Thermophysical and Electrochemical Properties of Ethereal Functionalised Cyclic Alkylammonium-based Ionic Liquids as Potential Electrolytes for Electrochemical Applications

Article abstract of DOI:10.1002/cphc.201700246

A series of hydrophobic room temperature ionic liquids (ILs) based on ethereal functionalised pyrrolidinium, piperidinium and azepanium cations bearing the bis[(trifluoromethyl)sulfonyl]imide, [TFSI]^?, anion were synthesized and characterized. Their physicochemical properties such as density, viscosity and electrolytic conductivity, and thermal properties including phase transition behaviour and decomposition temperature have been measured. All of the ILs showed low melting point, low viscosity and good conductivity and the latter properties have been discussed in terms of the IL fragility, an important electrolyte feature of the transport properties of glass-forming ILs. Furthermore, the studied [TFSI]^?-based ILs generally exhibit good electrochemical stabilities and, by coupling electrochemical experiments and DFT calculations, the effect of ether functionalisation at the IL cation on the electrochemical stability of the IL is discussed. Preliminary investigations into the Li-redox chemistry at a Cu working electrode are also reported as a function of ether-functionality within the pyrrolidinium-based IL family. Overall, the results show that these ionic liquids are suitable for electrochemical devices such as battery systems, fuel cells or supercapacitors.

Full text of DOI:10.1002/cphc.201700246

Accepted Article
Title: Thermophysical and Electrochemical Properties of Ethereal
Functionalised Cyclic Alkylammonium-based Ionic Liquids as
Potential Electrolytes for Electrochemical Applications
Authors: Alex Ryan Neale, Sinead Murphy, Peter Goodrich,
Christopher Hardacre, and Johan Jacquemin
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Thermophysical and Electrochemical Properties of Ethereal
Functionalised Cyclic Alkylammonium-based Ionic Liquids as
Potential Electrolytes for Electrochemical Applications
Alex R. Neale,*[a] Sinead Murphy,[a] Peter Goodrich, [a] Christopher Hardacre[a, b] and Johan
Jacquemin*[a, c]
Abstract: A series of hydrophobic room temperature ionic liquids (ILs)
based on ethereal functionalised pyrrolidinium, piperidinium and
azepanium cations bearing the bis{(trifluoromethyl)sulfonyl}imide,
[TFSI]-, anion were synthesized and characterized. Their
physicochemical properties such as density, viscosity and electrical
conductivity and thermal properties including phase transition
behaviour and decomposition temperature have been measured. All
of the ILs showed low melting point, low viscosity and good
conductivity and the latter properties have been discussed in terms of
the IL fragility, an important electrolyte feature of the transport
properties of glassy-forming ILs. Furthermore, the studied [TFSI]--
based ILs generally exhibit good electrochemical stabilities and, by
coupling electrochemical experiments and DFT calculations, the
effect of ether functionalisation at the IL cation on the electrochemical
stability of the IL is discussed. Preliminary investigations into the Li-
redox chemistry at a Cu working electrode are also reported as a
function of ether-functionality within the pyrrolidinium-based IL family.
Overall, the results show that these ionic liquids are suitable for
electrochemical devices such as battery systems, fuel cells or
supercapacitors.
stability and, therefore, have been proposed as new electrolytes
for energy devices[7-18] and as solvents for electrodeposition of
metals.[19, 20]
Within the field of energy storage devices, a number of ILs have
been developed with specific physiochemical applications. To
date, however, most of these are based on imidazolium,[21]
23]
quaternary ammonium,[22,
pyridinium[24] and quaternary
phosphonium cations.[25, 26] Therein, the phosphonium, aliphatic
and cyclic ammonium cations show a higher resistance to
electrochemical reduction than corresponding imidazolium and
pyridinium analogues and are, therefore, considered to be more
promising as electrolytes. Within the cyclic aliphatic ammonium
cations, typically pyrrolidinium[27-29], piperidinium[28, 30] and more
32]
recently azepanium[31,
cationic based structures have been
chosen for a large number of investigations for potential
electrochemical devices. However, the downside to this improved
stability is an increase in viscosity which results in significant
drops in performance at mid-to-high current rates.[12, 13]
Attempts to reduce the viscosity of ILs based on cyclic
alkylammonium-[TFSI] based ILs have been achieved by
blending with appropriate organic solvents,[33-35] utilising
alternative anion structures,[14, 36, 37] and functionalisation of the
alkyl chains of the cyclic aliphatic ammonium cations with different
group functionalities. In this regard, some studies have
highlighted the positive influence of ethereal functionalisation on
the transport properties of acyclic and cyclic ammonium-based
ionic liquids.[37-41] In contrast, ionic liquids containing two or more
ether groups appended to the cation have been somewhat
understudied. Functionalised quaternary ammonium ILs based on
di-, tri- and tetra- ethereal-based cations have recently been
reported as promising candidates for energy applications.[42-45] In
addition, a series of novel pyrrolidinium and piperidinium ILs
based on cations with two identical mono-ether groups have been
developed which showed promising applications for energy
devices.[28]
In order to develop ILs with good physiochemical properties for
possible energy applications, we have investigated the effect of
cation structural modifications for a family of ILs based on
pyrrolidinium, piperidinium and azepanium cations functionalised
with two or four ether groups paired with the benchmark [TFSI]-
anion. The [TFSI]- anion was utilised herein owing to its,
reasonable cost and commercial availability and also due to the
very good promotion of hydrophobicity and thermal, chemical and
electrochemical stability of the resulting ILs. The physiochemical
and electrochemical properties of these new ILs have been
compared against the corresponding non-functionalised ionic
liquids.
1. Introduction
Ionic liquids have gained considerable interest over the past few
decades due to their numerous attractive properties such as
extremely low vapour pressure, low flammability, high thermal
stability, and large liquid range. This has led to a number of
groups researching ionic liquid applications particularly in the
fields of catalysis [1-3] and separations[4, 5] and nanotechnology.[6]
In comparison with molecular solvents, ionic liquids also possess
good ionic conductivity coupled with good electrochemical
[a]
A. R. Neale, S. Murphy, Dr. J. Jacquemin, Dr. P. Goodrich, Prof C.
Hardacre
School of Chemistry and Chemical Engineering
Queen’s University Belfast
David Keir Building
Stranmillis Road, Belfast, BT9 5AG (UK)
E-mail: aneale01@qub.ac.uk
johan.jacquemin@qub.ac.uk or jj@univ-tours.fr
[b]
[b]
Prof C. Hardacre
School of Chemical Engineering & Analytical Science
The University of Manchester
The Mill, Sackville Street, Manchester, M13 9PL (UK)
Dr. J. Jacquemin
Laboratoire PCM2E, Université François Rabelais
Parc de Grandmont, 37200, Tours (France)
Supporting Information and the ORCID identification number(s) for
the author(s) of this article can be found under http://dx.doi.org/....
This article is protected by copyright. All rights reserved.

10.1002/cphc.201700246
ChemPhysChem
ARTICLE
2. Results and Discussion
2.1. Synthesis of ether-functionalised cyclic-alkylammonium
ILs
Several synthetic strategies were undertaken to develop these ILs
with the maximum yield and purity. Initial attempts to synthesise
these materials using conventional alkylation procedures
employed for imidazolium-based ILs proved problematic due to
the low alkylation yield observed between the alkylamine and the
ethereal precursors. Therein, low to moderate yields of deeply
coloured ILs were achieved. Therefore, the proposed strategy
was to firstly synthesise the corresponding aminoethereal based
species and then perform the alkylation as the final step. These
aminoethereal compounds were easily synthesised using the
conventional Williamson procedure under aqueous conditions
(Scheme 1).
Initially, the secondary cyclic amine (pyrrolidine, piperidine or
azepane) was reacted with one molar equivalent of the desired
alkylating agent (RBr, where R = butyl, 2-methoxyethyl or 2-
(methoxyethoxy)ethyl). The reactions were carried out in water
which upon completion resulted in the formation of a monophasic
system. The corresponding hydrobromide salt was then treated
with potassium hydroxide and after stirring for ca. 24 h at room
temperature a biphasic system formed. The resulting crude amine
was then fractionally distilled in vacuo resulting in good yields
(61–71%). Mono-functionalised ILs were obtained by reacting
these tertiary amine compounds with dimethylsulfate generating
the corresponding methylsulfate ILs in almost quantitative yields.
Di-functionalised ILs were obtained by reacting the tertiary amine
compounds with either 2-methoxyethyl or 2-(methoxyethoxy)ethyl
bromide.
Figure 1. Chemical structures, and respective abbreviations, of the ether-
functionalised
cyclic
alkylammonium
cations
and
the
bis{(trifluoromethyl)sulfonyl}imide, [TFSI]-, anion.
These reactions proceeded with lower reactivity due to the
reduced nucleophilicty of the N-centre of the glyme-functionalised
tertiary amines. Metathesis of the bromide or methylsulfate anion-
containing ILs with lithium bis{(trifluoromethyl)sulfonyl}imide
(Li[TFSI]) in biphasic water/dichloromethane systems yielded the
corresponding [TFSI]- salts in good yields (92-97%). The prepared
[TFSI]- salts were dried in vacuo for at least 48 h and stored in a
glove box.
2.2. Characterisation of ether-functionalised cyclic-
alkylammonium ILs
This family of ILs were functionalised with different combinations
of alkyl and ethereal chains in order to introduce ethereal-like
properties to the liquid. The abbreviations [Pyrr]+, [Pip]+ and [Aze]+
are used to express the cations with 5, 6 and 7 membered
alkylammonium rings, respectively. The subscript text is used to
express two additional groups bonded to the nitrogen of the ring;
1 = methyl-, 4 = butyl-, (2o1) = methoxyethyl-, (2o2o1) =
methoxyethoxyethyl-. The chemical structure of all synthesised
ether-functionalised cations, and their respective abbreviations
are shown in Figure 1. Table 1 summarises the thermal behaviour
physiochemical properties at 298.15 K of all the ILs.
Thermal properties
The decomposition temperatures (Td) of the studied ILs was
determined by dynamic thermogravimetric analysis (TGA)
wherein the temperature of the sample is increased at a rate of
10 K·min-1. The Td values for range of studied ILs, determined as
the temperature at which a 5 % mass loss is recorded, were in the
range of 560 – 679 K. The TGA mass loss traces for each IL are
shown in Figure S1 in the electronic supplementary information
(ESI). For the appending alkyl group an increase in the number of
ether linkages generally resulted in a decrease in thermal stability.
The exception to this was the [Aze(2o2o1)2][TFSI] which showed
Scheme 1. Strategic routes for the synthesis of ether-functionalised cyclic-
alkylammonium ILs.
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Table 1. The thermal properties and physiochemical properties at 298.15 K of all the [TFSI]- ionic liquids.
Ionic liquid
Mr / g·mol-1
Tg / K
Td / K
ρ / g·cm-3
η / mPa·s
σ / mS·cm-1
Λ / S·cm2·mol-1
[Pyrr14][TFSI]
422.83
424.40
468.46
468.46
556.56
436.86
438.43
482.49
482.49
570.59
450.89
452.46
496.52
496.52
584.62
679
636
643
561
560
639
630
606
631
604
649
641
625
605
639
1.4054
1.4547
1.4137
1.4308
1.4010
1.3775
1.4543
1.3956
1.4087
1.3531
1.3729
1.4179
1.3949
1.3393
1.3527
77.15
49.22
2.72
3.69
2.80
1.72
1.24
1.04
1.74
1.62
1.63
0.98
0.60
1.10
1.09
0.70
0.56
0.818
1.076
0.929
0.563
0.494
0.329
0.526
0.560
0.557
0.411
0.198
0.350
0.388
0.259
0.241
[Pyrr1(2o1)][TFSI]
[Pyrr1(2o2o1)][TFSI]
[Pyrr(2o1)2][TFSI]
[Pyrr(2o2o1)2][TFSI]
[Pip14][TFSI]
187
51.19
198
199
198
200
193
206
201
205
198
195
207
201
74.58
124.49
170.71
104.46
83.64
[Pip1(2o1)][TFSI]
[Pip1(2o2o1)][TFSI]
[Pip(2o1)2][TFSI]
[Pip(2o2o1)2][TFSI]
[Aze14][TFSI]
108.01
137.42
307.15
134.87
127.80
135.74
209.99
[Aze1(2o1)][TFSI]
[Aze1(2o2o1)][TFSI]
[Aze(2o1)2][TFSI]
[Aze(2o2o1)2][TFSI]
the presence of the cis-conformer of [TFSI]- anions in favour of a
system predominantly containing the lower energy trans-[TFSI]-
conformer. All of the other ILs investigated showed no observable
Tf or Tm. The majority of the ILs, with the exception of
[Pyrr14][TFSI] and [Pyrr1(2o2o1)][TFSI], showed glass transition
temperatures (Tg) between 187 K and 207 K. Figure 3 shows the
DSC thermogram of [Aze1(2o1)][TFSI] as a representative example
of a typical trace observed for these ILs, with the exception of
[Pyrr14][TFSI] and [Pyrr1(2o2o1)][TFSI]. The Tg value was estimated
from the half-maximum on the rise of the broad endothermic peak.
The DSC heating traces of the remaining ILs are presented in
Figure S2 in the ESI. Functionalisation of ionic liquid alkyl chains
with ether groups has been previously reported to result in a
decrease in melting.[49] Therein, the reason for the observed Tm
and Tg depression was thought to be due to the increased
rotational freedom and subsequent reduction in lattice energy as
has been described elsewhere.[38] No such clear trend is observed
for this family of ILs. For the two ILs which did not show any
obvious glass transition, it is likely that the actual Tg falls outside
of the measured temperature range. Notably, slight evidence for
a peak associated with a glass transition can be seen in the DSC
trace of [Pyrr14][TFSI] (Figure 2).
increased stability compared to the corresponding [Aze1(2o1)]+,
[Aze1(2o2o1)]+, [Aze(2o1)2]+-based ILs.
A similar study involving imidazolium and morpholinium cations
also observed that the replacement of the alkyl group with an
ether functionalised group also decreased the thermal stability of
the corresponding ionic liquid.[46]
The thermal behaviour of the synthesised ILs was also
investigated by DSC within the temperature range of 183.13 K to
323.15 K. Only the [Pyrr14][TFSI] IL showed observable
transitions associated with crystallisation (Tf) and melting (Tm)
(Figure 2). The exotherm (Tf) observed on the heating step is
associated with the crystallization of the supercooled liquid at 219
K. Further warming of the sample results in the observation of a
small endothermic process (Tcc) at 243.7 K prior to the main
endotherm associated with melting (Tm) of the IL at 253 K. The
endothermic feature Tcc is associated with a crystal-crystal phase
transition between two metastable phases. The existence of
these polymorphs in [Pyrr14][TFSI] has been previously reported
using DSC[47] and, more recently, by temperature-dependent
Raman spectroscopy investigations of the phase changes
occurring in this IL.[48] The results of the latter study suggested
that the solid-solid phase transition occurs with a suppression of
Figure 3. Exemplary DSC thermogram of [Aze1(2o1)][TFSI]. Glass transition
temperature, Tg, is determined at the midpoint of the curve as highlighted by the
dashed lines. Inset graph shows the DSC trace across the full temperature
range.
Figure 2. Differential scanning calorimeter (DSC) thermogram of [Pyrr14][TFSI].
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Unfortunately, the position of this feature is too close to the
minimum temperature of the DSC measurement, ca. 183 K, to be
accurately distinguished.
Density
Density measurements of the alkyl and ether-functionalised ILs
were performed between 293 – 363 K and the experimental data
is summarised in Figure 4. The numerical raw data is presented
in Table S1 in the ESI. As is typical for ionic liquids, the density
was observed to decrease linearly as a function of temperature.
This data has, therefore, been fitted to Equation 1;
ꢀ = 퐴 + 퐵푇
1
where ρ is the density (g.cm-3), T is the temperature (K) and A and
B are the linear fitting parameters, see Table 2. Regardless of the
cationic scaffold, the density distinctly increases in the order of
[X14][TFSI] < [X1(2o2o1)][TFSI] < [X1(2o1)][TFSI] (where X represents
either pyrrolidinium, piperidinium or azepanium). This observation
is in accordance with previous literature values concerning the
study of ether and alkyl functionalised pyrrolidinium, piperidinium
[50]
and azepanium [TFSI]-based ILs.[31] Additionally, the specific
density values at 298 K for the pyrrolidinium- and azepanium-
based ILs are in reasonable agreement with the equivalent
literature values (Table 2). In addition, where the alkyl functional
groups remain constant, the observed density generally
decreases as the cation size increases; in the order [Pyrr1X][TFSI]
> [Pip1X][TFSI] > [Aze1X][TFSI]. However, for the ILs functionalised
with the methoxyethyl group, a minimal difference is observed
between the density of the pyrrolidinium (1.4594 g·cm-3 at 293 K)
and piperidinium (1.4588 g·cm-3 at 293 K). The values in Table 2
also show that, with the exception of [Aze1(2o2o1)][TFSI], the fitting
parameters obtained for linear correlation are also in reasonable
agreement with the literature values.
Table 2. Density values, ρ, of the [TFSI]--based ILs at 298.15 K and the fitting
parameters, A and B, for the linear fitting equation ρ = A + BT.
IL
cation
[Pyrr14]+
ρ
A
104 B
Adj. R2
Figure 4. Density variation as a function of temperature for (a) pyrrolidinium, (b)
piperidinium and (c) azepanium-based [TFSI]--based ILs. Solid lines represent
the linear fit of each dataset based on the equation ρ = A + BT. Data point
symbols represent the functional groups of a given cyclic alkylammonium
cation;  = 14,  = 1(2o1),  = 1(2o2o1),  = (2o1)2,  = (2o2o1)2.
/ g·cm-3
/ g·cm-3
/ g·cm-3·K-1
1.4054 (1.3950)a
1.4547 (1.4546)a
1.4137
1.6709
1.7268
1.6856
1.7091
1.6762
1.6358
1.7316
1.6633
1.6865
1.6104
-8.9107
-9.1390
-9.1160
-9.3433
-9.2370
-8.6690
-9.3024
-8.9843
-9.3173
-8.6300
0.99993
0.99990
0.99883
0.99989
0.99988
0.99989
0.99994
0.99993
0.99997
0.99994
+
[Pyrr1(2o1)
]
+
+
[Pyrr1(2o2o1)
]
]
Furthermore, the isobaric coefficient of thermal expansion, αp, for
the studied ILs was calculated using equation 2;
+
1.4308 (1.39)b
1.4010
[Pyrr(2o1)2
]
[Pyrr(2o2o1)2
1
휕푉
1 휕ꢀ
( )
푚
2
[Pip14]+
훼푝
=
(
)
푝
= −
1.3775 (1.3786)b
1.4543 (1.4355)b
1.3956
푉
푚
휕푇
ꢀ 휕푇
푝
+
[Pip1(2o1)
]
where Vm, the quotient of molar mass (g·mol-1) divided by the
density (g·cm-3), is the molar volume (cm3·mol-1) of the IL. For all
of the ILs reported in this work, the value of αp lies within the range
of 5.7 x 10-4 – 7.5 x 10-4 K-1 across the studied temperature range,
with the exception of several seemingly anomalous data points.
Notably, αp of [Aze1(2o2o1)][TFSI] increases to ca. 9 x 10-4 - 9.9 x
10-4 K-1 in the higher temperature region 353 – 363 K. This
significant deviation in αp is consistent with the apparent lack of
linearity observed for the high temperature density of this IL, as
represented by the red filled diamonds in Figure 4c. However, the
observed range of coefficients for this set of ILs is in reasonable
accordance with previously reported ILs with similar structures
+
+
[Pip1(2o2o1)
]
]
+
1.4087 (1.36)c
1.3531
[Pip(2o1)2
]
[Pip(2o2o1)2
[Aze14]+
1.3729 (1.3661)d 1.6284 (1.6183)d -8.5737 (-8.4670)d0.99990
1.4179 (1.4156)d 1.6851 (1.6801)d -8.9633 (-8.8730)d0.99994
1.3949 (1.3815)d 1.6947 (1.6404)d -10.036 (-8.6893)d0.99709
+
[Aze1(2o1)
]
+
[Aze1(2o2o1)
]
]
+
[Aze(2o1)2
]
1.3393
1.3527
1.6092
1.6108
-9.0517
-8.6610
0.99997
0.99987
+
[Aze(2o2o1)2
Literature values shown in brackets: a. Ref [39], b. Ref [50], c. Ref [28], d. Ref [31]
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
(examples given at 298 K): [Pyrr14][TFSI], αp = 6.56 x 10-4 K-1 [51]
;
;
[51]
[1-butyl-3-methylimidazolium][TFSI], αp = 7.24 x 10-4 K-1
[Pyrr14][thiocyanate], αp = 5.08 x 10-4 K-1 [52]; [Pip14][thiocyanate],
αp = 5.5 x 10-4 K-1 [52]; [Pyrr14][dicyanimide], αp = 5.56 x 10-4 K-1 [53]
.
The calculated values for αp are given for all ILs across the full
studied temperature range in Table S2 in the ESI.
Viscosity
The viscosity of an electrolyte is an important parameter which
contributes to the transport capability of the active substances
during electrochemical processes. To study the effect of the
cation structure on this property, the viscosity of the ILs was
measured between the temperature range of ca. 293 – 363 K. The
temperature dependences of the viscosity are presented in the
form of Arrhenius-type plots in Figure 5. The numerical raw
viscosity data is presented in Table S3 in the ESI. These plots
show that the majority of the ILs do not follow Arrhenius-type
behaviour. As is common for ILs, the temperature dependence of
viscosity, η, can be well described using the Vogel-Tammann-
Fulcher (VTF) equation 3:
퐵휂
푇 − 푇휂
휂 = 휂표exp⁡(
)
3
표
η
where ηo, Bη and To are the fitting parameters and T is the
temperature. The best fit parameters for each IL are listed in Table
3, where Bη is related to the pseudo activation energy barrier
linked to the energy barrier needed to be overcome for ions to
move past each other (i.e. Bη = Ea /R, where R is the gas
constant).[54] To is related to the glass transition temperature of the
ionic liquid and, by its position in the VTF equation, the
temperature at which the viscosity would approach infinity. Good
correlation of the temperature dependence of viscosity of the ILs
was achieved with the VTF equation within the measured range.
Table 3. Viscosity, η, data at 298.15 K and VTF fitting parameters for the ether
and alkyl functionalised ammonium [TFSI]--based ILs.
IL
η
ηo
Bη
To
Adj. R2
η
cation
/ mPa·s
/ mPa·s
/ K
/ K
[Pyrr14]+
Figure 5. Viscosity variation as a function of temperature for (a) pyrrolidinium,
(b) piperidinium and (c) azepanium-based [TFSI]--based ILs. Solid lines
represent the correlation of the data based on the VTF equation (Equation 3).
77.15 (79.6)a
49.22 (54.8)a
51.19
0.154 846.38 161.98 0.99999
0.269 633.58 176.56 0.99999
0.237 654.17 176.46 0.99995
0.081 933.84 161.31 0.99989
0.364 510.70 210.61 0.99973
0.159 815.32 181.29 0.99999
0.215 718.86 181.92 0.99997
0.220 694.30 181.27 0.99992
0.255 548.19 207.51 0.99997
0.342 579.01 201.60 0.99963
0.129 862.49 187.22 0.99999
0.206 720.29 187.09 0.99998
0.174 746.08 185.10 0.99995
0.208 650.48 197.80 0.99997
0.186 724.94 195.01 0.99996
+
[Pyrr1(2o1)
]
Data point symbols represent the functional groups of
a given cyclic
+
+
[Pyrr1(2o2o1)
]
]
alkylammonium cation;  = 14,  = 1(2o1),  = 1(2o2o1),  = (2o1)2,  =
(2o2o1)2.
+
74.58 (55)c
[Pyrr(2o1)2
]
Further discussion of the coefficients of the VTF correlation is
detailed in a following section.The results demonstrate that the
viscosity of the IL tends to increase as the size of the alkyl ring of
the cation is increased and in general the viscosity of the IL
follows the trend [PyrrX][TFSI] < [PipX][TFSI] < [AzeX][TFSI]
(where X represents a particular set of functional groups).
Comparing ILs of the same cation ring size, the data also shows
that substituting the butyl group with a methoxyethyl group
reduces the measured viscosity of the IL. The ether group has
been previously reported to reduce the viscosity of ILs relative to
similarly sized alkyl-functionalised equivalents.[38, 39, 50] This effect
has been associated with the flexibility of ether moiety due to the
free rotation of the C•O-C bond.[39] Alternatively, the results of a
molecular dynamics study of similarly sized butyl- and 2-
methoxyethyl-functionalised acyclic ammonium ILs suggested
that the reduction in IL viscosity is related to changes in cation-
[Pyrr(2o2o1)2
124.50
[Pip14]+
170.71 (182)b
104.46 (102)b
83.64
+
[Pip1(2o1)
]
+
+
[Pip1(2o2o1)
]
]
+
108.01 (122)c
[Pip(2o1)2
]
[Pip(2o2o1)2
[Aze14]+
137.42
307.15 (315)d
134.87 (160.3)d
127.80 (136.5)d
135.74
+
[Aze1(2o1)
]
+
[Aze1(2o2o1)
]
]
+
[Aze(2o1)2
]
+
[Aze(2o2o1)2
209.99
Literature values shown in brackets: a. Ref [39], b. Ref [50], c. Ref [28],d. Ref [31]
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
cation interactions; i.e. short range van der Waals interaction
between alkyl chains are more efficient than for the alkoxy-group
equivalent.[55] However, increasing the length of this ether chain
appears to have a varying effect; for pyrrolidinium and azepanium
ILs the viscosity decreases in the order of [X14][TFSI] >
[X1(2o2o1)][TFSI] > [X1(2o1)][TFSI]. This behaviour suggests there is
a balance between increasing the flexibility of the ether chain
(reducing viscosity) and increasing the cation size (increasing
viscosity). However, for the piperidinium class of ILs, the
introduction of the methoxyethoxyethyl chain yields a very
significant drop in the measured viscosity.
The ILs for which the cations have been functionalised with two
ether groups ([X(2o1)2][TFSI] and [X(2o2o1)2][TFSI]) exhibit different
behaviours depending on the size of the ring. For the
pyrrolidinium-based ILs, the additional groups lead to a large
increase in the IL viscosity from 58.8 mPa·s for [Pyrr14][TFSI] to
72.9 mPa·s and 124.4 mPa·s for [Pyrr(2o1)2][TFSI] and
[Pyrr(2o2o1)2][TFSI], respectively. For the piperidinium and
azepanium-based ILs, the effect of the additional ether chains (in
place of the original methyl group) does not follow a simple trend.
Firstly, the viscosity of [Pip(2o1)2][TFSI] and [Pip(2o2o1)2][TFSI] is
lower than the analogous [Pip14][TFSI]. Secondly, compared to
[Aze14][TFSI], the viscosity of [Aze(2o1)2][TFSI] is lower while the
viscosity of [Aze(2o2o1)2][TFSI] is higher. These observed variations
between the ILs with different ring sizes could possibly arise again
from a combination of opposing factors which positively or
negatively impact the viscosity. Increasing ring size and ether
chain length may increase the number of available low energy
conformers thus increasing the potential void space promoting
mass transport of the ions, i.e. decreasing the viscosity. However,
increasing the molecular weight of the cation increases the van
der Waals interactions, inhibiting the mass transport leading to an
increased viscosity.
Table 4. Conductivity, σ, data at 298.15 K and VTF fitting parameters for the
ether and alkyl functionalised ammonium [TFSI]--based ILs.
Adj. R2
σ
IL
cation
σ
σo
Bσ
/ K
To
/ mS·cm-1 / mS·cm-1
/ K
2.72 (2.62)a
3.69 (3.72)a
2.80
574.90
593.50
465.61
286.63
324.48
695.27
535.11
390.93
385.28
360.71
736.69
384.32
384.79
442.69
501.72
672.95
666.73
658.85
496.88
622.46
747.58
661.16
623.07
618.76
647.83
770.50
576.29
623.40
713.74
756.22
172.47 0.99999
166.94 0.99999
169.28 0.99999
201.04 0.99990
186.28 0.99997
183.30 0.99985
182.68 0.99996
184.59 0.99996
184.98 0.99996
188.59 0.99995
189.77 0.99989
199.80 0.99969
191.88 0.99998
187.50 0.99994
186.96 0.99993
[Pyrr14]+
+
[Pyrr1(2o1)
]
+
+
[Pyrr1(2o2o1)
]
]
1.72 (2.29)c
1.24
+
[Pyrr(2o1)2
]
[Pyrr(2o2o1)2
1.04 (1.1)b
1.74 (2.0)b
1.62
[Pip14]+
+
[Pip1(2o1)
]
+
+
[Pip1(2o2o1)
]
]
1.63 (1.08)c
0.98
+
[Pip(2o1)2
]
[Pip(2o2o1)2
[Aze14]+
0.60 (0.56)d
1.10 (1.12)d
1.09 (0.97)d
0.70
+
[Aze1(2o1)
]
+
[Aze1(2o2o1)
]
]
+
[Aze(2o1)2
]
+
0.56
[Aze(2o2o1)2
Literature values shown in brackets: a. Ref [39], b. Ref [50], c. Ref [28], c. Ref [31]
parameters for the conductivity data of the studied ILs are
presented in Table 4.
The reported data shows that increasing the alkyl ring size of the
cation has a negative effect on the resulting conductivity. This
observation corresponds with the increase in viscosity (Table 3,
Figure 5) and the fact that increasing the overall cation size is
typically linked to a drop in conductivity. Since conductivity is
dependent on ionic mobilities (as well as the number of charge
carriers), the inverse relationship between the two quantities
(conductivity and viscosity) is expected. The effect of substituting
the butyl group for the methoxyethyl group results in an increase
in conductivity regardless of the alkylammonium ring size. These
observations correlate well with the observed reduction in IL
viscosity and is attributed to the increased ionic mobility.
Increasing the size of the ether chain, from methoxyethyl to
methoxyethoxyethyl, has a negative effect on the conductivity
across the measured temperature range. Furthermore, increasing
the overall cation size by substituting the methyl group for a
second ether functional group yields a small drop in the observed
conductivity for the piperidinium-based ILs but a significant drop
in the measured conductance of the pyrrolidinium and
azepanium-based ILs. This observation is likely to be the result of
a trade-off between increasing the degrees of rotational freedom
of the flexible ether chains and increasing the size of the cation
which reduces the mobility of the cation and hinders the
conductivity. As discussed in terms of the measured IL viscosities,
there will exist a balance between factors either positively or
negatively impacting the conductivity of the IL. In the case of the
majority of the cations functionalised with two ether chains, the
increased cation size appears to dominate these factors leading
to lower conductance than the equivalent IL functionalised with
only one ether chain.
Conductivity
The electrolytic conductivity, σ, of the ether-functionalised was
measured across the temperature range of 293 – 363 K. The
experimental results, organised by the size of the alkyl ring of the
cations, are presented as Arrhenius-type plots in Figure 6. The
numerical conductivity data is presented in Table S4 in the ESI.
The solid lines represent the non-linear fitting of the data with
respect to the VTF equation
−퐵휎
ꢁ = ꢁ표exp⁡(
)
4
푇 − 푇휎
표
where T is the absolute temperature and σo, Bσ and Toσ represent
the fitting parameters. Like viscosity, the VTF fitting equation is
commonly used to accurately fit the temperature dependency of
the conductivity of ionic liquids which do not follow the Arrhenius
Law.[56]
Analogous to the parameters of the VTF equation used to
describe IL viscosity in the previous section, Bσ is related to a
pseudo-activation energy for conduction (Bσ = Ea /R, where R is
the gas constant) and To is related to the glass-transition
temperature of the IL, theoretically representing the temperature
at which the conductivity of the system would disappear. Good
correlation of the temperature dependence of the conductivity
was achieved by correlation with Equation 4. The VTF correlation
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10.1002/cphc.201700246
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ARTICLE
viscosities of certain ILs did not directly translate to higher
conductivities. For example, the viscosity of [Pip1(2o2o1)][TFSI]
(83.6 mPa·s at 298 K) was lower than its single ether-group
analogue, [Pip1(2o1)][TFSI] (108.0 mPa·s at 298 K), yet the
conductivity of the latter example was found to be higher (1.62
mS·cm-1 and 1.74 mS·cm-1 at 298 K for [Pip1(2o2o1)][TFSI] and
[Pip1(2o1)][TFSI], respectively). Some of these variations are
highlighted graphically in Figure 7. The lack of a direct trend
between viscosity and conductivity shows other factors and
properties intrinsic to each ionic liquid are clearly contributing to
the viscosity and conductivity in different ways.
The discrepancies between these quantities have been discussed
previously in terms of the ionicity, or degree of ionic association,
of the ILs and the Walden rule; Λη = c. This can be
straightforwardly represented with the use of a Walden plot; a plot
of log(η -1) vs. log(Λ), where the reciprocal viscosity (or fluidity, η-
1) is given in Poise-1 and the molar conductivity, Λ, is given in
S·cm2·mol-1. The derived Λ values for each IL at 298.15 K is
provided in Table 5. The Walden plots within the temperature
range of 298 – 343 K for the 15 ILs are shown in Figure 8. Herein
the viscosity and conductivity of each IL is calculated by accurate
interpolation of VTF correlation discussed previously and the
molar conductivity is calculated by division of electrolytic
conductivity by the IL concentration [Λ = σ (S·cm-1)/concentration
(mol·cm-3)]. The IL concentration was calculated by linear
interpolation of the temperature-dependency of the IL density and
the IL molecular weight [concentration = ρ (g·cm-3) / Mr (g·mol-1)].
The calculated values for Λ and η-1 over a temperature range of
293 K – 343 K are presented in Table S5 in the ESI. The solid
lines in Figure 8(a-c) is the ideal KCl line representing the ideal
Walden behaviour taken from extrapolation of the behaviour of a
0.01 mol·dm-3 KCl aqueous solution, a strong electrolyte where
the ions are known to be fully dissociated and considered equally
mobile.[57] Deviations below this calibration line, i.e. where the
conductivity of the IL is lower than may be expected at a given
viscosity according to the Walden rule, imply the existence of
significant ionic association between the cation and anion with the
tendency to behave as ion pairs. The magnitude of this deviation,
ΔW (where ΔW represents the vertical displacement from the
ideal KCl line) has been used to classify the ionicity of the liquid.[57]
Figure 6. Conductivity variation as a function of temperature for (a) pyrrolidinium,
(b) piperidinium and (c) azepanium-based [TFSI]--based ILs. Solid lines
represent correlation of the data based on the VTF equation (Equation 4). Data
point symbols represent the functional groups of a given cyclic alkylammonium
cation;  = 14,  = 1(2o1),  = 1(2o2o1),  = (2o1)2,  = (2o2o1)2
This is also consistent with previously reported observations
which demonstrate the positive effect of the small flexible alkoxy
50]
chains.[39,
A further comparison of the effect of the ether
functionalisation versus cation size within this class of ILs could
be made by measuring the respective properties of dibutyl or
diheptyl functionalised cations. Such IL cations would be of similar
size to the studied diether cations, [X(2o1)2][TFSI] and
[X(2o2o1)2][TFSI].
Ionicity
A typical observation made when characterising ILs is the
aforementioned inverse relationship between the viscosity and
conductivity of the liquid. The cooperative dependency of these
behaviours is unsurprising given that viscosity describes a liquids
resistance to flow while the conductance of an ionic liquid relies
on the mobility of the ions carrying the charge. However,
considering the data discussed in the previous sections, the lower
Figure 7. Variations in the fluidity (η-1, reciprocal viscosity, left y-axis),
conductivity (σ, right inner y-axis) and molar conductivity (Λ, right outer y-axis)
in the studied [TFSI]--based ILs at 298 K. The x-axis markers highlight the
abbreviations of the functional groups X and Y attached to the labelled cyclic
alkylammonium cation.
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Table 5. The derived molar conductivity, Λ, at 298.15 K for each of the alkyl-
and ether-functionalised [TFSI]- ILs. Also shown are the linear fitting
coefficients, Log10(c’) and αW, of Equation 6 for the temperature dependence of
Walden behaviour and the % ionicity derived from the fractional Walden rule.
Λ
%
Log10(c’)
αW
Adj. R2
IL cation
/ S·cm2·mol-1 ionicity
[Pyrr14]+
0.818
64.17
-0.193
-0.257
-0.302
-0.375
-0.229
-0.261
-0.261
-0.327
-0.228
-0.260
-0.244
-0.325
-0.310
-0.462
-0.319
0.935 0.99998
0.938 0.99992
0.929 0.99995
0.991 0.99907
0.809 0.99920
0.959 1.00000
0.947 1.00000
0.961 1.00000
0.780 0.99936
0.911 0.99978
0.943 1.00000
1.004 0.99989
0.948 0.99998
0.939 0.99986
0.927 0.99992
+
[Pyrr1(2o1)
]
1.076
0.929
0.563
0.494
0.329
0.526
0.560
0.557
0.411
0.198
0.350
0.388
0.259
0.241
55.28
49.85
42.18
58.97
54.89
54.81
47.14
59.10
54.92
56.96
47.27
48.96
34.52
47.98
+
+
[Pyrr1(2o2o1)
]
]
+
[Pyrr(2o1)2
]
[Pyrr(2o2o1)2
[Pip14]+
+
[Pip1(2o1)
]
+
+
[Pip1(2o2o1)
]
]
+
[Pip(2o1)2
]
[Pip(2o2o1)2
[Aze14]+
+
[Aze1(2o1)
]
+
[Aze1(2o2o1)
]
]
+
[Aze(2o1)2
]
+
[Aze(2o2o1)2
For example, ILs for which ΔW is greater than one order of
magnitude below ideal Walden behaviour (ΔW > 1, where ΔW =
1 is represented by the diagonal dashed lines in Figure 8) lie
towards the bottom-right of the Walden plot and can be classified
as poor ionic liquids since the molar conductivity is much lower
than possibly expected from a given low viscosity (high η-1).[58]
Conversely, if the Walden behaviour of an IL lies closer to the
ideal KCl line, where ΔW is small, it can be understood that these
liquids exhibit more dissociated ionic character in which the
constituent ions are more independently mobile.
The data in Figure 8 shows all the ILs in this study lie below the
ideal calibration line, highlighting a certain degree of ionic
association or aggregation between the constituent ions of the IL.
However, the deviation is not so strong as to classify any of the
ILs as poor and ΔW < 0.4 (% ionicity > 40 %) for all studied ILs at
298 K, with the single exception of [Aze(2o1)2][TFSI] (for which ΔW
= 0.45 and % ionicity = 35 %). Using the example discussed
earlier in this section concerning [Pip1(2o2o1)][TFSI] and
[Pip1(2o1)][TFSI], it can be seen that the IL functionalised with the
longer ether chain lies further from the ideal line. This indicates
that slightly stronger ionic interactions may be present in
[Pip1(2o2o1)][TFSI] and possibly account for the fact that this IL
exhibits a lower conductivity even though it is less viscous than
the analogous piperidinium IL functionalised with the shorter
chain.
Furthermore, the temperature dependant Walden behaviour of
the ILs (Figure 8) shows that the apparent degree of dissociation
is not independent of the temperature according to the Walden
rule (i.e. the slope of the data points is not parallel to the ideal KCl
line). To account for these variations as a function of the
temperature an additional exponent, αW, can be introduced to give
the fractional Walden rule
Figure 8. Walden plots for a temperature range of 293 – 343 K for (a)
pyrrolidinium, (b) piperidinium and (c) azepanium-based [TFSI]--based ILs. Data
point symbols represent the functional groups of a given cyclic alkylammonium
cation;  = 14,  = 1(2o1),  = 1(2o2o1),  = (2o1)2,  = (2o2o1)2. The solid
lines represent the ideal Walden behaviour of an aqueous KCl solution. The
dashed lines represent a ΔW of 1 (% ionicity = 10 %), beneath which ILS are
classed as poorly ionic (or poor ionic liquids).
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
훬휂훼 = 푐′ = 푐표푛푠푡푎푛푡
5
6
푊
small deviations within the measurements and, as such, direct
comparison can become problematic. Nevertheless, in many
cases, both To values show a reasonable match for the studied
ILs and the most significant discrepancies appear with the diether-
functionalised ILs. Furthermore, due to the intricate relationship
between IL viscosity and conductivity (or more specifically molar
conductivity), differences in the activation energy barriers for the
two processes should also be related to changes in the apparent
ionicity of the ILs with temperature (i.e. αW).[61] Since typical
Arrhenius plots (ln(η) or ln(Λ) vs. 1/T) of the viscosity and molar
conductivity of ILs yield non-linear behaviour, it is inferred that the
apparent activation energy derived from Arrhenius equations
(Equation 7 and 8) is temperature dependant.
log10(훬) = log10(푐′) + 훼푊log10(휂−1
)
where c’ is analogous to the aforementioned Walden product.
The linear fitting coefficients of Equation 6 and the calculated %
ionicity as calculated from the fractional Walden rule for each IL
are presented in Table 5.
As a result of extrapolation of the seemingly ideal behaviour of
aqueous KCl solutions, αW is considered to be unity for the ideal
KCl line. Conversely, the ionicity of majority of the ILs reported in
this work appears temperature dependant and αW values between
0.9-0.96 are exhibited, showing that, at increasing temperatures,
ΔW increases and the IL viscosity decreases at greater rate than
the increase in conductivity that may be expected from so-called
ideal Walden behaviour. However, comparison between the αW
values for ILs and dilute electrolytic salt-in-solvent solutions has
been previously described as largely meaningless; particularly
since real αw values relating the viscosity and conductivity of
aqueous KCl solutions are found to roughly equal 0.87.[59]
Nevertheless, relative comparison of the temperature-dependant
Walden behaviour of ILs shows that the relationship between IL
conductivity and fluidity, whereby αw values roughly equate to 0.9-
0.95, is in reasonable agreement with previously reported ILs; e.g.
퐸푎휂
7
휂 = 퐴휂푒푥푝 (
)
푅푇
−퐸푎훬
푅푇
8
훬 = 퐴훬푒푥푝 (
)
where R represents the gas constant (8.3145 J·mol-1·K-1) and
η
Λ
Aη/AΛ and Ea /Ea represent the pre-exponential factor and the
activation energy barrier, respectively for the two quantities,
viscosity and molar conductivity. As such, following a procedure
described previously,[61] the Arrhenius plots for each respective
quantity were firstly fitted by a third-order polynomial within the
range of 293 – 343 K. By subsequent differentiation of the derived
polynomial equation, the gradient (Ea/R) was calculated at a
series of tangents to the curve. The activation energy (i.e. Ea = R
· gradient) was then taken as an average of the calculated values
within the stated temperature range. The coefficients of the
polynomial expansions and the derived average Ea values are
shown in Tables S6 and S7 in the ESI. Interestingly, the ratio
[EMIm][TFSI], αw = 0.906 [59], [EMIm][dicyanamide], αw = 0.95 [59]
[BMIm][TFSI], αw = 0.998 [51]; and [Pyrr14]][TFSI], αw = 0.971.[51]
;
The effects contributing to ion pair formation or aggregation will
relate partly to delocalisation of the ionic charge surrounding the
cation/anion centre and structural contributions of bulky side
chains contributing to steric hindrance of coulombic interaction.
Within this family of ILs, the negative charge of the bulky [TFSI]-
anion is highly delocalised via the electron withdrawing
(trifluoromethane)sulfonyl groups, reducing affinity for ion pair
formation. Within the cyclic alkylammonium cations, generally the
apparent % ionicity is greatest for the alkyl-functionalised cations.
Where ether functionality is introduced and the cation remains of
similar size (i.e. butyl groups substituted by the 2-methoxyethyl
group), electron donation from the lone pairs of the ethereal
oxygen atoms would be expected to contribute somewhat to the
localisation of the positive charge and, in turn, increase the
propensity for ionic interaction with the counter ion. We reported
a similar reduction in apparent ionic dissociation for ether vs. alkyl
functionalised cyclic sulfonium [TFSI]- based ILs based on the
tetrahydrothiophenium cation.[60] For larger degrees of ether
functionalisation, the tendency for ion-ion pair formation would be
further dependant on additional factors including steric hindrance,
organisation/freedom and interaction of flexible ether chains,
varying ring conformers and anion solvation/interaction in
competition with ethereal groups and cation centres. As such, a
complex variation in the % ionicity is not unexpected.
Λ
between the two activation energies, Ea /Eaη (detailed in Table 6)
closely represent the derived values for the exponent of the
fractional Walden rule, αW. In fact, the majority of the values for
the studied ILs differ by less than 2 % (bracketed values in Table
6). This is in agreement with the observations, and the original
hypothesis of this relation, detailed by Schreiner et al.[61]
As mentioned previously, the To values derived from VTF
correlation of viscosity and conductivity data are related to the IL’s
glass transition temperature, Tg (as measured by DSC). Typically
for many ILs, To usually falls approximately between 10 and 50 K
62]
below the Tg of the IL.[61,
However, the larger difference
between Tg, and the theoretical To, are considered to depend on
the liquid fragility or strength. Furthermore, while lower values of
Tg would be expected to yield lower viscosity, or higher
conductivity, ILs under ambient conditions, it has been shown that
these properties are dependent on both fragility and Tg.[58, 63] The
concept of fragility, introduced by Angell,[64, 65] describes the rate
at which the transport properties and relaxation dynamics of a
glassy forming material changes as the temperature approaches
the Tg. Materials described as highly fragile, or low strength,
exhibit higher susceptibility to structural change and, in turn, large
changes in transport properties over small temperature variations
close to the Tg.
VTF parameters and fragility
Table 6 shows a comparison of the Tg derived from DSC
measurements and the VTF derived parameter, Toη/σ, from
temperature dependant viscosity and conductivity measurements.
The VTF parameters should agree if the both charge transport
properties, viscosity and conductivity are closely coupled.
However, the derived VTF fitting parameters are very sensitive to
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10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Table 6. Comparison of measured Tg and the VTF correlation parameters, Toη/σ, from the viscosity and conductivity measurements for each of the alkyl-/ether-
functionalised [TFSI]--based ILs. The ratio between Arrhenius-type activation energies, EaΛ/Eaη, for viscosity and molar conductivity is shown and the value in
f
brackets represent the percentage deviation between EaΛ/Eaη and αW. The calculated fragility indices, mη/σ, and the ratios between measured Tg and the To
correlation parameters of the modified VTF equation (Equations 9 and 10) are also shown.
Toηf / K
T oσf / K
EaΛ/ Ea
mη
mσ
Tg/To
Tg/To
η
ηf
σf
IL cation
Tg / K
[Pyrr14]+
[Pyrr1(2o1)
[Pyrr1(2o2o1)
[Pyrr(2o1)2
[Pyrr(2o2o1)2
[Pip14]+
[Pip1(2o1)
[Pip1(2o2o1)
[Pip(2o1)2
[Pip(2o2o1)2
[Aze14]+
[Aze1(2o1)
[Aze1(2o2o1)
186b
187
161.98
176.56
176.46
161.31
210.61
181.29
181.92
181.27
207.51
210.92
187.22
187.09
185.10
197.80
195.01
172.47
166.94
169.28
201.04
186.28
183.30
182.68
184.59
184.98
188.59
189.77
199.80
191.88
187.50
186.96
0.94 (-0.1)a
0.93 (0.4)a
0.93 (0.3)a
1.01 (-2.0)a
0.79 (2.9)a
0.95 (0.8)a
0.95 (0.2)a
0.99 (-2.8)a
0.78 (0.0)a
0.92 (-1.1)a
0.93 (1.4)a
0.97 (3.5)a
0.94 (0.7)a
0.93 (1.2)a
0.90 (2.5)a
124c
147
126c
142
1.15c
1.12
1.13c
1.12
+
]
+
+
]
]
+
]
198
199
198
200
193
206
201
205
198
195
207
201
151
134
126
153
138
153
139
125
133
125
154
121
141
141
119
125
129
170
134
119
118
117
139
115
1.12
1.14
1.15
1.11
1.13
1.12
1.13
1.15
1.14
1.15
1.12
1.15
1.12
1.12
1.14
1.14
1.13
1.09
1.12
1.15
1.15
1.15
1.12
1.15
+
]
+
+
]
]
+
]
+
]
+
]
]
+
[Aze(2o1)2
[Aze(2o2o1)2
]
+
(a) the percentage deviation between EaΛ/ Eaη and αW shown in Table 5, (b) literature value[47], (c) calculated using literature Tg value.
Conversely, strong materials are considered more impervious to
structural changes upon heating above their Tg.[61] Ionic liquids,
materials known for undergoing glass transitions and commonly
inadequately described by normal Arrhenius-type temperature
dependence (in terms of transport properties), have been
frequently described as exhibiting intermediate to high fragility,
such that their transport and relaxation dynamics change several
orders of magnitude over relatively small temperature deviations
close to their Tg.[58, 61, 63, 66, 67]
To assess the apparent fragility of the studied ionic liquids,
quantified by m (the fragility index), the viscosity and conductivity
data was further correlated as a function of temperature by a
modified VTF equation
a steeper approach, greater changes in viscosity/conductivity as
T approaches Tg and a higher fragility). Full derivation of Equation
11 is explained in reference.[61]
11
푇
퐷푇
푔
표
푚 =
·
(푇 − 푇 )2
푙푛10
푔
표
Within the Angell plots of Figure 10, an approximation of ideal
Arrhenius-type behaviour of the transport properties of a strong
(or non-fragile) material is represented by the diagonal dash and
dash-dot-dot straight lines. For viscosity, this represents a
minimum fragility of 16 determined by the logarithmic ratio of the
viscosity at Tg and an approximation of the limiting viscosity (mmin
f
f
= Log10[η(Tg)/ηo ] = 16 where ηo ≈ 0.1 mPa·s). This trend is
supported by the behaviour of strong glass formers, e.g. SiO2 and
GeO2.[58] For the conductivity plot, since the limiting conductivity
is generally considerably more varied, the y-axis intercept of the
hypothetical Arrhenius line on the Angell plots was approximated
퐷휂푇휂푓
표
휂 = 휂표푓exp⁡(
)
푇 − 푇휂푓
9
표
−퐷휎푇휎푓
표
ꢁ = ꢁ표푓exp⁡(
)
f
10
from the average σo correlation parameter for the studied ILs
푇 − 푇휎푓
표
(2.86, Equation 10). This yielded an approximate value of 15 for
the fragility index of the ideal Arrhenius line for conductivity.
The data presented in Figure 9 shows the studied ILs all exhibit
very similar, high fragility behaviour. This similarity is expected
considering the shared anion and the closely related cationic
where Dη/σ represents the strength parameter, inversely
proportional to the fragility index, m, and the remaining
parameters are equivalent to the previously described VTF
equation (Equations 3 and 4; additional notation of a superscript f
is included here to differentiate between parameters used in these
and the previous equations). For each dataset, an additional point
of η(Tg) = 1015 mPa·s or σ(Tg) = 10-12 mS·cm-1 was included in the
correlation leading to slightly different correlation parameters.
These values represent an approximation of the sample viscosity
and conductivity, respectively, at the Tg of the material.[61, 63] The
correlation parameters of these fits, detailed in Table S8, are then
used to extrapolate the temperature dependant behaviour of
viscosity/conductivity towards the Tg. This is presented in the form
of an Angell-type plot (Figure 9), where the temperature along the
x-axis is scaled by the materials Tg. The fragility index, m,
calculated using Equation 11 (and by differentiation of the fitted
slope as T  Tg), is representative of the gradient of approach of
each respective property towards the Tg (higher m corresponds to
structures. The values for mη σ derived from the viscosity and
/
conductivity data are shown in Table 6 and, in ranging from 115 –
170, are of a similar order to values previously reported for fragile
ILs.[61, 63, 68] Furthermore, the values for the fragility indices derived
from the two different quantities, viscosity (mη) and conductivity
(mσ), show reasonable agreement between their magnitudes and
respective trends. Within the different ring-size families of ILs
studied, no clear trend is observed. However, for all cation ring
sizes, the bis(2-methoxyethyl) functionalised ILs (e.g.
[Pip(2o1)2][TFSI]) appear to exhibit the highest fragilities; explaining
why the ambient temperature fluidity/conductivity of these ILs is
not too drastically low, despite possessing the some of the highest
Tg values (particularly the piperidinium and azepanium ILs).
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10.1002/cphc.201700246
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Conversely, the fragility indices of the butyl/methyl functionalised
ILs (e.g. [Pyrr14][TFSI]) tend to lie at the lower end of the
measured values. Though the deviation of derived values is
relatively small, and all studied ILs would be classed as highly
fragile, the lower apparent fragility of the dialkyl functionalised
cations could be attributed to more efficient van der Waals
interactions between alkyl substituents which are effectively
dampened by substitution with ether-groups.
Following the analysis of the data by Equations 9 and 10, the
ηf
σf
derived coefficients To and To for viscosity and conductivity,
respectively, values consistently lie within a range of 17.6 – 27.4
K below the measured Tg. Furthermore, the ratios between these
coefficients and the measured Tg of the studied ILs (Tg/Toηf and
Tg/Toσf) are shown in Table 6. An average value for Tg/Toηf of 1.13
± 0.004 is obtained from both viscosity and conductivity derived
coefficients. All of these ratios satisfy the original observation of
Angell and described by Equation 12
푇
퐷
푔
= 1 +
12
푇푓
푚푚푖푛 · ln⁡(10)
표
where mmin refers to the minimum fragility index for each
respective property (mmin = 16 and 15 for the viscosity and
conductivity data, respectively).[65]
Electrochemical stability
The electrochemical stability of the alkyl and ether functionalised
cyclic alkylammonium [TFSI]- ILs was determined by cyclic
voltammetry (CV) at a 3 mm diameter glassy carbon (GC)
macrodisk electrode. Under dry, and inert conditions of an Ar-filled
glovebox, the potential of the working electrode cycled at 2 mV·s-
1 until a current density boundary of ± 0.5 mA·cm-2 was recorded,
at which point the potential sweep was reversed. The temperature
of all CV measurements was ca. 303 ± 1 K as dictated by the
internal atmosphere of the glovebox. The electrochemical
windows of the studied ILs are presented in full in Figure S3 in the
ESI. The CVs of the studied ILs all show the onset of massive
electrolyte decomposition by reduction and oxidation in the
regions of (-3.4 to -3) V vs. Fc+/Fc and (1.7 to 2.6) V vs. Fc+/Fc.
Furthermore, additional small oxidation peaks are observed in
many of the scans at ca. (0.4 to 0.6) V vs. Fc+/Fc. These are
generally attributed to the oxidation of small quantities of protic
impurities (possibly residual water). However, by varying the
voltage/current limits of the initial reducing sweep (which is always
completed before the oxidising sweep), we were able to show that
the existence of this small additional current response is related
to the oxidation of products formed during reductive
decomposition of the IL. Figure S4 in the ESI shows an exemplary
CV in [Pip1(2o2o1)][TFSI] wherein the reductive current limits are
sequentially increased. The CV traces show that the magnitude
of the impurity peak at ca. 0.7 V vs. Ag[NO3]/Ag increases as more
charge is passed on reduction. In fact, there appears to be a
proportionality between the two processes (as shown in Figure
S4a, inset graph) and the oxidative integral charge of the impurity
peak corresponds to ca. 22-25 % of the total charge passed on
the reductive sweep.
Figure 9. Tg/T scaled Angell-plots for viscosity (left y-axis) and conductivity
(right y-axis) behaviour of the (a) pyrrolidinium, (b) piperidinium and (c)
azepanium [TFSI]--based ILs. Data point symbols represent the functional
groups of a given cyclic alkylammonium cation;  = 14,  = 1(2o1),  =
1(2o2o1),  = (2o1)2,  = (2o2o1)2. Solid, coloured lines represent correlation
by the modified VTF equation (Equations 9 and 10). Diagonal dashed and dash-
dot-dot straight lines represent ideal Arrhenius type dependence of a strong
materials and additional curves in (c) represent the effect of varying m/D: i) D =
15, m = 54; ii) D = 50, m = 27; iii) m = 15.
The electrochemical stability window, defined as the potential
range between the reductive and oxidative decomposition
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10.1002/cphc.201700246
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potentials of the electrolyte, is indicative of the stable operating
voltage of the material in electrochemical systems. The wide
potential window exhibited by many different IL structures is one
of the most attractive features of applying ILs as electrolytes in
electrochemical energy storage devices like Li-batteries and high-
voltage electrochemical double layer capacitors (EDLCs).[16, 69-72]
For EDLC devices in particular, extension of the maximum
operative voltage (Umax) range, typically limited by the electrolyte
formulation, can result in direct enhancement in device specific
energy (ESp) and specific power (P) capabilities due to the
proportionalities; ESp = 0.5·C·Umax2 and P = Umax2/(4ESR) where C
and ESR represent the capacitance and equivalent series
resistance, respectively.[70] The onset potentials of reductive and
oxidative decomposition (Ea and Ec, respectively) of the IL
electrolytes studied herein are approximated using a cut-off
current density of ± 0.1 mA·cm-2. As such, the electrochemical
stability window (ΔE) is defined by the difference between these
upper and lower potential limits (i.e. ΔE = (Ea - Ec) V). The
numerical potential limits determined by cyclic voltammetry are
presented in Table 7 and represented graphically in Figure 10.
With the exception of the bis(methoxyethoxyethyl) functionalised
ILs (e.g. [Pyrr(2o2o1)2][TFSI]), the studied ILs show electrochemical
windows greater than 5 V, wherein reductive decomposition
occurs close to the Li+/Li reduction potential and oxidative
decomposition occurs at ca. 4.9 – 5.9 V vs. Li+/Li. As described in
the experimental section, an internal ferrocene redox couple was
used to normalise the reference potential vs. the Li+/Li reduction
potential following determination of electrochemical windows
(based on the approximation EFc+/Fc ≈ 3.2 V vs. ELi+/Li). An
exemplary CV of the ferrocene redox couple versus the
Ag[NO3]/Ag reference and versus a Li-metal reference electrode
is shown in Figure S5a in the ESI to justify this approximation.
The data presented in Table 7 and Figure 10 straightforwardly
shows some apparent trends in the electrochemical stabilities of
the studied ILs.
Firstly, as may be expected, the dialkyl functionalised ILs,
[Pyrr14][TFSI], [Pip14][TFSI] and [Aze14][TFSI], exhibit the largest
reductive and oxidative stabilities and, in turn, the largest
electrochemical windows (ca. 5.9 V). Beyond this, increasing the
degree of ethereal group functionalisation on the three
alkylammonium rings appears to yield subsequent reductions in
the magnitude of the electrochemical windows, ΔE. Specifically,
substitution of butyl groups for the 2-methoxyethyl group reduces
Ec slightly but affects a larger reduction in the oxidative stability,
+
Ea. Lengthening of this ethereal group (e.g. from [Pyrr1(2o1)
]
to
[Pyrr1(2o2o1)]+) for all three ring sizes promotes a further significant
reduction in the oxidative stabilities of the studied ILs. This, and
the large degree of variation in oxidative stability of the remaining
ILs, is interesting since the more traditional consensus aligns that
the oxidative decomposition is generally attributed to oxidation of
the anion (and vice versa for the cathodic reductive limits).[73, 74]
However, molecular dynamics (MD) simulation and density
functional theory (DFT) calculations have suggested this may not
always be so straightforward.[75-77] To compare the trends in
oxidative and reductive stabilities of the studied ILs as a function
of the variation in cation structure, the HOMO/LUMO energy
levels of many geometrically optimized conformers of cations,
anions and ion pairs were determined from DFT calculations. An
extensive summary of the calculated HOMO/LUMO energy levels
and surface diagrams of single ions and ion pairs is provided in
Table S8 in the ESI.
Within the cation ring size families, increasing ether functionality
has almost identical effects on the computed HOMO and LUMO
values based on ion calculations; i.e. HOMO: [X14]+ < [X1(2o1)]+ ≈
[X(2o1)2]+ < [X1(2o2o1)]+ ≈ [X(2o2o1)2]+; and LUMO: [X14]+ ≈ [X1(2o1)]+ <
[X(2o1)2]+ < [X1(2o2o1)]+ < [X(2o2o1)2]+.
Table 7. Onset potentials of anodic (Ea) and cathodic (Ec) decomposition and
the respective electrochemical stability windows, ΔE, for the [TFSI]--based ILs
as determined by cyclic voltammetry at a GC working electrode.
IL cation
Ec
Ea
ΔE
/ V
/ V vs. Li+/Li
/ V vs. Li+/Li
[Pyrr14]+
-0.11
0.03
-0.19
0.10
0.15
-0.15
0.17
-0.13
-0.07
0.12
-0.20
-0.09
-0.24
0.12
0.13
5.76
5.40
4.97
5.26
4.96
5.72
5.50
4.96
4.99
5.01
5.72
5.49
4.92
5.32
4.99
5.87 (5.8)a, f
5.37 (5.0)b, f
5.16
+
[Pyrr1(2o1)
]
+
+
[Pyrr1(2o2o1)
]
]
+
[Pyrr(2o1)2
]
5.16 (5.0)c, e
[Pyrr(2o2o1)2
[Pip14]+
4.81
5.87g
+
[Pip1(2o1)
]
5.33
+
+
[Pip1(2o2o1)
]
]
5.09
5.06 (5.5)c, e
+
[Pip(2o1)2
]
[Pip(2o2o1)2
[Aze14]+
4.89
5.92 (6.5)d, e
5.58 (5.5)d, e
5.16 (5.5)d, e
5.2
+
[Aze1(2o1)
]
+
[Aze1(2o2o1)
]
]
+
[Aze(2o1)2
]
+
[Aze(2o2o1)2
4.86
Literature values shown in brackets: a. Ref [78], b. Ref [79], c. Ref [28], d. Ref [31]
.
e) measured at a GC working electrode, f) measured at a platinum carbon
Figure 10. Graphical representation of the electrochemical stability limits of the
[TFSI]--based ILs at a GC electrode vs. Fc+/Fc (top x-axis) and Li+/Li (bottom x-
axis) reference potentials. The abbreviations on each horizontal line represent
the functional groups, X and Y, attached to the labelled cyclic alkylammonium
cation.
working electrode, g) analogous IL 1-methyl-1-propylpiperidinium [TFSI]-
reported window of 5.9 V in Ref [12]
.
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However, the functional group modifications described affect a
contributes to reductions in the partial positive charge of the alpha
methyl/methylene protons and reductions in the computed cation
LUMO values. However, in terms of oxidative stabilities, the
increased Coulombic sharing of available orbitals between cation
and anion trough ion pair formation would be expected to increase
HOMO energy levels relative to single ion computed levels, i.e.
reducing the anodic potential barrier of HOMO oxidation and
reducing oxidative stability of the IL. Additionally, in our
observations reported for the ether-functionalised sulfonium
[TFSI]- ILs,[77] the DFT calculations of ion pairs indicated that when
increasing the level of ether-functionality on the sulfonium cation,
the cis-conformer of the [TFSI]- anion becomes more favoured
over the trans-conformer. Based on the single ion DFT
calculations, cis-[TFSI]- should exhibit a lower electrochemical
oxidative stability (by 0.22 eV) than the trans alternative. However,
while the comparative ion-pair simulations of the alkyl and ether
functionalised ammonium-based ILs did not reveal a clear-cut
change in the global minimum energy conformer of the [TFSI]-
anion, the relative energy difference between trans- and cis-
conformers was lower for [Pyrr1(2o2o1)][TFSI] than for [Pyrr14][TFSI]
(see Table S9 in the ESI). We relate this observation, in
conjunction with the reduced apparent ionicity through increased
cation-anion interaction, to the restriction of the anion positioning
around the cation centre as a result of the ether chain presence;
i.e. interaction between anion and ethereal oxygen atoms is
unfavourable and, thus, the freedom of ion coordination around
the cation centre is reduced. Consequently, the probability of the
less electrochemically stable cis-[TFSI]- conformers existing in the
ether functionalised IL is higher than may be expected for the alkyl
functionalised analogue. In the chosen example, the calculated
HOMO values for [Pyrr14][trans-TFSI] and [Pyrr1(2o2o1)][cis-TFSI]
were -6.626 eV and -6.191 eV, respectively. Coupling this set of
calculations with the observed reduction in ionicity of the ether
functionalised [Pyrr1(2o2o1)][TFSI] IL (15 % less than [Pyrr14][TFSI])
can possibly account for the reduced oxidative stabilities
observed. A third consideration relates to the possibility of the
ether-functionalised cations readily reacting with any radical
much larger change in HOMO values (e.g. ≈3 eV range between
lowest cation, [Pyrr14]+
= -12.39 eV, and highest cation,
+
[Pyrr(2o2o1)2
]
= -9.46 eV) than the computed single ion LUMO
values (where the range is ≈0.7 eV). For the analogous functional
group pairing, the effect cation ring size on the computed values
is generally small but the LUMO follows the trend of [AzeX]+ ≳
[PipX]+ ≳ [PyrrX]+, in approximate agreement with trends in
observed Ec values. Furthermore, the agreements in the
computed LUMO values and Ec limits can be further extended to
include our previously reported experimental and computational
observations for a series [TFSI]- ILs based on ether-functionalised
acyclic sulfonium cations.[77] These sulfonium based ILs undergo
reductive decomposition at ca. 1 V higher under identical
conditions and the computed LUMO values for the sulfonium
cations are ca. 1eV more negative than calculated for the cyclic
alkylammonium cations reported here.
Additionally, the trend in cation HOMO values is the same
approximate trend observed for Ea. A plot of the computed cation
HOMO values vs. the experimentally observed oxidative limits
(see Figure S6 in the ESI) for each IL displays good linear
correlation; wherein only [Pip(2o1)2][TFSI] appears an outlier.
However, given that the electrochemical window of an IL is limited
by the least stable component in both oxidative or reductive
potential regions, the variation in cation HOMO values does not
alone justify the observed Ea trends since the single ion HOMO
values computed for the [TFSI]- anion component are much
higher (i.e. less stable towards oxidation; -4.19 eV, -3.97 eV and
-3.89 eV for the three [TFSI]- conformers). However, we recently
reported similar experimental and computational observations for
the series of ether-functionalised acyclic sulfonium [TFSI]- ILs.[77]
There are several possible explanations for the prominent effect
of the cation structure on the oxidative stability of the IL. Firstly,
the cation may undergo direct oxidation at the electrode interface.
Given the well-structured nature of IL electrical double layers, the
presence of cationic species directly at highly positive electrode
surface appears counter-intuitive. However, the long and flexible
2-(methoxyethoxy)ethyl group will somewhat change the
organisation of any IL double layer with this ether group. Further,
the observed Ea of these ILs is similar to the oxidative stability of
analogous glyme type solvents (e.g. tri/tetra ethylene glycol
dimethyl ether; 4.6-4.9 V vs. Li+/Li)[80, 81] indicating the possibility
of the terminal ethereal oxygen (where the cation HOMO resides,
see surface diagrams in Table S8 in the ESI) behaving more as a
free ether group.
products of [TFSI]- oxidation. Such
a process would be
autocatalytic and somewhat enable a reduction in the oxidative
potential barrier required to proceed the reactions. Ultimately, the
overall processes may be combinations of the discussed ideas
and we are currently devising experimental procedures to
understand the reactions taking place during decomposition to
make sense of the prominent effects of cation structure.
A second rationale relates to many possible changes in cation-
Li-redox in pyrrolidinium ILs
anion interactions as
a
function of the increasing ether
One of the most desirable applications for electrochemical stable
ILs, in addition to EDLC electrolyte components, is as electrolytes
for alkali metal batteries, notably Li and Na. To utilise Li/Na-metal
anodes in place of traditional graphitic anodes, the electrolyte
must facilitate deposition and stripping of the alkali metal ions with
a high cycling stability. In the context of electrodeposition and
stripping of Li, the effect of the cation ether-group functionality on
the Li-redox chemistry in the family of pyrrolidinium-based ILs is
explored in preliminary CV experiments at a Cu working electrode.
Herein, binary mixtures (by mole fraction) of (0.9)IL – (0.1)Li-salt
functionality. In a previous section, we demonstrated using the
Walden behaviour that, with a few exceptions, the ionicity (or
degree of dissociation) of the ILs decreased with the introduction
of more ethereal groups. Electron donation from the ethereal
oxygen atoms into the nitrogen centre of the cation decreases the
delocalisation of the positive charge and, in turn, the increased
tendency for the formation of ion pairs and the associated
restricted freedom of the ions may increase the possibility of direct
participation of the cation during oxidative decomposition.
Additionally, increased localisation of the cation charge
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Figure 11. Cyclic voltammograms of Li electrodeposition/stripping in (0.9)IL – (0.1)Li-salt binary mixtures at a Cu working electrode for the 5 pyrrolidinium-based
ILs containing (a) Li[TFSI] (left panel) and (b) Li[FSI] (right panel) as the Li-salt component, respectively. The 1st (▬), 10th (▬), 20th (▪▪▪) and 30th (▬) cycles are
shown and the IL component is given above each respective trace. The arrows show the scan direction. The scan rate was 10 mV·s-1 and the potentials are
displayed vs. Fc+/Fc (top x-axis) and vs. Li+/Li (bottom x-axis).
compositions were prepared with Li[TFSI] and lithium
bis(fluorosulfonyl)imide (Li[FSI]) as the Li-salt component. These
molar fraction compositions, corresponding to ca. 0.25 – 0.18
mol·kg-1 Li-salt for the different ILs, were selected to limit the
impact of the Li-salt component on the transport properties of the
IL. Further, Li[FSI] was selected as a comparative salt as it
displays improved conductivity in IL/salt mixtures relative to
Li[TFSI] in these ILs,[82] and the [FSI]- also shows evidence of
good SEI (solid electrolyte interphase) formation during Li anode
plating/stripping processes[12, 36] and also during intercalation into
Li-ion graphitic anodes.[30, 82] Figure 11 shows the 1st, 10th, 20th,
and 30th cycles of CV experiments at a Cu macrodisk working
electrode performed at 10 mV·s-1 under the inert atmosphere of
the Ar-filled glovebox with (a) Li[TFSI] and (b) Li[FSI] as the salt
component. All ILs studied herein display cathodic Li metal
deposition at potentials more negative than -3.2 V vs. Fc+/Fc
which is succeeded by an anodic stripping current peak on the
reverse scan. As for the measurement of the electrochemical
windows at the GC electrode, the reference potentials were
normalised initially vs. an internal Fc+/Fc couple in each sample
and further normalised using the approximation EFc+/Fc ≈ 3.2 V vs.
ELi+/Li. Likewise, exemplary CVs of the Fc+/Fc couple vs. the
Ag[OTf]/Ag and Li-metal reference electrodes at the Cu working
electrode are presented in Figure S5b in the ESI to support this
approximation.
onsets of the cathodic deposition and anodic stripping currents
shift to more negative potentials vs. the Fc+/Fc potential by ca.
100 mV. This shift is consistent with a previous analogous system
containing Li[TFSI] in [Pip1(2o2o1)][TFSI],[41] and is more
pronounced for both examples containing to two ether chains,
[Pyrr(2o2o1)2][TFSI]. This behaviour prompts the consideration that
the secondary ethereal oxygen in the cation functional groups (i.e.
the oxygen furthest from the positive nitrogen centre) is involved
significantly in the Li+-ion solvation complex in these ILs.
Raman/NMR investigations have shown previously that
substitution of the butyl group for the short methoxyethyl group
enables direct IL-cation-Li+ interactions (via the weak Lewis basic
ethereal oxygen atom), reducing the coordination number of
[TFSI]- anions which form the Li+-solvation sphere.[83] Furthermore,
the presence of the longer ether groups in a N-ethyl-N,N,N-tri[2-
(2-methoxyethoxy)ethyl]ammonium-[TFSI] IL has been recently
reported to promote the solubility of Na[TFSI] and display
correlated relaxation dynamics between IL and Na+ ions,
indicative of significant interactions, and potential cation-cation
cross-linking, between alkali metal cation and IL cation.[43] As
such based on these interesting observations, and the apparent
increased cycling stability of the di[2-(2-methoxyethoxy)ethyl]-
functionalised IL (relative to [Pyrr1(2o2o1)][TFSI]), further
investigations in the IL/Li-salt interactions and the effects of IL-salt
concentrations in these ILs are ongoing.
During the first cycles, independently of the Li-salt component, the
Beyond the first cycle in all studied ILs, the onset of Li
onset of Li electrodeposition occurs at ca. -30 to -60 mV vs. Li+/Li electrodeposition shifted to less cathodic potentials slightly but the
+
in the [Pyrr14]+, [Pyrr1(2o1)
]
and [Pyrr(2o1)2]+-based ILs. On the
potential to drive the electrodeposition process at higher currents
shifted to more cathodic potentials, associated with evolution of
higher overpotentials required to drive the process. Similarly, with
progressive cycling, anodic shifts of the stripping peak can be
observed for most of the IL formulations. However, with both
reverse scan, following a characteristic nucleation loop, anodic
oxidation associated with stripping of the new Li surfaces begins
at ca. 0 – 50 mV vs. Li+/Li for the same ILs. However, one of the
most obvious differences between the electrolyte mixtures can be
seen for the ILs functionalised with the methoxyethoxyethyl-group, Li[TFSI] and Li[FSI] salts, the [Pyrr(2o1)2][TFSI] IL displayed
[Pyrr1(2o2o1)][TFSI] and [Pyrr(2o2o1)2][TFSI]. In these mixtures, the uniform cycling up to 30 cycles; albeit with very low charge
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10.1002/cphc.201700246
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utilisation. This performance, and the apparent reversibility of the
[Pyrr(2o2o1)2][TFSI] IL, is interesting given these di-ether
functionalised ILs display some of the poorest cathodic stability
limits (see Figure 10, Table 7) at the GC electrode. Evidence for
the apparent stabilisation of some [TFSI]--based ILs towards
reductive decomposition in the presence of dissolved Li-salts has,
however, been reported numerous times previously.[28, 41, 79]
The best cycling performance, taking into consideration stability
on repeat cycling and the maximum utilised current densities,
appears to be the combination of [Pyrr1(2o1)][TFSI] with Li[FSI] salt.
This may be expected given this IL exhibits the best transport
properties from the all ILs described in this work, the positive Li+
conducting SEI film forming properties of the [FSI]- anion,[12, 36]
and the good stripping/plating performance reported previously
for this IL.[79] However, the applicability of these ILs in for Li-metal
systems (and possibly Na systems also) could be further
optimised by investigation of alternative salts, salt concentration
ranges and SEI film forming additive compounds.
and further lengthening of the ether group yields greater
reductions. Most interestingly, despite the ILs sharing the
equivalent [TFSI]- anion, the modifications of the cation structure
affect the greatest changes in the oxidative stability of the IL,
contrary to the classical understanding. Calculations of the
HOMO/LUMO energies of single ion conformers reveal the least
oxidatively stable orbital of the cation lies on the terminal oxygen
of the ether groups but the calculated magnitudes still indicate the
lowest thermodynamic pathway towards IL oxidation should be
removal of the electron from the HOMO of the [TFSI]- anion.
However, increasing ether functionalisation at the cation, firstly,
increases the propensity for ion pair formation (Walden
behaviour) and, secondly, increases the geometric stability of the
less electrochemically stable cis-[TFSI]- conformer (DFT
calculations). These factors may contribute to the observed
reduction in the oxidative stability limits of the ILs but further
experimental work is required to fully understand the sequence of
events occurring at the electrode/electrolyte interface during
electrochemical decomposition of the IL. Preliminary
investigations also revealed, or further supported, the applicability
of these ILs for plating/stripping processes at Li-metal anodes.
Some of these results were indicative of interesting IL-cation – Li-
cation interaction effects observed with higher degree of ether-
group functionalisation, prompting further investigations into
these IL/Li-salt compositions.
Conclusions
The
synthesis
and
systematic
thermophysical
and
electrochemical investigations of 15 cyclic alkylammonium [TFSI]-
based ILs with and without varying degrees of ether functionality
is reported. All studied ILs exhibit low meting temperatures and
high thermal stabilities. Ether functionality on the cation slightly
reduces upper limits of thermal stabilities but all ILs exist in the
liquid phase under wide range of potential operating temperatures
for electrochemical devices (e.g. -20 to >100 °C). The transport
properties, relating to the IL viscosity and conductivity, are
strongly effected by cation structure. Initial substitution of alkyl
groups for like-sized ether groups affords important reductions in
IL viscosity and enhances IL conductivities observed, in line with
previous observations. However, increasing the degree of ether
functionality engages more complex trends in viscosity and
conductivity depending on cation ring size. For example, the
longer flexible ether chains may increase degrees of rotational
freedom and reduce cation-cation interactions, increasing IL
fluidity but, in most cases, reducing IL conductivity. The bulky side
chains obviously increase cation size, likely reducing ionic
mobilities (most prominent for the small pyrrolidinium cations), but
also increase the propensity for ion pair formation, as evidenced
by the assessment of the Walden-type behaviour for the studied
ILs. Herein, all but one of the studied ILs exhibited >40 % ionicity
and may be classed as good ionic liquids but, as stated, the
introduction of ether functionality at the cation generally affected
reductions in the observed ionicity. Furthermore, estimations of
the IL fragilities were completed using extrapolation of both the
viscosity and conductivity measurements towards glass transition
temperatures. These results further highlighted the coupled
nature of both physical properties and revealed the studied ILs to
be highly fragile materials; a desirable quality for IL electrolytes.
Furthermore, all of the studied ILs exhibited wide electrochemical
windows (>4.5 V), an important feature for potential IL electrolytes.
Nevertheless, initial substitution of the alkyl groups for ether
groups reduces the magnitude of the observed windows slightly
Experimental Section
Materials
The materials used to synthesise the ionic liquids were sourced as follows:
pyrrolidine (99 %), piperidine (99 %) N-butylpyrrolidine (99 %), dimethyl
sulfate (>99 %), 1-bromobutane (98 %), ferrocene (98 %) ex Sigma-
Aldrich; azepane (99 %) ex INVISTA; 2-bromoethyl methyl ether (99 %)
and 1-bromo-2-(2-methoxyethoxy)ethane (95 %) ex Fluorochem; lithium
bis{(trifluoromethyl)sulfonyl}imide (battery grade) ex 3M All synthesis
reagents were used as received. Lithium bis(fluorosulfonyl)imide
(<99.5 %) used for Li-redox experiments was sourced from Kanto
Chemical Co. Inc.
Ionic liquid synthesis
Full description of the synthesis of the 15 studied ILs is provided in the ESI
in conjunction with results of NMR, Microanalysis, and Li-content analysis.
Preparation of materials for physiochemical measurements
Prior to all physical and electrochemical measurements, all ILs were dried
under vacuum (ca. 10-3 mbar) at 353 – 373 K with stirring for at least 2
days. This treatment was found to reduce the water content to below 100
ppm, as measured by Karl Fischer coulometric titration. The resolution of
the water content measurements was 0.001 wt/wt % (or 10 ppm) and
measurements were completed in duplicate. Once dried, the IL samples
were stored in the Ar-filled glovebox with moisture levels less than 3 ppm
water. Lithium content of the ILs was analysed by inductively coupled
plasma optical emission spectroscopy (ICP-OES) on an Agilent 5100 ICP-
OES, and along with Microanalysis, was performed by Analytical Services
at Queen’s University, Belfast. 1H and 13C NMR spectra were recorded at
293 K on a Bruker Avance DPX spectrometer at 300 MHz and 75 MHz
respectively. 1H and 13C NMR spectra for all [TFSI]- ILs are shown in
Figures S7 to S36 in the ESI.
This article is protected by copyright. All rights reserved.

10.1002/cphc.201700246
ChemPhysChem
ARTICLE
Physical measurements
Li[TFSI] was dissolved in the ferrocene-containing IL and a small strip of
fresh Li-metal was used as the reference electrode for cyclic voltammetry
measurements of the ferrocene couple.
Density measurements were performed using a DM40 oscillating tube
density meter (Mettler Toledo, ± 1 x 10-4 g·cm-3) in the range of 293.15 –
363.15 K (± 0.1 K). The instrument was cleaned using acetone and dried
using dehumidified air prior to any measurements. The viscosity of the ILs
was measured using a Bohlin Gemini Rotonetic Drive 2 cone and plate
rheometer (±1 %) from 293 – 363 K (± 0.01 K) at atmospheric pressure.
The rheometer was calibrated using ultra-pure water and an oil viscosity
standard (ASTM Oil Standard S600, Cannon, 1053 mPa·s at 298.15 K).
Dynamic thermal decomposition profiles were collected by
thermogravimetric analysis (TGA) using a TGA Q5000 (TA Instruments).
Heating profiles were performed at a rate of 10 K·min-1 under nitrogen flow
from room temperature to 773 K (± 1 K). Thermal phase transitions were
recorded using differential scanning calorimetry (DSC) analysis on a DSC
Q2000 (TA Instruments) with ca. 5 mg IL samples in hermetically sealed
Al pans. A heating gradient of 5 K·min-1 was used and 183 K and 323 K (±
0.1 K) were used as the lower and upper temperature limits, respectively.
Conductivity measurements were conducted using a sensION+ EC71
benchtop meter with a 3-pole platinum sensION+ 5070 conductivity probe
(< 0.5 % of range) with an in-built Pt1000 temperature probe (Hach Lange).
Regular calibration of the probe was completed using aqueous KCl
standard conductivity solutions (147 μS·cm-1, 1413 μS·cm-1, and 12.88
mS·cm-1 at 298 K). The immersion and sealing of the conductivity probe in
the liquid sample was carried out in an Ar-filled glovebox. The conductivity
probe, when disconnected from the meter, was immersed in the liquid
sample (ca. 3 cm3) inside a glass sample tube also containing a small
magnetic stirrer bar. The sample was sealed with the probe in the glass
tube using an O-ring seal and Parafilm. The conductivity of the sample was
then recorded with stirring as a function of temperature (using the
temperature reading built into the conductivity probe). The temperature of
the sample was varied from 293 – 363 K (± 0.2 K) using a small oil bath
and a hot-plate with a thermocouple control. The sample conductivity and
temperature were recorded when the observed values were stable for ca.
1 min.
For the measurements of Li-redox chemistry, both Li[TFSI] and Li[FSI]
salts were firstly dried at 373 K under vacuum (ca. 10-3 mbar followed by
ca. 10-5 mbar, turbomolecular vacuum pump station) for several days.
IL/Li-salt binary mixtures were then prepared accurately by mass inside
the Ar-filled glovebox with molar fractions of (0.9)IL – (0.1)Li-salt and left
to stir overnight to ensure complete dissolution. The cells were then
prepared using the Cu working electrode, the Pt-wire counter electrode
and the Ag[OTf]/[Pyrr14][TFSI] reference electrode and CV was used at a
10 mV·s-1 scan rate to assess Li plating/stripping with repeated cycling.
Therein, the cathodic current density was limited to -3 mA·cm-2 or,
alternatively, to a portion (ca. ¾ to ⅔) of the maximum peak current
observed on the first cycle. Following completion of the cycling, the
reference potential was normalised vs. the ferrocene couple as described
previously.
Computational methods
Calculations of single-ion and ion pair HOMO and LUMO energies were
completed using the Turbomole 7.0 program package,[84] as described by
our group previously.[77] Before the visualization of the HOMO/LUMO
orbitals using TmoleX (version 4.1.1), single ion structures were optimized
in the gas phase, with a convergence criterion of 10-8 Hartree, using DFT
calculations combining the Resolution of Identity (RI) approximation[85, 86]
within the Turbomole 7.0 program package using the B3LYP function with
the def-TZVP basis set.[87-89] The resulting optimized structures were then
used as inputs in the COSMOconfX program (version 4.0) to generate the
conformers of each species. Single point energy calculations
(DFT/B3LYP/def-TZVP + RI approximation) within Turbomole were then
utilized to determine the HOMO/LUMO orbitals of the predefined
conformers of each species.
Electrochemical measurements
Acknowledgements
All electrochemical measurements of the electrochemical windows and the
Li-redox chemistry were completed using cyclic voltammetry using either
an Autolab PGSTAT 302 workstation (Metrohm) or a VMP3 multi-channel
workstation (BioLogic). A custom 3 port glass cell was prepared in-house
which allowed completion of all electrochemical measurements with ca. 1
cm3 of electrolyte. Cell preparation and electrochemical measurements
were conducted inside the Ar-filled glovebox. The temperature of the
electrochemical measurements was ca. 303 K, as determined by the
internal atmosphere of the glovebox. For the determination of
electrochemical windows, the working electrode was a 3 mm diameter
glassy carbon macrodisk electrode (ALS Co., Ltd). For Li-redox chemistry,
the working electrode was a 1.6 mm diameter copper macrodisk electrode
(ALS Co., Ltd). Prior to all measurements, the working electrodes were
cleaned with acetonitrile and ultra-pure water and then polished using
alumina slurries of decreasing grain size (1.0 μm, 0.3 μm and 0.05 μm,
Buehler) in ultra-pure water on soft microcloth lapping pads (Kemet Ltd.).
Following the polishing, the working electrodes were sonicated in ultra-
pure water for 2 minutes. The GC electrode was dried in an oven at ca.
353 K and the Cu electrode was dried under a nitrogen stream. The
counter electrode was a platinum wire coil heat sealed in a glass capillary
and the reference electrodes used were based on a silver wire in a 0.1
mol·dm-3 solution of either Ag[NO3] in 1-butyl-3-methylimidazolium nitrate
or silver trifluoromethanesulfonate (Ag[OTf]) in [Pyrr14][TFSI] separated
from the bulk solution by a glass frit. The reference potentials were
normalized vs. an internal ferrocene redox couple wherein a small amount
of ferrocene was dissolved in the IL following completion each
measurement. Cyclic voltammetry was then used to approximate the
formal potential of the ferrocene/ferrocenium redox couple. To verify the
approximation of the reference potential vs. Li+/Li, a small quantity of
The authors would like to gratefully acknowledge the funding from
EPSRC (EP/L505262/1) and Innovate UK for the Practical
Lithium-Air Batteries project (project number: 101577).
Additionally, the authors would like thank Keith Whiston and
INVISTA Intermediates for supplying the samples of azepane for
this work.
Keywords: Ammonium cations, electrochemistry, electrolytes
ether functionalisation, ionic liquids
[1]
[2]
Q. Zhang, S. Zhang, Y. Deng, Green Chem. 2011, 13, 2619-2637.
V. I. Pârvulescu, C. Hardacre, Chem. Rev. (Washington, DC, U. S.).
2007, 107, 2615-2665.
[3]
J. Dupont, R. F. de Souza, P. A. Z. Suarez, Chem. Rev. (Washington,
DC, U. S.). 2002, 102, 3667-3692.
[4]
[5]
[6]
[7]
[8]
X. Han, D. W. Armstrong, Acc. Chem. Res. 2007, 40, 1079-1086.
M. D. Joshi, J. L. Anderson, RSC Adv. 2012, 2, 5470-5484.
M. Palacio, B. Bhushan, Tribol. Lett. 2010, 40, 247-268.
M. Gorlov, L. Kloo, Dalton Trans. 2008, 2655-2666.
G. P. S. Lau, H. N. Tsao, S. M. Zakeeruddin, M. Grätzel, P. J. Dyson,
ACS Appl. Mater. Interfaces. 2014, 6, 13571-13577.
P. Wang, S. M. Zakeeruddin, J.-E. Moser, M. Grätzel, J. Phys. Chem.
B. 2003, 107, 13280-13285.
[9]
[10] S. Liu, L. Zhou, P. Wang, F. Zhang, S. Yu, Z. Shao, B. Yi, ACS Appl.
Mater. Interfaces. 2014, 6, 3195-3200.
[11] A. Lewandowski, A. Świderska-Mocek, J. Power Sources. 2009, 194,
601-609.
[12] H. Matsumoto, H. Sakaebe, K. Tatsumi, M. Kikuta, E. Ishiko, M. Kono,
J. Power Sources. 2006, 160, 1308-1313.
This article is protected by copyright. All rights reserved.

10.1002/cphc.201700246
ChemPhysChem
ARTICLE
[13] A. Brandt, S. Pohlmann, A. Varzi, A. Balducci, S. Passerini, MRS Bull.
2013, 38, 554-559.
[51] M. Vranes, S. Dozic, V. Djeric, S. Gadzuric, J. Chem. Eng. Data. 2012,
57, 1072-1077.
[14] C. Wolff, S. Jeong, E. Paillard, A. Balducci, S. Passerini, J. Power
Sources. 2015, 293, 65-70.
[15] A. J. R. Rennie, N. Sanchez-Ramirez, R. M. Torresi, P. J. Hall, J. Phys.
Chem. Lett. 2013, 4, 2970-2974.
[52] U. Domańska, M. Królikowska, J. Solution Chem. 2012, 41, 1422-1445.
[53] M. C. Cumicheo, L. C. S. Nobre, A. F. Santos, I. M. S. Lampreia, M. S.
C. S. Santos, F. J. V. Santos, H. Segura, C. A. Nieto de Castro, J.
Chem. Eng. Data. 2015, 60, 3766-3775.
[16] G. A. Elia, J. Hassoun, W. J. Kwak, Y. K. Sun, B. Scrosati, F. Mueller,
D. Bresser, S. Passerini, P. Oberhumer, N. Tsiouvaras, J. Reiter, Nano
Lett. 2014, 14, 6572-6577.
[54] E. Couadou, J. Jacquemin, H. Galiano, C. Hardacre, M. Anouti, J.
Phys. Chem. B. 2013, 117, 1389-1402.
[55] L. J. A. Siqueira, M. C. C. Ribeiro, J. Phys. Chem. B. 2009, 113, 1074-
[17] J. Wang, S. Y. Chew, Z. W. Zhao, S. Ashraf, D. Wexler, J. Chen, S. H.
Ng, S. L. Chou, H. K. Liu, Carbon. 2008, 46, 229-235.
[18] D. R. MacFarlane, N. Tachikawa, M. Forsyth, J. M. Pringle, P. C.
Howlett, G. D. Elliott, J. H. Davis, M. Watanabe, P. Simon, C. A. Angell,
Energy Environ. Sci. 2014, 7, 232-250.
1079.
[56] M. Galiński, A. Lewandowski, I. Stępniak, Electrochim. Acta. 2006, 51,
5567-5580.
[57] K. Ueno, H. Tokuda, M. Watanabe, Phys. Chem. Chem. Phys. 2010,
12, 1649-1658.
[19] A. P. Abbott, K. J. McKenzie, Phys. Chem. Chem. Phys. 2006, 8, 4265-
[58] W. Xu, E. I. Cooper, C. A. Angell, J. Phys. Chem. B. 2003, 107, 6170-
6178.
4279.
[20] S. Zein El Abedin, F. Endres, ChemPhysChem. 2006, 7, 58-61.
[21] J. C. Forgie, S. El Khakani, D. D. MacNeil, D. Rochefort, Phys. Chem.
Chem. Phys. 2013, 15, 7713-7721.
[22] L. Lombardo, S. Brutti, M. A. Navarra, S. Panero, P. Reale, J. Power
Sources. 2013, 227, 8-14.
[59] C. Schreiner, S. Zugmann, R. Hartl, H. J. Gores, J. Chem. Eng. Data.
2010, 55, 1784-1788.
[60] A. R. Neale, S. Murphy, P. Goodrich, C. Schütter, C. Hardacre, S.
Passerini, A. Balducci, J. Jacquemin, J. Power Sources. 2016, 326,
549-559.
[23] G. H. Lane, A. S. Best, D. R. MacFarlane, M. Forsyth, P. M. Bayley, A.
F. Hollenkamp, Electrochim. Acta. 2010, 55, 8947-8952.
[24] M. Diaw, A. Chagnes, B. Carré, P. Willmann, D. Lemordant, J. Power
Sources. 2005, 146, 682-684.
[61] C. Schreiner, S. Zugmann, R. Hartl, H. J. Gores, J. Chem. Eng. Data.
2010, 55, 4372-4377.
[62] S. Ferrari, E. Quartarone, C. Tomasi, D. Ravelli, S. Protti, M. Fagnoni,
P. Mustarelli, J. Power Sources. 2013, 235, 142-147.
[63] P. Sippel, P. Lunkenheimer, S. Krohns, E. Thoms, A. Loidl, Sci. Rep.
2015, 5, 13922.
[25] K. Tsunashima, M. Sugiya, Electrochem. Solid-State Lett. 2008, 11,
A17-A19.
[26] E. Frackowiak, G. Lota, J. Pernak, Appl. Phys. Lett. 2005, 86, 164104.
[27] R. S. Kühnel, N. Böckenfeld, S. Passerini, M. Winter, A. Balducci,
Electrochim. Acta. 2011, 56, 4092-4099.
[28] S. Fang, Z. Zhang, Y. Jin, L. Yang, S.-i. Hirano, K. Tachibana, S.
Katayama, J. Power Sources. 2011, 196, 5637-5644.
[29] H. Yoon, P. C. Howlett, A. S. Best, M. Forsyth, D. R. MacFarlane, J.
Electrochem. Soc. 2013, 160, A1629-A1637.
[64] R. Böhmer, K. L. Ngai, C. A. Angell, D. J. Plazek, J. Chem. Phys. 1993,
99, 4201-4209.
[65] C. A. Angell, Science. 1995, 267, 1924-1935.
[66] T. L. Greaves, A. Weerawardena, C. Fong, I. Krodkiewska, C. J.
Drummond, J. Phys. Chem. B. 2006, 110, 22479-22487.
[67] K. R. Harris, M. Kanakubo, L. A. Woolf, J. Chem. Eng. Data. 2006, 51,
1161-1167.
[30] J. Reiter, M. Nádherná, R. Dominko, J. Power Sources. 2012, 205,
402-407.
[31] T. Belhocine, S. A. Forsyth, H. Q. N. Gunaratne, M. Nieuwenhuyzen, P.
Nockemann, A. V. Puga, K. R. Seddon, G. Srinivasan, K. Whiston,
Green Chem. 2011, 13, 3137.
[68] M. C. C. Ribeiro, J. Chem. Phys. 2010, 133, 024503.
[69] E. Simonetti, G. Maresca, G. B. Appetecchi, G. T. Kim, N. Loeffler, S.
Passerini, J. Power Sources. 2016, 331, 426-434.
[70] F. Béguin, V. Presser, A. Balducci, E. Frackowiak, Adv. Mater.
(Weinheim, Ger.). 2014, 26, 2219-2251.
[32] N. Salem, L. Nicodemou, Y. Abu-Lebdeh, I. J. Davidson, J.
Electrochem. Soc. 2012, 159, A172-A176.
[33] V. Ruiz, T. Huynh, S. R. Sivakkumar, A. G. Pandolfo, RSC Adv. 2012,
2, 5591-5598.
[71] A. Balducci, R. Dugas, P. L. Taberna, P. Simon, D. Plée, M.
Mastragostino, S. Passerini, J. Power Sources. 2007, 165, 922-927.
[72] V. Borgel, E. Markevich, D. Aurbach, G. Semrau, M. Schmidt, J. Power
Sources. 2009, 189, 331-336.
[34] S. Pohlmann, T. Olyschläger, P. Goodrich, J. A. Vicente, J. Jacquemin,
A. Balducci, Electrochim. Acta. 2015, 153, 426-432.
[35] C. Schütter, A. R. Neale, P. Wilde, P. Goodrich, C. Hardacre, S.
Passerini, J. Jacquemin, A. Balducci, Electrochim. Acta. 2016, 220,
146-155.
[36] E. Paillard, Q. Zhou, W. A. Henderson, G. B. Appetecchi, M.
Montanino, S. Passerini, J. Electrochem. Soc. 2009, 156, A891-A895.
[37] J. Reiter, S. Jeremias, E. Paillard, M. Winter, S. Passerini, Phys. Chem.
Chem. Phys. 2013, 15, 2565-2571.
[73] P. Bonhôte, A.-P. Dias, N. Papageorgiou, K. Kalyanasundaram, M.
Grätzel, Inorg. Chem. 1996, 35, 1168-1178.
[74] M. C. Buzzeo, C. Hardacre, R. G. Compton, ChemPhysChem. 2006, 7,
176-180.
[75] S. P. Ong, O. Andreussi, Y. Wu, N. Marzari, G. Ceder, Chem. Mater.
2011, 23, 2979-2986.
[76] P. C. Howlett, E. I. Izgorodina, M. Forsyth, D. R. Macfarlane, Z. Phys.
Chem. (Muenchen, Ger.). 2006, 220, 1483.
[77] E. Coadou, P. Goodrich, A. R. Neale, L. Timperman, C. Hardacre, J.
Jacquemin, M. Anouti, ChemPhysChem. 2016, 17, 3992-4002.
[78] M. Montanino, F. Alessandrini, S. Passerini, G. B. Appetecchi,
Electrochim. Acta. 2013, 96, 124-133.
[79] J. Reiter, E. Paillard, L. Grande, M. Winter, S. Passerini, Electrochim.
Acta. 2013, 91, 101-107.
[80] K. Yoshida, M. Nakamura, Y. Kazue, N. Tachikawa, S. Tsuzuki, S.
Seki, K. Dokko, M. Watanabe, J. Am. Chem. Soc. 2011, 133, 13121-
13129.
[38] Z.-B. Zhou, H. Matsumoto, K. Tatsumi, Chem-Eur. J. 2005, 11, 752-
766.
[39] Z. J. Chen, T. Xue, J.-M. Lee, RSC Adv. 2012, 2, 10564-10574.
[40] A. R. Neale, P. Li, J. Jacquemin, P. Goodrich, S. C. Ball, R. G.
Compton, C. Hardacre, Phys. Chem. Chem. Phys. 2016, 18, 11251-
11262.
[41] H. Matsumoto, H. Sakaebe, K. Tatsumi, J. Power Sources. 2005, 146,
45-50.
[42] H.-B. Han, K. Liu, S.-W. Feng, S.-S. Zhou, W.-F. Feng, J. Nie, H. Li, X.-
J. Huang, H. Matsumoto, M. Armand, Z.-B. Zhou, Electrochim. Acta.
2010, 55, 7134-7144.
[43] C. R. Pope, M. Kar, D. R. MacFarlane, M. Armand, M. Forsyth, L. A.
O'Dell, ChemPhysChem. 2016, 17, 3187-3195.
[44] C. Pozo-Gonzalo, M. Kar, E. Jónsson, P. C. Howlett, D. R. MacFarlane,
M. Forsyth, Electrochim. Acta. 2016, 196, 727-734.
[45] M. Kar, B. Winther-Jensen, M. Armand, T. J. Simons, O. Winther-
Jensen, M. Forsyth, D. R. MacFarlane, Electrochim. Acta. 2016, 188,
461-471.
[81] H. G. Jung, J. Hassoun, J. B. Park, Y. K. Sun, B. Scrosati, Nat. Chem.
2012, 4, 579-585.
[82] M. Nádherná, J. Reiter, J. Moškon, R. Dominko, J. Power Sources.
2011, 196, 7700-7706.
[83] M. Kunze, E. Paillard, S. Jeong, G. B. Appetecchi, M. Schonhoff, M.
̈
Winter, S. Passerini, J. Phys. Chem. C. 2011, 115, 19431-19436.
[84] R. Ahlrichs, M. Bär, M. Häser, H. Horn, C. Kölmel, Chem. Phys. Lett.
1989, 162, 165-169.
[85] F. Weigend, M. Häser, Theor. Chem. Acc. 1997, 97, 331-340.
[86] F. Weigend, M. Häser, H. Patzelt, R. Ahlrichs, Chem. Phys. Lett. 1998,
294, 143-152.
[87] A. D. Becke, Phys. Rev. A. 1988, 38, 3098-3100.
[88] C. Lee, W. Yang, R. G. Parr, Physical Review B. 1988, 37, 785-789.
[89] S. Grimme, J. Comput. Chem. 2006, 27, 1787-1799.
[46] K.-S. Kim, S. Choi, D. Demberelnyamba, H. Lee, J. Oh, B.-B. Lee, S.-J.
Mun, Chem. Commun. 2004, 828-829.
[47] D. R. MacFarlane, P. Meakin, J. Sun, N. Amini, M. Forsyth, J. Phys.
Chem. B. 1999, 103, 4164-4170.
[48] F. Capitani, F. Trequattrini, O. Palumbo, A. Paolone, P. Postorino, J.
Phys. Chem. B. 2016, 120, 2921-2928.
[49] K. Tsunashima, M. Sugiya, Electrochem. Commun. 2007, 9, 2353-
2358.
[50] Z.-B. Zhou, H. Matsumoto, K. Tatsumi, Chem-Eur. J. 2006, 12, 2196-
2212.
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Entry for the Table of Contents
ARTICLE
Alex R. Neale,* Sinead Murphy, Peter
Goodrich, Christopher Hardacre and
Johan Jacquemin*
Page No. – Page No.
Thermophysical and Electrochemical
Properties of Ethereal Functionalised
Cyclic Alkylammonium-based Ionic
Liquids as Potential Electrolytes for
Electrochemical Applications
Balancing stabilities and mobilities: Trends in physical and electrochemical
characteristics of a variety of alkyl/ether-functionalised cyclic ammonium
bis{(trifluoromethyl)sulfonyl}imide ionic liquids highlight balances and trade-offs
between chemical structures and important electrolyte properties.
This article is protected by copyright. All rights reserved.