Complex Nature of Ionic Coordination in ... - ACS Publications

Apr 24, 2014 - Complex Nature of Ionic Coordination in Magnesium Ionic Liquid-. Based Electrolytes: Solvates with Mobile Mg2+ Cations. Guinevere A. Gi...
2 downloads 0 Views 3MB Size
Download PDF
Article pubs.acs.org/JPCC

Complex Nature of Ionic Coordination in Magnesium Ionic LiquidBased Electrolytes: Solvates with Mobile Mg2+ Cations Guinevere A. Giffin,*,† Arianna Moretti,† Sangsik Jeong,† and Stefano Passerini*,†,‡ †

Institute of Physical Chemistry and MEET Battery Research Center, University of Muenster, Corrensstrasse 28, 48149 Muenster, Germany ‡ Helmholtz Institute (HIU), Karlsruhe Institute of Technology, Albert-Einstein-Allee 11, 89081 Ulm, Germany S Supporting Information *

ABSTRACT: The Raman shifts of the TFSI− expansion-contraction mode in N-butyl-Nmethylpyrrolidinium bis(trifluoromethanesulfonyl)imide ionic liquid (IL) electrolytes were analyzed to compare the ionic coordination of magnesium with lithium and sodium. In the Mg2+-IL electrolytes, the TFSI− anions are found in three different potential energy environments, while only two populations of TFSI− are evident in the Na+- and Li+-IL electrolytes. For Mg2+, the high frequency peak component is associated with a TFSI− that is in a bidentate coordination with a single metal cation and can therefore be considered a contact ion pair (CIP) solvate. The mid frequency component is attributed primarily to bridging aggregate (AGG) TFSI− solvate or a weakly bound monodentate CIP TFSI−. The low frequency peak is well-known to be associated with “free” TFSI− anions. The average number of TFSI− per Mg2+ cation (n) is 3 to 4. In comparison, the value of n is 4 at very low concentrations and decreases with increasing salt mole fraction to 2 for Li+ and Na+, where n of Na+ is larger than that of Li+ at any given concentration. The results imply the existence of anionic magnesium solvates of varying sizes. The identity of the Mg2+ charge-carrying species is complex due to the presence of bridging AGG solvates in solution. It is likely that there is a combination of single Mg2+ solvate species and larger complexes containing two or more cations. In comparison, the primary Li+ and Na+ charge-carrying species are likely [Li(TFSI)2]− and [Na(TFSI)3]2− in the concentration range successfully implemented in IL-based electrolyte batteries. These solvates result in Mg2+ cations that are mobile in the IL-based electrolytes as demonstrated by the reversible magnesiation/ demagnesiation in V2O5 aerogel electrodes.

■

INTRODUCTION Magnesium is a metal that has been considered to hold significant promise as an alternative to lithium-based battery systems. It has a higher specific volumetric capacity than lithium, which could allow it to be used in high energy-density batteries, and it is highly abundant in the earth’s crust.1 Current electrolytes for rechargeable magnesium cells are based on Mg organohaloaluminate complex solutions.1 Ionic liquids (ILs) have been investigated as alternative electrolyte components for magnesium-based batteries, but always in the presence of a Grignard or Lewis acid−base type salt.2−4 Although electrolytes containing non-nucleophilic salts with ILs for the reversible deposition and dissolution of magnesium have been previously studied, the results could not be reproduced.2,3 However, there are significant advantages in terms of safety if the nucleophilic salt could be replaced by one like magnesium bis(trifluoromethanesulfonyl)imide. From the fundamental point of view, there seem to be few, if any, studies on the ionic interactions of ILs with magnesium cations. ILs have received significant attention as an alternative to conventional electrolytes particularly in lithium battery applications due to their many favorable properties, but most notably their low volatility and flammability.5 One of the drawbacks of IL-based electrolytes is their low conductivity and high viscosity as compared to conventional organic liquid-based © 2014 American Chemical Society

electrolytes. It has been reported that the conductivities of saltin-IL electrolytes are lower than those of the ionic liquids themselves, which is a result of the increased viscosity of the system due to interactions between the metal cations and the IL anions.6,7 These interactions are generally classified by several solvate structures such as solvent separated ion pairs (SSIP), contact ion pairs (CIP), or aggregates (AGG), where the anions are coordinated to zero, one and two or more metal cations, respectively.8−11 The formation of such solvates, particularly in the case of AGGs, increases the size and decreases the mobility of the charge-carrying species. In addition, these interactions account for the low transference numbers of metal cations in IL electrolytes.12,13 The ionic interactions must be better understood to determine strategies to improve the properties of IL-based electrolytes, particularly the conductivity and the metal cation mobility. Vibrational spectroscopy is a very good method to examine the ionic interactions between the IL anions and metal cations. When an anion interacts with a metal cation, there is change in the potential energy environment of the anion. This induces a redistribution of the electron density toward the cation that Received: March 7, 2014 Revised: April 23, 2014 Published: April 24, 2014 9966

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

Figure 1. Raman spectra of Li+-, Na+-, and Mg2+-IL electrolyte solutions in the spectral region containing the TFSI− expansion-contraction normal mode of vibration.

IL electrolytes. Differences in the ionic interactions are expected to arise from the varying valence states and the charge to size ratios of the three cations. The results will be used to deduce possible charge-carrying species in the IL electrolytes. Additionally, the feasibility of using nonnucleophilic salts in ILs as electrolytes for magnesium batteries is demonstrated.

results in a change in the internal force constant and therefore the characteristic frequency of a particular vibrational mode.14 Generally, some normal modes are more sensitive to coordination than others. In the case of bis(trifluoromethanesulfonyl)imide (TFSI−), the normal mode associated with the expansion and contraction of the entire anion is particularly sensitive to ionic coordination.10 It has been well-documented that in the “free” anion, which corresponds to an SSIP solvate, this expansion−contraction normal mode is located at ca. 742 cm−1.10,15 When the TFSI− anion is coordinated to Li+, the mode is found at 748 cm−1.10 The ionic coordination of Mg2+ with TFSI− in ILs has not been reported. Decomposition of the TFSI− expansion-contraction Raman peak allows the ratio of coordinated/“free” TFSI− and, subsequently, the average number of TFSI− coordinated to each metal cation to be determined.16,17 It has been previously reported that the number of TFSI− coordinated to each Li+ and Na+ ion in ILs is two and three, respectively. This would imply the existence of [Li(TFSI)2]− and [Na(TFSI)3]2− species.17−20 In the case of Li+, there have also been several publications that have suggested the existence of solvates such as [Li(TFSI)4]3− and [Li(TFSI)3]2−.12,13,21 The general consensus is that lithium solvation in these larger anionic forms is extremely unfavorable for charge transport, and therefore attempts were made to reduce the solvation of Li+.12,13 Theoretical studies combined with spectroscopy are one of the best ways to probe the ionic coordination and deduce the charge-carrying species. The validity of this approach has been previously established by several other groups.14,19,22 Quantum chemical calculations can be used to distinguish possible ionic cluster geometries, and thus the ionic coordination, that would produce normal modes with appropriate frequency shifts. It is through such studies that the conformations and normal modes with frequencies of “free” TFSI− have been well established.23,24 In this work, the ionic coordination of Mg2+ with TFSI− in IL-based electrolytes is investigated using Raman spectroscopy and quantum chemical calculations. The combined theoretical and spectroscopic studies are used to assign the spectral features of the TFSI− expansion-contraction peak to various solvate species. These results are compared to Li+-IL and Na+-

■

EXPERIMENTAL METHODS Electrolyte Preparation. The ionic liquid N-butyl-Nmethylpyrrolidinium bis(trifluoromethanesulfonyl) imide (Pyr 14 TFSI) was synthesized via a two-step synthesis procedure, direct alkylation of N-methylpyrrolidine (>99%, Fluka) with 1-bromobutane (99%, Acros) followed by anion exchange with LiTFSI in aqueous solution, as previously described.25,26 The lithium, sodium and magnesium bis(trifluoromethanesulfonyl) imide salts (LiTFSI, 99.9 wt % battery grade, 3M; NaTFSI, 99.5%, Solvionic; Mg(TFSI)2, 99.5%, Solvionic) were dried at 120 °C for 24 h under vacuum. The TFSI− salts and the ionic liquids had less than 20 ppm water as measured by Karl Fischer titration. The IL-based electrolytes were prepared by mixing appropriate amounts of the salt and ILs and stirring until all of the salt was dissolved. The mixtures were then dried under vacuum. All the sample preparation was performed in a dry room (dew point < −50 °C). Vibrational Spectroscopy. The Raman measurements were recorded on a RAM II FT-Raman module of a Bruker Vertex70 FT-IR spectrometer with a laser wavelength of 1064 nm. The collected spectra are the average of 500 scans at an optical resolution of 2 cm−1. The samples were sealed in glass tubes under vacuum. The resultant spectrum were fit in the spectral region from 860 to 680 cm−1 with the multipeak fitting package in IGOR PRO 6.22A using a Voigt function with a fixed Lorentzian:Gaussian ratio (shape factor = 1.2). Electrochemical Measurements. V2O5 hydrogel was synthesized by reacting crystalline V2O5 with hydrogen peroxide. The red gel was aged for 3 days, washed with acetone several times to remove the water, and finally dried using supercritical CO2 (32 °C, 82 bar).27 Electrodes were 9967

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

prepared by mixing the V2O5 powder with super C65 and polyvinylidene difluoride (PVDF) in a weight ratio 80:10:10. The slurry was cast on aluminum foil current collector. Circular electrodes (1.13 cm2) were cut and dried at 120 °C under vacuum overnight. The average mass loading of the electrodes was 2.5 mg·cm−2. Three-electrode Swagelock cells were assembled in a glovebox (MBraun) with an argon atmosphere (H2O and O2 < 1 ppm). Magnesiated-V2O5 gel was produced via chemical intercalation by soaking the electrodes in 1 mL of 1 M dibutylmagnesium solution in heptane (Sigma-Aldrich). Pristine V2O5 and magnesiated-V2O5 were used as working and counter electrodes, respectively. The cell potential was controlled using an aluminum quasi-reference electrode. The electrolyte was Pyr14TFSI:Mg(TFSI)2 (9:1). Electrochemical tests were conducted by cyclic voltammetry (CV) at scan rate of 0.1 mV·s−1 at 20 °C in a climatic chamber (Binder).

■

THEORETICAL CALCULATIONS

Ionic complexes containing varying numbers of Mg2+ and TFSI− ions were constructed in Avogadro28 (an open-source molecular builder and visualization tool, Version1.1.0. http:// avogadro.openmolecules.net/). From these starting geometries, the complexes were optimized, and vibrational frequencies were calculated using a 6-31G** basis set with a B3LYP density functional theory (DFT) functional. This basis set was chosen so that the geometries and vibrational frequencies would be directly comparable to those published by Herstedt et al.23 All calculations were made using Gaussian 09.29 The vibrational mode assignments were made by animating each mode in Gaussview.30

Figure 2. Fits of the Raman spectra of Li+-, Na+-, and Mg2+-IL electrolyte solutions in the spectral region containing the TFSI− expansion-contraction normal mode of vibration. In all cases, the salt mole fraction in the electrolyte is 0.16. The blue circle markers correspond to the experimental data, the red lines are the fit curves and the gray lines are the fit peaks attributed to the different TFSI− species. The fit peaks are vertically shifted down for clarity.

coordinated TFSI− peak position in the IL electrolytes occurs at slightly lower frequencies (ca. 746 cm−1) than in the Li+-IL electrolytes. This implies that the interaction of Na+ with the TFSI− anion is slightly weaker than that of Li+ with TFSI−. This result, which has been previously noted for electrolytes containing the trifluoromethanesulfonate (Tf−) anion and in other Na+-IL electrolytes,20,31 is not particularly surprising due to the smaller charge to size ratio of the Na+ cation. It is interesting to note that in the solid LiTFSI and NaTFSI, the Raman shift of the TFSI− peak is the same in both salts. It is well-known that the lattice energy and melting point of salts generally decreases with increasing cation size in a given group of the periodic table. While this trend holds for the higher molecular weight alkali metal TFSI− salts, the melting point of LiTFSI is lower than that of NaTFSI.32 It was suggested that this irregularity could be caused by differences in the lattice energies of salts with large organic anions where the electrostatic interactions are more complex than in simple inorganic salts.32 When considering both the melting points and the spectral shift of the TFSI− peak, it seems that the TFSI− potential energy environments are similar in the solid LiTFSI and NaTFSI despite the different charge to size ratios of the cations. The spectra of the Mg(TFSI)2-Pyr14TFSI electrolytes, shown in Figure 1c, are quite different than those of the Na+-IL and Li+-IL electrolytes. There are three components evident in the spectra located at 752, 746, and 742 cm−1. The presence of three components with different frequency shifts implies that there are three populations of TFSI− anions in different potential energy environments. The lowest frequency component, which is dominant at the lowest Mg2+ contents, likely corresponds to “free” TFSI− anions. The features at 746 and 752 cm−1 are likely associated with Mg2+-coordinated TFSI−

■

RESULTS AND DISCUSSION Ionic Coordination. The spectral region containing the TFSI− normal mode of vibration associated with the expansion and contraction of the entire anion is shown in Figure 1 for electrolyte mixtures containing Li+, Na+, or Mg2+ cations. From this point forward, any reference to the TFSI− peak/normal mode will refer to the TFSI− expansion-contraction normal mode of vibration, which is located at 742 cm−1 in neat Pyr14TFSI. The contributions to the spectra discussed below can be more clearly seen in Figure 2, where examples of the fit are shown. The spectra of the LiTFSI-Pyr14TFSI electrolytes (Figure 1a) are consistent with those published for other TFSI−-containing IL electrolytes.16,18 The “free” TFSI− anion is located at 742 cm−1 as seen in the bottom curve (neat Pyr14TFSI). “Free” signifies that the anion is not coordinated or very weakly interacting with any Mn+ ions and therefore is considered to be spectroscopically free. As the Li+ ion content increases, a new peak appears at 748 cm−1, which is associated with the Li+coordinated TFSI−.10 As a reference, the spectrum of LiTFSI is shown at the top of Figure 1a. The frequency shift of the peak in solid LiTFSI is the same as that in the ionic liquid solutions. This suggests that the change in the force constant of the normal mode and therefore the potential energy environment of the coordinated TFSI− anions is approximately the same in the liquid solution and the pure salt. The comparable Na+-IL electrolyte spectra are shown in Figure 1b. There is no distinct separation of the two bands in the presence of sodium, as seen in the Li+-IL spectra. Instead, there is a broad peak encompassing both bands. The Na+9968

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

anions. The presence of two different populations of coordinated TFSI− anions makes it tempting to attribute these components to CIPs and AGG solvate species. However, for reasons elaborated in the following discussion, the peak assignments are not that straightforward. The mid frequency feature at 746 cm−1 is most evident at the highest salt concentrations, but is also discernible from the band asymmetry in most of the electrolytes. The 4 cm−1 shift of this peak with respect to the “free” anion is approximately the same as that reported by Bakker et al. from IR spectra of PEOnMg(TFSI)2 electrolytes (PEO refers to poly(ethylene oxide)). The authors attributed this peak to the formation of CIPs between the Mg2+ ions and TFSI−.33 However, the coordination environment of the Mg2+ cation in the PEOnMg(TFSI)2 system is not comparable to that of the Mg2+-IL electrolytes studied here. Here TFSI− is the only metal cationcoordinating species in the solution, whereas in the PEOnMg(TFSI)2 electrolytes the PEO ether oxygens also can act as ligand molecules. The coordination of the Mg2+ ions by the PEO ether oxygens likely reduces the interaction of the Mg2+ with the TFSI−. The high frequency peak found in this work (752 cm−1) was not seen in the case of the PEOnMg(TFSI)2 electrolytes. This TFSI− peak has a larger frequency shift than that of either of the monovalent alkali cation electrolytes, which indicates a significant shift of the TFSI− electron density toward the divalent magnesium ion. The results of a recent work on glyme-NaTFSI electrolytes further complicates the process of assigning the vibrational peaks.34 In this study, the AGG TFSI− peak is located on the low frequency side of the CIP band in the crystalline phase and is not detected in the liquid state. However, as mentioned above for the PEOnMg(TFSI)2 system, the Mg2+-IL electrolytes studied here are quite different from the glyme-NaTFSI system due to the coordinating ability of the ether oxygen atoms in the glyme molecules. Based on this information, it is still not immediately clear which solvate species can be associated with the high and low frequency TFSI− peaks in the Mg2+-IL electrolytes. However, quantum chemical calculations can be used to identify ionic complexes, and thus the ionic coordination, that would produce appropriate frequency shifts. As a starting point, simple Mg(TFSI)2 CIPs were constructed and optimized. These CIPs are similar to those determined for a variety of divalent cations, including magnesium, by Li and Nie.35 The initial TFSI− starting geometries were equivalent to the C1 and C2 conformers as reported by Johansson et al.24 The calculated C1 and C2 TFSI− geometries had vibrational frequencies of 713.2 and 715.7 cm−1, respectively, which are in good agreement with those given by Herstedt et al.23 The Mg(TFSI)2 structures, which are shown in Figure 3, feature a tetrahedral coordinatedMg2+ cation surrounded by two bidentate TFSI− anions. Analogous CIPs were also determined for [Li(TFSI)2]− complexes by two different groups.16,19 The [Li(C2) 2]− complex ions were reported to have the lowest enthalpies and the largest binding energies.19 Here, the Mg(C2)2 complex also had the lowest enthalpy and the largest binding energy. The addition of C1 TFSI− in the complex adds about 4.5 kJ· mol−1 to the enthalpy per anion. The differences in the binding energies differed by only 2 kJ·mol−1 in the three lowest enthalpy structures. The binding energy was calculated as the difference between the calculated complex enthalpy and the enthalpies of the optimized free ions (C2-TFSI−, C1-TFSI− and Mg2+).

Figure 3. Optimized geometries for Mg(TFSI)2 complexes obtained from DFT B3LYP/6-31G** calculations. The Raman activities are given in parentheses in units of Å4·amu−1.

In all three structures, the vibrational motion of the two TFSI− anions is coupled. As a result, there are two TFSI− normal modes where the motion of the two anions is (1) inphase and therefore the normal mode has a nonzero Raman activity; and (2) out-of-phase and thus the Raman activity is effectively zero. The calculated frequencies given in Figure 3 are too low to account for those seen in the experimental spectra in Figures 1 and 2. These structures also do not give two distinguishable populations of TFSI− anions. Therefore, larger structures with varied TFSI− coordination geometries must be considered. In addition, hexa-coordinate magnesium should be examined as many crystal structures containing Mg2+ ions have octahedrally coordinated cations.36−39 Larger anionic complexes containing between three and six TFSI− and one penta- or hexa- coordinated Mg2+ ion are shown in Figure 4. The maximum binding energies occur in the complexes with three or four TFSI− anions. This would imply that penta- or hexa-coordinated Mg2+ ions surrounded by three or four TFSI− ions are energetically favored. Of the Mg[(TFSI)4]2− complexes, those with the six-coordinated Mg2+ have larger binding energies than the five-coordinated Mg2+ complex. However, the difference between the binding energies of these three complexes is small in comparison to the difference between the complexes that do not have the same number of coordinating TFSI − . The Mg[(C 2 ) 2 (C 1 ) 2 ]2− structure has the lowest enthalpy of Mg[(TFSI)4]2− complexes; the ΔH is 4 and 20 kJ·mol−1 for the six- and five-coordinated Mg[(C2)4]2− clusters, respectively. The calculated vibrational frequencies associated with the TFSI− normal modes are also given for the complexes in Figure 4. The frequencies fall into two general ranges as compared to those of the “free” TFSI− anions: (1) shifted to higher frequencies by two cm−1 or less or to lower frequencies and occur in TFSI− coordinated to the Mg2+ ion in a monodentate fashion; and (2) shifted to higher frequencies by more than five cm−1 and arise from anions with a bidentate coordination. The small shift of the monodentate coordinated TFSI− suggests a small redistribution of the electron density that is mostly localized to the part of the anion interacting with the magnesium cation. In contrast, bidentate coordination results in a more substantial redistribution of the electron density throughout the entire ion. This idea is supported by the calculated Mulliken charges. For the bidentate TFSI−, there is a decrease of the electron density on the CF3 groups and the SO2 oxygen atoms not interacting with the Mg2+ cation, while there is an increase in the charge of the oxygen atoms interacting with the Mg2+ ion. In the case of the monodentate TFSI−, similar changes are only seen on the half of the anion interacting with the Mg2+ ion. The Mulliken charges on the other half of the 9969

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

from the “free” TFSI− to solely account for the population at 746 cm−1. It is possible that these monodentate anions cannot be distinguished from the remaining “free” TFSI− at 742 cm−1 or are distributed between the “free” and 746 cm−1 populations. Additionally, the structures in Figure 4 do not model any TFSI− that can be considered AGG solvates. Therefore, complexes containing TFSI− coordinated to two Mg2+ ions were constructed. The optimized geometries are shown in Figure 5.

Figure 5. Optimized geometries for several Mg2[(C2)n(C1)m](n+m−4)− complexes obtained from DFT B3LYP/6-31G** calculations. The Raman activities are given in parentheses in units of Å4·amu−1. The subscripts “mono” and “bi”, given with the calculated frequencies, indicate that the TFSI− anions are coordinated in a monodentate or bidentate configuration, respectively. The subscript “bridge” indicates that the TFSI− is a bridging ligand coordinated between two Mg2+ ions. The subscript “c” indicates that the normal mode is the result of the coupled motion from more than one TFSI− anion.

Figure 4. Optimized geometries for several Mg[(C2)n(C1)m](n+m−2)− complexes obtained from DFT B3LYP/6-31G** calculations. The Raman activities are given in parentheses in units of Å4·amu−1. The subscripts “mono” and “bi”, given with the calculated frequencies, indicate that the TFSI− are coordinated in a monodentate or bidentate configuration, respectively. The subscript “c” indicates that the normal mode is the result of the coupled motion from more than one TFSI− anion.

The structures in Figure 5 all feature two Mg2+ ions coordinated by two bidentate anions and one TFSI− that acts as a bridge between the two cations. The difference between the structures is the coordination number of the Mg2+ cations and therefore the number of monodentate anions. The complexes have the lowest enthalpy geometries for each anionic Mg2[(TFSI)n](n−4)− complex. It is interesting to note that in all three cases, the bridging TFSI− has a C2 conformation. The vibrational frequency of the normal mode of the bridging TFSI− is in all cases located at lower frequencies than the normal mode of bidentate coordinated TFSI−. However, the specific shift in each complex depends on the coordination number of the Mg2+ ions and the conformation of the TFSI−. In the Mg2[(TFSI)6]2− complexes, the Mg2[(C2)3(C1)3]2− complex has a bridging TFSI− that has a gauche-like conformation when looking down the SNS axis, as opposed

anion, i.e., the −SO2CF3 not interacting with the Mg2+, are largely unchanged. A comparison of the binding energies of the Mg(TFSI)2 and Mg[(TFSI)3]− complexes suggests that the interaction of the bidentate TFSI− with the magnesium is stronger than that of the monodentate TFSI− as only a small increase of the binding energy occurs when adding the additional monodentate TFSI− anion. While the complexes in Figure 4 give information about the interaction of the Mg2+ ion with the TFSI− and its subsequent influence on the vibrational normal modes and corresponding frequencies, they do not account for the three different TFSI− populations seen in Figures 1 and 2. The frequencies of the monodentate coordinated TFSI− is not significantly different 9970

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

to the eclipsed geometry in the “free” C1 anion, and has a lower vibrational frequency than that in the Mg2[(C2)4(C1)2]2− complex. The bidentate ligands are also sensitive to conformation. In both Figures 4 and 5, the C2 bidentate TFSI− have higher frequencies than the C1 bidentate TFSI−. However, even the C1 bidentate ions have frequencies that are higher than the bridging ligands. Based on the results presented here, it can be concluded that the frequencies of all the TFSI normal modes, but in particular the bridging TFSI− frequencies, depend on several factors including but not exclusive to the Mg2+ coordination number and geometry and the TFSI− conformation. The outcome of the DFT calculations can be used to make assignments of the high (752 cm−1) and mid (746 cm−1) frequency peaks. Given that the normal mode of the bidentate TFSI− is consistently found at frequencies higher than the bridging, monodentate or “free” anions, it seems reasonable to conclude that the high frequency peak is associated with these species. The bidentate TFSI− interact with only one Mg2+ and therefore can be classified as CIPs. The mid frequency peak is attributed primarily to the bridging TFSI− anions, but the contribution of monodentate anions cannot be excluded. The frequencies of monodentate TFSI − ions in the DFT calculations encompass an eight cm−1 frequency range. Therefore, it is safe to assume that the electron density distribution of these anions is sensitive to the strength of the interaction with the Mg2+ ion, which in turn depends on the various factors discussed above. As a result, it is likely that in some cases these monodentate anions will have frequencies close to those of the “free” TFSI−, and, in other cases, they will approximate those of the bridging TFSI−. The mid frequency peak can therefore be assigned to the TFSI− that can be classified as AGG solvates (bridging TFSI−) or weakly bound CIPs (monodentate anions). All of the spectra in Figure 1 have been fit using Voigt functions with a fixed Gaussian−Lorenzian ratio as shown in the examples in Figure 2. Using the area of the peaks (A) determined from the fit, the average number of TFSI− coordinated to each metal cation (n) is calculated from eq 1: n=

fTFSI−coordinated xmol salt

=

Figure 6. Average number of TFSI− coordinated to the metal cation in the IL electrolytes. The error bars are derived from the error in the area from the fitting procedure.

of some monodentate TFSI− anions are close to those of the “free” TFSI−. The number of TFSI− per cation decreases more for Li+ and Na+ than for Mg2+. In addition, Na+ coordinates a larger number of anions than Li+ at comparable concentrations. This result may seem counterintuitive given that sodium has a smaller charge to size ratio than lithium. However, the Raman shift of the TFSI− peak is smaller in the spectra of sodium electrolytes than in lithium electrolytes. As mentioned above, a smaller shift suggests a weaker interaction between Na+ and TFSI− than between Li+ and TFSI−, which could explain the higher number of TFSI− per Na+. The lowest average number of metal cations per TFSI− anion occurs in the Li+-IL electrolytes, but as in the case of Mg2+ and Na+, n is not constant as a function of the salt mole fraction. It has been previously published that the number of TFSI− coordinated to each Li+ in IL solutions is 2.17−19 This would imply that the lithium charge carriers in solution are primarily [Li(TFSI)2]− anions. However, based on the results presented here, these are not the only lithium containing species in the IL electrolytes particularly at low salt concentrations. It is likely that there are a mix of [Li(TFSI)4]3−, [Li(TFSI)3]2− and [Li(TFSI)2]− solvates existing in solution. These results are in good agreement with previously published studies that have concluded that solvates such as [Li(TFSI)4]3− and [Li(TFSI)3]2− are likely to be found in IL solutions.12,13,21 However, exchange of TFSI− between the various species is likely to be very fast. In fact, PFSE-NMR studies are not able to distinguish between TFSI− bound to Li+ and “free” TFSI−.40 The presence of larger aggregates is probably one of the factors that account for the decreased conductivity and increased viscosity of the IL electrolytes as compared to the neat ILs.6,7 In the concentration range where Li+-IL electrolytes have been successfully implemented in prototype IL batteries (9:1 IL:salt),41 the predominant lithium charge-carrying species is likely [Li(TFSI)2]−. In the same concentration range, the primary sodium complex is likely [Na(TFSI)3]2−. A deduction of the principle magnesium solvate is more difficult due to the presence of bridging AGG solvates in solution. It is likely that there is a combination of single Mg2+ species and larger complexes containing two or more cations as illustrated in Figure 5. As a result, it is possible to imagine a large number of solvates that would give a value of n that is between 3 and 4. Exchange of ions between these solvates is likely rapid, but in

ATFSI−coordinated A total

xmol salt

(1)

where xmol is the mole fraction of the salt and f is the fraction of coordinated TFSI−, which includes both the 752 and 746 cm−1 spectral components. The average number of TFSI− coordinated to each metal cation ion is plotted as a function of the salt mole fraction in Figure 6. It is interesting to note that at low concentrations, the average number of TFSI− per metal cation is approximately the same (n ≈ 4) for all the cations. This result implies the number of anions surrounding the cation is high and is likely only limited by the preferential coordination number and steric effects at very low concentrations when the amount of TFSI− anions is much larger than that of the metal cations. In all cases, the number of the TFSI− per cation decreases as the mole fraction of the salt increases. Unsurprisingly, the Mg2+ electrolytes have the highest n, which decreases from 4 to 3 over the investigated concentration range. This value of n is consistent with the trend in the binding energies, discussed with respect to Figure 4 above, which showed that 3 to 4 TFSI− per Mg2+ is energetically favorable. It is possible that calculated average values for Mg2+ may be slightly low as the frequencies 9971

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

solution the conductivity is likely primarily due to the TFSI− anions. Studies on the physical and transport properties of Mg2+-IL electrolytes will be reported in a future publication. Magnesium Mobility. Electrochemical tests were conducted to demonstrate the feasibility of using the Mg2+-IL electrolyte in a magnesium rechargeable battery. It has been previously shown that V2O5 aerogels can achieve a high intercalation capacity of Mg2+ ions.27 Therefore, this material is an excellent host to demonstrate the mobility of the Mg2+ ions in the IL-based electrolytes. The CV curve of the electrochemical cells is shown in Figure 7.

may still quite good as the interaction between Na+ and TFSI− is weaker than that between Li+ and TFSI−. In the case of the Mg2+-IL electrolytes, the coordination number does not tell the whole story. The existence of both contact ion pair and aggregate solvates further increases the size of the chargecarrying species and likely leads to a lower mobility of the Mg2+ ions in the electrolyte as compared to the mobility of Li+. Despite the size of the charge-carrying species, the Mg2+ ions are mobile in the IL-based electrolytes as demonstrated by the reversible magnesiation/demagnesiation of V2O5 electrodes in Pyr14TFSI:Mg(TFSI)2 (9:1) electrolyte.

■

ASSOCIATED CONTENT

* Supporting Information S

Cartesian coordinates and Mulliken charges of the optimized complexes from the DFT calculations (B3LYP/6-31G**); full author list for ref 29. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Authors

*E-mail: gwen.giffi[email protected]. *E-mail: [email protected]; stefano. [email protected]. Notes

Figure 7. Cyclic voltammetry for the magnesiation/demagnesiation of Mg2+ in V2O5 electrodes at 20 °C. Scan rate 0.1 mV·s−1. Working electrode: V2O5 aerogel. Counter electrode: Magnesiated V2O5 aerogel. Electrolyte:Pyr14TFSI:Mg(TFSI)2 (9:1).

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported by the Helmholtz Gemeinschaft within the project “Helmholtz-Energie-Allianz - Stationäre elektrochemische Feststoff-Speicher und -Wandler” (Förderkennzeichen: HA-E-0002).

The shape of the CV is similar to that reported for Mg2+ insertion into xerogel V 2O 5 using Mg(ClO 4) 2 salt in acetonitrile.42 Three broad peaks are visible during the cathodic scan at 3.2 V, 2.6 V, and 2.1 V, respectively. As already observed by other authors43 the insertion/deinsertion process into V2O5 seems to proceed via different steps, probably related to the insertion of Mg2+ ions into different crystallographic sites.42 In the anodic scan the two sharp peaks, located respectively at around 2.7 and 3.3 V, correspond to the demagnesiation process. These results demonstrate that the Mg2+ ions are mobile in the IL-based electrolytes.

■

REFERENCES

(1) Yoo, H. D.; Shterenberg, I.; Gofer, Y.; Gershinsky, G.; Pour, N.; Aurbach, D. Mg Rechargeable Batteries: An On-going Challenge. Energy Environ. Sci. 2013, 6, 2265−2279. (2) Amir, N.; Vestfrid, Y.; Chusid, O.; Gofer, Y.; Aurbach, D. Progress in Nonaqueous Magnesium Electrochemistry. J. Power Sources 2007, 174, 1234−1240. (3) Cheek, G. T.; O’Grady, W. E.; El, A. S. Z.; Moustafa, E. M.; Endres, F. Studies on the Electrodeposition of Magnesium in Ionic Liquids. J. Electrochem. Soc. 2007, 155, D91−D95. (4) Kakibe, T.; Hishii, J.-y.; Yoshimoto, N.; Egashira, M.; Morita, M. Binary Ionic Liquid Electrolytes Containing Organo−Magnesium Complex for Rechargeable Magnesium Batteries. J. Power Sources 2012, 203, 195−200. (5) MacFarlane, D. R.; Tachikawa, N.; Forsyth, M.; Pringle, J. M.; Howlett, P. C.; Elliott, G. D.; Davis, J. H.; Watanabe, M.; Simon, P.; Angell, C. A. Energy Applications of Ionic Liquids. Energy Environ. Sci. 2014, 7, 232−250. (6) Pitawala, J.; Kim, J.-K.; Jacobsson, P.; Koch, V.; Croce, F.; Matic, A. Phase Behaviour, Transport Properties, and Interactions in Li-Salt Doped Ionic Liquids. Faraday Discuss. 2012, 154, 71−80. (7) Zhou, Q.; Boyle, P. D.; Malpezzi, L.; Mele, A.; Shin, J.-H.; Passerini, S.; Henderson, W. A. Phase Behavior of Ionic Liquid−LiX Mixtures: Pyrrolidinium Cations and TFSI- Anions - Linking Structure to Transport Properties. Chem. Mater. 2011, 23, 4331−4337. (8) Han, S.-D.; Allen, J. L.; Jónsson, E.; Johansson, P.; McOwen, D. W.; Boyle, P. D.; Henderson, W. A. Solvate Structures and Computational/Spectroscopic Characterization of Lithium Difluoro(oxalato)borate (LiDFOB) Electrolytes. J. Phys. Chem. C 2013, 117, 5521−5531. (9) Seo, D. M.; Borodin, O.; Han, S.-D.; Boyle, P. D.; Henderson, W. A. Electrolyte Solvation and Ionic Association II. Acetonitrile−Lithium

■

CONCLUSIONS The TFSI− expansion-contraction normal mode of vibration has been analyzed to investigate the ionic coordination of the Mg2+ ions in IL-based electrolytes. Three spectral features are found at 752, 746, and 742 cm−1 in the Raman spectra of the Mg2+-IL electrolytes. In comparison, the TFSI− peak is located at 748/742 and 746/742 cm−1 in the Li+-IL and Na+-IL electrolytes, respectively. DFT calculations have been used to assign the three features in the Mg2+-IL system. The high frequency peak (752 cm −1 ) is associated with TFSI − coordinated to a single metal cation in a bidentate geometry and can therefore be considered CIPs. The mid frequency component (746 cm−1) is attributed primarily to bridging AGG TFSI− or weakly bound monodentate CIP TFSI−. The values of n, which are not constant as a function of the salt mole fraction, suggest that in all three electrolyte systems a mix of anionic solvates exist. In general, the largest anionic solvates likely occur with Mg2+, while the smallest occur with Li+. A comparison of the n values and the Raman shifts of the cation-coordinated TFSI− peak for Na+ and Li+ suggests that while Na+ likely exists in larger solvates, the mobility of the Na+ 9972

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973

The Journal of Physical Chemistry C

Article

Salt Mixtures: Highly Dissociated Salts. J. Electrochem. Soc. 2012, 159, A1489−A1500. (10) Brouillette, D.; Irish, D. E.; Taylor, N. J.; Perron, G.; Odziemkowski, M.; Desnoyers, J. E. Stable Solvates in Solution of Lithium Bis(trifluoromethylsulfone)imide in Glymes and Other Aprotic Solvents: Phase Diagrams, Crystallography and Raman Spectroscopy. Phys. Chem. Chem. Phys. 2002, 4, 6063−6071. (11) Henderson, W. A. Glyme−Lithium Salt Phase Behavior. J. Phys. Chem. B 2006, 110, 13177−13183. (12) Saito, Y.; Umecky, T.; Niwa, J.; Sakai, T.; Maeda, S. Existing Condition and Migration Property of Ions in Lithium Electrolytes with Ionic Liquid Solvent. J. Phys. Chem. B 2007, 111, 11794−11802. (13) Umecky, T.; Saito, Y.; Okumura, Y.; Maeda, S.; Sakai, T. Ionization Condition of Lithium Ionic Liquid Electrolytes under the Solvation Effect of Liquid and Solid Solvents. J. Phys. Chem. B 2008, 112, 3357−3364. (14) Huang, W.; Frech, R.; Wheeler, R. A. Molecular Structures and Normal Vibrations of Trifluoromethane Sulfonate (CF3SO3−) and Its Lithium Ion Pairs and Aggregates. J. Phys. Chem. 1994, 98, 100−110. (15) Rey, I.; Johansson, P.; Lindgren, J.; Lassègues, J. C.; Grondin, J.; Servant, L. Spectroscopic and Theoretical Study of (CF3SO2)2N− (TFSI−) and (CF3SO2)2NH (HTFSI). J. Phys. Chem. A 1998, 102, 3249−3258. (16) Umebayashi, Y.; Mitsugi, T.; Fukuda, S.; Fujimori, T.; Fujii, K.; Kanzaki, R.; Takeuchi, M.; Ishiguro, S.-I. Lithium Ion Solvation in Room-Temperature Ionic Liquids Involving Bis(trifluoromethanesulfonyl) Imide Anion Studied by Raman Spectroscopy and DFT Calculations. J. Phys. Chem. B 2007, 111, 13028− 13032. (17) Lassegues, J.-C.; Grondin, J.; Talaga, D. Lithium Solvation in Bis(trifluoromethanesulfonyl)imide-Based Ionic Liquids. Phys. Chem. Chem. Phys. 2006, 8, 5629−5632. (18) Duluard, S.; Grondin, J.; Bruneel, J.-L.; Pianet, I.; Grélard, A.; Campet, G.; Delville, M.-H.; Lassègues, J.-C. Lithium Solvation and Diffusion in the 1-Butyl-3-methylimidazolium Bis(trifluoromethanesulfonyl)imide Ionic Liquid. J. Raman Spectrosc. 2008, 39, 627−632. (19) Lassègues, J.-C.; Grondin, J.; Aupetit, C.; Johansson, P. Spectroscopic Identification of the Lithium Ion Transporting Species in LiTFSI-Doped Ionic Liquids. J. Phys. Chem. A 2009, 113, 305−314. (20) Monti, D.; Jónsson, E.; Palacín, M. R.; Johansson, P. Ionic Liquid Based Electrolytes for Sodium-Ion Batteries: Na+ Solvation and Ionic Conductivity. J. Power Sources 2014, 245, 630−636. (21) Borodin, O.; Smith, G. D.; Henderson, W. Li+ Cation Environment, Transport, and Mechanical Properties of the LiTFSI Doped N-Methyl-N-alkylpyrrolidinium+TFSI− Ionic Liquids. J. Phys. Chem. B 2006, 110, 16879−16886. (22) Gejji, S. P.; Suresh, C. H.; Babu, K.; Gadre, S. R. Ab Initio Structure and Vibrational Frequencies of (CF3SO2)2N−Li+ Ion Pairs. J. Phys. Chem. A 1999, 103, 7474−7480. (23) Herstedt, M.; Smirnov, M.; Johansson, P.; Chami, M.; Grondin, J.; Servant, L.; Lassègues, J. C. Spectroscopic Characterization of the Conformational States of the Bis(trifluoromethanesulfonyl)imide Anion (TFSI−). J. Raman Spectrosc. 2005, 36, 762−770. (24) Johansson, P.; Gejji, S. P.; Tegenfeldt, J.; Lindgren, J. The Imide Ion: Potential Energy Surface and Geometries. Electrochim. Acta 1998, 43, 1375−1379. (25) Appetecchi, G. B.; Scaccia, S.; Tizzani, C.; Alessandrini, F.; Passerini, S. Synthesis of Hydrophobic Ionic Liquids for Electrochemical Applications. J. Electrochem. Soc. 2006, 153, A1685−A1691. (26) Montanino, M.; Alessandrini, F.; Passerini, S.; Appetecchi, G. B. Water-Based Synthesis of Hydrophobic Ionic Liquids for High-Energy Electrochemical Devices. Electrochim. Acta 2013, 96, 124−133. (27) Le, D. B.; Passerini, S.; Coustier, F.; Guo, J.; Soderstrom, T.; Owens, B. B.; Smyrl, W. H. Intercalation of Polyvalent Cations into V2O5 Aerogels. Chem. Mater. 1998, 10, 682−684. (28) Hanwell, M.; Curtis, D.; Lonie, D.; Vandermeersch, T.; Zurek, E.; Hutchison, G. Avogadro: An Advanced Semantic Chemical Editor, Visualization, and Analysis Platform. J. Cheminf. 2012, 4, 17.

(29) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.et al.; Gaussian 09, revision A.02; Gaussian Inc.: Wallingford, CT, 2009. (30) Dennington, R.; Keith, T.; Millam, J. GaussView, version 5; Semichem Inc.: Shawnee Mission, KS, 2009. (31) Sanders, R. A.; Frech, R.; Khan, M. A. Structural Investigation of Crystalline and Solution Phases in N,N,N′,N′-Tetramethylethylenediamine (TMEDA) with Lithium Triflate (LiCF3SO3) and Sodium Triflate (NaCF3SO3). J. Phys. Chem. B 2003, 107, 8310−8315. (32) Hagiwara, R.; Tamaki, K.; Kubota, K.; Goto, T.; Nohira, T. Thermal Properties of Mixed Alkali Bis(trifluoromethylsulfonyl)amides. J. Chem. Eng. Data 2008, 53, 355−358. (33) Bakker, A.; Gejji, S.; Lindgren, J.; Hermansson, K.; Probst, M. M. Contact Ion Pair Formation and Ether Oxygen Coordination in the Polymer Electrolytes M[N(CF3SO2)2]2PEOn for M = Mg, Ca, Sr and Ba. Polymer 1995, 36, 4371−4378. (34) Mandai, T.; Nozawa, R.; Tsuzuki, S.; Yoshida, K.; Ueno, K.; Dokko, K.; Watanabe, M. Phase Diagrams and Solvate Structures of Binary Mixtures of Glymes and Na Salts. J. Phys. Chem. B 2013, 117, 15072−15085. (35) Li, X. Y.; Nie, J. Density Functional Theory Study on Metal Bis(trifluoromethylsulfonyl)imides: Electronic Structures, Energies, Catalysis, and Predictions. J. Phys. Chem. A 2003, 107, 6007−6013. (36) Aurbach, D.; Gizbar, H.; Schechter, A.; Chusid, O.; Gottlieb, H. E.; Gofer, Y.; Goldberg, I. Electrolyte Solutions for Rechargeable Magnesium Batteries based on Organomagnesium Chloroaluminate Complexes. J. Electrochem. Soc. 2002, 149, A115−A121. (37) Haas, A.; Klare, C.; Betz, P.; Bruckmann, J.; Krueger, C.; Tsay, Y. H.; Aubke, F. Acyclic Sulfur-Nitrogen Compounds. Syntheses and Crystal and Molecular Structures of Bis((trifluoromethyl)sulfonyl)amine ((CF3SO2)2NH), Magnesium Hexaaquo Bis((trifluoromethyl)sulfonyl)amide Dihydrate ([Mg(H2O)6][(CF3SO2)2N]2·2H2O), and Bis(bis(fluorosulfonyl)amino)sulfur ((FSO2)2NSN(SO2F)2. Inorg. Chem. 1996, 35, 1918−1925. (38) Muldoon, J.; Bucur, C. B.; Oliver, A. G.; Sugimoto, T.; Matsui, M.; Kim, H. S.; Allred, G. D.; Zajicek, J.; Kotani, Y. Electrolyte Roadblocks to a Magnesium Rechargeable Battery. Energy Environ. Sci. 2012, 5, 5941−5950. (39) Muldoon, J.; Bucur, C. B.; Oliver, A. G.; Zajicek, J.; Allred, G. D.; Boggess, W. C. Corrosion of Magnesium Electrolytes: Chlorides The Culprit. Energy Environ. Sci. 2013, 6, 482−487. (40) Park, J.-W.; Yoshida, K.; Tachikawa, N.; Dokko, K.; Watanabe, M. Limiting Current Density in Bis(trifluoromethylsulfonyl)amideBased Ionic Liquid for Lithium Batteries. J. Power Sources 2011, 196, 2264−2268. (41) Balducci, A.; Jeong, S. S.; Kim, G. T.; Passerini, S.; Winter, M.; Schmuck, M.; Appetecchi, G. B.; Marcilla, R.; Mecerreyes, D.; Barsukov, V.; et al. Development of Safe, Green and High Performance Ionic Liquids-Based Batteries (ILLIBATT Project). J. Power Sources 2011, 196, 9719−9730. (42) Imamura, D.; Miyayama, M. Characterization of MagnesiumIntercalated V2O5/Carbon Composites. Solid State Ionics 2003, 161, 173−180. (43) Gershinsky, G.; Yoo, H. D.; Gofer, Y.; Aurbach, D. Electrochemical and Spectroscopic Analysis of Mg2+ Intercalation into Thin Film Electrodes of Layered Oxides: V2O5 and MoO3. Langmuir 2013, 29, 10964−10972.

9973

dx.doi.org/10.1021/jp502354h | J. Phys. Chem. C 2014, 118, 9966−9973