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Structural and Ion Dynamics in Fluorine-Free Oligoether
Carboxylate Ionic Liquid-Based Electrolytes
Faiz Ullah Shah,* Oleg I. Gnezdilov, Inayat Ali Khan, Andrei Filippov, Natalia A. Slad,
and Patrik Johansson*
Cite This: J. Phys. Chem. B 2020, 124, 9690−9700
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sı Supporting Information
*
ABSTRACT: Here, we investigate the physicochemical and electrochemical
properties of fluorine-free ionic liquid (IL)-based electrolytes with two different
cations, tetrabutylphosphonium, (P4,4,4,4)+, and tetrabutylammonium, (N4,4,4,4)+,
coupled to a new anion, 2-[2-(2-methoxyethoxy)ethoxy]acetate anion (MEEA)−,
for both neat and (P4,4,4,4)(MEEA) also doped with 10−40 mol % of Li(MEEA).
We find relatively weaker cation−anion interactions in (P4,4,4,4)(MEEA) than in
(N4,4,4,4)(MEEA), and for both ILs, the structural flexibility of the oligoether
functionality in the anion results in low glass transition temperatures, also for the
electrolytes made. The pulsed field gradient nuclear magnetic resonance (PFG
NMR) data suggest faster diffusion of the (MEEA)− anion than (P4,4,4,4)+ cation
in the neat IL, but the addition of a Li salt results in slightly lower mobility of the
former than the latter and lower ionic conductivity. This agrees with the
combined 7Li NMR and attenuated total reflection−Fourier transform infrared
(ATR−FTIR) spectroscopy data, which unambiguously reveal preferential
interactions between the lithium cations and the carboxylate groups of the IL anions, which also increased as a function of the
lithium salt concentration. In total, these systems provide a stepping stone for further design of fluorine-free and low glass transition
temperature IL-based electrolytes and also stress how crucial it is to control the strength of ion−ion interactions.
(OCH2CH2)n−OR], that often are abbreviated as Gn where
“n” represents the number of repeating units and denoted as
monoglyme (G1), diglyme (G2), triglyme (G3), tetraglyme
(G4), etc.5 As compared to carbonate-based organic solvents,
glymes have higher flash points and better solubility for lithium
salts6they can even dissolve an equimolar amount of lithium
salt via an ether multidentate coordination functionality
making complexes known as solvate ionic liquids (SILs) as
they mimic ionic liquids (ILs) with high thermal and
electrochemical stabilities and high ionic conductivities.7,8
Such alkali ion−glyme complexes can even be reversibly
intercalated inside graphite electrodes to enable charge storage
based on solvent cointercalation.9,10 Tetraglymes have been
suggested to be stable against nucleophilic attack by superoxides and therefore are used in electrolytes for lithium−air
batteries.11 As they also form very thin and flexible solid
electrolyte interphase (SEI) layers on highly reactive alkali
metal anode surfaces, they are also considered for next-
INTRODUCTION
The liquid electrolytes used in conventional lithium-ion
batteries (LIBs) are most often composed of the salt lithium
hexafluorophosphate (LiPF6), or sometimes lithium tetrafluoroborate (LiBF4), salts that dissolve in flammable carbonatebased organic solvents.1,2 Such electrolytes have a number of
drawbacks adversely affecting the battery performance when
used and also cause difficulties at the recycling stage. The most
problematic properties are they are intrinsically flammable,
have unsatisfactory electrochemical stability windows (ESWs),
not the least for future high-voltage LIBs, and have limited
thermal stability. A particularly problematic aspect of the latter
is the decomposition of LiPF6 at elevated temperatures
producing lithium fluoride (LiF) and phosphorous pentafluoride (PF5), the latter forming toxic hydrofluoric acid
(HF)a serious hazard that must be dealt with.3 Taken
altogether, this urges for extensive research to develop
thermally and electrochemically stable, but still performant,
electrolytes to replace the conventional organic solvent-based
electrolytes.
Among the organic solvent-based electrolytes studied,
different oligoethers (also known as “glymes”) have been
extensively studied over the past decades, much due to their
desirable physicochemical properties.4 Oligoethers is a class of
organic compounds with the general chemical formula, [R−
■
© 2020 American Chemical Society
Received: May 26, 2020
Revised: September 21, 2020
Published: October 20, 2020
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everything from thermal, phase, and electrochemical stabilities
as well as the local interactions and dynamics.
A wide range of characterization techniques are employed
here to study the physicochemical and electrochemical
properties to enable a better understanding of the fundamental
potential of these systems as LIB electrolytes.
generation batteries (NGBs) such as lithium−air and lithium−
sulfur batteries,12 but they suffer from instability at high redox
potentials.13
Here, we combine the features of glyme solvents with those
of ILs. ILs have for long been used as promising electrolyte
solvents for LIBs14−16 and NGBs.17,18 ILs are salts made
entirely of cations and anions with a melting point below 100
°C, and most are liquid at room temperature: room
temperature ionic liquids (RTILs). As touched upon briefly
above, ILs/RTILs possess a combination of properties making
them excellent bases for electrolytes such as a nonflammability,
negligible vapor pressures, high chemical and thermal
stabilities, inherent high ionic conductivities, and wide
electrochemical stability windows (ESWs)sometimes
quoted up to 6.0 V.19−21
IL-based LIB electrolytes can thus arguably overcome some
of the drawbacks of conventional liquid electrolytes. The most
commonly studied IL-based LIB electrolytes use heavily
fluorinated anions, bis(trifluoromethanesulfonyl)imide
(TFSI) and more recently bis(fluorosulfonyl)imide (FSI).22
Typically, these electrolytes are made by dissolving LiTFSI/
LiFSI in ILs such as pyrrolidinium TFSI or imidazolium
TFSI,23 pyrrolidinium FSI,24 morpholinium FSI,25 piperidinium TFSI,26 or phosphonium FSI.27 The use of fluorinated
anions, however, makes these electrolytes less stable and
sensitive to moisture,28 and thus fluorine-free IL-based
electrolytes are highly desired.
We exploit a combination of oligoethers and ILs by creating
ILs with tetraalkyl phosphonium, (PR4)+, and ammonium,
(NR4)+, cations coupled to a new common fluorine-free
oligoether-based anion, 2-[2-(2-methoxyethoxy)ethoxy]acetate, (MEEA)−, and the corresponding Li-salt, Li(MEEA)
(Figure 1). The (MEEA)− anion is designed aiming at IL-
EXPERIMENTAL SECTION
Electrolyte Preparation. The fluorine-free oligoether
carboxylate-based Li-salt and ILs were synthesized by a slight
modification of the procedure described earlier.32 The
synthesis protocols and material characterizations including
multinuclear (1H, 13C, 7Li, and 31P) nuclear magnetic
resonance (NMR) spectroscopy, mass spectrometry, and
elemental analysis are described in detail in the Supporting
Information. The electrolytes were prepared by mixing 10−50
mol % of Li(MEEA) in (P4,4,4,4)(MEEA) (Table 1), but phase
■
Table 1. Compositions and Abbreviations of the
Electrolytes Made
electrolyte acronym
[Li(MEEA)]0.1[(P4,4,4,4)
(MEEA)]0.9
[Li(MEEA)]0.2[(P4,4,4,4)
(MEEA)]0.8
[Li(MEEA)]0.3[(P4,4,4,4)
(MEEA)]0.7
[Li(MEEA)]0.4[(P4,4,4,4)
(MEEA)]0.6
Li(MEEA)
(mol %)
(P4,4,4,4)(MEEA)
(mol %)
molality
(mol kg−1)
10
90
0.254
20
80
0.572
30
70
0.981
40
60
1.526
separation was observed for the highest salt concentration and
this is therefore not further included in this study. All the
samples were kept in a vacuum oven at 80 °C for 3−6 days
until the water content was less than 200 ppm as determined
by the Karl Fischer titration using a 917 coulometer
(Metrohm).
Nuclear Magnetic Resonance Spectroscopy. The
structure and purity of the synthesized Li-salt and ILs were
characterized using a Bruker Ascend Aeon WB 400 (Bruker
BioSpin AG, Fällanden, Switzerland) NMR spectrometer.
CDCl3 was used as a solvent. The working frequencies were
400.21 MHz for 1H, 100.64 MHz for 13C, 162.01 MHz for 31P,
and 155.53 MHz for 7Li. The 7Li spectra of the neat
electrolytes were recorded by placing the samples in a 5 mm
standard NMR tube, which was further placed inside a 10 mm
standard NMR tube containing CDCl3. The 7Li NMR spectra
were indirectly referenced to 1.0 M LiCl(aq). Data were
processed using Bruker Topspin 3.5 software.
Pulsed Field Gradient Diffusometry. Pulsed field
gradient (PFG) NMR measurements were performed on a
Bruker Avance III (Bruker BioSpin AG, Fällanden, Switzerland) NMR spectrometer. The working frequencies for 1H and
7
Li were 400.27 and 155.56 MHz, respectively. Data were
processed using Bruker Topspin 3.1 software. NMR selfdiffusion measurements were performed on 1H and 7Li with a
PFG NMR probe Diff50 (Bruker) with a maximum amplitude
of the magnetic field gradient pulse of 29.73 T m−1. The
sample was placed in a standard 5 mm glass sample tube and
closed with a plastic stopper to avoid contact with air. Prior to
measurements, each sample was equilibrated at a specific
temperature for 30 min.
Figure 1. Chemical structures and acronyms for the ILs cations and
anions.
based electrolytes with low glass transition temperatures, due
to the presence of ethylene oxide units, which might enhance
the flexibility of the alkyl chain, as previously demonstrated for
IL cations.29 The reason for using both (N4,4,4,4)+ and (P4,4,4,4)+
IL cations is that although the cationic central atoms are largely
shielded by the alkyl chains, they still have a significant effect
on the IL properties mainly due to the differences in charge
density of the nitrogen (higher) and phosphorous (lower)
atoms.30 This was analyzed by Coutinho and co-workers using
density functional theory (DFT) calculations for (N4,4,4,6)(NTf2) and (P4,4,4,6)(NTf2) ILs, and they found stronger
Coulomb interactions and more positively charged connecting
methylene groups, that is, a larger charge delocalization for the
former cation.31 This might have profound effects on
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Figure 2. (a) TGA thermograms and (b) DSC traces for the neat (N4,4,4,4)(MEEA) and (P4,4,4,4)(MEEA) ILs and the [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes. DSC traces are shifted along the y axis for clarity.
Electrochemical Characterization. The ESWs and the
ionic conductivities were measured using a Metrohm Autolab
PGSTAT302N electrochemical workstation with an FRA32M
module for impedance measurements, all controlled using
Nova 2.02 software. About 70 μL of the sample was placed in a
sealed Microcell HC from RHD instruments. Cyclic
voltammetry (CV) and linear sweep voltammetry (LSV)
experiments were carried out at 20 °C temperature using a
three electrode setup: 2 mm diameter glassy carbon (GC) and
a platinum wire with a diameter of 0.25 mm as working
electrodes (WEs), Pt crucible as a sample container as well as a
counter electrode (CE), and an Ag wire coated with AgCl was
used as a pseudoreference electrode (RE). The cyclic
voltammograms were recorded at 100 mV s−1, while the linear
sweeps used were 1 mV s−1. The electrochemical potentials
were recorded with ferrocene as an internal reference and
shifted using ELi/Li+ ≈ EFc/Fc+ + 3.2 V.37 The ESW limits were
determined using 0.1 and 0.2 mA cm−2 cutoff current
densities.38
The ionic conductivities were determined from impedance
measurements performed in a frequency range of 1 Hz to 1
MHz with an AC voltage amplitude of 10 mVrms and from −20
to 100 ± 0.1 °C; while the heating and cooling cycles match
very well (Figure S17), the data presented are taken from the
heating cycles. A two-electrode setup was used with a 2 mm
diameter GC WE and a 70 μL Pt crucible as a sample
container as well as a CE. Both the electrodes were polished
with a Kemet diamond paste 0.25 μm prior to each
measurement. The cell constant was determined using a 100
μS cm−1 KCl standard solution from Metrohm (Kcell = 1.486
cm−1). The cell was thermally equilibrated for at least 10 min
before each measurement.
The details of the PFG NMR technique for measuring
molecular diffusion coefficients can be found elsewhere.33 A
primary source of information about the diffusivity of a
molecule is the diffusion decay (DD) of amplitude A of the
NMR spectral line, obtained by Fourier transformation of the
descending half of stimulated echo (STE), as a function of the
amplitude of the applied pulsed field gradient. For a stimulated
echo pulse sequence used, the diffusion decay of A in the case
of a simple nonassociating molecular liquid can be described
by eq 134
A(g , δ , td) = A(0)exp( −γ 2g 2δ 2Dtd)
(1)
where A(0) is the factor proportional to the proton content in
the system and to spin−lattice and spin−spin relaxation times;
γ is the gyromagnetic ratio for a used nucleus; g and δ are the
amplitude and duration of the gradient pulse; td is the diffusion
time; and D is the self-diffusion coefficient. In our experiments,
td was in the range of 4 to 100 ms for 1H diffusion, while td was
in the range of 200 to 700 ms for 7Li diffusion. No diffusion
time dependence was observed in these measurements.
Infrared Spectroscopy. Attenuated total reflection−Fourier transform infrared (ATR−FTIR) spectra were recorded on
a Bruker IFS 80v spectrometer equipped with a deuterated
triglycine sulphate (DTGS) detector and a diamond ATR
accessory. All spectra were recorded at room temperature
(∼22 °C) using the double-side forward−backward acquisition
mode. A total number of 256 scans were coadded and signalaveraged at an optical resolution of 4 cm−1.
Thermal Analysis. Thermogravimetric analysis (TGA) was
performed using PerkinElmer 8000 TGA apparatus. The
dynamic TGA experiments used a heating rate of 10 °C min,
nitrogen gas as the inert carrier gas, and 2−4 mg of the sample.
The onset of decomposition temperature, Tonset, was calculated
from the intersection of the baseline weight and the tangent of
the weight versus temperature curve using Pyris software.35,36
Differential scanning calorimetry (DSC) was performed using
PerkinElmer DSC 6000 apparatus. About 2−5 mg of the
sample was packed in an aluminum pan for each experiment.
DSC data were recorded during both cooling and heating
traces from −100 to 100 °C at a scanning rate of 5 °C min−1.
The glass transition temperature, Tg, was determined as the
onset of the transition. An inert nitrogen gas was supplied to
the instrument at a constant flow of 20 mL min−1 in order to
preserve a dry environment inside the sample chamber.
RESULTS AND DISCUSSION
For both ILs and electrolytes, we start by first assessing the
phase, thermal, and electrochemical stabilities before addressing more direct LIB relevant performance parameters such as
ionic conductivity and diffusion together with ion−ion
interactions. A more or less complete picture of these
electrolytes as promising LIB electrolytes is hereby created.
Thermal Properties. As thermal stability is one of the key
properties where IL-based electrolytes are superior to conventional organic solvent-based electrolytes, all systems, (N4,4,4,4)(MEEA), (P4,4,4,4)(MEEA), and [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x), were assessed by TGA and all remain stable
at >200 °C and most are stable at >300 °C (Figure 2a and
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Table 2. Cathodic Limiting Potentials, Anodic Limiting Potentials, ESWs, Glass Transition Temperatures, and Thermal
Decomposition Temperatures of the Neat ILs and the IL-Based Electrolytes
0.1 mA cm−2
IL/electrolyte
(N4,4,4,4)(MEEA)
(P4,4,4,4)(MEEA)
[Li(MEEA)]0.1[(P4,4,4,4)(MEEA)]0.9
[Li(MEEA)]0.2[(P4,4,4,4)(MEEA)]0.8
[Li(MEEA)]0.3[(P4,4,4,4)(MEEA)]0.7
[Li(MEEA)]0.4[(P4,4,4,4)(MEEA)]0.6
+
0.2 mA cm−2
+
EC (V vs Li/Li ) EA (V vs Li/Li )
2.20
1.11
1.08
1.07
0.98
1.06
5.07
4.75
4.73
4.87
4.74
4.81
+
ESW (V) EC (V vs Li/Li ) EA (V vs Li/Li+) ESW (V)
2.87
3.64
3.65
3.80
3.76
3.75
1.74
0.88
0.84
0.82
0.70
0.69
5.38
5.0
4.99
5.08
5.0
5.20
3.64
4.12
4.15
4.26
4.30
4.51
Tg (°C) Tonset (°C)
−67
−80
−78
−77
−77
−76
204
317
330
334
337
340
ESWs of the electrolytes as the difference can arise only due to
the IL cation. However, the widths of the ESWs (ca. 4.1−4.5
V) are anyhow quite remarkable, especially when compared to
the standard LIB electrolyte, 1 M LiPF6 in EC:DMC (EC 1.3 V
vs Li/Li+, EA 5.0 V vs Li/Li+, and ESW 3.7 V).48 These ESWs
are comparable with the commonly studied imidazolium- and
phosphonium-based ILs containing TFSI and other fluorinated
anions.49 The ESWs of all these electrolytes are narrower on
the surface of the platinum electrode as compared with the GC
electrode (Figure S14 and Table S1). Yet, all these ESWs
should be seen as upper limits as we are using a sweep
technique and the ILs are viscous; and furthermore, in a real
LIB system, the electrodes might catalyze reactions earlier and
hence the ESWs become narrower. By using CV, the long-term
electrochemical stability and reversibility were confirmed for
any electrochemical events in the neat (P4,4,4,4)(MEEA) IL and
the electrolytes (Figures S11−S13). No significant changes are
observed in the CV curves of (P4,4,4,4)(MEEA) IL even after
100 cycles on both GC and platinum electrode surfaces. The
shoulder observed at ca. 4.8 V versus Li/Li+ in the CVs is due
to oxidation of a previous oxidation or reduction product of
the IL.50
Ionic Conductivity. While the ionic conductivities are
always high for both neat ILs and IL-based electrolytes, it is
significantly higher for (P4,4,4,4)(MEEA) than for (N4,4,4,4)(MEEA) throughout all temperatures (Figure 4a, Table S2). It
is expected because (P4,4,4,4)(MEEA) has lower viscosity and
density as compared with (N4,4,4,4)(MEEA) (Figures S15 and
S16). The impedance spectra of some representative samples
are shown in Figures S18 and S19. For the electrolytes, the
ionic conductivity, as expected, decreases as a function of the
salt concentration as increasing the Li+ concentration reduces
the free volume and increases the ion−ion Coulombic
interactions, primarily between the lithium ion and the
(MEEA)− anion (Figure 4b, Table S2).51,52 In addition, the
Li+ ion will most likely be coordinated with multiple (MEEA)−
anions, carboxylate, and/or ether groups, leading to dynamic
cross-linking and thus enhanced ion association. A similar
cross-linking effect was observed when Na-salt was added to an
alkoxy-ammonium-based IL.53
The ionic conductivities can be further analyzed by fitting
the data to the Vogel−Fulcher−Tammann (VFT) equation
Table 2). While these stabilities probably are somewhat
overestimated, as for more exact thermal stability determination of IL-based electrolytes, isothermal TGA must be
applied,39,40 they are anyhow highly encouraging in comparison with the standard LIB electrolyte “LP30”, that is, 1 M
LiPF6 in EC:DMC that starts to decompose already at
temperatures <100 °C with the formation of LiF and PF5.41 As
all systems have the (MEEA)− anion in common, it is clear
that the phosphonium cation is thermally more stable than the
ammonium cation, which is consistent with the literature.42,43
The addition of Li(MEEA) increases the thermal stability,
more or less linearly as a function of composition and by a
maximum of +23 °C for 40 mol % added.
The corresponding DSC traces reveal all systems to be glassforming liquids, that is, they all have glass transitions (Figure
2b and Table 2). It is also clear that (N4,4,4,4)(MEEA) has a
higher ionic strength and/or stronger ion−ion interactions,
and thus, a higher thermal energy would be required to reach
the same ionic mobility as for (P4,4,4,4)(MEEA). As expected,44
increasing Tg values are observed as a function of Li(MEEA)
addition, but they show very small increasesmaximum of +3
°C for 40 mol % of Li(MEEA) (Table 2). The overall low Tg
values can be attributed to the ethylene oxide units present in
the (MEEA)− anion, allowing low energy rotations,45 and for
all the electrolytes, this seems to dominate over the increasing
content of “hard”, dynamically cross-linking Li+ ions, which
most often would cause much larger increases in Tg.46,47
Electrochemical Stability. The ESW of each sample is
obtained from cathodic and anodic LSV experiments. The
ESWs are revealed to be wider for (P4,4,4,4)(MEEA) than for
(N4,4,4,4)(MEEA) (Figure 3 and Table 2). This also affects the
ij −B yz
zz
σ = σ0expjjj
j (T − T0) zz
k
{
(2)
where σ0, B, and T0 are the fitting parameters, a preexponential factor, a factor related to the activation/
pseudoactivation energy, and the dynamic glass transition
temperature, respectively. The latter indicates the temperature
at which the free volume is reduced so that molecular mobility
Figure 3. Linear sweep voltammetry of the neat ILs and the
[Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes using GC as the
WE.
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Figure 4. Ionic conductivity as a function of temperature for (a) neat ILs and (b) [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes. The solid lines
indicate the best fit of data using the Vogel−Fulcher−Tammann (VFT) equation.
Figure 5. Diffusion coefficients of ions in (a) neat (P4,4,4,4)(MEEA) IL and [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes: (b) the (P4,4,4,4)+
cation, (c) the (MEEA)− anion, and (d) Li+ ion. The solid lines indicate the best fits of diffusion data using the VFT equation.
is arrested.54 The activation energy for ionic conductivity (Eσ)
is related to B as Eσ = B·R. The fitting procedure was carried
out over the full temperature range in two steps. In the first
step, we plot ln(σ) versus 1/(T−T0) and selected T0 to have a
linear dependence. In the second step, we fit the dependence
by a linear regression to obtain the fitting parameters (σ0, B).
The resulting VFT parameters (Table S3) show that the Eσ
for (N4,4,4,4)(MEEA) is slightly higher compared to that of
(P4,4,4,4)(MEEA), which agrees with the DSC data and again
indicates that a higher thermal energy is required to reach the
same ion mobility. Similarly, as expected, the Eσ increases as a
function of the salt concentration for the [Li-
(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes. Building further
on the comparison with the DSC data, the experimental Tg
values are higher than T0, but Tg − T0 is in the range of 51−57
K and T0/Tg is in the range of 0.71−0.75, which agree well
with the empirical approximation for ILs being Tg − T0 ≈ 50 K
and T0/Tg ≈ 0.7555 and also for previously observed IL-based
electrolytes.53
NMR Diffusometry. Having established the overall
mobilities by DSC and impedance spectroscopy at a more
macroscopic level, PFG NMR diffusometry was employed to
better understand the mobility at the molecular level, and
especially, the influence of Li+ addition in the IL-based
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Figure 6. Diffusion coefficients of the (P4,4,4,4)+ cation, the (MEEA)− anion, and the Li+ ion in the (a) [Li(MEEA)]0.1[(P4,4,4,4)(MEEA)]0.9 and (b)
[Li(MEEA)]0.4[(P4,4,4,4)(MEEA)]0.6 electrolytes. The solid lines indicate the best fits of diffusion data using the VFT equation.
The apparent transference numbers of the individual ions
are determined from their diffusion coefficients in neat
(P4,4,4,4)(MEEA) and the electrolytes using the following
equation56,57
electrolytes. The PFG NMR diffusion coefficients of all species,
(P4,4,4,4)+, (MEEA)−, and Li+ show monotonous increases as a
function of temperature (Figures 5 and 6, and Table S4) and
show the VFT behavior. In the case of neat (P4,4,4,4)(MEEA),
the anion diffusivity is faster than the cation diffusivity in the
whole temperature range (Figure 5a), following the relative ion
sizes. Upon Li-salt doping, however, while both the IL cation
and anion diffusivities decrease linearly as a function of the salt
concentration (Figure 5b,c), the diffusivity of the (MEEA)−
anion is affected to a much greater extent, suggesting Li+−
(MEEA)− interactions. This is corroborated by the decreased
diffusivity of the Li+ ion with the increasing salt concentration
(Figure 5d).
Comparing the [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) x = 0.1
and x = 0.4 electrolytes directly (Figure 6, and Figures S20 and
S21) show that the diffusivities of all ions decrease with the salt
concentration. For the less concentrated electrolyte, the
diffusivity of (MEEA)− is still slightly higher than for
[P4,4,4,4]+ (Figure 6a), but for the most concentrated
electrolyte, it is vice versa (Figure 6b). Clearly, the Li+ ions
and the (MEEA)− anions move in a correlated fashion.
The temperature dependence of the diffusivity of ILs can be
described by another VFT equation
ji −B zyz
D = D0expjjj
z
j (T − T0) zz
k
{
ti =
xiDi
ΣixiDi
(4)
where ti is the apparent transference number, xi is the molar
fraction of each individual ion, and D is the self-diffusion
coefficient of an ion in m2 s−1. The transference number as
measured by PFG NMR is usually referred to as a transport
number as the diffusion coefficients measured are averages of
individual charged species, neutral species (ion pairs), and
aggregates. The latter diffuse slower due to their larger radii,
and this leads to lower diffusion coefficients and transference
numbersone of the reasons that transference numbers
measured by electrochemical techniques are usually twice as
large.58
The apparent transference numbers (ti) of Li+, (P4,4,4,4)+, and
(MEEA)− ions as a function of the Li-salt concentration and
temperature show that the (MEEA)− anion moves faster than
the (P4,4,4,4)+ cation (Figure 7, Table S6). The transference
number of Li+ increases linearly with the increasing Li-salt
concentration, while it decreases for the (P4,4,4,4)+ cation, which
is attributed to the increasing/decreasing number densities.59
However, no significant changes are seen for the (MEEA)−
anion.
(3)
where D0, T0, and B are adjustable parameters. The energy of
activation for the diffusion of ions is associated with B as ED =
B·R and as for the ionic conductivity, the fitting procedure was
carried out in two stepswith the same layout. The VFT
parameters (Table S5) show that the pre-exponential factor,
D0, for the (MEEA)− anion is five times larger than for the
(P4,4,4,4)+ cation, indicating it to be the main contributor for
ionic diffusivity. Although D0 monotonically decreases for all
ions with the addition of Li-salt, the effect is more prominent
for the (MEEA)− anion, confirming increasing interactions
between these ions. The values of T0 estimated from the
diffusion data are comparable to those calculated from the
ionic conductivity data and do not change significantly with
the addition of Li-salt. Similarly, the apparent activation energy
for diffusion, ED, changes slightly with the addition of Li-salt,
which reveals a minor contribution in the change of ionic
diffusion as compared with D0.
Figure 7. Apparent transfer numbers of ions in neat (P4,4,4,4)(MEEA)
and [Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes.
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(MEEA)−, we expect relevant features. In the former region,
there is, for (P4,4,4,4)(MEEA), a main band at 1673 cm−1 with
distinct satellite bands on both sides at 1713 and 1609 cm−1,
indicating several different types of coordination/interaction
active sites, while for (N4,4,4,4)(MEEA), there is a single, albeit
broad, feature at 1725 cm−1, which is more similar to a clear
carbonyl group interaction (Figure 9a).66 The symmetric
stretching band of the carboxylate group appears as a single
and symmetrical peak at 1463 cm−1 for (P4,4,4,4)(MEEA) and
for (N4,4,4,4)(MEEA) also with a shoulder at 1489 cm−1, again
suggesting stronger interactions with the carboxylate group.
The stronger and single mode of interaction also agrees well
with the lower ionic conductivity of (N4,4,4,4)(MEEA).
Looking further at the spectral regions sensitive to the
oligoether functional groups of the (MEEA)− anion, a C−O
stretching band appears at 1099 cm−1 for (P4,4,4,4)(MEEA) and
it is slightly upshifted for (N4,4,4,4)(MEEA) (Figure 9b). At the
lower wavenumbers in the range of 900 to 800 cm−1, the
infrared absorptions represent a mixture of CH2 rocking and
CO stretching motions.67 These vibrations are rather sensitive
to any structural changes and shift upon interactions with a
counter cation.68 The absorption bands at 850 and 817 cm−1
for (P4,4,4,4)(MEEA) IL are shifted toward higher wavenumbers
at 879 and 853 cm−1 for (N4,4,4,4)(MEEA) IL, suggesting much
stronger interactions of the (MEEA)− anion with the (N4,4,4,4)+
cation than the (P4,4,4,4)+ cation. A comparison of the full range
FTIR spectra of neat (P4,4,4,4)(MEEA) and (N4,4,4,4)(MEEA)
ILs is shown in Figure S27.
The addition of Li(MEEA) to (P4,4,4,4)(MEEA) changes the
interactions; the bands at 1713 and 1673 cm−1 both disappear
and the peak at 1609 cm−1 increases, which also is
approximately the region of the pure Li(MEEA) salt feature
(Figure 9c). This change clearly shows that the Li+ ion
interacts strongly with the carboxylate group of the (MEEA)−
anion, but the symmetric stretching band region is much
difficult to analyze. In contrast, neither the C−O bond
stretching region nor the 800−1000 cm−1 region changes to
any large extent (Figure 9d) although significant changes have
previously been observed in this region for mixtures of various
oligoethers and alkali metal salts.69−71 Here the Li+ ions seem
to be only weakly interacting with the oxygen atoms of the
ethylene oxide units of the (MEEA)− anion. The only major
differences are increases in the intensity of the C−O stretching
band at around 1100 cm−1 as a function of the Li(MEEA)
concentration, quite naturally as it induces relatively higher
(MEEA) − concentrations. This picture of the cation
coordination is in support of the DSC analysis, suggesting
that the addition of Li(MEEA) has a negligible effect on the
flexibility of the oligoether backbone within the (MEEA)−
anion, which is also in great accordance with the findings for
the alkali metal salts of the 2,5,8,11-tetraoxatridecan-13-oate
anion, (TOTO)−, which is a rather similar anion, showing
strong coordination with the carboxylate group and moderate
interactions with the oligoether groups.72
The ratio of the molar conductivity obtained from the
impedance measurements (Λimp) and the molar conductivity
calculated from the diffusion coefficients of ions (ΛNMR) is
commonly used to investigate the ionic association and
dissociation under equilibrium in RTILs. The Λimp/ΛNMR
ratio indicates the percentages of ions responsible for the ion
conduction within the diffusing unit, that is, the ionicity.60
Here, the Λimp/ΛNMR ratio for neat (P4,4,4,4)(MEEA) is
investigated as a function of temperature (Figure S26). The
molar conductivity is calculated from the self-diffusion
coefficients using the Nernst−Einstein equation:61
ΛNMR = F 2(Dcation + Danion)/RT
Article
(5)
where F is the Faraday constant and R is the universal gas
constant.
The Λimp/ΛNMR ratio being lower than unity indicates the
presence of ionic association and that only a fraction of the
diffusing species contributes to the ionic conduction.62 The
ratio decreases as a function of temperature showing more
ionic association, which is sensitive to temperature, as
expected.63
7
Li NMR Spectroscopy. 7Li NMR spectroscopy was
employed to get further insights into the local environment
of the Li+ ions and the interactions with the (MEEA)− anions.
Both the chemical shift and the line width change upon the
addition of Li-salt, suggesting that the local environment
changes as well (Figure 8). The downshifts in the 7Li NMR
Figure 8. 7 Li NMR spectra of the [Li(MEEA)] x [(P 4,4,4,4 )(MEEA)](1−x) electrolytes.
spectra agree with previous observations for LiFSI-doped
(C3mpyr)(FSI).64 In contrast, the 31P NMR spectra show only
a single resonance line for the (P4,4,4,4)+ cation, suggesting no
significant changes (Figures S22−S25), most likely due to the
fact that the alkyl chains wrap around and shield the cation
positive center from the anion. The line widths of the 7Li
NMR spectra increase with the salt concentration and as the
homogeneous NMR line broadening is determined by the
molecular motions,65 it reveals that the lithium ions become
less mobile in the electrolytes with higher salt concentrations,
which is in accordance with both the ionic conductivity and
the ion diffusivity data.
Infrared Spectroscopy. To get an additional and deeper
insight into the ion−ion interactions, FTIR spectroscopy was
applied, and especially in the regions of the asymmetric and the
symmetric stretching bands of the carboxylate group of
CONCLUSIONS
Among the neat ILs, (P4,4,4,4)(MEEA) has more beneficial
properties than (N4,4,4,4)(MEEA) such as much better thermal
and electrochemical stabilities, a lower glass transition
temperature, and higher ionic conductivity. As is not
uncommon to IL-based electrolytes, both the total ionic
conductivity and the diffusivity of the Li+ ions decrease as a
function of the salt concentration. The local coordination
■
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Article
Figure 9. FTIR spectra of the (a and b) neat (N4,4,4,4)(MEEA) and (P4,4,4,4)(MEEA) ILs and (c and d) (P4,4,4,4)(MEEA), Li(MEEA), and the
[Li(MEEA)]x[(P4,4,4,4)(MEEA)](1−x) electrolytes.
points to strong ion−ion interactions, albeit to a very varying
degree for the three different cations, with the (MEEA)−
anionbut all primarily with the carboxylate group.
Altogether, while the strong cation−anion interactions might
be less beneficial in terms of charge carrier concentration and
in terms of creating a dynamically cross-linked structure, they
render the high flexibility of the oligoethylene glycol moiety of
the (MEEA)− anion intactand hence more globally dynamic
electrolytes, as seen by the small changes induced in the glass
transition temperatures. The combination of the latter and
FTIR data revealed that ion−ion interactions can even be used
to look at how the different IL cations interact with the
(MEEA)− anion. Overall, the here created IL-based electrolytes have promising potential to be used in batteries operating
over a wide range of temperature and electrochemical
potential.
■
spectra of the neat electrolytes; ratio of molar
conductivities measured by impedance and NMR
spectroscopies, FTIR spectra of neat (P4,4,4,4)(MEEA)
and (N4,4,4,4)(MEEA) ILs; Tables of cathodic limiting
potentials, anodic limiting potentials, and ESWs; ionic
conductivity; VFT equation parameters; diffusivity data;
and transference numbers of the ionic liquids and the
electrolytes (PDF)
■
AUTHOR INFORMATION
Corresponding Authors
Faiz Ullah Shah − Chemistry of Interfaces, Luleå University of
Technology, SE-971 87 Luleå, Sweden; orcid.org/00000003-3652-7798; Email: faiz.ullah@ltu.se
Patrik Johansson − Department of Physics, Chalmers University
of Technology, SE-412 96 Gothenburg, Sweden; ALISTOREEuropean Research Institute, 80039 Amiens, France;
orcid.org/0000-0002-9907-117X;
Email: patrik.johansson@chalmers.se
ASSOCIATED CONTENT
sı Supporting Information
*
The Supporting Information is available free of charge at
https://pubs.acs.org/doi/10.1021/acs.jpcb.0c04749.
Authors
Oleg I. Gnezdilov − Institute of Physics, Kazan Federal
University, 420008 Kazan, Russia
Inayat Ali Khan − Chemistry of Interfaces, Luleå University of
Technology, SE-971 87 Luleå, Sweden; orcid.org/00000002-7940-7297
Andrei Filippov − Chemistry of Interfaces, Luleå University of
Technology, SE-971 87 Luleå, Sweden; Medical and Biological
Synthesis of the lithium salt and the ionic liquids; 1H,
13
C, 31P, and 7L NMR spectra; CV and LSV curves;
dynamic viscosity and density of neat (P4,4,4,4)(MEEA)and (P4,4,4,4)(MEEA) ILs; heating and cooling cycles of
the ionic conductivity of (P4,4,4,4)(MEEA); Nyquist
plots; comparison of the diffusivity of ions; 31P NMR
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Physics, Kazan Medical University, 420012 Kazan, Russia;
orcid.org/0000-0002-6810-1882
Natalia A. Slad − Institute of Polymers, Kazan National
Research Technological University, 420015 Kazan, Russia
Complete contact information is available at:
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Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS
The financial support from the Swedish Research Council
(project number: 2018-04133) is gratefully acknowledged. P.J.
is grateful for the continuous support from the Chalmers Areas
of Advance Materials Science, Energy, and Transport. O.I.G.
acknowledges the subsidy allocated to the Kazan Federal
University for the state assignment in the sphere of scientific
activities (project number: 0671-2020-0051).
■
■
Article
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