Physicochemical Properties of Fluoride-Based Eutectic Electrolytes Composed of Tetramethylammonium Fluoride and N-Methyltrifluoroacetamide

Abstract

Development of F⁻-containing electrolytes with high electrochemical oxidation stability is crucial for electrochemical fluorination. Although various combinations of fluoride salts and solvents have been investigated for preparing the electrolytes, there have been few studies on eutectic systems using fluoride salts and solid hydrogen bond donors, especially amide compounds. In this study, fluoride-based eutectic electrolytes (F-EEs) composed of tetramethylammonium fluoride ([TMA]F) and N-methyltrifluoroacetamide (MTFAA) as an amide compound are developed, and their physicochemical properties are evaluated including comparison with that composed of [TMA]F and N-methylacetamide (MAA). The F-EEs which are in liquid state at room temperature are successfully obtained for [TMA]F·x[MTFAA] with x = 8.0 and 4.0 owing to melting-point depression by hydrogen bonds between N–H and F⁻. The F-EEs with MTFAA and MAA exhibits the favorable ion transport properties because they have only one N–H bond per molecule and hydrogen bonds of N–H···F⁻ are moderate. The F-EEs with MTFAA exhibit superior electrochemical oxidation stability due to the CF₃ group compared to that with MAA and will be promising as electrolytes for electrochemical fluorination applications.

1. Introduction

Fluorination imparts unique properties to compounds owing to distinctive properties of fluorine such as a small atomic radius and the highest electronegativity and is crucial for various fields such as materials science, energy storage, and medicine.^1,2 In particular, the development of F⁻-containing electrolytes with high electrochemical oxidation stability is significant for the electrochemical synthesis of fluorinated compounds.^3–9

In the preparation of fluoride-containing electrolytes, organic fluoride salts are known to exhibit better solubility than inorganic fluoride salts. Generally, aprotic solvents have favorable electrochemical stability but low solubility of fluoride salts; therefore, it has been attempted to improve the solubility by exploring suitable solvents and designing the structure of cations.^4,10 On the other hand, protic solvents tend to exhibit superior solubility of fluoride salts but poor electrochemical stability. Recent studies have shown that the electrochemical stability of fluoride salts can be improved by dissolving them in high concentrations, which represents prospects of the electrolytes using protic solvents.^7,8

Another possible approach to prepare electrolytes in the protic system is to obtain eutectic electrolytes (EEs) by combining fluoride salts and protic solids, accompanied by melting-point depression. In general, the combination of hydrogen bond acceptors (HBAs) and hydrogen bond donors (HBDs) is known to significantly lower the melting point, and the resulting liquids around room temperature are referred to as deep eutectic solvents (DESs).^11–15 They attract attention as an environmentally-friendly and low-cost material with properties similar to ionic liquids, such as low volatility and flammability. A wide variety of combinations have been studied, whereas there are few studies on the fluoride system, especially those using amide HBDs rather than alcohol HBDs.^16–22

Recently, we reported a fluoride-based DES (F-DES) of [TMA]F·3.5[1,3-DMU], which is composed of tetramethylammonium fluoride ([TMA]F) as a HBA and 1,3-dimethylurea (1,3-DMU) as an amide HBD.²³ It exhibits superior electrochemical stability rather than that composed of ethylene glycol as an alcohol HBD. This motivated us to develop fluoride-based EEs (F-EEs) and DESs with high electrochemical oxidation stability required for the application to electrochemical fluorination by using another amide HBD. In this study, we develop F-EEs composed of [TMA]F and N-methyltrifluoroacetamide (MTFAA) (Fig. 1 for their chemical structures). Here, we call them as F-EEs, not F-DESs, because the melting point of MTFAA is 323 K (see the following Results and discussion), which is comparatively near room temperature, and the melting point can be below room temperature without its deep reduction. The application of MTFAA, which contains CF₃ group, is expected to provide the high electrochemical oxidation stability to the resulting F-EEs, similarly to F⁻-containing electrolytes with 1,1,1,3,3,3-hexafluoro-2-propanol of a CF₃-containing alcohol.^3,24,25 MTFAA is relatively inexpensive and readily available; moreover, it is useful to investigate the effect of CF₃ group in physicochemical properties of F-EEs by comparison with N-methylacetamide (MAA). The F-EEs composed of [TMA]F and MTFAA are prepared, and their physicochemical and electrochemical properties are investigated.

Figure 1.

Schematics of chemical structures of [TMA]F, MTFAA, and MAA.

2. Experimental

Materials were handled under a dry air (dew point < −233 K) or Ar atmosphere in a glove box (dew point < −203 K). Molecular sieves 4A (Fujifilm Wako Chemicals) was washed with distilled water, followed by being dried under vacuum at 513 K for 12 hours. [TMA]F (Sigma Aldrich Co. LLC, > 97.0 %) was used as purchased. MTFAA (Tokyo Chemical Industry Co., Ltd., > 98.0 %) and MAA (Tokyo Chemical Industry Co., Ltd., > 99.0 %) were dried over the dried molecular sieves for one day, as necessary, while being melted on a hotplate at 343 K and 323 K, respectively. Subsequently, the molecular sieves were removed, and the resulting MTFAA and MAA were sublimated at 353 K under reduced Ar atmosphere (≈ 0.01 atm) in a sealed glass container. Silver trifluoromethanesulfonate (Sigma Aldrich Co. LLC, ≥ 99 %) was used as purchased, and 1-methyl-1-propylpyrrolidinium bis(trifluoromethanesulfonyl)amide was prepared as mentioned in a previous study.²⁶

F-EEs composed of [TMA]F and MTFAA/MAA were prepared by mixing the two compounds in the designated ratios and stirring them in the glove box at room temperature.

Phase transition behavior was investigated by differential scanning calorimeter (SII NanoTechnology Inc., DSC6220) at a scan rate of 5 K min⁻¹ in sealed Al pans. Melting points were determined based on intersection points of baselines and tangents of endothermic peaks. Infrared (IR) spectra were recorded using Cary 630 (Agilent Technologies) by the single reflection attenuated total reflection method under the dry air atmosphere. Densities were estimated at 298 K by weighing samples in 1 mL of a volumetric cylinder whose volume had been calibrated with ultrapure water beforehand. Viscosities were measured using an electro-magnetically spinning viscometer (EMS-1000S, Kyoto Electronics Manufacturing Co., Ltd.).

Electrochemical measurements were performed using an electrochemical analyzer (HZ-7000, Meiden Hokuto Corp.). Ionic conductivities of the electrolytes were determined by electrochemical impedance spectroscopy with an AC perturbation of 10 mV in the frequency range of 1 kHz–200 kHz. An airtight four-probe conductivity cell consisting of two inner platinum wire electrodes (1.0 mm diameter) for monitoring the potential difference and two outer platinum disk-electrodes (10.0 mm diameter) was used. The ionic conductivities of samples (σ) were obtained based on an equation: σ (mS cm⁻¹) = 1000 × k (cm⁻¹)/ρ (ohm), where k is the cell constant and ρ is the solution resistance. The cell constant was obtained by measuring the conductivity of a standard KCl aqueous solution (Kishida Chemical Co., Ltd., 1.41 mS cm⁻¹ at 298 K), and the solution resistances were determined by estimating where data at the high frequencies intercepts the real resistance axis in Nyquist plots. Linear sweep voltammograms (LSVs) were obtained at room temperature at a scan rate of 1 mV s⁻¹ in the glove box with the aid of the electrochemical analyzer in a three-electrode conic cell made of polypropylene. Oxidation and reduction limit potentials were determined based on the current density of 0.1 mA cm⁻². Pt disk electrodes (diameter: 0.3 cm and surface area: 0.0707 cm²) were used as a working electrode, and Pt wire electrode (surface area: 2.35 cm²) was used as a counter electrode, respectively. The surface of the working electrodes was polished using alumina suspension before measurements. A reference electrode comprised Ag wire immersed in 0.1 mol dm⁻³ silver trifluoromethanesulfonate in 1-methyl-1-propylpyrrolidinium bis(trifluoromethanesulfonyl)amide, which was separated from the targeted electrolytes by a porous glass frit.

3. Results and Discussion

The F-EEs [TMA]F·x[MTFAA] (x: the molar ratio of MTFAA to [TMA]F) were prepared by mixing them in various ratios. The presence of solid precipitates in the liquid was visually confirmed for x = 9.0, whereas increasing [TMA]F corresponding to x = 8.0 and 4.0 resulted in the formation of transparent and colorless liquid products. Too much increase of [TMA]F corresponding to x = 2.0 formed a product containing solids. Figure 2 shows DSC curves of MTFAA and [TMA]F·x[MTFAA] (x = 8.0, 4.0, and 2.0). The product with x = 9.0 was not measured because the liquid and solid were separated heterogeneously. The melting point of [TMA]F·8.0[MTFAA] is 279 K, which is lower than 323 K for that of MTFAA and 443 K for the decomposition temperature of [TMA]F,²⁷ indicating that a liquid F-EE is obtained at room temperature by melting-point depression. On the other hand, [TMA]F·4.0[MTFAA] shows only the glass transition at −200 K without endothermic peaks for determining its melting point although it was confirmed to be liquid visually with the density of 1.32 g cm⁻³ and the F⁻ concentration of 2.2 mol dm⁻³ at 298 K. For [TMA]F·2.0[MTFAA], two endothermic peaks are observed at 309 K and 320 K above room temperature. Thus, the F-EEs [TMA]F·x[MTFAA] were obtained with x = 8.0 and 4.0.

Figure 2.

DSC curves of (i) MTFAA, (ii) [TMA]F·8.0[MTFAA], (iii) [TMA]F·4.0[MTFAA], (iii) [TMA]F·2.0[MTFAA], (iv) MAA, and (v) [TMA]F·4.0[MAA] at a scan rate of 5 K min⁻¹ measured by using sealed Al pans.

Figure 3(i–iv) shows IR spectra of [TMA]F, MTFAA, and [TMA]F·x[MTFAA] (x = 8.0 and 4.0). A band ascribed to N–H stretching mode is observed around 3300 cm⁻¹ in the spectrum of pure MTFAA.²⁸ As shown in the spectra of the [TMA]F·x[MTFAA], its band intensity is weakened by combining with [TMA]F and a broad band appears in the broad range of 2500–3100 cm⁻¹. Such behavior was observed for [TMA]F·3.5[1,3-DMU] and [TMA]F·x[1,3-DMTU] (1,3-DMTU: 1,3-dimethylthiourea, x = 6.0, 3.5, and 2.0) in the previous study, where bands attributed to N–H stretching mode of 1,3-DMU and 1,3-DMTU shifted to lower wavenumbers by combining with [TMA]F due to NH···F⁻ hydrogen bonds (HBs).²³ This demonstrates that NH···F⁻ HBs are formed in [TMA]F·x[MTFAA], which results in breaking the intermolecular HBs in MTFAA and the melting-point depression. For [TMA]F·8.0[MTFAA], a band attributed to N–H stretching mode also remains at the same wavenumber as that for MTFAA, whereas the band for [TMA]F·4.0[MTFAA] is distributed from lower wavenumbers to slightly higher wavenumbers than that for pure MTFAA. The blue-shift could reflect the strengthened N–H bonds caused by weakened N–H···O=C HBs between different MTFAA molecules by the increased F⁻ ratio,²⁸ and such behavior was also observed for [TMA]F·x[1,3-DMTU] in the previous study.²³

Figure 3.

IR spectra of (i) [TMA]F, (ii) MTFAA, (iii) [TMA]F·8.0[MTFAA], (iv) [TMA]F·4.0[MTFAA], (v) MAA, and (vi) [TMA]F·4.0[MAA].

For the comparison of the red-shift due to the NH···F⁻ HBs, another F-EE of [TMA]F·4.0[MAA] was prepared using MAA with the chemical structure where CF₃ in MTFAA is replaced by CH₃ as shown in Fig. 1. It was visually confirmed to be liquid with the density of 0.98 g cm⁻³ and the F⁻ concentration of 2.5 mol dm⁻³ at 298 K by melting-point depression from that of MAA (302 K) although no endothermic peaks are observed in its DSC curve in Fig. 2. As shown in Fig. 3(v), a band ascribed to N–H stretching mode is observed around 3300 cm⁻¹ in the spectrum of MAA.²⁸ After combining with [TMA]F, a broad peak corresponding to NH···F⁻ appears around 2900 cm⁻¹ for [TMA]F·4.0[MAA], demonstrating that the red-shift by NH···F⁻ HBs is larger for MTFAA than MAA. Such the larger red-shift for fluorinated HBDs agrees with the previous study about the comparison of the red-shift of O–H stretching mode by OH···F⁻ between 2-propanol and 1,1,1,3,3,3-hexafluoro-2-propanol.²⁴

Figure 4a shows the viscosities of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA]. Those of [TMA]F·8.0[MTFAA] and [TMA]F·4.0[MTFAA] are 16 mPa s and 65 mPa s at 298 K, respectively, indicating that the viscosity is increased by increased F⁻ ratio. [TMA]F·4.0[MAA] has the viscosity of 59 mPa s at 298 K, which is similar to [TMA]F·4.0[MTFAA]. The activation energy E_η can be obtained using the following Arrhenius equation:

  

\begin{equation} \eta = \eta_{0}\exp (E_{\eta }/\text{R}T) \end{equation} (1)

where η is the viscosity, η₀ is a constant related to the viscosity at infinite temperature, T is the absolute temperature, and R is the gas constant. The E_η of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] in 298–323 K are 28 kJ mol⁻¹, 41 kJ mol⁻¹, and 37 kJ mol⁻¹, respectively. The lower F⁻ ratio provides lower E_η and there is not significant difference between E_η of F-EEs with MTFAA and MAA in the same molar ratio. However, the viscosities of the F-EEs are 1 orders and 3–4 orders of magnitude lower than those of [TMA]F·3.5[1,3-DMU] (η = 4.4 × 10² mPa s at 298 K and E_η = 75.6 kJ mol⁻¹ in 293–313 K) and [TMA]F·3.5[1,3-DMTU] (η = 1.0 × 10⁵ mPa s at 298 K and E_η = 176 kJ mol⁻¹ in 293–313 K) in the previous study,²³ respectively, as well as the much smaller activation energies although they are not at exactly the same molar ratio. These derive from different chemical structures of the amide compounds as discussed in the following discussion.

Figure 4.

(a) Viscosities and (b) ionic conductivities of (i) [TMA]F·8.0[MTFAA], (ii) [TMA]F·4.0[MTFAA], and (iii) [TMA]F·4.0[MAA]. The lines represent fitting results calculated based on the Arrhenius equation. Coefficients of determination for (a-i), (a-ii), (a-iii), (b-i), (b-ii), and (b-iii) are 0.998, 0.999, 0.998, 0.994, 0.999, and 0.999, respectively. (c) Walden plots of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] (closed symbols) as well as those of [TMA]F·3.5[1,3-DMU], [TMA]F·6.0[1,3-DMTU], and [TMA]F·3.5[1,3-DMTU] in the previous study (open symbols).²³

Figure 4b shows the ionic conductivities of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA]. The ionic conductivities of [TMA]F·8.0[MTFAA] and [TMA]F·4.0[MTFAA] are 2.7 mS cm⁻¹ and 1.4 mS cm⁻¹ at 298 K, respectively, indicating that the ionic conductivity is increased by the decreased F⁻ ratio. That of [TMA]F·4.0[MAA] is 2.0 mS cm⁻¹ at 298 K, indicating that MAA provides the similar ionic conductivity to that for MTFAA to the resulting F-EEs although the value for the F-EE with MAA is slightly higher in the same molar ratio. The activation energy E_σ can be obtained using the following Arrhenius equation:

  

\begin{equation} \sigma = \sigma_{0}\exp (-E_{\sigma }/\text{R}T) \end{equation} (2)

where σ is the ionic conductivity, σ₀ is a constant related to the ionic conductivity at infinite temperature. The E_σ of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] in 298–323 K are 27 kJ mol⁻¹, 38 kJ mol⁻¹, and 35 kJ mol⁻¹, respectively. The lower F⁻ ratio provides lower E_σ similarly to the viscosity, which suggests that the increase of the neutral species moderates the Coulombic interactions between cations and anions and improves mobilities of the dissolved species. There are not significant differences between E_σ of F-EEs with MTFAA and MAA in the same molar ratio; however, the ionic conductivities of the F-EEs are 1 order and 3 orders of magnitude higher than those of [TMA]F·3.5[1,3-DMU] (σ = 0.42 mS cm⁻¹ at 298 K and E_σ = 55.2 kJ mol⁻¹ in 293–333 K) and [TMA]F·3.5[1,3-DMTU] (η = 3.9 × 10⁻³ mPa s at 298 K and E_σ = 155 kJ mol⁻¹ in 293–313 K) in the previous study,²³ as well as the much smaller activation energies. In [TMA]F·x[1,3-DMU] and [TMA]F·x[1,3-DMU], F⁻ is strongly stabilized with two N–H bonds.²³ Extensive HB network is formed in [TMA]F·x[1,3-DMTU] due to cis-trans conformation of 1,3-DMTU with regard to the position of two N–H bonds, and F⁻ is stabilized locally in [TMA]F·x[1,3-DMU] by HBs with two N–H bonds of 1,3-DMU with trans-trans conformation although the HB network is limited. For [TMA]F·x[MTFAA] and [TMA]F·x[MAA], on the other hand, MTFAA and MAA have only one N–H bond per molecule; therefore, the HB network in them is limited and the dissociation of F⁻ from the HBs is facilitated owing the moderate interactions, leading to the improved physicochemical properties. Differences in the dissolved state of F⁻ attributed to the different numbers and positions of N–H bonds in an amide compound have more significant influences on the physicochemical properties and their activation energies rather than the different strengths of each NH···F⁻ HB based on the groups of CF₃ or CH₃ in the amides.

The obtained viscosities and molar ionic conductivities were plotted based on the Walden plot. Walden rule (or fractional Walden rule) is shown as below:^29–32

  

\begin{equation} \eta {\cdot} \varLambda = \text{const.}\ (\text{or}\ \eta^{\alpha }{\cdot} \varLambda = \text{const.}) \end{equation} (3)

where η is viscosity, Λ is molar ionic conductivity, and α is the decoupling constant corresponding to the Walden plot gradient. The Walden plots show the charge mobility against the frictional resistance. Figure 4c shows the Walden plots of the F-EEs with MTFAA and MAA as well as those with 1,3-DMU and 1,3-DMTU in the previous study at 298 K.²³ The dotted line is an ideal line corresponding to 1 mol dm⁻³ KCl aqueous solution where each ion behaves as a charged ion without forming ion pairing due to the full dissociation of the salt.^29,33 The plots for [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] are slightly below the ideal line in the similar distance, suggesting that there would not be significant differences on the dissociation states by the different HBD ratios and the functional groups of CF₃ or CH₃ in the amide compounds. Those for [TMA]F·3.5[1,3-DMU], [TMA]F·6.0[1,3-DMTU], and [TMA]F·3.5[1,3-DMTU], where 1,3-DMU and 1,3-DMTU have two N–H bonds per molecule, are closer to or slightly above the ideal line compared to [TMA]F·x[1,3-DMU] and [TMA]F·x[1,3-DMTU]; however, their difference is not very significant and the difference in viscosities and ionic conductivities by the different numbers of N–H bonds per molecule is rather noticeable.

Figure 5 shows LSVs of [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA]. Reduction limit potentials for [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] are −1.93 V, −1.99 V, and −2.68 V vs. Ag⁺/Ag, respectively, based on the current density of 0.1 mA cm⁻². The reduction of MAA and MTFAA would proceed at the reduction limit potentials because the reduction of TMA⁺ is not considered to occur at these potentials according to the more negative reduction limit potential of −2.95 V vs. Ag⁺/Ag (determined based on ±0.05 mA cm⁻²) for [TMA]F·3.5[1,3-DMU].²³ [TMA]F·8.0[MTFAA] and [TMA]F·4.0[MTFAA] exhibit the similar values, indicating that there is not significant difference by the different compositions. The comparison of those for [TMA]F·4.0[MTFAA] and [TMA]F·4.0[MAA] indicates that MAA provides better electrochemical reduction stability to the resulting F-EEs than MTFAA. This would be attributed to the lower lowest unoccupied molecular orbital (LUMO) level of MTFAA than that of MAA caused by the substitution of CF₃ for CH₃.^34,35 Meanwhile, oxidation limit potentials for [TMA]F·8.0[MTFAA], [TMA]F·4.0[MTFAA], and [TMA]F·4.0[MAA] are +0.34 V, +0.22 V, and −0.04 V vs. Ag⁺/Ag, respectively. [TMA]F·8.0[MTFAA] and [TMA]F·4.0[MTFAA] exhibit the close values despite the different compositions similarly to the result for the reduction limit potentials. This could suggest that there is not significant difference in species and bonds determining the reduction/oxidation limits in this composition range, which could be characterized by computational chemical approaches in future studies. [TMA]F·4.0[MTFAA] exhibits the higher oxidation stability than [TMA]F·4.0[MAA] because of the higher oxidation stability of MTFAA than that of MAA resulting from the lowered highest occupied molecular orbital (HOMO) level attributed to CF₃ for MTFAA instead of CH₃ for MAA.^34,35 The oxidation limit potentials for [TMA]F·x[MTFAA] are as high as that for [TMA]F·3.5[1,3-DMU] corresponding to a high potential of approximately 2 V vs. PbF₂/Pb.²³ It demonstrates the superior electrochemical oxidation stability of [TMA]F·x[MTFAA] enough to be used as electrolytes for fluorinating noble metals and organic compounds.

Figure 5.

LSVs of the Pt disc electrode in (i) [TMA]F·8.0[MTFAA], (ii) [TMA]F·4.0[MTFAA], and (iii) [TMA]F·4.0[MAA] at a scan rate of 1 mV s⁻¹.

4. Conclusions

In this study, F-EEs composed of [TMA]F and MTFAA were developed, and their physicochemical properties were evaluated including comparison with that composed of [TMA]F and MAA. The formation of F-EEs which were in liquid state at room temperature was confirmed for [TMA]F·x[MTFAA] with x = 8.0 and 4.0 by melting-point depression. IR spectroscopy revealed that HBs of NH···F⁻ caused the melting point depression and the HBs are stronger for MTFAA than MAA. The viscosity and ionic conductivity measurements showed that MTFAA and MAA of amide compounds with one N–H bond per molecule provided better ion transport properties to the resulting F-EEs than 1,3-DMU with two N–H bonds per molecule by comparison with the previous study.²³ The LSVs demonstrated that the F-EEs with MTFAA possessed superior electrochemical oxidation stability owing to the CF₃ group compared to that with MAA and would be promising as electrolytes for electrochemical fluorination applications in future studies.

Acknowledgment

This research was financially supported by Japan Society for the Promotion of Science (JSPS) KAKENHI Grant Numbers 21K20564 (H. Y.) and 23K13834 (H. Y.).

CRediT Authorship Contribution Statement

Hiroki Yamamoto: Conceptualization (Lead), Data curation (Lead), Funding acquisition (Lead), Investigation (Lead), Resources (Lead), Writing – original draft (Lead), Writing – review & editing (Lead)

Kazuki Yoshii: Resources (Equal), Writing – review & editing (Supporting)

Conflict of Interest

The authors declare no competing financial interests.

Funding

Japan Society for the Promotion of Science: 21K20564

Japan Society for the Promotion of Science: 23K13834

Footnotes

H. Yamamoto and K. Yoshii: ECSJ Active Members

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