TFSA]/HFE solution, the Raman peaks of HFE overlap with the breathing mode of [Na(glyme)]+ as shown in Figure S5. To analyze the coordination between glymes and Na+ in solution, the Raman spectra were normalized by the peak intensity of HFE at 670 cm−1 (Figure S6) and the HFE spectrum was subtracted from each [Na(glyme)][TFSA]/HFE spectrum. As shown in Figure 1, the obtained difference spectra are similar to those of previously reported molten [Na(glyme)][TFSA] complexes.9 The difference spectra are deconvoluted into 7 peaks using the Gaussian-Lorentzian (pseudo-Voigt) function as shown in Figure S7. Intense Raman peaks are observed at 868 and 742 cm−1 and are assignable to the breathing mode of [Na(glyme)]+ and the S−N symmetric stretching vibration coupled with the CF3 bending of [TFSA]−, respectively.1,9,44 In addition to the breathing mode, weaker bands at 850, 835, 813, and 795 cm−1 are also observed, and these peaks are also assigned to the vibration modes of glyme in the [Na(glyme)]+ complex.9 As shown in Figure 2, the integral intensities of the breathing mode (Ibreath) and the peak of [TFSA]− (ITFSA) increase linearly with cNa/cHFE, where cNa and cHFE are the concentrations of Na[TFSA] and HFE, respectively. Clearly, the Raman scattering coefficient of [TFSA]− is independent of the concentration. Assuming that the Raman scattering coefficient of the breathing mode is independent of the concentrations of glyme, Na[TFSA], and HFE, we can consider that the linear relationship between Ibreath and cNa/cHFE indicates that the [Na(glyme)]+ coordination structure is kept intact in the [Na(glyme)][TFSA]/HFE solutions. If the complex structure of [Na(glyme)]+ had been destroyed and uncoordinated glyme had been liberated by the addition of HFE, the Ibreath would not be proportional to cNa/cHFE. Ibreath/ITFSA for the [Na(glyme)][TFSA] complexes with and without HFE is ~0.4 (Figure S8). This result also supports that the molar ratio of [TFSA]− to [Na(glyme)]+ is always 1 in solution and the liberation of uncoordinated glymes does not occur. Thus, we confirm that the complex cation structure [Na(G4 or G5)]+ is retained by the addition of HFE to ACS Paragon Plus Environment

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[Na(G4 or G5)][TFSA]. Conversely, the interaction between HFE and Na+ is negligible. We previously reported liquid structures of [Li(glyme)][TFSA]/HFE solutions.25,32,45 As mentioned before, glymes have multiple ether oxygen atoms, and the electron-pair donating ability of glymes is strong due to the lone pairs of oxygen atoms resulting in the strong attractive force for Li+ owing to electrostatic and induction interactions.14 On the other hand, the solvation ability of HFE is rather weak because the electron-pair donating ability of HFE’s ether oxygen is weak due to electron withdrawing by the fluorine atoms.45 The Gutmann donor number (DN), which is a good metric for solvent donating ability, of HFE is as low as 2 kcal mol−1.32 In contrast, the DNs of glymes are much higher and in the range of 14−20 kcal mol−1.46 Therefore, the interaction between Li+ and glyme is much stronger than that of Li+-HFE, and glyme coordinates to Li+ preferentially while HFE hardly participates in the solvation. Similarly, the interaction between Na+ and glyme should be stronger than that of Na+-HFE, and HFE negligibly takes part in the solvation of Na+.

Figure 1. Raman spectra for (a) [Na(G4)][TFSA]/HFE and (b) [Na(G5)][TFSA]/HFE. The HFE spectrum was subtracted from each spectrum.


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Figure 2. Integral intensities for Raman bands assigned to the breathing mode (868 cm–1) and [TFSA]− (742 cm–1) as a function of the Na/HFE molar ratio (cNa/cHFE) in (a) [Na(G4)][TFSA]/HFE and (b) [Na(G5)][TFSA]/HFE.

The relative dielectric constant (εr) of HFE at 25 °C is 6.21,25,32 which is lower than those of other solvents such as propylene carbonate (εr: 64.9), water (εr: 78.4), and glymes (εr: 7−8).46 As the relative permittivity of a medium (solvent) becomes lower, the electrostatic interaction between the cation and the anion becomes stronger. The [TFSA]– anions interact with complex [Na(glyme)]+ cations and form ion-pairs in the solution. The Raman band for [TFSA]− at around 740 cm−1 is known to be sensitive to the ion-pair formation.1,47−49 The Raman band for [TFSA]− appears at ~745−747 cm−1 when [TFSA]− coordinates Na+ and forms contact ion-pairs (CIPs) and/or aggregates (AGGs). On the other hand, the Raman band for [TFSA]− appears at ~739−742 cm−1 when the interaction between [TFSA]− and Na+ is weak and [TFSA]− is not bound to Na+, i.e., [TFSA]− is free and/or in the solvent shared ion-pair (SSIP). In this study, we could not distinguish the bound [TFSA] from the free- and/or SSIP-[TFSA] clearly by Raman spectroscopy; however, the peak position of [TFSA]− shifted depending


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on the solution composition (Figure S9). Figure 3 represents the peak position for the [TFSA]− band as a function of [Na(glyme)][TFSA] concentration. In the entire concentration range, the peak position of the [TFSA]− band for [Na(G4)][TFSA]/HFE is higher than that for [Na(G5)][TFSA]/HFE, indicating that the interaction between the complex cation and [TFSA]− is stronger and the population of CIP and/or AGG is higher in the former solution. This is because the attractive electrostatic and induction interactions of [Na(G4)]+−[TFSA]− are stronger than those of [Na(G5)]+−[TFSA]− (vide infra). It should be noted that for [Na(G5)][TFSA], the [TFSA]− band at 741 cm−1 shifted to higher wavenumber side with deceasing [Na(G5)][TFSA] content (Figure 3). Similarly, a blue shift of the [TFSA] peak due to the addition of HFE is observed for [Na(G4)][TFSA], however the shift is less significant. The reason for the blue shift is thoroughly discussed later in the “Ionicity” section.

Figure 3. Peak position of the [TFSA]− band as a function of [Na(glyme)][TFSA]/HFE concentration.

Transport Properties Figure 4 shows the concentration dependence of ionic conductivity (σ) and viscosity (η) for the [Na(glyme)][TFSA]/HFE solution. The viscosity was lowered upon the addition of HFE and ionic ACS Paragon Plus Environment

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conductivity reached a maximum at ~1 mol dm−3. The ionic conductivity is proportional to the number density and mobility of the charge carriers. While the mobility of charge carriers increased upon the addition of low viscous solvent HFE, the number density decreased resulting in the above maximum. This behavior is similar to conventional electrolytes used in lithium and sodium secondary batteries.50,51 The ionic conductivity of the G4 electrolyte is lower than that of the G5 electrolyte for the same concentration despite their similar viscosities. This can be attributed to the difference in the number density of charge carriers contributing to the ionic conduction. The cation and anion form ion-pairs and AGGs in the solution (vide supra). If the lifetime of electrically neutral ion-pairs is relatively long, the ion-pairs contribute less to the ionic conduction in the solution. Based on the Raman measurements (Figure 3), the CIP population of the G4 system is higher than that of the G5 system, leading to the relatively lower ionic conductivity of the G4 electrolyte.

Figure 4. Concentration dependencies of ionic conductivity (σ) and viscosity (η) for (a) [Na(G4)][TFSA]/HFE and (b)[Na(G5)][TFSA]/HFE at 30 °C. The molar ratio of [Na(glyme)][TFSA] to HFE is 1 : x.


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Figure 5 shows the concentration dependence of self-diffusion coefficients measured by PFG-NMR for glyme, [TFSA]−, and HFE in [Na(glyme)][TFSA]/HFE. Unfortunately, the self-diffusion coefficient for Na+ could not be measured using our apparatus. The self-diffusion coefficient for each species is increased as the [Na(glyme)][TFSA] concentration is decreased owing to the lowering viscosity. The self-diffusion coefficients for glyme and [TFSA]− are similar, indicating that [Na(glyme)]+ and [TFSA]− have comparable hydrodynamic radii. In contrast, HFE diffuses faster than other species. Similar results were observed for [Li(G3)][TFSA]/HFE and [Li(G4)][TFSA]/HFE.25,32

Figure 5. Concentration dependencies of self-diffusion coefficients for glyme, [TFSA]−, and HFE in (a) [Na(G4)][TFSA]/HFE and (b) [Na(G5)][TFSA]/HFE measured at 30 °C.

Ionicity From the self-diffusion coefficients for the cation and the anion in the solution, we can calculate the molar conductivity (ΛNMR) using the Nernst-Einstein equation. =

# " !" !"


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where F is the Faraday constant, R is the gas constant, T is the absolute temperature, and Dcation and Danion are the self-diffusion coefficients for the cation and the anion, respectively. From the ionic conductivity of the solution measured by the impedance method (Figure 4), the molar conductivity (Λimp) can be evaluated as follows. $% =

& 'Na

The ΛNMR is a molar conductivity of the electrolyte where all cations and anions in the system are assumed to contribute to the ionic conduction. In contrast, the Λimp corresponds to the net migration of charged species under an electric field. Therefore, the ratio of these molar conductivities, Λimp/ΛNMR, so-called ionicity, represents the apparent dissocitativity of the salt.34 The ΛNMR can be calculated from the self-diffusion coefficients for [TFSA]− (DTFSA) and glyme (DG) (Figure 5). Although the diffusion coefficient for Na+ could not be measured, we can assume that Na+ and glyme diffuse together as [Na(glyme)]+ and the diffusion coefficient for Na+ is identical to DG. This assumption is supported by the results of Raman measurements for [Na(glyme)][TFSA]/HFE (Figure 1 and Figure 2). Figure 6 shows the concentration dependence of ionicity for [Na(glyme)][TFSA]/HFE. The ionicity decreases upon decreasing the concentration of [Na(glyme)][TFSA] (or addition of HFE) in both the G4 and G5 systems (Figure 6). The decrease in ionicity indicates that the association of [Na(glyme)]+ and [TFSA]− is enhanced by the addition of HFE. It is interesting to note that the ionicity of Li salt dissolved in high polar solvents such as propylene carbonate increases as the concentration of the salt decreases.52 However, the present system, [Na(glyme)][TFSA]/HFE, demonstrates the opposite trend.


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Figure 6. Concentration dependence of ionicity (Λimp/ΛNMR) for [Na(glyme)][TFSA]/HFE at 30 °C.

As shown in Figure 6, the dissociativity of [Na(G5)][TFSA] is the greatest in the molten [Na(G5)][TFSA] complex (without HFE). [TFSA]− in the molten [Na(G5)][TFSA] complex is surrounded by several [Na(G5)]+ cations, and the electrostatic and induction interactions between [TFSA]− and [Na(G5)]+ is weakened by other [Na(G5)]+ cations. Therefore, the attractive interaction of [Na(G5)]+−[TFSA]− in the molten complex is weaker than that of the isolated ion-pair in gas phase where other ions and molecules do not interact with the ion-pair. In the molten [Na(G5)][TFSA] complex, the complex cation and the anion are adjacent to other anions and cations, respectively, and the ion-pairs exchange the ions at a relatively high rate (e.g., [Na(G5)]+−[TFSA]− + [Na(G5)]*+ → [Na(G5)]*+−[TFSA]− + [Na(G5)]+). A similar ion exchange possibly occurs between the AGGs such as triple ions ([Na(G5)]+−[TFSA]−−[Na(G5)]+ and [TFSA]−−[Na(G5)]+−[TFSA]−). As a result, the molten [Na(G5)][TFSA] complex has a relatively high dissociativity (ionicity) in spite of the high salt concentration. When HFE is added to [Na(G5)][TFSA], HFE is hardly involved in the first solvation shell of Na+ and the complex structure of [Na(G5)]+ is retained (vide supra). In addition, the attractive ACS Paragon Plus Environment

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interaction between [Na(G5)]+ and [TFSA]− is enhanced by the addition of HFE, as revealed by Raman spectroscopy (Figure 3). This suggests that HFE hardly has the ability to break the ion-pair of [Na(G5)]+−[TFSA]− into separate ions, [Na(G5)]+ and [TFSA]−. However, HFE is miscible with [Na(G5)][TFSA] and produces a homogenous solution macroscopically. A possible explanation is that the HFE molecule gets between two ion-pairs ([Na(G5)]+−[TFSA]−···[Na(G5)]+−[TFSA]− + HFE → [Na(G5)]+−[TFSA]−···HFE···[Na(G5)]+−[TFSA]−). HFE may also get between AGGs such as triple ions. This model is similar to that proposed by Dupont53 and Hunger et al.54 Ion exchange between the ion-pairs (and/or AGGs) takes place by collision, therefore, decreasing the complex concentration and increasing the distance between ion-pairs (and/or AGGs) by the addition of HFE should slow down the ion exchange rate. In other words, the lifetime of an ion-pair should be increased by the addition of HFE. As the distance between ion-pairs in [Na(G5)][TFSA]/HFE increases, the electrostatic interaction between [Na(G5)]+ and [TFSA]− of an ion-pair is enhanced because the electric fields of other ions against the cation and the anion become weaker. Therefore, the Raman peak of [TFSA]− shifts toward higher frequency as HFE increases in the solution (Figure 3). In addition, the ion-pairs with long lifetimes contribute less to ion conduction, resulting in a decrease in dissciativity (ionicity). The ionicity of the G4 system is lower than that of the G5 one and the Raman results suggest that the population of [TFSA]− bound to Na+ in [Na(G4)][TFSA]/HFE is higher than that in [Na(G5)][TFSA]/HFE (Figure 3). To further understand the effect of glyme chain length on the interaction between [Na(glyme)]+ and [TFSA]−, the stabilization energies for complex formation from isolated species (∆Eform) were calculated. The stabilization energies (∆Eform) calculated for the formation of the [Na(glyme)]+ and [Na(glyme)][TFSA] complexes are listed in Table 1. The ∆Eform for [Na(G5)]+ is greater (more negative) than that for [Na(G4)]+, indicating that the interaction between Na+ and G5 is ACS Paragon Plus Environment

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stronger than that between Na+ and G4, while the difference between the ∆Eform for [Na(G4)][TFSA] and that for [Na(G5)][TFSA] is small. The difference between the ∆Eform for [Na(gylme)][TFSA] and that for [Na(glyme)]+ (∆Ebind) corresponds to the binding energy of [Na(glyme)]+ and [TFSA]−. The ∆Ebind for [Na(G4)][TFSA] is greater than that for [Na(G5)][TFSA], indicating that the interaction between [Na(G4)]+ and [TFSA]− is stronger than that between [Na(G5)]+ and [TFSA]−. We previously reported the ∆Ebind between [Li(glyme)]+ and [TFSA]− and the contributions of electrostatic and induction energies (∆Ees and ∆Eind).14 The ∆Ebind calculated for [Li(G3)][TFSA] and [Li(G4)][TFSA] are −82.8 and −70.0 kcal mol−1, respectively. The ∆Ees for the two complexes are −87.4 and −74.1 kcal mol−1. The ∆Eind for the two complexes are −19.7 and −12.6 kcal mol−1, respectively. The ∆Ebind decreased with an increase in the glyme chain length because the attractive electrostatic and induction interactions between the [Li(glyme)]+ complex cation and [TFSA]− anion are weakened by this increase. The weaker interaction between [Na(G5)]+ and [TFSA]− is the reason behind the higher solubility of [Na(G5)][TFSA] in HFE and the smaller population of CIP and/or AGG in [Na(G5)][TFSA]/HFE compared with that of [Na(G4)][TFSA]. The weaker interaction will reduce the lifetime of the ion-pair in [Na(G5)][TFSA]/HFE, leading to greater dissociativity (ionicity) and the higher ionic conductivity.

Table 1. Stabilization energies for the formation of [Na(glyme)]+ and [Na(glyme)][TFSA] complexes.[a] ∆Eform[b]

Glyme

∆Ebind[c]

[Na(glyme)]+

[Na(glyme)][TFSA]

G4

−77.1

−149.1

−72.0

G5

−84.9

−146.8

−61.9


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[a] Energy in kcal mol−1. [b] Stabilization energy for complex formation from isolated glyme and ions. [c] Difference between the ∆Eform for the [Na(glyme)][TFSA] complex and that for the [Na(glyme)]+ complex. ∆Ebind corresponds to the binding energy between the [Na(glyme)]+ complex and [TFSA]− anion.

Electrochemical Windows The electrochemical stability of the [Na(glyme)][TFSA]/HFE electrolytes were elucidated. Hereafter, the electrolytes with a [Na(glyme)][TFSA] concentration of ca. 1 mol dm−3, which have the highest ionic conductivities in both the G4 and G5 systems, are investigated. Figure 7(a) shows the cyclic voltammograms (CVs) for [Na(G4)][TFSA]/HFE and [Na(G5)][TFSA]/HFE electrolytes. The reduction currents corresponding to the deposition of Na metal are observed at lower than 0.3 V vs Li/Li+ in both electrolytes during the cathodic scan. However, the oxidation currents corresponding to the dissolution of Na metal during the anodic scan are smaller than the reduction current, indicating the low reversibility of the Na deposition/dissolution and the irreversible reductive decomposition of the electrolytes at lower than 0.3 V. At present, the decomposition mechanism of the electrolytes is not clear. According to a recent report by Hosokawa et al.,55 [TFSA]− decomposes gradually on the Na metal. In addition, the glymes (G4 and G5) and HFE may also decompose. However, glymes, [TFSA]−, and HFE are stable on the Li metal, and Li batteries using [Li(glyme)][TFSA]/HFE can be operated for prolonged charge-discharge cycles.25,28 For the suppression of irreversible decomposition of the electrolytes on the Na metal, the stabilization of the passivation layer (solid electrolyte interphase) on Na metal may be necessary.


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Figure 7(b) shows the linear sweep voltammograms (LSVs) for [Na(G5)][TFSA] (without HFE) and [Na(glyme)][TFSA]/HFE at 30 °C. The onset of anodic current at around 4.2 V and the rapid increase at greater than 5 V vs. Li/Li+ are observed for each electrolyte. As we previously reported, the free glyme, which does not coordinate to Na+, decomposes at greater than 4.0 V vs. Li/Li+.12 The present electrolytes are stable up to 4.2 V because free glymes hardly exist in the solutions. The oxidation of glymes is caused by the extraction of electrons from the lone pairs of ether oxygen atoms. The lone pair electrons are attracted by the positive charge of Na+, and, therefore, the HOMO energy levels of glymes are lowered by the complexation with Na+. The HOMO energy levels of G4 and G5 of [Na(G4)][TFSA] and [Na(G5)][TFSA] are −11.91 and −11.98 eV, respectively, while those of free G4 and G5 are −11.46 and −11.47 eV, respectively.10 The lowering of HOMO levels results in the shift of the anodic limit for the glymes.10,12,22,29 Therefore, the oxidative stability of [Na(glyme)][TFSA] is enhanced owing to the absence of free glymes. In addition, HFE is not oxidized up to 5 V vs. Li/Li+28 because the electrons of the HFE ether oxygen is withdrawn by the fluorine atoms.45

Figure 7. (a) Cyclic voltammograms with Cu working electrode and (b) linear sweep voltammograms with Pt working electrode for [Na(glyme)][TFSA]/HFE=1:4 and [Na(G5)][TFSA]/HFE =1:4 measured at a scan rate of 1 mV s−1 at 30 °C. ACS Paragon Plus Environment

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Battery Applications The [Na(glyme)][TFSA]/HFE solutions are stable in the 0.3−4.2 V (vs. Li/Li+) range. Although the reductive stability of the electrolytes or the passivation layer on the Na metal should be improved before the practical use of the electrolytes in sodium batteries, preliminary battery tests were carried out. The [Na(glyme)][TFSA]/HFE solutions were used as electrolytes for batteries with a Na metal anode and a Na0.44MnO2 cathode. The charge and discharge capacities of the cell were limited by the Na0.44MnO2 cathode and an excess amount of Na metal was used to make up for the loss of Na due to the irreversible reductive decomposition at the anode. In the initial charging of the [Na | [Na(glyme)][TFSA]/HFE] | Na0.44MnO2] cell, the Na+ is released from Na0.44MnO2: Na0.44MnO2 → Na0.18MnO2 + 0.26Na+ + 0.26e−.56 The theoretical capacity of this reaction is 72 mA h g−1 based on the weight of Na0.44MnO2. After the initial charging, the Na+ insertion and extraction (Na0.18MnO2 + 0.46Na+ + 0.46e− ⇆ Na0.64 MnO2) take place reversibly. The theoretical capacity of this reaction is 127 mA h g−1. Figure 8 shows the results for the [Na | [Na(G4 or G5)][TFSA]/HFE] | Na0.44MnO2] cells. The galvanostatic charge-discharge tests were performed in the 2−4 V range at a current density of 12.7 mA g−1. Multiple voltage plateaus in the charge and discharge curves are observed. According to the literature, the crystal structure changes of NaxMnO2 (0.18 < x < 0.64) occur during the insertion and extraction of Na+ at each plateau.56 The initial charge capacity was ca. 55 mA h g−1 and the discharge capacity was ca 100 mA h g−1. In subsequent cycles, the reversible behavior was observed for both [Na(G4)][TFSA]/HFE and [Na(G5)][TFSA]/HFE for more than 50 cycles, and the charge and discharge capacities of ca. 110 mA h g−1 were obtained. The Coulombic efficiency following the 2nd cycle was over 99%, indicating that severe side reactions such as irreversible oxidation of electrolyte on the ACS Paragon Plus Environment

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cathode and the corrosion of Al current collector of the cathode do not take place. If free glymes existed in the electrolyte, the decomposition of glymes and dissolution of Al corrosion would occur.28

Figure 8. Charge and discharge curves, capacities, and Coulombic efficiencies of [Na | [Na(glyme)][TFSA]/HFE | Na0.44MnO2] cells with a current density of 12.7 mA g−1 (70 µA cm−2) at 30 °C; (a, c) [Na(G4)][TFSA]/HFE = 1:3 and (b, d) [Na(G5)][TFSA]/HFE = 1: 4.

To elucidate the effect of transport properties on battery performance, charge-discharge tests were carried out at various current densities (Figure S10). Figure 9 shows the dependences of charge and discharge capacities on the current density. The 1 C rate is defined as the gravimetric current density of 127 mA g−1 based on the mass of Na0.44MnO2, corresponding to the geometric current density of ca. 0.7 mA cm−2. The discharge and charge capacities of the cells with [Na(G4)][TFSA]/HFE and ACS Paragon Plus Environment

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[Na(G5)][TFSA]/HFE do not change greatly at a current density lower than 1/3 C rate. However, the discharge and charge capacities of the cell with [Na(G4)][TFSA]/HFE greatly decreased at a current density higher than 1/2 C rate and the cell delivered a discharge capacity of only 15 mA h g−1 at 1 C rate. In contrast, the cell with [Na(G5)][TFSA]/HFE showed a better rate capability. The cell with [Na(G5)][TFSA]/HFE delivered a discharge capacity of 80 mA h g−1 at 1 C rate, and the capacity decreased at a higher current density. Apparently, the transport properties of the electrolyte have a significant effect on the rate capability of the cell. The decrease in charge and discharge capacities of a cell at high current density is attributed to the limitation of the mass transport process in the electrolyte.23 At higher current density, the concentration gradient of [Na(glyme)][TFSA] is formed in the separator and the porous Na0.44MnO2 composite cathode, and the cell cannot be charged and discharged at current densities greater than the diffusion limiting current density. Although the diffusion coefficients of [Na(G4)]+ and [Na(G5)]+ are not so different (Figure 5 and Table S1), [Na(G4)][TFSA]/HFE has a lower ionic conductivity than [Na(G5)][TFSA]/HFE (Figure 4). As the ionic conductivity decreases, the iR drop during charge and discharge increases, and the cell reaches the cut-off voltage before achieving the full capacity of active material. The [Na(G5)][TFSA]/HFE electrolyte has a higher ionic conductivity and lowers the iR drop of the operating voltage, leading to the higher rate capability, i.e., the power density of the cell. Note that, in this study, a glass filter (ca. 200 µm thick) was used as a separator between the anode and the cathode. The rate capability of a cell is affected significantly by the diffusion of the electrolyte in the separator.23 Therefore, if we use a thinner separator, the cell will show an improved rate capability.


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Figure 9. Charge and discharge capacities of [Na | [Na(glyme)][TFSA]/HFE | Na0.44MnO2] cells measured at various current densities.

Finally, we applied the [Na(G5)][TFSA]/HFE electrolyte to a sodium-sulfur (Na-S) battery. The elemental sulfur (S8) can be converted to Na2S electrochemically (S8 + 16Na+ + 16e− → 8Na2S), and the theoretical discharge capacity is 1672 mA h g−1 based on the mass of S8. If we can use the full capacity of S8 and the cell can be operated stably at room temperature, the Na-S cell is attractive as a next generation storage device having a high energy density. Previously, we reported the stable charge and discharge of lithium-sulfur (Li-S) cells using [Li(glyme)][TFSA] electrolytes.24−26 The Li-S cells exhibit high discharge capacities of ca. 900 mA h g−1. In addition, the cell could be charged and discharged reversibly for more than 400 cycles with a high Coulmbic efficiency (>98%). This encouraged us to test the Na-S battery using [Na(G5)][TFSA]/HFE. Figure 10 shows the charge and discharge curves and capacities of a Na-S cell with the [Na(G5)][TFSA]/HFE electrolyte. The reduction reaction of sulfur


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occurred below 2.1 V and the discharge voltage decreased gradually as the depth of discharge increased. The charging reaction took place in the voltage region of 1.8−2.5 V with flat plateaus at 2 V and 2.3 V. The initial discharge and charge capacities were 975 mA h g−1 and 920 mA h g−1, respectively. Although the discharge capacity was smaller than the theoretical one, the sulfur cathode exhibited a much higher discharge capacity than other cathode materials studied as candidates for the Na-ion batteries.57 However, the discharge and charge capacities sharply decreased in the subsequent 5 cycles and the capacity of the 5th cycle was 500 m A h g−1. After that, the capacity decreased gradually as the cycle number increased and the capacity of the 50th cycle was 400 mA h g−1. Sodium polysulfides, Na2Sx (2 ≤ x ≤ 8), which are reaction intermediates of the sulfur cathode, easily dissolve in ether solvents.58 The solubility of elemental sulfur and Na2Sx in [Na(G5)][TFSA]/HFE was analyzed (Figure S11). We reported that the solubility of lithium polysulfides Li2Sx (2 ≤ x ≤ 8) in molten [Li(glyme)][TFSA] complexes and [Li(glyme)][TFSA]/HFE solutions is also very low.25 For the dissolution of Li2Sx, the Li2Sx should be solvated, however, almost all glymes coordinate to Li+ and uncoordinated glymes hardly exist in these electrolytes. The absence of free glyme results in the very low solubility of Li2Sx. In addition, the solvation ability of HFE is also very weak, and the solubility of Li2Sx in [Li(glyme)][TFSA]/HFE is also very small.25 Similarly, the solubility of Na2Sx is lower than 10 mM in the [Na(G5)][TFSA]/HFE electrolyte (Figure S11). From the amounts of electrolyte and S8 contained in the cathode, the dissolved sulfur species in the electrolyte of Na-S cell was calculated to be only 1.8% and the remaining 98.2% should be in the solid phase. Consequently, the conversion of S8 ↔ Na2Sx proceeds mainly in the solid phase during charge and discharge. The detailed reaction mechanisms of discharge and charge are not clear at present and need further investigations.


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The capacity degradation of the sulfur cathode is considered to originate from the volume change of active material in the composite cathode. The conversion of S8 into Na2S is accompanied by 200% volume expansion of the active material. This large volume change during charge and discharge reactions possibly leads to loose contact between the active material and the conductive carbon support. If the active material is isolated electrically in the composite cathode, the discharge capacity is decreased. In addition, the electronic conduction path consisting of carbon particles in the composite cathode may be deteriorated by the volume change, and this increases the resistance of the cathode resulting in its gradual degradation.

Figure 10. (a) Galvanostatic charge–discharge curves and (b) capacities and Coulombic efficiencies of the [Na | [Na(G5)][TFSA]/HFE (1: 4) | S] cell measured at a current density of 139 mA g−1, based on the mass of S8 (34.8 µA cm−2) at 30 °C.

CONCLUSIONS The physicochemical properties of [Na(glyme)][TFSA] dissolved in HFE were investigated. The 1:1 complex structures of glyme (G4 or G5) and Na+, [Na(glyme)]+ are maintained in solution, while free (uncoordinated) glymes are not generated by the addition of HFE. The detailed analysis of


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Raman spectra revealed that HFE molecules are scarcely involved in the solvation of Na+ because of their low electron-pair donating ability. The attractive interaction between the complex [Na(glyme)]+ cation and [TFSA]− anion is enhanced by the addition of HFE, resulting in the formation of CIPs and/or AGGs in the solution. The viscosity of the solution is lowered with increasing HFE, and the ionic conductivity shows a maximum at around 1 mol dm−3 concentration of [Na(glyme)][TFSA]. The self-diffusion coefficients for [Na(glyme)]+, [TFSA]−, and HFE are increased as the amount of HFE increased. The ionicity of [Na(glyme)][TFSA] decreased with increasing HFE. This is because the rate of ion exchange between ion-pairs (and/or AGGs) slows as the [Na(glyme)][TFSA]concentration decreases. The attractive interaction in the [Na(G5)]+−[TFSA]− complex is weaker than that in the [Na(G4)]+−[TFSA]− complex. This results in the smaller population of CIP and/or AGG, greater dissociativity (ionicity), and the higher ionic conductivity of the [Na(G5)][TFSA]/HFE solution. The Na metal deposition and dissolution is possible in the [Na(glyme)][TFSA]/HFE solution at the electrode potential of 0.3 V vs. Li/Li+, however, the irreversible decomposition of the electrolytes occurs simultaneously. The solutions are stable up to 4.2 V vs. Li/Li+ and decompose at electrode potential higher than 4.2 V. The [Na | [Na(glyme)][TFSA]/HFE | Na0.44MnO2] cells can be charged and discharged reversibly in the voltage range of 2−4 V at relatively low current densities. The charge and discharge capacities of the cells are decreased as increasing the current density. The cell with [Na(G5)][TFSA]/HFE shows better rate performance than one with [Na(G4)][TFSA]/HFE. This is because the ionic conductivity of [Na(G5)][TFSA]/HFE is higher owing to the higher dissociativity of [Na(G5)][TFSA]. The [Na(G5)][TFSA]/HFE electrolyte was also tested in a Na-S cell. The Na-S cell shows a high capacity of over 950 mA h g−1 at initial discharge. The cell can be recharged with a Coulombic efficiency of 95%, suggesting that severe side reactions are suppressed in the cell and the ACS Paragon Plus Environment

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cell can be charged and discharged repeatedly at room temperature, in principle. During the charge and discharge of the S cathode, the conversion of S8 to Na2Sx proceeds mainly in the solid phase because of the low solubility of the reaction intermediates (Na2Sx) in [Na(G5)][TFSA]/HFE. The capacity of Na-S cell decays upon cycling. This degradation is caused by the large volume change of active material in the cathode during charging and discharging. The reductive decomposition of [Na(glyme)][TFSA]/HFE on the negative electrode should also be suppressed for practical use in Na batteries. Finally, the stabilization of the passivation layer on the negative electrode requires further investigation.

ASSOCIATED CONTENT

Supporting Information. Viscosity, density, molar concentration, ionic conductivity, self-diffusion coefficients, and ionicity of [Na(glyme)][TFSA]/HFE (Table S1), Optimized geometries of the complexes (Figures S1−S4), Coordinates of the optimized geometries (Table S2), Calculated interaction energies (Table S3), Calculated change of the Gibbs free energy by the complex formation (Table S4), Raman spectra for glyme, HFE, and [Na(glyme)][TFSA] (Figure S5), Raman spectra for [Na(glyme)][TFSA]/HFE (Figure S6), Deconvolution of Raman spectrum (Figure S7), Intensity ratio of Raman bands at 868 and 742 cm−1 (Figure S8), Raman spectra of TFSA (Figure S9), Charge-discharge curves (Figure S10), and Solubility of Na2Sx (Figure S11). This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION ACS Paragon Plus Environment

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Corresponding Author *[email protected]

Present Addresses ¶

Present address: Department of Applied Chemistry, Graduate School of Urban Environmental Sciences,

Tokyo Metropolitan University, 1-1 Minami-Ohsawa, Hachioji, Tokyo 192-0397, Japan #

Present address: Graduate School of Sciences and Technology for Innovation, Yamaguchi University,

2-16-1 Tokiwadai, Ube 755-8611, Japan

Notes The authors declare no conflict of interest.

ACKNOWLEDGEMENTS This study was supported in part by the MEXT program “Elements Strategy Initiative to Form Core Research Center” of the Ministry of Education, Culture, Sports, Science, and Technology (MEXT) of Japan, the JSPS KAKENHI (No. 15H03874, No. 15K13815, and No. 16H06368) from the Japan Society for the Promotion of Science (JSPS), and the Advanced Low Carbon Technology Research and Development Program (ALCA) of the Japan Science and Technology Agency (JST).


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