Article pubs.acs.org/JPCC
Role of Mixed Solvation and Ion Pairing in the Solution Structure of Lithium Ion Battery Electrolytes Daniel M. Seo, Stefanie Reininger, Mary Kutcher, Kaitlin Redmond, William B. Euler, and Brett L. Lucht* Department of Chemistry, University of Rhode Island, Kingston, Rhode Island 02881, United States S Supporting Information *
ABSTRACT: The solution structures of organic carbonate solvents (ethylene carbonate (EC), propylene carbonate (PC), dimethyl carbonate (DMC), and diethyl carbonate (DEC)) as electrolyte solutions of LiPF6 were investigated with FTIR and NMR spectroscopy and DFT computational methods. Both coordinated and uncoordinated solvents are observed by IR spectroscopy, allowing the determination of solvent coordination numbers, which a range from 2 to 5. The predominant species in solution changes as a function of LiPF6 concentration. At low salt concentrations (2.0 M) the predominant species in solution is the contact ion pair. In mixed solvent systems (PC−DMC, PC−DEC, EC−DMC, or EC−DEC), the mixed solvated cations are observed in the presence of high concentrations of uncoordinated cyclic carbonate despite the much larger dielectric constant of the cyclic carbonates than dielectric constant of linear carbonate. and ethyl methyl carbonate (EMC).22,23 The most widely used salt in lithium ion battery electrolytes is LiPF6 due to a high degree of dissociation, high conductivity, ability to passivate the Al current collector, and relatively low cost. However, LiPF6 is not ideal due to its poor thermal and hydrolytic stabilities.24 Herein, we report an investigation of molecular interactions of carbonate solutions of LiPF6 by FTIR and NMR spectroscopies aided by density functional theory (DFT) computational simulations. The results suggest that both solvent-separated ion pairs and contact ion pairs are present in the electrolyte solutions and that cooperative solvation effects may be more important than previously realized.
1. INTRODUCTION The development of novel electrode materials for advanced lithium ion batteries has increased interest in developing an understanding of electrolytes for lithium ion batteries.1−3 Optimization of lithium ion batteries requires appropriate electrolyte properties including conductivity, electrochemical stability window, thermal and hydrolytic stability, and ability to generate a stable anode solid electrolyte interface (SEI).4,5 The properties of the electrolyte are dependent upon the electrolyte solution structure. Thus, the development of a better understanding of the electrolyte solution structure is critical for the development of improved electrolytes for lithium ion batteries. The generation of a stable SEI is one of the most critical aspects of lithium ion battery electrolytes, and the electrolyte solution structure has been reported to be integral to this process.4−8 The electrolyte solution structure is primarily determined by two different interactions of the electrolyte components: the interaction between the cation and the anion and the interaction between the cation and the polar aprotic organic solvent. Many different analytical methods have been used to study the solution structure including FTIR, Raman, NMR spectroscopy, conductivity, and computational simulations.9−21 Cyclic carbonate solvents, such as ethylene carbonate (EC) and propylene carbonate (PC), are the primary organic solvents for state-of-the-art lithium ion battery electrolytes. The high dielectric constant of the carbonate solvents suggests that the carbonates coordinate to the Li+ cation strongly, improving the ion transport properties of electrolyte. However, EC is a solid at room temperature, resulting in electrolytes with high viscosity. Alternatively, PC has a lower melting point and lower viscosity but does not form a stable anode SEI on graphite. Therefore, binary solvent systems are utilized with blends of the cyclic carbonates, primarily EC, and linear alkyl carbonate solvents such as dimethyl carbonate (DMC), diethyl carbonate (DEC), © XXXX American Chemical Society
2. EXPERIMENTAL SECTION EC, PC, DMC, and DEC were used as-received (Sigma-Aldrich, anhydrous). LiPF6 was also used as-received (Sigma-Aldrich). Samples were prepared in a Vacuum Atmospheres inert atmosphere (N2) glovebox (20 times larger than the dielectric constant for DMC, but the solvation numbers are very similar at all concentrations of LiPF6.
Figure 3. FTIR spectra of carbonyl group in binary solvent mixtures with LiPF6: (a) EC−DMC; (b) EC−DEC; (c) PC−DMC; (d) PC− DEC.
Table 1. Selected Properties of Carbonate Solvents at 25 °C4 solvent
MW (g/mol)
density (g/mL)
permittivity (ε)
mol no. in 1 L
EC PC DMC DEC
88 102 90 118
1.321a 1.200 1.063 0.969
89.78 64.92 3.1 2.8
15.00 11.80 11.89 8.25
a
Li cation in all forms, Li+(DAC) is the concentration of dialkyl carbonate coordinated to the Li cation in all forms, PC is the concentration of uncoordinated PC, and DAC is the concentration of uncoordinated dialkyl carbonate. This assumption depends upon the similarity of the IR absorption for all forms of coordinated PC and DAC (Li+PC4, Li+PC3DAC, Li+PC2DAC2, etc..), which is consistent with the experimental data.
The density of EC is from 40 °C.
To better understand competitive solvation between linear and cyclic carbonates for the Li cation, binary solvent mixtures were prepared (cyclic carbonate/linear carbonate) with different concentrations of LiPF6. FTIR spectra of the carbonyl bands in different binary carbonate solvent mixtures with various concentrations of LiPF6 are presented in Figure 3. In the electrolyte with a mixture of PC and DMC, six carbonyl bands are observed: uncoordinated PC (1805 and 1790 cm−1), coordinated PC (1772 and 1752 cm−1), uncoordinated DMC (1755 cm−1), and coordinated DMC (1724 cm−1). Very similar results are observed for mixtures of PC and DEC, except that the carbonyl absorptions for the uncoordinated and coordinated DEC are observed at 1724 and 1715 cm−1, respectively. Similar results are also observed with EC−DMC and EC−DEC mixtures, except that the absorptions for coordinated and uncoordinated EC overlap at 1805 and 1755 cm−1. Deconvolution of the coordinated and uncoordinated solvents leads to the determination of the average solvation number for the different carbonate solvents in direct competition (Figure 4; Supporting Information). The average solvation number for PC is higher than the average solvation number of either DMC or DEC at all LiPF6 concentrations. At 1.2 M LiPF6 the solvation numbers for DMC or DEC are 1.8 and 1.7, whereas the solvation numbers for PC are 2.2 and 2.4, respectively. Thus, competitive solvation appears to be preferred despite the much higher dielectric constant for PC and the large excess of uncoordinated PC at low salt concentrations. The solvation equilibrium constants and solvation free energies can be obtained using eq 2,32 where Li+(PC) is the concentration of PC coordinated to the
Li+(PC) + DAC ⇌ Li+(DAC) + PC
(2)
The results suggest the difference in solvation free energies (ΔG°) at 1.2 M LiPF6 are small, 1.3 and 0.0 (± 0.2) kJ/mol for DMC and DEC, respectively. Thus, PC solvates the Li cation slightly better than DMC, but the solvation energies of PC and DEC are comparable. The small differences in solvation free energies are surprising considering the very large differences in dielectric constants (Table 1) and solvation spheres analyzed by ESI-MS.6 Although there is significant interest in directly comparing the solvation affinities of EC and PC, the overlap of the carbonyl absorptions for PC and EC does not allow a direct competition by IR spectroscopy. However, the relative solvation affinities can be determined via an indirect competition between PC and dialkyl carbonates compared to EC and dialkyl carbonates (Figure 5). At all LiPF6 concentrations investigated, the relative amounts of coordinated DMC and DEC are slightly lower with EC than with PC, suggesting that EC has a slightly greater binding affinity for the lithium cation than PC. This is in agreement with computational investigations, as discussed below, but in contradiction to the results of ESI-MS experiments.6 In addition to analysis of the IR absorption of the CO group of the carbonate solvents, the analysis of the PF absorption of the PF6− anion can also provide insight into the solution structure of the electrolyte. The PF absorption has been observed at 844 cm−1 for the uncoordinated anion in solution, but upon formation of an ion pair the symmetry of the anion is C
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Figure 4. Relative area of coordinated solvent in binary PC−DMC and PC−DEC solvent mixtures with LiPF6 mixtures and calculated solvation number in changing concentration of LiPF6.
system has an anion coordination structure similar to that observed in PC. The results are consistent with the solventseparated ions being the dominant species in solution at low LiPF6 concentration (2.0 M). 3.2. NMR Study on Solvation. The rapid time scale of IR spectroscopy allows detection of separate absorptions for uncoordinated and coordinated solvent. However, the time scale of NMR spectroscopy is slower, and solvent exchange frequently occurs more rapidly than the NMR time scale for solvents coordinated to the lithium cation.32,34 Thus, instead of observing separate peaks for coordinated and uncoordinated solvents, a single weighted average peak is typically observed. NMR spectra of carbonate solvents with different concentrations of LiPF6 are depicted in Figure 7. Pure PC has a 13C resonance at 154.6 ppm. Upon increasing the concentration of LiPF6, the 13C resonance shifts to higher frequency. The shift is consistent with a weighted average peak resulting from rapid exchange between uncoordinated PC and PC coordinated to the lithium cation. As the LiPF6 concentration is increased, the chemical shift continues to shift to higher frequency, consistent with more of the PC coordinated to the Li+ cation. Pure DMC and DEC have 13C resonances at 155.7 and 154.4 ppm, respectively. Similar to the results observed for PC, incorporation of LiPF6 shifts the 13C resonances to higher frequency (Figure 7). Related results have been observed with 17O NMR spectroscopy.35 Unfortunately, the solidification of EC/LiPF6 mixtures prevented solution NMR spectroscopy at room temperature. Because the chemical shift of the NMR resonance is a weighted average, the chemical shift of the resonance should correlate with the mole fraction of the solvent molecules coordinated to the lithium cation as determined by IR spectroscopy. A plot of the chemical shift versus the mole fraction of coordinated solvent provides a linear fit (Figure 8). The linearity of the fit provides additional support for the concentrations of coordinated and uncoordinated solvents as determined by IR spectroscopy. Extrapolation of the linear fit to the chemical shift of the pure solvent provides an approximate value for the chemical shift of the pure solvent, but small differences are observed that result from medium effects of the high concentrations of LiPF6 in the solvents used for the experiments. This can also be observed below where the chemical shift of PC in DMC (154.5 ppm) is
Figure 5. Relative amount of coordinated linear carbonate solvent with EC and PC solvent in LiPF6 mixtures: (a) DMC; (b) DEC.
disrupted and the absorption is altered into two bands at 877 and 834 cm−1.16,33 The IR absorption of the PF6− anion at different concentrations of LiPF6 in PC or DMC is depicted in Figure 6. EC−LiPF6 mixtures are solid at room temperature at high LiPF6 concentrations, resulting in difficulty in obtaining homogeneous material for FTIR spectroscopy, whereas DEC has an IR absorption at ∼860 cm−1, which overlaps with the absorptions of the PF6− anion, making analysis of the PF absorption problematic. The peak at 845 cm−1 has a strong intensity in dilute PC−LiPF6 solutions, consistent with most PF6− anions being uncoordinated. The relative intensity of the peak is diminished at higher concentrations of LiPF6, whereas new peaks are observed at 877 and 834 cm−1, consistent with the formation of contact ion pairs or aggregates. The change in the absorption spectra is consistent with the conversion of a solvent-separated ion pair at low LiPF6 concentration to a contact ion pair at high LiPF6 concentration. A similar trend is observed for DMC− LiPF6 solutions. However, higher relative intensities for peaks at 877 and 834 cm−1 suggest that the concentration of ion pairs is higher in the DMC−LiPF6 mixture than in PC−LiPF6 mixtures. Additionally, the IR absorptions for the PF6− anion were monitored in a binary solvent mixture (1:1 EC−DMC, Figure 7). The change in the IR spectra of the PF6− anion is similar to that observed for LiPF6 in PC, suggesting that the binary solvent D
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Figure 7. 13C NMR resonance for carbonyl group in LiPF6 mixtures with (a) PC, (b) DMC, and (c) DEC with changing concentration of LiPF6.
Figure 6. FTIR spectra of PF6− anion in (a) PC, (b) DMC, and (c) EC− DMC mixtures with changing concentration of LiPF6.
slightly different from the chemical shift of PC in pure PC (154.7 ppm).36 The small differences are consistent with medium effects for the mixed solvents compared to the pure solvents. NMR spectra of binary solvent mixtures (PC−DMC and PC− DEC) with and without added LiPF6 are depicted in Figure 9. In binary mixtures without lithium salt, the 13C carbonyl resonance for PC is located at 154.5 ppm, whereas the carbonyl resonances for DMC and DEC are at 155.5 and 154.4 ppm, respectively. Upon addition of LiPF6, all carbonyl resonances shift to higher frequency. In agreement with the investigations of single-solvent electrolytes, the change of the chemical shift of the cyclic carbonate upon Li+ coordination is greater than that observed for the linear carbonates (DMC and DEC). The chemical shift of the 13 C carbonyl resonance is plotted against the mole fraction of coordinated solvent as determined by IR spectroscopy in Figure 10. The linearity of the fit is very good, supporting the
Figure 8. Correlation between 13C NMR resonances for carbonyl group and relative amount of coordinated solvents with changing concentration of LiPF6 in (a) PC, (b) DMC, and (c) DEC.
concentrations of coordinated and uncoordinated solvents and solvation free energies as determined by IR spectroscopy. 3.3. Computational Analysis. EC, PC, and DMC and their lithium ion adducts were modeled using DFT with the B3LYP E
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were not considered. Table 2 shows the calculated energies, HOMO, LUMO, and dipole moments for each of the modeled Table 2. Calculated Quantities for EC, PC, and DMC and Li+ Coordinated Species
EC PC aa-DMC as-DMC ss-DMC Li(EC)4+ Li(PC)4+ Li(aa-DMC)4+ Li(as-DMC)4+ Li(EC)3(aaDMC)+ Li(EC)3(asDMC)+ Li(PC)3(aaDMC)+ Li(PC)3(asDMC)+ Li(EC)3PF6 Li(PC)3PF6 Li(aa-DMC)3PF6 Li(as-DMC)3PF6 Li(EC)2(asDMC)PF6 Li(EC)(asDMC)2PF6 Li(EC)2(aaDMC)PF6 Li(EC)(aaDMC)2PF6
Figure 9. 13C NMR resonance for carbonyl group in binary solvent mixtures with changing concentration of LiPF6: (a) PC−DMC; (b) PC−DEC. Asterisk (∗) indicates resonance for PC.
energy (hartrees)
HOMO (eV)
LUMO (eV)
dipole moment (D)
−342.488387 −381.817886 −343.699786 −343.695123 −343.665455 −1377.45761 −1534.77839 −1382.27345 −1382.28613 −1378.66595
−8.22 −8.14 −7.98 −7.94 −7.80 −11.14 −10.97 −11.61 −10.99 −10.88
0.63 0.63 0.99 0.99 0.46 −2.37 −2.13 −3.02 −2.34 −2.56
5.35 5.50 0.29 3.93 5.91 0.26 0.22 0.48 0.11 5.73
−1378.66477
−10.99
−2.39
2.66
−1496.65617
−10.81
−2.37
5.18
−1496.65506
−10.91
−2.22
3.10
−1975.94574 −2093.93009 −1979.56750 −1979.56205 −1977.15346
−8.60 −8.64 −8.71 −8.50 −8.63
0.23 0.17 −0.45 0.16 0.19
−1978.35930
−8.51
0.22
−1977.15610
−8.53
0.14
−1978.36434
−8.43
−0.11
solvents. The dipole moments for EC and PC are similar and sizable, consistent with the high observed dielectric constants, whereas the low dipole moment for DMC is also consistent with the low observed dielectric constant. Tetra-coordinated Li+ species were calculated, and the structures of the Li(solv)4+ species are shown in Scheme 1. Table 2 gives the calculated energies, HOMO, LUMO, and dipole moments for each of the solvated ions. In the case of PC, which is optically active, all of the coordinated PC molecules were chosen to be in the R conformation. The calculated structure for Li(as-DMC)4+ is lower in energy than Li(aaDMC)4+. This may be due to steric repulsion because the aaDMC conformation is more crowded around the lithium ion but may also arise for electronic reasons, as the binding of a single DMC to Li+ follows the same pattern.38 This suggests that upon coordination of DMC to Li+ there is an isomerization of the solvent molecule. This affects the HOMO and LUMO of the Li(DMC)4+ significantly: both the HOMO and LUMO are more negative for Li(aa-DMC)4+, whereas in Li(as-DMC)4+ the HOMO and LUMO are closer to those found for Li(EC)4+ and Li(PC)4+. This suggests that the oxidation and reduction potentials for Li(EC)4+, Li(PC)4+, and Li(as-DMC)4+ are similar and significantly different from Li(aa-DMC)4+. Four different mixed solvated species were modeled, as well, where a DMC molecule was substituted for either EC or PC in the coordination sphere. These substitutions had small effects on the HOMO and LUMO energies but increased the dipole moment of the solvated lithium ion considerably.39 This suggests that DMC can help the
Figure 10. Correlation between 13C NMR resonances for carbonyl group and relative amount of coordinated solvents with changing concentration of LiPF6 in binary solvent mixtures: (a) PC−DMC; (b) PC−DEC.
functional and the 6-311G* basis set. Whereas the B3LYP functional does not give as accurate results as some other methods,37 the computations are fast and the relative comparisons between species are expected to be valid. PC is optically active, and the results for the all-R Li(PC)4+ cluster are reported for all calculations. A calculation of the all-S structure gave the same results as the all-R configuration, whereas all calculations with mixed R and S enantiomers gave higher energies and always had one or more imaginary frequencies. DMC also has multiple isomers, depending upon the dihedral angles of the O−C−O−C bonds: if both of the O−C−O−C dihedral angles are 0°, the molecule has a “w” shape and is termed the anti,anti-isomer (aa-DMC); if one O−C−O−C angle is 0° and the other is 180°, the molecule has a “z” shape and is termed the anti,syn-isomer (as-DMC); and if both O−C−O−C angles are 180°, the molecule as a “u” shape and is termed the syn,synisomer (ss-DMC). The lowest energy conformation is aa-DMC, whereas the ss-DMC isomer is significantly higher in energy than either aa-DMC or as-DMC, so is not considered further in this discussion. Li(DMC)4+ species with mixed DMC geometries F
DOI: 10.1021/acs.jpcc.5b03694 J. Phys. Chem. C XXXX, XXX, XXX−XXX
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earlier work. Binding of DMC to Li+ is weaker than either EC or PC, but isomerization of the DMC solvent upon complexation is significantly favored, as estimated both by enthalpy and by free energy. If the as-DMC were the predominant conformer in the solvent, the calculations predict that binding to Li+ would be comparable to that to PC. It should be noted that the predicted differences in the calculated binding energies are small and are likely within the error of the method (DFT/B3LYP is expected to have uncertainties on the order of 25 kJ/mol compared to coupled cluster methods37). To model the situation in mixed solvents, the thermodynamic quantities for substitution of EC or PC by DMC were calculated. In no case is the substitution reaction favored at room temperature. In contrast to the case for the tetra-coordinated Li+, upon coordination of a single DMC molecule to Li+ the aaDMC isomer is slightly favored, which is the reverse trend previously reported using the M05-2X functional.37 Substitution of PC by DMC is slightly more favored than substitution of EC. The B3LYP calculations predict that the free energy for the substitution of PC or EC by DMC should be slightly disfavored, on the order of 8−12 kJ/mol. The calculation gives the correct sign compared to the experimental value estimated to be 1.2 kJ/ mol using IR spectroscopy, but overestimates the magnitude. This discrepancy is probably caused by multiple sources: the computation considers only the first solvation shell, whereas the experiment samples all solvation shells; likewise, the experimental IR areas do not distinguish the possible intensity differences for ions with different coordination spheres, which affects the quantitative estimate of ΔG°. The role of ion pairing was also considered. The calculated energies and HOMO and LUMO energies of ion pairs consisting of three solvent molecules and PF6− are also given in Table 2. Formation of the ion pairs has a significant effect on the HOMO and LUMO energies. In all cases when Li(sol)4+ is compared to Li(sol)3PF6, the orbital energies are lowered by about 2 eV. This implies that as ion pairing becomes important there should be an observable shift in the electrochemical potentials of the solvent decomposition. This general prediction follows for both pure and mixed solvents. The calculated thermodynamic quantities for several ion pair formation reactions are given in Table 3. All ion pair formation reactions are found to be strongly favorable, consistent with the stabilization expected from the electrostatic attraction between the cation and the anion.
net solvation by making the primary coordinated ion more polar, which then can increase the attraction of the second solvation sphere. Also note that in the mixed solvated species, the aa-DMC conformation is the lower energy species. Thermodynamic quantities were also calculated to allow modeling of reaction energies. Selected mixed solvated species were also modeled in an attempt to better understand the solution composition in mixed solvents. Table 3 gives the Table 3. Calculated Thermodynamic Values for Selected Solvation Reactions at 298.15 Ka reaction Li+ + 4 EC → Li(EC)4+ Li+ + 4 PC → Li(PC)4+ Li+ + 4 aa-DMC → Li(aa-DMC)4+ Li+ + 4 aa-DMC → Li(as-DMC)4+ Li+ + 4 as-DMC → Li(as-DMC)4+ Li(EC)4+ + aa-DMC → Li(EC)3(aaDMC)+ + EC Li(EC)4+ + aa-DMC → Li(EC)3(asDMC)+ + EC Li(PC)4+ + aa-DMC → Li(PC)3(aaDMC)+ + PC Li(PC)4+ + aa-DMC → Li(PC)3(asDMC)+ + PC Li(EC)4+ + PF6− → Li(EC)3PF6 + EC Li(PC)4+ + PF6− → Li(PC)3PF6 + PC Li(aa-DMC)4+ + PF6− → Li(aaDMC)3PF6 + aa-DMC Li(as-DMC)4+ + PF6− → Li(asDMC)3PF6 + as-DMC Li(as-DMC)4+ + PF6− → Li(asDMC)3PF6 + aa-DMC Li(EC)3(aa-DMC)+ + PF6− → Li(EC)2(aa-DMC)PF6 + EC Li(EC)3(as-DMC)+ + PF6− → Li(EC)2(as-DMC)PF6 + EC Li(EC)3(aa-DMC)+ + PF6− → Li(EC)3PF6 + aa-DMC
ΔH° (kJ/mol)
ΔS° (J/mol·K)
ΔG° (kJ/mol)
−569.7 −576.2 −490.6 −523.8 −572.5 9.7
−715.5 −763.4 −730.3 −725.1 −751.2 −2.9
−356.4 −348.6 −272.9 −307.6 −348.6 10.6
11.5
−2.3
12.1
11.0
9.7
8.2
13.4
8.6
10.8
−290.5 −272.3 −337.1
5.6 17.0 16.9
−292.2 −277.4 −342.2
−278.0
18.1
−283.4
−290.2
11.6
−293.6
−297.0
5.3
−298.6
−292.2
6.4
−294.2
−300.3
8.5
−302.8
a
The enthalpy corrections were found using the frequency calculation with the B3LYP function and the 6-311G* basis set.
calculated reaction energies for several different modes of binding. The binding energies for all carbonates to generate the tetra-coordinated lithium ion are strongly favorable and enthalpy driven. For the reaction of binding of the solvent to Li+, Li(PC)4+ is favored over Li(EC)4+ enthalpically, whereas Li(EC)4+ has a slightly larger binding free energy compared to Li(PC)4+.40 These results agree with results of Bhatt et al.41 for enthalpy and free energy for Li(EC)4+ and free energy for Li(PC)4+. However, the results here are the opposite trend to those of Bhatt et al.41 for the enthalpy of binding for Li(PC)4+. This may be related to the choice of enantiomer for PC, which was not indicated in the
4. CONCLUSION A combination of FTIR spectroscopy, NMR spectroscopy, and DFT computational methods has been utilized to investigate the solvation sphere of the lithium cation in common electrolyte compositions for lithium ion batteries. The quantity of carbonate solvents uncoordinated and coordinated to the lithium cation has G
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been analyzed using the CO IR absorption band for different solvents (PC, DMC, and DEC) with various concentrations of LiPF6. The solvent coordination numbers are higher for DMC and PC than for DEC. At low salt concentrations (0.8 M LiPF6) the solvent coordination numbers approach 5 for PC and DMC, whereas at high salt concentrations (3 M) the solvent coordination numbers are decreased to ∼3. The high solvent coordination numbers at low salt concentration may result from secondary solvation effects because the lithium cation has typically been reported to have a primary coordination sphere of 4. The IR absorptions for the PF6− anion suggest that at low salt concentrations the anion and cation are a solvent-separated ion pair, whereas at high salt concentrations the changes in the absorption are consistent with generation of a contact ion pair. The presence of the PF6− anion in the coordination sphere of the cation compensates for the reduced number of solvent molecules in the coordination sphere. The 13C NMR resonances of the carbonyl group of carbonate solvents are also shifted with Li+ coordination. The chemical shift of the NMR resonance has a linear correlation with the relative amount of coordinated solvent from the FTIR investigation. This correlation shows NMR data can provide solvation information and complement FTIR studies. The DFT computational investigations support the stabilization of the lithium cation by carbonate solvent coordination and support the stability of contact ion pairs. The competitive solvation of the lithium cation by different carbonate solvents was also investigated. Mixtures of PC and DMC and PC and DEC reveal significant involvement of the linear carbonates in the coordination sphere of the lithium cation at all salt concentrations. Interestingly, DMC and DEC coordinate to the lithium cation in the presence of high concentrations of uncoordinated PC. This is surprising because the dielectric constant of PC is an order of magnitude greater that the dielectric constants of DMC and DEC. However, the computational investigations suggest that the DMC undergoes a conformational change upon binding to the lithium cation, which significantly increases the dipole moment of DMC. The results are also in contrast with experimental ESI-MS results, which suggest significant preferential binding for the cyclic carbonates, EC and PC, compared to dialkyl carbonates, DMC and DEC. However, the difference with ESI-MS may be due to the boiling point of the carbonate solvents. Lower boiling solvents DMC and DEC evaporate more readily in ESI-MS, and the observed ratio of solvents in the cation coordination sphere may be influenced by this difference in evaporation rate. One of the most important aspects of understanding the solution structure of standard lithium ion battery electrolytes relates to the formation of the anode SEI.42 The formation reactions of the SEI are closely related to the desolvation of the lithium cation. Whereas this investigation confirms that dialkyl carbonates have an important contribution to the bulk solution structure of standard electrolytes for lithium ion batteries, previous reports suggest that the last solvent to coordinate to the lithium cation during desolvation is a cyclic carbonate (EC).43,44 Thus, the solvation of the lithium cation at the graphite anode interface may be dominated by cyclic carbonate desolvation, which can explain the higher relative concentration of the reduction products of EC compared to dialkyl carbonates in the graphite anode SEI.42,45
Article
ASSOCIATED CONTENT
S Supporting Information *
Deconvolution analysis of the FT-IR absorption bands. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcc.5b03694.
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AUTHOR INFORMATION
Corresponding Author
*
[email protected] Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS We gratefully acknowledge funding from a Department of Energy Office of Basic Energy Sciences EPSCoR Implementation award (DE-SC0007074).
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