Article pubs.acs.org/JPCB
Effect of Structure on Transport Properties (Viscosity, Ionic Conductivity, and Self-Diffusion Coefficient) of Aprotic Heterocyclic Anion (AHA) Room-Temperature Ionic Liquids. 1. Variation of Anionic Species Liyuan Sun, Oscar Morales-Collazo, Han Xia, and Joan F. Brennecke* Department of Chemical and Biomolecular Engineering, University of Notre Dame, Notre Dame, Indiana 46556, United States S Supporting Information *
ABSTRACT: A series of room temperature ionic liquids (RTILs) based on 1-ethyl-3-methylimidazolium ([emim]+) with different aprotic heterocyclic anions (AHAs) were synthesized and characterized as potential electrolyte candidates for lithium ion batteries. The density and transport properties of these ILs were measured over the temperature range between 283.15 and 343.15 K at ambient pressure. The temperature dependence of the transport properties (viscosity, ionic conductivity, self-diffusion coefficient, and molar conductivity) is fit well by the Vogel−Fulcher−Tamman (VFT) equation. The best-fit VFT parameters, as well as linear fits to the density, are reported. The ionicity of these ILs was quantified by the ratio of the molar conductivity obtained from the ionic conductivity and molar concentration to that calculated from the self-diffusion coefficients using the Nernst−Einstein equation. The results of this study, which is based on ILs composed of both a planar cation and planar anions, show that many of the [emim][AHA] ILs exhibit very good conductivity for their viscosities and provide insight into the design of ILs with enhanced dynamics that may be suitable for electrolyte applications.
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INTRODUCTION Ionic liquids (ILs) are organic salts with melting points below 373.15 K. Many ILs have melting points around room temperature and are known as room-temperature ionic liquids (RTILs). RTILs have attracted great interest for practical applications in recent years due to their unique properties such as negligible vapor pressure, low flammability, high thermal, chemical, and electrochemical stability, and good solvating properties. Additionally, physicochemical properties of an IL can be easily tuned by changing the structure of the cation and anion.1−4 Because ILs are comprised entirely of ions, their intrinsic ionic conductivity and wide electrochemical windows make them attractive electrolyte candidates for electrochemical devices, such as lithium ion batteries, electric double layer capacitors, dye-sensitized solar cells, and fuel cells.5−10 On the other hand, the viscosity of ILs is usually 1 to 3 orders of magnitude higher than that of water and conventional organic solvents,11 which significantly reduces ion transport and thus hinders the applications of ILs. As such, understanding how transport properties of ILs are correlated with the ion structure is very important and has been the focus of many studies.12−22 Aprotic heterocyclic anion (AHA) ILs have been recently developed and found to be promising for CO2 capture due to their ability to react stoichiometrically with CO2.16 They may also be promising electrolyte candidates because of their moderate viscosity and conductivity in addition to the key favorable properties of most ILs. However, previous studies and reviews that consider the use of ILs as electrolytes have © 2015 American Chemical Society
primarily focused on some well-studied anions, including [(C2F5SO2)2N]−, [(CF3SO2)2N]−, [CF3SO3]−, [CF3CO2]−, [BF4]−, [PF6]−, and halides,6,13,19−21,23,24 although with various cations based on imidazolium, pyridinium, pyrrolidinium, and ammonium. Only limited experimental studies are available on the physicochemical properties of the AHA RTILs15−18 and data on the correlation between self-diffusion coefficients of ions and their structures is particularly lacking in the literature for AHA RTILs. In the present work, we report a systematic study of the correlation between ion structures and the physicochemical properties with an emphasis on the ion transport behavior of a family of AHA RTILs. Specifically, variation of anionic species with a fixed cation, 1-ethyl-3metylimidazolium ([emim]+), was considered.
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EXPERIMENTAL SECTION Synthesis. The structures, names, and abbreviations of all the ILs discussed in this study are listed in Table 1. 1H-1, 2, 3triazole (97% purity), 1,2,4-triazole (98% purity), 4-nitro-1Hpyrazole (97% purity), 3-nitro-1H-pyrazole, 3(trifluoromethyl)pyrazole (99% purity), and 3-methyl-5(trifluoromethyl)pyrazole (99% purity) were purchased from Sigma-Aldrich. Tetrazole (0.45 M in acetonitrile) and pyrrole2-carbonitrile (99% purity) were purchased from Alfa Aesar and Received: September 20, 2015 Revised: October 27, 2015 Published: October 27, 2015 15030
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
Article
The Journal of Physical Chemistry B Table 1. Structures, Names, and Abbreviations of ILs Discussed in This Work
was less than 0.05% by weight (or 500 ppm). The structures of the ILs were verified by 1H NMR (Varian INOVA-600) spectroscopy with dried ILs dissolved in deuterated dimethyl sulfoxide (DMSO-d6 , 99.9 atom % D, Sigma-Aldrich) containing 1% v/v TMS. Any residual halide in the aqueous phase in contact with the ILs was not detectable by adding an AgNO3 aqueous solution. Thermal Analysis. Thermal stabilities of the ILs were determined by using a Mettler Toledo TGA/SDTA thermal gravimetric analyzer (TGA). Approximately 20 mg samples were kept in a standard 40 μL aluminum crucible and were heated to 773.15 K at a heating rate of 10 K min−1 after drying in situ at 383.15 K for 45 min. An onset temperature was used to evaluate the relative thermal stability of each IL and was
used without further purification. ILs were prepared by a twostep procedure. First, 1-ethyl-3-methylimidazolium hydroxide ([emim][OH]) was prepared by treating 1-ethyl-3-methylimidazolium bromide, [emim][Br] (97% purity, Iolitec) with Amberlite IRN78 (Sigma-Aldrich), a [OH] ion-exchange resin in methanol (ACS grade, Fischer Scientific). An acid−base reaction takes place between [emim][OH] and an equimolar amount of corresponding heterocyclic precursor. Methanol and other volatiles were removed at 323.15 K under reduced pressure (10 mbar). Complete removal of volatiles was confirmed by 1H NMR. Byproduct water was then removed by further drying at 323.15 K under reduced pressure for approximately 3 days. The water content of each IL was determined by a Brinkman 831 Karl Fischer coulometer and 15031
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
Article
The Journal of Physical Chemistry B
measurements were performed over the frequency range from 20 Hz to 1 MHz with a perturbation value of 10 mV, and the uncertainty is approximately ±3%. Self-Diffusion Coefficient. The self-diffusion coefficients of the ILs were determined by pulsed gradient spin−echo (PGSE) NMR on a Bruker Avance III HD 400 MHz NMR spectrometer. DMSO-d6 containing 1% v/v TMS was used as the external solvent reference at temperatures from 298.15 to 343.15 K and deuterated chloroform (CDCl3) containing 1% v/v TMS was used at 283.15 and 293.15 K. 1H nucleus was used to measure both the cationic and anionic self-diffusion coefficients of the ILs studied in this work. The 1H NMR chemical shifts were determined by use of a double tube with pure and dried ILs tightly sealed in a PTFE liner inner tube and the external solvent reference in the outer 5 mm NMR tube. The sample filling height was approximately 3 cm in the inner tube. A 100% ethylene glycol standard solution was used for temperature calibration, and the gradient strength was calibrated with a lab built “doped water” sample containing 1% H2O in D2O plus 0.1% CuSO4 at 298.15 K.26 The free diffusion echo signal attenuation, E, is related to the experimental parameters and the self-diffusion coefficient by the following well-known Stejskal equation27,28
defined as the intersection of the baseline weight (after the drying step) and the tangent of the weight versus temperature curve as decomposition occurs. The TGA measurement has an accuracy of ±0.25 K but the uncertainty of the onset temperature is ±2 K due to the manual determination of the tangent point. The melting point (Tm) and the glass transition (Tg) temperature were determined by using a Mettler-Toledo differential scanning calorimeter (DSC). The samples were kept in a standard 40 μL aluminum crucible and were cooled from room temperature (298.15 K) to 153.15 K at a cooling rate of −10 K min−1 and were held at 153.15 K for 3 min and then heated from 153.15 to 323.15 K at a heating rate of 10 K min−1 under a N2 atmosphere. The midpoint and onset point were used to determine Tg and Tm, respectively. The uncertainties of the reported T m and T g values are approximately ±1 K. Both the TGA and DSC were calibrated using indium and zinc. Electrochemical Window. The electrochemical windows (ECWs) of the ILs were measured in an undivided threeelectrode cell with a VoltaLab50 potentiostat at a scanning rate of 100 mV/s. The cell was kept in a glovebox (Labmaster SP MB 20, H2O < 0.1 ppm, O2 < 30 ppm) under N2 atmosphere at ambient temperature. Glassy carbon and platinum wire were used as the working and counter electrode, respectively. The reference electrode was constructed from a silver wire immersed in 0.01 M silver nitrate dissolved in a 0.1 M [emim][(CF3SO2)2N]/acetonitrile solution. The reference electrode was calibrated against the ferrocene/ferrocenium (Fc/Fc+) redox couple. Density. The densities of the ILs were measured with a DMA 4500 Anton Paar oscillating U-tube densitometer at ambient pressure with a precision of ±0.00001 g/cm3. The temperature was controlled by two integrated Pt 100 platinum thermometers with a precision of ±0.01 K. The uncertainty of the density measurement is approximately ±0.0005 g/cm3 when taking the sample impurities into account. Viscosity. The viscosities of the ILs below 230 mPa·s were measured with an Anton Paar automated microviscometer at ambient pressure. A capillary of 1.6 mm or 3.0 mm in diameter was used depending on the viscosity range measured with an inclination angle of 40°. Quartic measurements of each sample at each temperature were performed and the results were averaged. The experimental uncertainty is ±3%. Viscosities above 230 mPa·s were measured with an ATS Rheosystems viscometer equipped with a cone-and-plate spindle. The samples were kept under a N2 atmosphere by continuously purging the gas over the samples. The uncertainty of the viscosity measurements with the ATS system is ±5%. Ionic Conductivity. The ionic conductivities were measured with a computer controlled electrochemical impedance spectroscopy (EIS) system, which consists of a Solartron SI 1287 electrochemical interface and an SI 1260 impedance/gainphase analyzer. All samples were loaded and sealed into airtight conductivity cells under N2 in a glovebox to avoid their exposure to moisture in the atmosphere. The conductivity cells were constructed with two platinized platinum electrodes with the cell constants of about 1.0 cm−1. The actual cell constants were calibrated with a standard KCl solution with known conductivity of 10 mS/cm at 298.15 K, as described by Barthel et al.25 The cells were kept in a Binder refrigerated incubator for temperature control and were thermally equilibrated at each temperature for at least 45 min before measurements. The
⎛ S ⎞ ⎛ ⎞ ⎟⎟ = −Dγ 2g 2δ 2⎜Δ − δ ⎟ ln(E) = ln⎜⎜ ⎝ 3⎠ ⎝ Sg = 0 ⎠
(1)
where S is the spin−echo signal intensity, D is the self-diffusion coefficient, γ is the gyromagnetic ratio, g is the magnitude of the field gradient, δ is the duration of the field gradient, and Δ is the interval between the two gradient pulses. In the present measurements, the g value used was constant (5.35 Tm−1), Δ was in the range of 0.2−0.65 s, δ was in the range of 3−13 ms, and γ is a preset value of the program. The uncertainty of the self-diffusion coefficient measurement is ±5%.
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RESULTS Thermal Properties, Electrochemical Window, Density, and Molar Concentration. As Table 2 shows, all of the Table 2. Molecular Weight and Thermal Properties ILs
Mw/g mol−1
Tg/Kc
Td/Kd
[emim][2CNPyr]a [emim][4triz]a [emim][3triz]b [emim][4NO2Pyra] [emim][3NO2Pyra] [emim][3CF3Pyra] [emim][3CH35CF3Pyra] [emim][tetz]b
202.26 179.22 179.22 223.23 223.23 246.23 261.27 180.21
212.15 205.15 201.15 217.15 209.15 201.15 216.15 194.15
479.15 481.15 491.15 494.15 497.15 468.15 482.15 466.15
a From ref 17. bFrom ref 15. cGlass transition temperature determined by differential scanning calorimetry. dDecomposition temperature determined by thermal gravimetric analysis.
ILs discussed here exhibit a glass transition temperature between 194.15 and 217.15 K and a decomposition temperature above 482.15 K at a scanning rate of 10 K min−1. No endothermic peaks, corresponding to melting points, were observed in the scanning cycles. These results indicate a good possibility for these [emim][AHA] ILs to work at low temperatures. 15032
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
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The Journal of Physical Chemistry B Figure 1 shows the electrochemical windows (ECWs) of the [emim][AHA] ILs measured by cyclic voltammetry at ambient
Figure 2. Density of [emim][AHA] ILs as a function of temperature. a From ref 17. bFrom ref 15.
[emim][3triz] due to their structural similarities. The density of [emim][3CH35CF3Pyra] is noticeably lower than that of [emim][3CF3Pyra]. This should be attributed to increased dispersion interaction caused by the introduction of the alkyl substitution on the heterocyclic ring, which is in agreement with previous studies.29,30 The densities of [emim][AHA] ILs studied here are in general 18−30% lower than that of [emim][(CF3SO2)2N] (the density of [emim][(CF3SO2)2N] is 1.51 g/cm3 at 303.15 K).24 The molar concentration at 303.15K (M30) was calculated from the molecular weight and density of each IL and is also listed in Table 3. M30 of the [emim][AHA] ILs studied in this work ranges from 4.62 to 6.53 mol L−1, which is considerably higher than that of many imidazolium, pyridinium, pyrrolidinium and ammonium RTILs with the [(CF 3 SO 2 ) 2 N] − anion (2.77−4.15 mol L−1 ) reported.20,21 This is primarily because the AHA anions have lower molecular weights than the [(CF3SO2)2N]− anion. Viscosity. Figure 3 presents the temperature dependence of the viscosities for the [emim][AHA] ILs over the temperature range from 283.15 to 343.15 K. The full collection of the measured viscosity data can be found in the Supporting Information (Table S3). Consistent with previous studies,19−21 the experimental viscosity data are fit well by the Vogel− Fulcher-Tammman (VFT) equation31,32
Figure 1. Electrochemical windows of [emim][AHA] ILs measured at ambient temperature at a scanning rate of 100 mV/s on a glassy carbon working electrode. aFrom ref 17.
temperature at a scanning rate of 100 mV/s. The values of the ECWs, as well as the cathodic reduction potential (Epc) and anionic oxidation potential (Epa) of the ILs can be found in the Supporting Information (Table S1). As shown by Figure 1, Epc of the [emim][AHA] ILs is very close to each other, indicating that the variation of anions does not noticeably affect the oxidation potential of the cation. As expected, Epa of the [emim][AHA] ILs differ considerably with the various anions. The ECW values of the ILs follow the order of [tetz]− > [3CF3Pyra]− > [3triz]− ≥ [3CH35CF3Pyra]− > [4triz]− > [4NO2Pyra]− > [2CNPyr]− > [3NO2Pyra]−. It is interesting to notice that while the ECW of [emim][3triz] is noticeably larger than that of [emim][4triz], the ECW of [emim][3NO2Pyra] is smaller than that of [emim][4NO2Pyra]. Figure 2 depicts the temperature dependent densities of the [emim][AHA] ILs measured over the temperature range from 283.15 to 343.15 K. The original density data can be found in the Supporting Information (Table S2). The densities of the ILs have been found to decrease linearly with increasing temperature as described by a linear equation ρ = a + bT
⎤ ⎡ B η = η0 exp⎢ ⎥ ⎣ (T − T0) ⎦
(3)
where η0 (mPa·S), B (K), and T0 (K) are fitting parameters. The best fit parameters for eq 3 are tabulated in Table 4, and the solid lines in Figure 3 represent the profiles from the VFT equation. T0 values have been found to be close to Tg for each IL (−42K to 17 K deviation from Tg), which agrees well with earlier researchers’ observation (e.g., 49 K below Tg).33 The viscosities of [emim][AHA] ILs with [4triz]−, [3triz]−, [tetz]−, [3CF3Pyra]−, and [2CNPyr]− anions are close to each other in the temperature range of the measurements and are considerably lower than those of the other three ILs in the order of [3NO2Pyra]− < [4NO2Pyra]− < [3CH35CF3Pyra]−. Interestingly, the viscosity of [emim][4NO2Pyra] is noticeably higher than that of [emim][3NO2Pyra], while the viscosities of [emim][4triz] and [emim][3triz] are similar. It was found by simulation that the location of substituent groups on the imidazolium ring of ILs can alter the dynamics in unpredictable ways,22 which is most likely the same case here. The viscosity of [emim][3CH35CF3Pyra] is substantially higher than that of
(2)
where T is temperature and a and b are fitting parameters with the best fit values summarized in Table 3. The densities of the [emim][AHA] ILs decrease in the order of [3CF3Pyra]− > [4NO2Pyra]− ≥ [3NO2Pyra]− > [3CH35CF3Pyra]− > [tetz]− > [4triz]− ≥ [3triz]− > [2CNPyr]−. The densities of [emim][4NO2Pyra] and [emim][3NO2Pyra] are very close to each other, which is the same case for [emim][4triz] and 15033
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
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The Journal of Physical Chemistry B Table 3. Density Equation Parameters and Molar Concentration at 303.15 K (M30) ρ = a + bT a/g cm−3
ILs [emim][2CNPyr]a [emim][4triz]a [emim][3triz]b [emim][4NO2Pyra] [emim][3NO2Pyra] [emim][3CF3Pyra] [emim][3CH35CF3Pyra] [emim][tetz]b a
1.268 1.321 1.312 1.427 1.418 1.474 1.435 1.360
± ± ± ± ± ± ± ±
b/10−4 g cm−3 K−1 −6.069 −6.098 −6.000 −6.802 −6.755 −7.640 −7.496 −6.198
0.000 0.001 0.000 0.001 0.001 0.000 0.001 0.002
± ± ± ± ± ± ± ±
0.010 0.017 0.000 0.023 0.020 0.012 0.028 0.062
R2
M30/10−3 mol cm−3
1.0000 1.0000 1.0000 0.9999 0.9999 1.0000 0.9999 0.9993
5.357 6.339 6.302 5.466 5.436 5.046 4.620 6.526
Fitted from ref 17. bFitted from ref 15.
⎡ ⎤ B σ = σ0 exp⎢ ⎥ ⎣ (T − T0) ⎦
(4)
where σ0 (S/cm), B (K), and T0 (K) are fitting parameters, wherein T0 once again was found to be close to Tg (30 K below Tg).33 Figure 4 shows the ionic conductivities of the
Figure 3. Temperature-dependent viscosity of [emim][AHA] ILs. a From ref 17. bFrom ref 15.
Table 4. VFT Equation Parameters of Viscosity Data η = η0 exp[B/(T − T0)] η0/m Pas
ILs a
[emim][2CNPyr] [emim][4triz]a [emim][3triz]b [emim][4NO2Pyra] [emim][3NO2Pyra] [emim][3CF3Pyra] [emim] [3CH35CF3Pyra] [emim][tetz]b a
0.63 0.73 1.32 0.16 0.14 0.09 0.38
± ± ± ± ± ± ±
0.14 0.09 0.23 0.06 0.06 0.02 0.09
1.24 ± 0.22
B/102 K 3.75 3.92 3.29 6.91 6.89 7.76 5.23
± ± ± ± ± ± ±
0.34 0.20 0.28 0.77 0.89 0.49 0.33
3.48 ± 0.30
R2
T0/K 218 214 215 200 195 181 220
± ± ± ± ± ± ±
3 2 3 5 7 4 2
0.9998 0.9999 0.9998 0.9999 0.9999 1.0000 1.0000
211 ± 4
0.9998
Figure 4. Temperature-dependent conductivity of [emim][AHA] ILs. a From ref 17. bFrom ref 15.
[emim][AHA] ILs as a function of temperature on a logarithmic scale and the solid lines in Figure 4 represent the profiles from the VFT equation. The original data can be found in the Supporting Information (Table S4). Table 5 lists the best fit parameters of eq 4 for the measured conductivity data. The ionic conductivities of [emim][AHA] ILs with [2CNPyr]−,
Fitted from ref 17. bFitted from ref 15.
Table 5. VFT Equation Parameters of Ionic Conductivity Data
[emim][3CF3Pyra] because the introduction of the methyl group on the anion induces additional van der Waals interactions, which subsequently increases the friction between the ions. The viscosities of the [emim][AHA] ILs studied here are generally 1.9−7.9 times that of [emim][(CF3SO2)2N] at 303.15 K (viscosity of [emim][(CF3SO2)2N] at 303.15 K is 27 mPa·s24) and 2.1−9.7 times that of [emim][BF4] at 298.15 K (viscosity of [emim][BF4] at 298.15 K is 32 mPa·s34). Ionic Conductivity. Conductivity is one of the most important aspects to evaluate an IL for electrochemical applications. As such, conductivities of the [emim][AHA] ILs were measured over the same temperature range as other transport properties. Analogous to the viscosity, the temperature dependence of the ionic conductivities for each of the [emim][AHA] ILs exhibits convex curved profiles and also can be well described by the VFT equation28,32
σ = σ0 exp[−B/(T − T0)] ILs [emim][2CNPyr]a [emim][4triz]a [emim][3triz]b [emim] [4NO2Pyra] [emim] [3NO2Pyra] [emim][3CF3Pyra] [emim] [3CH35CF3Pyra] [emim][tetz]b a
15034
σ0/102 mS/cm 8.04 8.41 10.08 7.69
± ± ± ±
0.77 2.41 0.65 0.27
B/102 K 4.47 4.82 5.72 5.19
± ± ± ±
0.22 0.67 0.17 0.08
R2
T0/K 207 203 187 209
± ± ± ±
3 8 2 1
0.9999 0.9995 1.0000 1.0000
7.91 ± 0.37
5.01 ± 0.11
204 ± 1
1.0000
6.77 ± 0.20 8.99 ± 2.07
5.09 ± 0.07 6.81 ± 0.58
195 ± 1 198 ± 5
1.0000 0.9999
11.23 ± 0.25
6.21 ± 0.08
178 ± 1
1.0000
Fitted from ref 17. bFitted from ref 15. DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
Article
The Journal of Physical Chemistry B [tetz]−, [3triz]−, and [4triz]− anions are close to each other, slightly higher than that of [emim][3CF3Pyra], and are considerably higher than those of the other three ILs in the order of [3NO2Pyra]− > [4NO2Pyra]− > [3CH35CF3Pyra]−, which coincides with the reverse order of the viscosities. The ionic conductivities of [emim][AHA] ILs discussed here are generally 0.29−0.72 that of [emim][(CF3SO2)2N] at 303.15 K (the conductivity for [emim][(CF3SO2)2N] at 303.15 K is 11 mS/cm24) and 0.17−0.46 that of [emim][BF4] at 298.15 K (the conductivity of [emim][BF4] at 298.15 K is 14 mS/cm34), while the conductivity of [emim][3CF35CF3Pyra] is only 0.13 that of [emim][(CF3SO2)2N] at 303.15 K and 0.07 that of [emim][BF4] at 298.15 K, respectively. Walden Plot. A common way in the literature to qualitatively estimate the ionicity of ILs or other electrolytes is the empirical Walden Rule,15,35−37 which correlates the conductivity and viscosity, as shown in eq 5 Λη = constant
(5)
where Λ is the equivalent conductivity, calculated as Λ = σVE (VE is the volume containing one Faraday of positive charge, which could be calculated equivalently as VE = M/(ρn) where ρ is the density of the electrolyte, n is the moles of positive charges per mole of the electrolyte) and η is the viscosity. Figure 5 shows the Walden plot of the [emim][AHA] ILs
Figure 6. Temperature-dependent self-diffusion coefficients of [emim][AHA] ILs for (a) cations and (b) anions.
where D0 (cm2/s), B (K), and T0 (K) are adjustable fitting parameters with the best fit values listed in Table 6 for the diffusivities of the cations and anions, Dcation and Danion, and the simple summation of the cation and anion diffusivities, Dcation + Danion. The solid lines in Figure 6 are calculated with the best fit parameters and eq 6. The cationic self-diffusion coefficients of [emim][AHA] ILs follow the order of [3CF3Pyra]− > [3triz]− ≈ [2CNPyr]− ≈ [tetz]− > [4triz]− > [3NO2Pyra]− > [4NO2Pyra]− > [3CH35CF3Pyra]−. The anionic self-diffusion coefficients follow the order [2CNPyr]− ≈ [3CF3Pyra]− ≈ [tetz]− ≈ [3triz]− ≈ [4triz]− > [3NO2Pyra]− > [4NO2Pyra]− > [3CH35CF3Pyra]−. Dcation, Danioin, and Dcation + Danioin (not plotted) self-diffusion coefficients generally change in an order that coincides with the reverse order of the viscosities, as expected and consistent with previous studies.21,34 The ionic size effect19 on the self-diffusion coefficients does not hold for the present set of ILs with a same cation, as the sizes of [4triz]− and [3triz]− are the same while the [3triz]− diffuses noticeably faster than [4triz]−, as shown in Figure 6b. Therefore, the diffusion of the ions is affected not only by the size of the ions but other factors, such as the shapes of the ions, electron distribution, and local interactions between the cations and the anions, as well. The transference number, which directly describes the portion of charge carried by a specific ion species, is an important index to evaluate an electrolyte for electrochemical applications.12,37,38 By definition, the apparent cationic transference number, t+, in an IL can be calculated by using the selfdiffusion coefficients as follows Dcation t+ = (Dcation + Danion) (7)
Figure 5. Walden plot of [emim][AHA] ILs. aFrom ref 17. bFrom ref 15. cFrom ref 20.
discussed. As expected, all the [emim][AHA] ILs fall below the so-called ideal line that runs from corner to corner on the Walden plot and is determined by data for 1 M KCl aqueous solution except a few points for [emim][4NO2Pyra] and [emim][3CH35CF3Pyra] at elevated temperatures. This indicates that strong ion interactions exist in the ILs. However, the variation of the anions seems to only have a trivial effect on the ionicity of the ILs as the ILs all fall in the same range in relative to [emim][(CF3SO2)2N] on the Walden plot. Self-Diffusion Coefficient. Figures 6 show the temperature dependence of the self-diffusion coefficients of the cation (Dcation) and anion (Danion) for the [emim][AHA] ILs. The original data can be found in the Supporting Information (Table S5). Similar to viscosity and ionic conductivity, the diffusivity of the ILs also follows the VFT equation ⎡ ⎤ B D = D0 exp⎢ ⎥ ⎣ (T − T0) ⎦
(6) 15035
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
Article
The Journal of Physical Chemistry B Table 6. VFT Equation Parameters for Self-Diffusion Coefficient Data D = D0 exp[−B/(T − T0)] D0/10−4 cm2 s−1
ILs [emim][2CNPyr]
[emim][4triz]
[emim][3triz]
[emim][4NO2Pyra]
[emim][3NO2Pyra]
[emim][3CF3Pyra]
[emim][3CH35CF3Pyra]
[emim][tetz]
Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation Dcation Danion Dcation
+ Danion
+ Danion
+ Danion
+ Danion
+ Danion
+ Danion
+ Danion
+ Danion
1.86 4.26 4.78 1.10 2.33 3.13 1.42 2.47 3.63 0.65 0.69 1.34 2.21 2.18 4.39 1.99 1.84 4.65 0.79 2.10 2.14 1.10 2.88 3.33
± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±
B/102 K
0.35 1.19 1.02 0.07 0.67 0.32 0.81 1.57 2.16 0.50 0.58 1.07 2.09 1.95 4.05 0.97 0.62 2.17 0.23 0.65 0.54 0.33 1.07 1.06
8.40 10.89 9.13 7.84 9.88 8.76 8.89 10.84 9.75 6.68 6.84 6.76 8.98 9.06 9.02 9.02 9.63 9.86 6.45 9.84 7.56 8.34 11.52 9.67
± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±
0.51 0.87 0.61 0.18 0.87 0.29 1.72 2.09 1.87 1.85 2.03 1.93 2.65 2.49 2.57 1.47 1.03 1.45 0.67 0.86 0.61 0.87 1.27 1.00
T0/K
R2
± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±
1.0000 1.0000 1.0000 1.0000 1.0000 1.0000 0.9997 0.9997 0.9997 0.9991 0.9990 0.9991 0.9995 0.9995 0.9995 0.9998 0.9999 0.9998 0.9999 1.0000 1.0000 0.9999 0.9999 0.9999
178 161 173 180 165 173 163 151 158 197 196 196 178 178 178 165 162 159 205 180 196 168 145 158
4 6 5 2 7 2 15 16 15 16 18 17 20 19 20 12 8 11 6 6 5 8 9 8
[4triz]−. Note that the transference numbers for [emim][4NO2Pyra] and [emim][4triz] are very close to 0.5. By comparison, the transference numbers for [emim][(CF3SO2)2N] are between 0.62−0.64 over the same temperature range20 and are noticeably higher than those of the [emim][AHA] ILs studied here. This is largely because the planar AHAs diffuse faster than the chain-structured [(CF3SO2)2N]− anion. Molar Conductivity. The molar conductivity ratio, Λimp/ ΛNMR, defined as the molar conductivity calculated from the ionic conductivity and the molar concentration (Λimp) versus that estimated from the self-diffusion coefficients using the Nernst−Einstein equation (ΛNMR), is another commonly used index to quantify the ionicity of an electrolyte.
Apparent cationic transference numbers of the [emim][AHA] ILs are shown in Figure 7 as a function of temperature
ΛNMR =
NAe 2(Dcation + Danion) F 2(Dcation + Danion) = kT RT (8)
Figure 7. Apparent cationic transference number of [emim][AHA] ILs as a function of temperature.
where NA is the Avogadro number, e is the electric charge on each ion, Dcation and Danion are the self-diffusion coefficients of the cation and anion, respectively, k is the Boltzmann constant, F is the Faraday constant, R is the universal gas constant, and T is the absolute temperature. The assumption behind the ΛNMR value is that all of the diffusing species detected by NMR contribute to the molar conductivity, while the Λimp value originates from the migration of the charged species, or ions, in an electric field. Because PGSE NMR measurement detects a nucleus (i.e., 1H) and cannot distinguish between the ions and their charge neutral aggregates or clusters, the Λimp/ΛNMR ratio indicates the percentage of ions contributing to the conduction among the diffusing species on the time scale of the measurement and gives a measure of the degree of cation− anion aggregation in ILs at equilibrium. As shown in Figure 8a, the Λimp values for the present [emim][AHA] ILs follow the
to compare the self-diffusion coefficients between the cation and the anion. Interestingly, the cationic self-diffusion coefficients are at least a bit larger than the anionic selfdiffusion coefficients for all of the ILs over the entire temperature range of measurements making the apparent cationic transference numbers range between 0.5−0.6, which indicates that the cation, even though it has a larger radius, can diffuse faster than the anion. This is also reflected by the larger B values in the VFT fitting for the anions than the cations (Table 6), indicating that the activation energy for the diffusion of the anion is greater than that for the cation for an individual IL. The cationic transference numbers approximately follows the order of [3CH35CF3Pyra]− ≈ [3CF3Pyra]− > [3triz]− > [tetz]− ≈ [2CNPyr]− > [3NO2Pyra]− > [4NO2Pyra]− ≈ 15036
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
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Table 7. VFT Equation Parameters for Molar Conductivity Data ILs [emim][2CNPyr]a [emim][4triz]a [emim][3triz]b [emim] [4NO2Pyra] [emim] [3NO2Pyra] [emim] [3CF3Pyra] [emim] [3CH35CF3Pyra] [emim][tetz]b [emim][2CNPyr] [emim][4triz] [emim][3triz] [emim] [4NO2Pyra] [emim] [3NO2Pyra] [emim] [3CF3Pyra] [emim] [3CH35CF3Pyra] [emim][tetz]
Figure 8. Temperature-dependent molar conductivity of [emim][AHA] ILs (a) obtained from ionic conductivity and molar concentration and (b) calculated from ion self-diffusion coefficients and the Nernst−Einstein equation. aFrom ref 17. bFrom ref 15.
a
Λ0/102 S cm2 mol−1 Λimp= 1.71 ± 1.50 ± 1.99 ± 1.61 ±
B/102 K
Λ0 exp[−B/(T − T0)] 0.16 4.67 ± 0.22 0.44 5.01 ± 0.70 0.13 6.04 ± 0.18 0.05 5.39 ± 0.07
T0/K
R2
± ± ± ±
3 8 2 1
1.0000 0.9996 1.0000 1.0000
206 202 184 207
1.67 ± 0.07
5.22 ± 0.10
203 ± 1
1.0000
1.57 ± 0.04
5.36 ± 0.07
193 ± 1
1.0000
2.27 ± 0.52
7.07 ± 0.59
196 ± 5
0.9999
176 ± 1
1.0000
2.01 ± 0.06 6.52 ± 0.09 ΛNMR= Λ0 exp[−B/(T − T0)] 7.62 ± 1.35 7.52 ± 0.47 5.15 ± 0.45 7.25 ± 0.23 5.19 ± 2.52 7.72 ± 1.41 2.57 ± 1.85 5.76 ± 1.64
181 181 169 202
4 2 13 16
1.0000 1.0000 0.9997 0.9989
6.71 ± 3.88
7.33 ± 1.49
187 ± 13
0.9996
5.77 ± 1.88
7.47 ± 0.92
173 ± 9
0.9998
4.72 ± 1.14
6.91 ± 0.57
198 ± 5
0.9999
4.95 ± 1.31
7.76 ± 0.78
168 ± 7
0.9999
± ± ± ±
b
Fitted from ref 17. Fitted from ref 15.
order of [2CNPyr]− ≥ [tetz]− ≥ [3triz]− ≥ [4triz]− > [3CF 3 Pyra] − > [ 3NO 2 Py ra] − > [4NO 2 Pyra] − > [3CH35CF3Pyra]−, which agrees well with the ionic conductivity trend. Figure 8b demonstrates that the ΛNMR values for the ILs discussed follow the same order as for the simple summation of the cationic and anionic self-diffusion coefficients with [3CF3Pyra]− ≥ [3triz]− ≈ [2CNPyr]− ≈ [tetz]− > [4triz]− > [3NO2Pyra]− > [4NO2Pyra]− > [3CH35CF3Pyra]−. Both molar conductivities can be well described by the VFT equation ⎤ ⎡ B Λ = Λ 0 exp⎢ ⎥ ⎣ (T − T0) ⎦
Figure 9. Temperature-dependent molar conductivity ratio of [emim][AHA] ILs. aContains data from ref 17. bContains data from ref 15.
(9)
where Λ0 (S cm2 mol−1), B (K) and T0 (K) are fitting parameters, whose best fit values are summarized in Table 7. Figure 9 represents the molar conductivity ratios for the present ILs as a function of temperature. Consistent with previous studies,19−21 the molar conductivity ratios exhibit little dependence on temperature over the temperature range from 283.15 to 333.15 K and range between 0.38 and 0.93 for the ILs in this study. It is noteworthy that the Λimp/ΛNMR values appear to drop considerably from 333.15 to 343.15 K for all the ILs. We believe this is an artifact; it is most likely because convection becomes important in our system at this elevated temperature as has been observed by other researchers.39,40 We have attempted to avoid the convection effect by reducing the sample loading height to 2 cm and by reducing the delay time between the two gradient pulses (Δ), but without much success. Because this work is not dedicated to reducing convection effects in PGSE NMR measurements, we only used
the self-diffusion coefficient data over the temperature range of 283.15 to 333.15 K, neglecting the data point at 343.15 K for the VFT fitting discussed above. Nonetheless, all of the [emim][AHA] ILs give Λimp/ΛNMR values less than unity, indicating the presence of some ion pairing and aggregation in the ILs. The Λimp/ΛNMR value for [emim][2CNPyr] is considerably higher than the other seven [emim][AHA] ILs. This could be explained by the fact that [emim][2CNPyr] has moderate self-diffusion coefficients and a substantially lower molar concentration than the other five ILs. Also it ranks among the highest for the ionic conductivity, which means that the ions in [emim][2CNPyr] are more dissociated and therefore more effective in conduction. Interestingly, while the Λimp/ΛNMR value for [emim][4triz] is higher than that for 15037
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
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[emim][3triz], [emim][4NO2Pyra] exhibits a considerably lower Λimp/ΛNMR value compared with [emim][3NO2Pyra]. This again indicates that the location of substituent groups on the ions can alter the dynamics of an IL in unpredictable ways.22 The introduction of a methyl group to the [3CF3Pyra]− anion on one hand leads to a decrease in the molar concentration (Table 3), which reduces the Coulombic attractive forces between the cation and anion. On the other hand, the increase of hydrocarbon units increases the van der Waals interactions through both the alkyl chains-ion inductive forces and the hydrocarbon-hydrocarbon interactions, among which the former tends to be predominant in RTILs.20 The Λimp/ΛNMR value of [emim][3CH35CF3Pyra] is remarkably lower than that of [emim][3CF3Pyra], indicating a larger influence from the intermolecular inductive forces compared with the Coulombic attractions. By comparison, the Λimp/ΛNMR value for [emim][(CF3SO2)2N] is between 0.73−0.76 over the same temperature range.20 It is worth pointing out that in contrast to similar placement on the Walden plot in Figure 5, Figure 9 indicates that the choice of the anion significantly affects the ionicity of the IL. This reveals that the two ways (the Walden plot and the Λimp/ΛNMR value) that are commonly used in the literature to measure the ionicity of ILs actually report on different things. The Λimp/ΛNMR values appear to be somewhat more discriminating in identifying attractive IL electrolytes.
Article
ASSOCIATED CONTENT
S Supporting Information *
The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpcb.5b09175. Tables of electrochemical windows (ECWs), density, viscosity, ionic conductivity, and self-diffusion coefficient of the ILs at various temperature. (PDF)
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AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS This work is supported by the University of Notre Dame Incropera-Remick Endowment for Excellence. The authors acknowledge Dr. Jaroslav Zajicek from the Notre Dame Magnetic Resonance Research Center for help with the PGSE NMR measurements.
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REFERENCES
(1) Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. Solubility of CO2, CH4, C2H6, C2H4, O2, and N2 in 1-Hexyl-3-Methylpyridinium Bis(trifluoromethylsulfonyl)imide: Comparison to Other Ionic Liquids. Acc. Chem. Res. 2007, 40, 1208−1216. (2) Armand, M.; Endres, F.; MacFarlane, D. R.; Ohno, H.; Scrosati, B. Ionic-Liquid Materials for the Electrochemical Challenges of the Future. Nat. Mater. 2009, 8, 621−629. (3) Fernicola, A.; Scrosati, B.; Ohno, H. Potentialities of Ionic Liquids as New Electrolyte Media in Advanced Electrochemical Devices. Ionics 2006, 12, 95−102. (4) Plechkova, N. V.; Seddon, K. R. Applications of Ionic Liquids in the Chemical Industry. Chem. Soc. Rev. 2008, 37, 123−150. (5) Kawano, R.; Katakabe, T.; Shimosawa, H.; Nazeeruddin, M. K.; Graetzel, M.; Matsui, H.; Kitamura, T.; Tanabe, N.; Watanabe, M. Solid-state Dye-Sensitized Solar Cells Using Polymerized Ionic Liquid Electrolyte with Platinum-Free Counter Electrode. Phys. Chem. Chem. Phys. 2010, 12, 1916−1921. (6) Lewandowski, A.; Swiderska-Mocek, A. Ionic Liquids as Electrolytes for Li-Ion Batteries-An overview of Electrochemical Studies. J. Power Sources 2009, 194, 601−609. (7) Nanjundiah, C.; McDevitt, S. F.; Koch, V. R. Differential Capacitance Measurements in Solvent-Free Ionic Liquids at Hg and C Interfaces. J. Electrochem. Soc. 1997, 144, 3392−3397. (8) Sakaebe, H.; Matsumoto, H. N-Methyl-N-propylpiperidinium Bis(trifluoromethanesulfonyl)imide (PP 13-TFSI) - Novel Electrolyte Base for Li Battery. Electrochem. Commun. 2003, 5, 594−598. (9) Ue, M.; Takeda, M.; Toriumi, A.; Kominato, A.; Hagiwara, R.; Ito, Y. Application of Low-Viscosity Ionic Liquid to the Electrolyte of Double-Layer Capacitors. J. Electrochem. Soc. 2003, 150, A499−A502. (10) de Souza, R. F.; Padilha, J. C.; Goncalves, R. S.; Dupont, J. Room Temperature Dialkylimidazolium Ionic Liquid-Based Fuel Cells. Electrochem. Commun. 2003, 5, 728−731. (11) Endres, F. A.; MacFarlane, D. R.; Abbott, A. Electrodeposition from Ionic Liquids; Wiley-VCH: Weinheim, Germany, 2008. (12) Liao, C.; Shao, N.; Han, K. S.; Sun, X. G.; Jiang, D. E.; Hagaman, E. W.; Dai, S. Physicochemical Properties of ImidazoliumDerived Ionic Liquids with Different C-2 Substitutions. Phys. Chem. Chem. Phys. 2011, 13, 21503−21510. (13) Every, H. A.; Bishop, A. G.; MacFarlane, D. R.; Oradd, G.; Forsyth, M. Transport Properties in a Family of Dialkylimidazolium Ionic Liquids. Phys. Chem. Chem. Phys. 2004, 6, 1758−1765. (14) Moganty, S. S.; Chinthamanipeta, P. S.; Vendra, V. K.; Krishnan, S.; Baltus, R. E. Structure-Property Relationships in Transport and
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CONCLUSION Transport properties of a family of AHA RTILs consisting of eight different anions paired with the [emim]+ cation have been studied systematically. Measurements include electrochemical windows, densities, viscosities, ionic conductivities, and cation and anion self-diffusion coefficients from PGSE NMR spectroscopy. The main conclusion is that many of the [emim][AHA] ILs show very good conductivity for their viscosities. This is frequently described in the literature as ILs with high “ionicity”. Specifically, the molar conductivity ratio, which compares the ionic conductivity to the total transport of the cations and anions from PGSE NMR (Λimp/ΛNMR), is as high as 0.93 for [emim][2CNpyr]. This is much higher than values of 0.73−0.76 observed for the popular [emim][(CF3SO2)2N] IL in the same temperature range. In fact, five of the [emim][AHA] ILs have Λimp/ΛNMR values as high or higher than [emim][(CF3SO2)2N]. This is at least partly due to good diffusion of the anions. The fraction of the diffusivity coming from the cation (i.e., cation transference number, Dcation/(Dcation + Danion)) is between 0.5 and 0.6 for the [emim][AHA] ILs, compared to 0.62−0.64 for [emim][(CF3SO2)2N]. This means that in the [emim][AHA] ILs the diffusivities of the anions are almost as high as the [emim]+ cations, which is not the case for [emim][(CF3SO2)2N]. In addition, the [emim][AHA] ILs show good “ionicity” on the Walden Plot and have ionic conductivities as high as 6.4 mS/cm at room temperature. The interrelations among the microscopic ionic diffusivity, macroscopic viscosity, and ionic conductivity demonstrate that the macroscopic physicochemical properties of the ILs are significantly affected by the microscopic ion state and dynamics. Ionicity is determined not only by the ion size, but the ion shape, electron distribution, and local interactions between the cation and anion. The results of this study help to establish a relationship between the structures and transport properties of AHA RTILs and provide insight into rational design of new ILs for practical electrochemical applications. 15038
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and a Critique of the So-Called Ideal KCl Line for the Walden Plot. J. Chem. Eng. Data 2010, 55, 1784−1788. (36) Xu, W.; Cooper, E. I.; Angell, C. A. Ionic liquids: Ion Mobilities, Glass Temperatures, and Fragilities. J. Phys. Chem. B 2003, 107, 6170− 6178. (37) Zugmann, S.; Fleischmann, M.; Amereller, M.; Gschwind, R. M.; Winter, M.; Gores, H. J. Salt Diffusion Coefficients, Concentration Dependence of Cell Potentials, and Transference Numbers of Lithium Difluoromono(oxalato)borate-Based Solutions. J. Chem. Eng. Data 2011, 56, 4786−4789. (38) Petrowsky, M.; Frech, R. Concentration Dependence of Ionic Transport in Dilute Organic Electrolyte Solutions. J. Phys. Chem. B 2008, 112, 8285−8290. (39) Mohoric, A.; Stepisnik, J. Effect of Natural Convection in a Horizontally Oriented Cylinder on NMR Imaging of the Distribution of Diffusivity. Phys. Rev. E: Stat. Phys., Plasmas, Fluids, Relat. Interdiscip. Top. 2000, 62, 6628−6635. (40) Hayamizu, K.; Price, W. S. A New type of Sample Tube for Reducing Convection Effects in PGSE-NMR Measurements of SelfDiffusion Coefficients of Liquid Samples. J. Magn. Reson. 2004, 167, 328−333.
Thermodynamic Properties of Imidazolium Bistriflamide Ionic Liquids for CO2 Capture. Chem. Eng. J. 2014, 250, 377−389. (15) Seo, S.; DeSilva, M. A.; Brennecke, J. F. Physical Properties and CO2 Reaction Pathway of 1-Ethyl-3-Methylimidazolium Ionic Liquids with Aprotic Heterocyclic Anions. J. Phys. Chem. B 2014, 118, 14870− 14879. (16) Seo, S.; Quiroz-Guzman, M.; DeSilva, M. A.; Lee, T. B.; Huang, Y.; Goodrich, B. F.; Schneider, W. F.; Brennecke, J. F. Chemically Tunable Ionic Liquids with Aprotic Heterocyclic Anion (AHA) for CO2 Capture. J. Phys. Chem. B 2014, 118, 5740−5751. (17) Shi, C.; Quiroz-Guzman, M.; DeSilva, A.; Brennecke, J. F. Physicochemical and Electrochemical Properties of Ionic Liquids Containing Aprotic Heterocyclic Anions Doped With Lithium Salts. J. Electrochem. Soc. 2013, 160, A1604−A1610. (18) Shi, C. Development of Ionic Liquids for Li-ion Battery Applications. Ph.D. Dissertation, University of Notre Dame, Notre Dame, Indiana, 2013. (19) Tokuda, H.; Hayamizu, K.; Ishii, K.; Abu Bin Hasan Susan, M.; Watanabe, M. Physicochemical Properties and Structures of Rroom Temperature Ionic Liquids. 1. Variation of Anionic Species. J. Phys. Chem. B 2004, 108, 16593−16600. (20) Tokuda, H.; Hayamizu, K.; Ishii, K.; Susan, M.; Watanabe, M. Physicochemical Properties and Structures of Room Temperature Ionic Liquids. 2. Variation of Alkyl Chain Length in Imidazolium Cation. J. Phys. Chem. B 2005, 109, 6103−6110. (21) Tokuda, H.; Ishii, K.; Susan, M.; Tsuzuki, S.; Hayamizu, K.; Watanabe, M. Physicochemical Properties and Structures of RoomTemperature Ionic Liquids. 3. Variation of Cationic Structures. J. Phys. Chem. B 2006, 110, 2833−2839. (22) Zhang, Y.; Maginn, E. J. Direct Correlation between Ionic Liquid Transport Properties and Ion Pair Lifetimes: A Molecular Dynamics Study. J. Phys. Chem. Lett. 2015, 6, 700−705. (23) Wang, X. J.; Chi, Y. L.; Mu, T. C. A Review on the Ttransport Properties of Ionic Liquids. J. Mol. Liq. 2014, 193, 262−266. (24) Tokuda, H.; Tsuzuki, S.; Susan, M.; Hayamizu, K.; Watanabe, M. How Ionic are Room-Temperature Ionic Liquids? An Indicator of the Physicochemical Properties. J. Phys. Chem. B 2006, 110, 19593− 19600. (25) Barthel, J.; Feuerlein, F.; Neueder, R.; Wachter, R. Calibration of Conductance Cells at Various Temperatures. J. Solution Chem. 1980, 9, 209−219. (26) Zick, K. NMR Diffusion User Manual. Bruker BioSpin, version 003; Bruker Corporation: Billerica, MA, 2009 (27) Stejskal, E. O.; Tanner, J. E. Spin Diffusion Measurements: Spin Echoes in the Presence of a Time-Dependent Field Gradient. J. Chem. Phys. 1965, 42, 288−292. (28) Tanner, J. E. Use of Stimulated Echo in NMR-diffusion Studies. J. Chem. Phys. 1970, 52, 2523−2526. (29) Fredlake, C. P.; Crosthwaite, J. M.; Hert, D. G.; Aki, S.; Brennecke, J. F. Thermophysical Properties of Imidazolium-Based Ionic Liquids. J. Chem. Eng. Data 2004, 49, 954−964. (30) Rodriguez, H.; Brennecke, J. F. Temperature and Composition Dependence of the Density and Viscosity of Binary Mixtures of Water Plus Ionic Liquid. J. Chem. Eng. Data 2006, 51, 2145−2155. (31) Fulcher, G. S. Analysis of Recent Measurements of the Viscosity of Glasses. J. Am. Ceram. Soc. 1925, 8, 339−355. (32) Vogel, H. The Temperature Dependence Law of the Viscosity of Fluids. Phys. Z. 1921, 22, 645−646. (33) Hayamizu, K.; Aihara, Y.; Nakagawa, H.; Nukuda, T.; Price, W. S. Ionic Conduction and Ion Diffusion in Binary Room-temperature Ionic Liquids Composed of emim BF4 and LiBF4. J. Phys. Chem. B 2004, 108, 19527−19532. (34) Noda, A.; Hayamizu, K.; Watanabe, M. Pulsed-gradient Spinecho H-1 and F-19 NMR Ionic Diffusion Coefficient, Viscosity, and Ionic Conductivity of Non-Chloroaluminate Room-temperature Ionic Liquids. J. Phys. Chem. B 2001, 105, 4603−4610. (35) Schreiner, C.; Zugmann, S.; Hartl, R.; Gores, H. J. Fractional Walden Rule for Ionic Liquids: Examples from Recent Measurements 15039
DOI: 10.1021/acs.jpcb.5b09175 J. Phys. Chem. B 2015, 119, 15030−15039
