Article pubs.acs.org/JPCC
Structural and Electrochemical Properties of Li Ion Solvation Complexes in the Salt-Concentrated Electrolytes Using an Aprotic Donor Solvent, N,N‑Dimethylformamide Kenta Fujii,* Hideaki Wakamatsu, Yanko Todorov, Nobuko Yoshimoto, and Masayuki Morita Graduate School of Sciences and Technology for Innovation, Yamaguchi University, 2-16-1 Tokiwadai, Ube, Yamaguchi 755-8611, Japan S Supporting Information *
ABSTRACT: We report the relation between the structural and electrochemical properties of N,N-dimethylformamide (DMF)-based electrolytes containing lithium bis(trifluoromethanesulfonyl)amide (LiTFSA) in the concentration range cLi = 0−3.2 mol dm−3. Raman spectroscopy and DFT calculations indicate that Li+ ions are solvated by DMF molecules in the form of [Li(DMF)4]+ complexes at low cLi ( 2.5 mol dm−3. The high cLi solutions, in which all the DMF molecules solvate to Li+ ions (i.e., no DMF remains in the bulk), make the electrochemical window wider; the oxidative stability is enhanced owing to lower HOMO energy levels of solvated DMF molecules relative to those in the bulk. The salt concentration also controls the reductive stability; coordinated TFSA− anions within the Li-ion complexes formed in concentrated solutions affect the LUMO energy levels of the electrolyte. The LUMOs located on the TFSA− anions lead to a preferential reduction of the TFSA component rather than DMF to form a solid electrolyte interphase on graphite negative electrodes, resulting in the Li-ion insertion/desertion into/from graphite in the concentrated solutions.
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INTRODUCTION Electrolytes employing nonaqueous solvents are one of the key materials, as well as the electrode materials, required for the development of more stable, safer, and higher voltage lithiumion batteries (LIBs).1,2 It is well established that in conventional LIB systems with a popular graphite negative electrode, using carbonate-based solvents such as ethylene carbonate (EC) and dimethyl carbonate is necessary for graphite to perform its function reversibly and stably. This is because the reductive decomposition of the carbonate component (particularly, EC) occurs on the graphite anode during the first charging process to form a so-called solid−electrolyte interphase (SEI) film,3 which plays an essential role in reversible Li+ insertion into graphite.4−6 Hence, design and functionalization of electrolyte materials are still limited to the use of carbonate solvents, and there has been no breakthrough in commercial LIB electrolyte systems, in contrast to the active progress in positive/negative electrode materials. Recently, a new concept of functional electrolyte in practical LIBs has been proposed, i.e., a carbonate solvent-free electrolyte solution. The key of the concept is the high Li-salt concentration in the electrolyte solution, that is, solvent-in-salt and the choice of counteranion composing of Li salt.7−13 © XXXX American Chemical Society
Yamada et al. reported that highly concentrated solutions of lithium bis(trifluoromethanesulfonyl)amide (LiTFSA) in acetonitrile (AN) exhibit excellent electrochemical reductive stability, allowing for reversible Li-ion insertion into graphite electrodes without the need for EC or other carbonate solvents.9,14 Furthermore, they proposed that Li salts including the bis(fluorosulfonyl)amide (FSA) anion15−17 give highly concentrated electrolytes with relatively low viscosity due to the high ionic dissociation ability of LiFSA, enabling the development of superfast-charging LIBs using concentrated LiFSA electrolytes.8−10 The unusual electrochemical phenomena in LIB chemistry originate from the specific solution structure in the highly concentrated electrolyte system: (1) All of the solvent molecules solvate Li+ ions; in other words, there are no solvent molecules in the bulk. (2) Anion species, which are generally “noncoordinating anions” in solutions, such as TFSA− and FSA−, coordinate with Li+ ions to form a polymeric network-like structure comprising Li+ ions, solvents, and anions. (3) The coordinated anion (TFSA−or FSA−) is Received: May 5, 2016 Revised: July 23, 2016
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DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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The Journal of Physical Chemistry C
The sample solution was filled in a quartz cell. The obtained Raman spectra were deconvoluted using a nonlinear leastsquares curve-fitting procedure to extract single bands. A single band is assumed to be represented as a pseudo-Voigt function, f V(ν) = γf L(ν) + (1 − γ)f G(ν), where f L(ν) and f G(ν) represent the Lorentzian and Gaussian components, respectively, and the parameter γ (0 < γ < 1) represents the Lorentzian component. The intensity I of a single band is evaluated according to I = γIL + (1 − γ)IG, where IL and IG denote the integrated intensities of the Lorentzian and Gaussian components, respectively. The integrated intensity of DMF in the bulk (free DMF) is represented as If = Jfcf, where Jf and cf stand for the Raman scattering coefficient and concentration, respectively, of DMF. By considering the mass balance equations cf = cT − cb = cT − ncLi, where cT, cb, and cLi denote the concentrations of total DMF, DMF bound to Li+ ions (bound DMF), and Li+ ions, respectively, and n is the solvation number of DMF molecules around Li+ ions, we obtain the following equation:
preferentially reduced and decomposed on the graphite electrode during the first charging process to form a stable SEI on the electrode, though this is not the case for the corresponding dilute system in which decomposition of solvent molecules and SEI formation do not occur. The anion-based SEI shows high electrochemical resistance to reductive decomposition during further charging cycles. Watanabe et al. also reported a systematic study of a concentrated electrolyte system based on oligoether-type solvents, glymes.18−23 The highly concentrated Li salt/glyme electrolytes (molar ratio of Li salt:glyme = 1:1) are in the liquid state at room temperature in the form of solvated ionic liquids, [Li(glyme)][X] (X: anions such as TFSA) and can be used as electrolytes for 4 V-class LIBs. In this system, they pointed out that reversible Li-ion insertion into graphite in the electrolyte system is mainly dominated by the desolvation process at the graphite/ electrolyte interface, not by the SEI formation due to reductive decomposition of anions coordinated around Li+ ions.21,23 The Li-ion insertion into graphite electrodes has also been reported for concentrated solution systems using other nonaqueous solvents.7,24−28 However, knowledge of the mechanism of the Li-ion insertion/desertion process is still limited at present; it is not found whether SEI formation from the coordinated anions around Li+ ions is indispensable. In this study, the solution structures of dilute to highly concentrated nonaqueous LiTFSA electrolytes were investigated at the molecular level. In particular, we focused on two issues: (1) how the Li-ion solvation complexes depend on the salt concentration and (2) the relationship between the structural and electrochemical properties of their complexes. We used N,N-dimethylformamide (DMF) as a model solvent, which is a typical aprotic donor solvent. DMF has a much larger electron pair donating ability: Gutmann’s donor number DN = 26.6 kcal mol−1 among the aprotic donor solvents,29,30 whereas TFSA− is well-known as an essentially noncoordinating anion relative to the popularly used BF4−, PF6−, and halide anions. This large difference in coordination ability between the solvent and the counteranion allows us to easily control the solvation environment of Li+ ions in the solutions by adjusting the molar ratio of LiTFSA to DMF (i.e., LiTFSA concentration, cLi). We might expect that Li+ ions are completely solvated by DMF molecules (as opposed to being paired with the noncoordinating TFSA− anions) in the low cLi solutions (2.0 mol dm−3), resulting in an improvement in the electrochemical performance of the electrolyte for LIBs.
If /c T = −nJf (c Li /c T) + Jf
(1)
Plots of If/cT against cLi/cT yield a straight line with a slope α = −nJf and an intercept β = Jf, and the n value is thus obtained from n = −α/β. The detailed procedure for the analyses is described elsewhere.31 The same procedure was also applied to the analysis of the coordination of TFSA− anions with Li+ ions by assuming free and bound TFSA components in the electrolyte solutions. To quantitatively determine the n and J values, the measured Raman spectra were normalized by using the band intensities at 1339.0 cm−1 of the TFSA− as an internal standard (see Figure S1 in the Supporting Information). Ionic Conductivity and Viscosity. The ionic conductivity was measured by an ac impedance method in the frequency range from 100 kHz to 10 mHz using a frequency response analyzer (Solartron 1260). A cell with platinum electrodes was used for the measurements. The cell constant was determined by using KCl solution (0.01, 0.1, and 1.0 mol dm−3). The viscosity was measured using a DV-I Prime viscometer (Brookfield). Both ionic conductivity and viscosity measurements were conducted at 298 K. Electrochemical Measurements. The electrochemical windows of the electrolytes were determined by linear sweep voltammetry (LSV, HZ-5000; Hokuto Denko) on a Pt working electrode with Li foil as both counter and reference electrodes. The LSVs were performed by scanning the cell voltage from the open circuit potential toward more negative or positive potentials at a scan rate of 5.0 mV s−1. The measurements with cathodic and anodic scans were performed in the difference cells. Cyclic voltammetry (CV, HZ-5000; Hokuto Denko) measurements were conducted using a conventional three-electrode cell with graphite as the working electrode and Li foil as the counter and the reference electrodes. The scan rate was 0.1 mV s−1. For preparation of the graphite (TIMREX KS6, TIMCAL) electrode, N-methylpyrrolidone (NMP) was added to a mixture of graphite and a poly(vinylidene difluoride) (PVdF) binder in a mass ratio of 9:1. Here, note that in the highly concentrated electrolyte solutions, the electrode potential of reference electrode (Li metal) strongly depends on the salt concentration; the activity of the free solvent in the bulk is very low, which was reported by Dokko et al.21 Therefore, it is plausible that the potential V used in this study should be considered as the apparent value; V vs Li/Li+ in cLi solution (cLi = 1.0, 2.5, and 3.2 mol dm−3).
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EXPERIMENTAL SECTION Materials. LiTFSA salt (Kanto Chemical, Battery grade) was vacuum-dried at 100 °C for 100 h. Solvent DMF (Wako Chemicals, Battery grade) was used without further purification, and the water content was determined to be less than 50 ppm by Karl Fischer titration. The sample electrolyte solutions were prepared at the required molarities cLi in an Ar-filled glovebox. Raman Spectroscopy. Raman spectroscopic measurements were conducted at room temperature using an FTRaman/IR spectrometer (JASCO, FT/IR-6100) equipped with an Nd:YAG laser (1064 nm). Raman spectra were obtained with an optical resolution of 4.0 cm−1 and were accumulated 1024 times to achieve a sufficiently high signal-to-noise ratio. B
DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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The Journal of Physical Chemistry C DFT Calculations. Density functional theory (DFT) calculations were conducted using the Gaussian09 software package.32 The geometries of isolated DMF and TFSA (cis and trans conformers33,34) and Li-ion complexes were fully optimized at the B3LYP/6-311G** level, followed by normal frequency analyses.
shifted toward higher frequency upon further increasing cLi (solid red lines). This strongly suggests that the solvation structure of Li+ ions in the solutions changes at approximately cLi = 2.0 mol dm−3, which will be discussed in detail in the later sections. We performed a curve-fitting analysis of the observed Raman spectra to deconvolute them into their constituent single bands. The typical result for cLi = 1.0 mol dm−3 is shown in Figure 2a. The Raman spectra for cLi = 0 (neat DMF) to ∼2.0 mol dm−3 could be deconvoluted into two bands at 663.3 and 675.9 cm−1, which originate from free and bound DMF components. We can estimate the solvation number of the Li+ ions (nDMF) by analyzing the integrated intensity of the 663.3 cm−1 band (free DMF) as a function of cLi. Figure 2b shows a plot of If/cT against cLi/cT based on eq 1 described in the Experimental Section. For cLi ≤ 2.0 mol dm−3, the plot produced a straight line with a slope (−nJf) and intercept (Jf), which enables us to estimate the nDMF value according to n = −slope/intercept, yielding nDMF = 3.9 ± 0.1. This indicates that four DMF molecules solvate the Li+ ion to form an [Li(DMF)4]+ complex at lower cLi solutions ( 2.0 mol dm−3, resulting in no free solvent molecules in the bulk. In this case, it is plausible that the HOMO is located on the solvated DMF molecules (or TFSA− anions) around the Li+ ions, which is essentially different from those in the bulk owing to polarization induced by the strong electrostatic field around the Li+ ion. Therefore, we estimated the HOMO energy levels of DMF molecules in possible Li-ion complexes (complexes 1a−g and 2a−g in Figure S4) using DFT calculations; the results are listed in Table S2. For all complexes examined here, the HOMO energy levels tend to be in the following order: isolated DMF > mononuclear-type [Li(DMF)n(TFSA)m] > mononuclear [Li(DMF)4]+ ≈ binuclear-type [Li2(DMF)n(TFSA)m]. The values for DMF bound to Li+ ions are appreciably lower than that for isolated DMF. It is thus plausible that the oxidation current flowing in the range 2.8−4.0 V in lower cLi solutions (cLi = 1.0 mol dm−3) during LSV measurements originates from the decomposition of DMF molecules in the bulk and that at >4.0 V results from the [Li(DMF)4]+ complex. In the concentrated solutions (cLi = 2.5 and 3.2 mol dm−3), DMF molecules do not exist in the bulk; all of the DMF molecules bind to Li+ ions to form Li-ion complexes containing both DMF and TFSA. In the present LSV results, no oxidative decomposition occurs up to 4.8 V, implying that the Li-ion complexes formed in the concentrated solutions are more oxidatively stable than the [Li(DMF)4]+ complexes in dilute solutions. Dokko et al. reported that the coordination of TFSA− anions to Li+ ions (contact ion-pair formation) in the concentrated LiTFSA/ tetraglyme (G4) system increases the HOMO energy levels of the solvated [Li(G4)]+ complex.22 This is because the negative charge of the anion weakens the electric field around the metal ion. Indeed, in this work, the value for the mononuclear-type [Li(DMF)n(TFSA)m] is higher than that for [Li(DMF)4]+ (Table S2). In contrast, the binuclear-type [Li2(DMF)n(TFSA)m] examined here has a HOMO energy comparable with [Li(DMF)4]+. This result suggests that the formation of multinuclear complexes in the concentrated solutions plays a key role in enhancing the oxidative stability and widening the electrochemical window of electrolytes. In the present work, it was difficult to investigate multinuclear Li-ion complexes with molecular orbital calculations because there are too many possible combinations of complexes including Li+ ions, DMF, and TFSA. Studies of the concentrated systems using molecular dynamics simulations43 are needed to obtain detailed information about multinuclear complexes (> binuclear). With regard to the reductive stability, in Figure 7 (inset), it was found that the slight reductive currents in the range 0.4−0.8 V were observed in the concentrated solutions (cLi = 2.5 and 3.2 mol dm−3), whereas there was no current in the potential range in the diluted solution cLi = 1.0 mol dm−3. The result can be interpreted in terms of the LUMO energy levels of the Li ion complexes listed in Table S2: isolated DMF, [Li(DMF)4]+ > binuclear-type [Li2(DMF)n(TFSA)m], which is discussed later in detail (reductive decomposition on graphite negative electrode).
Figure 8. cLi dependences of ionic conductivity (filled circles) and viscosity (open circles) for LiTFSA/DMF solutions at 298 K.
exhibited a maximum at cLi = 1.0 mol dm−3, followed by a decrease in the ionic conductivity upon further increasing cLi. The η gradually increased with increasing cLi. The σ increase up to cLi = 1.0 mol dm−3 is due to an increase in ion concentration. In contrast, the decrease in σ observed for cLi > 1.0 mol dm−3 may be caused by a decrease in both mobility and concentration of ionic species owing to the formation of contact ion pairs or multinuclear Li complexes. Figure 9 shows
Figure 9. CVs for a graphite electrode in LiTFSA/DMF solutions at cLi = (a) 1.0, (b) 2.5, and (c) 3.2 mol dm−3.
typical CV profiles of a graphite electrode in cLi = (a) 1.0, (b) 2.5, and (c) 3.2 mol dm−3 electrolyte solutions. For the cLi = 1.0 mol dm−3 solution (Figure 9a), no significant redox peak was observed in the voltage range 0−0.3 V (vs Li/Li+ in cLi solution), although the ionic conductivity was the highest among the electrolytes examined here. This CV behavior is ascribed to there being no insertion/desertion of Li+ ions into/ from the graphite negative electrode. A weak and broad current was observed at approximately 0.5 V during the cathodic scan, F
DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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The Journal of Physical Chemistry C
Figure 10. (a) CVs for a graphite electrode over the limited range 0.3−2.8 V and (b) XRD patterns of a pristine graphite electrode and an electrode charged at 0.3 V in cLi = 2.5 mol dm−3 electrolyte solution.
Figure 11. LUMOs located on the optimized geometries of [Li(DMF)2(TFSA)1] and [Li2(DMF)2(TFSA)1]+ complexes.
solutions ([Li(G3)x][TFSA]: x = 1−4) giving rise to a redox peak at 0−0.3 V and (2) co-intercalation of solvated Li+ ions, that is, [Li(G3)1]+ complexes, occurs in solutions with x > 1.5, yielding a cathodic current at approximately 0.8 V. The experimental evidence for (2) was clearly provided by CV measurements in the narrow voltage range (1.5−0.4 V) and also by XRD measurements. To confirm the origin of the 0.5 V peak observed in this work, additional CV (in the narrow voltage range of 2.8−0.3 V) and XRD measurements were conducted, which are shown in Figures 10a and 10b, respectively. As can be seen in Figure 10a, the graphite negative electrode showed irreversible behavior; a small reductive current was observed, which is similar to that in Figure 9a (wide voltage range 2.8−0 V), whereas no oxidative current was observed during the anodic scan. This behavior is quite different from that exhibited in the LiTFSA/G3 system mentioned above; reductive and oxidative currents are observed in CVs during cathodic and anodic scans, respectively, in a narrow voltage range. Therefore, the 0.5 V peak observed for the LiTFSA/DMF system (this work) is not likely to be the coinsertion of solvated Li+ ion species into graphite. Figure 10b shows XRD patterns of a pristine graphite electrode and an electrode charged at 0.3 V. It is clear that there is no change in the (002) peak at 28.88° for both pristine graphite and graphite charged at 0.3 V. This indicates that the interlayer distance of graphite remains unchanged at 0.3 V, meaning that no coinsertion of solvated Li+ ion species could have occurred. From these results, we conclude that the current peak appearing in the present CVs in the range 0.3−0.5 V mainly originates from the decomposition of the DMF component of the solutions.
which is a result of the decomposition of the DMF component in the solution, as discussed in the subsequent sections. In the concentrated cLi solutions (2.5 and 3.2 mol dm−3), a significant reductive current was observed in the range 0.3−0 V during the cathodic scan, which originates from Li-ion insertion into the graphite, and an oxidative current due to Li-ion desertion was observed in the range 0−0.45 V during the anodic scan. The reductive current appreciably decreased during the second cycle compared with the first cycle, but no change in the CV profile was observed for subsequent cycles. The same was also observed in the LiTFSA/acetonitrile system: in the highly concentrated solution, the reductive current decreased after the second cycles (see CV data in Figure S5). This decrease in the reductive current after the second cycles is ascribed to a stable SEI formation on the graphite electrode in the first cycle, resulting in a suppression of further reductive decomposition. As described later, we proposed on the basis of DFT calculations that the SEI formed in the concentrated electrolytes originates from the reductive decomposition of the TFSA species bound to Li ion. For the cLi = 2.5 mol dm−3 solution (Figure 9b), small reductive currents appear in the range 1.5−0.4 V during the first cycle, a relatively sharp peak at 0.5 V, and a broad peak centered at 0.9 V, in addition to the Li-ion insertion/desertion process mentioned above. The same applies to the 3.2 mol dm−3 solution, although the 0.5 V peak seems to disappear. Here, we should note the previous study for the concentrated electrolyte system using triglyme (G3) reported by Dokko et al.21,44 They reported that (1) Li-ion insertion/desertion into/ from graphite electrodes occurs in LiTFSA/G3 electrolyte G
DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
The Journal of Physical Chemistry C
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SEI Formation by Reductive Decomposition of Bound TFSA. As mentioned in the Introduction, Yamada et al. reported a superconcentrated electrolyte for LIBs using acetonitrile (AN) and LiTFSA as a solvent and a Li salt, respectively.9 They pointed out that (1) the concentrated electrolyte solution (cLi > 4 mol dm−3), in which all of the AN molecules and TFSA− anions are bound to Li+ ions, enhances its reductive stability, (2) the LUMO energy levels of the TFSA− anions become lower than those of the AN molecules in the concentrated solution and the reverse is true for dilute solutions, and (3) the LUMO located on the TFSA leads to the predominate reduction reaction, rather than that on the AN, forming a stable TFSA-based SEI on graphite electrodes. Figure S5 shows the CV profile of the concentrated LiTFSA/AN solution (cLi = 3.7 mol dm−3) that was examined in this work. We confirmed that significant redox peaks corresponding to Li+ insertion/desertion into/from the graphite electrode are observed in the range 0−0.3 V as well as in the present LiTFSA/DMF system. Here, we note that a broad reductive current appears in the range 0.8−1.5 V, indicating the decomposition of the TFSA component in the concentrated electrolyte on the graphite electrode. On the basis of this experimental fact, we can conclude that the broad current peak observed at approximately 0.9 V for the concentrated DMFbased solutions (Figure 9b,c) is attributable to the decomposition of the TFSA− anions, resulting in SEI formation on the graphite. Figure 11 shows typical results of the LUMOs on (a) mononuclear [Li(DMF)(TFSA) 2 ] − and (b) binuclear [Li2(DMF)2(TFSA)2] complexes obtained by DFT calculations. It was found that the LUMOs are located on the DMF molecule in mononuclear [Li(DMF)(TFSA)2]−, whereas those on the TFSA− anion interacted with two Li+ ions in binuclear [Li2(DMF)2(TFSA)2]. As discussed above, a polymeric network-like structure based on multinuclear Li-ion complexes linked by anions and solvent molecules was suggested to form in concentrated AN-based solutions according to DFT-MD simulations.9,14 Thus, we can conclude that the complex structure in the concentrated solutions continuously changes from mononuclear to multinuclear as cLi is increased, and the multinuclear complex formation is important for shifting the LUMOs from DMF to TFSA.
Article
ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.6b04542. cLi, xLi, and density values (Table S1); HOMO/LUMO energy levels and binding energies of Li-ion complexes (Table S2); Raman intensities at 1339.0 cm−1 used as the normalization (Figure S1); theoretical Raman bands for Li+−DMF complexes (Figure S2); theoretical Raman bands for [Li(DMF)n(TFSA)m] complexes (Figure S3); optimized geometries of Li-ion complexes (Figure S4); CV profile for acetonitrile-based electrolyte solution (Figure S5) (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected] (K.F.). Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS
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REFERENCES
This work has been financially supported by Grant-in-Aids for Scientific Research from the Ministry of Education, Culture, Sports, Science, and Technology (No. 15K17877 to K.F.).
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CONCLUSIONS In this study, we investigated the structures of Li-ion complexes in dilute and highly concentrated DMF-based electrolyte solutions using Raman spectroscopy with the aid of DFT calculations. The coordination numbers of DMF and TFSA (nDMF and nTFSA) were estimated quantitatively in the solutions with cLi in the range 0−3.2 mol dm−3. In highly concentrated solutions, the TFSA− anion, which is well-known as a “noncoordinative anion” in conventional molecular solvents, coordinates to Li+ ions in both monodentate and bidentate modes, depending on cLi. Increasing the salt concentration to the point when no DMF molecules exist in the bulk (cLi > 2.0 mol dm−3) widens the electrochemical window of the electrolyte. We also observed that coordinated TFSA species in the concentrated solutions play a key role in the reversible Li-ion insertion/desertion into/from graphite electrodes; the reductive reaction on graphite electrodes that preferentially occurs for the coordinated TFSA around Li+ ions leads to the formation of a stable SEI, but this is not the case in dilute solutions in which reductive decomposition of the DMF component occurs. H
DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
Article
The Journal of Physical Chemistry C
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DOI: 10.1021/acs.jpcc.6b04542 J. Phys. Chem. C XXXX, XXX, XXX−XXX
