jo u r n al hom ep age: www.elsev ier .com/ locate /e lec tac ta
hysicochemical properties of ammonium-based deep eutecticolvents and their electrochemical evaluation using organometalliceference redox systems
aleh Bahadoria, Mohammed Harun Chakrabarti a,b,∗, Farouq Sabri Mjalli c,nas Muen AlNashefd, Ninie Suhana Abdul Manane, Mohd Ali Hashima
Department of Chemical Engineering, Faculty of Engineering, University of Malaya, Kuala Lumpur 50603, MalaysiaEnergy Futures Lab, Electrical Engineering Building, Imperial College London, South Kensington, London SW7 2AZ, United KingdomPetroleum & Chemical Engineering Department, Sultan Qaboos University, Muscat 123, OmanChemical Engineering Department, King Saud University, P.O. Box 800, Riyadh, Saudi ArabiaDepartment of Chemistry, University of Malaya, Kuala Lumpur 50603, Malaysia
r t i c l e i n f o
rticle history:eceived 24 July 2013eceived in revised form8 September 2013ccepted 18 September 2013vailable online 12 October 2013
a b s t r a c t
Seven deep eutectic solvents (DESs) containing ammonium based salts are prepared by means ofhydrogen bonding with acid, amine, amide and nitrate based compounds. The major physicochem-ical properties of the DESs in terms of density, viscosity, electrical conductivity, molar conductivityand pH are investigated prior to ascertaining their electrochemical characteristics by means of cyclicvoltammetry and chronoamperometry. Nitrate based DESs exhibit higher conductivities but lower vis-cosities than other DESs, whereas the amide based DES displays the widest electrochemical potential
window. Diffusion coefficient, D, of two organometallic redox couples, Fc/Fc+ (ferrocene/ferrocenium)and Cc/Cc+(cobaltocene/cobaltocenium) is found to be of the order of 10−9 to 10−8 cm2 s−1 in all stud-ied DESs while the heterogeneous rate constant for electron transfer across the electrode/DES interfaceis of the order of 10−4 cm s−1. The Stokes–Einstein products of Fc and Cc+ in the DESs have also beendetermined.
alden plot of DESs
. Introduction
Ionic liquids (ILs) have been intensively studied as electrolytesue to their attractive characteristics such as broad potential win-ows, intrinsic conductivities and high thermal stabilities [1–4] inomparison to common organic solvents like acetonitrile [5]. Theresence of an abundance of charge carriers abrogates the needor added supporting electrolytes thereby simplifying the experi-
ental setup [6–8]. ILs have been used in many electrochemicalpplications such as electrochemical sensors [9], solar cells [10],apacitors [11], fuel cells [12] and in lithium batteries [13]. Theyave already been found to provide several benefits during elec-
rochemical studies when used as either the solvent or electrolyte14–16]. For electrochemical applications, fundamental charac-eristics related to electrochemical and physical properties of ILs
∗ Corresponding author at: Department of Chemical Engineering, Faculty of Engi-eering, University of Malaya, Kuala Lumpur 50603, Malaysia. Tel.: +60 3 79677655;
include viscosity, electrochemical stability, conductivity and thediffusion coefficient of the electroactive species in the ILs [17–19].
A related class of ILs named deep eutectic solvents (DESs)have been revealed to overcome the high price and toxicity ofILs [20–22]. DESs are obtained by fusing quaternary ammonium-or phosphonium-based salts with hydrogen bond donors, whichare capable of forming a eutectic mixture [23–25]. DESs sharemany favorable characteristics of traditional ILs, for instance non-volatility, non-reactivity with water and biodegradability [21,26].
In the electrochemistry of ILs, the oxidation of Fc/Fc+
(ferrocene/ferrocenium), or reduction of Cc/Cc+ (cobal-tocene/cobaltocenium) exhibit a well-defined reversible redoxprocess that are widely used as voltammetric potential referencescales [27–31]. Fc has limited solubility in ILs whereas Cc+ isgenerally soluble [27,29]. The standard redox potential of Fc ishighly dependent on solvation effects of the solvent and support-ing electrolytes [32]. However, Fc has seen limited investigationsin DESs [33] and information on the dependence of DESs on the
Cc+/Cc couple has only seen recent light [34], although that workhas remained limited to a single DES only.
We prepared several DESs based on two different quater-nary ammonium salts and evaluated their physicochemical and
lectrochemical properties (viscosity, conductivity, electrochem-cal stability, diffusion coefficient, etc.) in a similar manner tohat for ILs. Within the knowledge of the authors, this may bene of the few reports that provides a detailed account on theetermination of electrochemical potential windows, internalotential references (based on the reduction of cobaltoceniumexafluorophosphate and oxidation of ferrocene), diffusion andate coefficients of Fc and Cc+ as well as the Stokes–Einsteinroducts of Fc and Cc+ in such DESs.
. Experimental
.1. Preparation and characterization of DESs
The chemical formulae of the salts and hydrogen bond donorshat make up the DESs studied in the present work are shownn Table 1. All chemicals are purchased from Merck (Germany)
ith high purity (≥98%) and have been stored in an inert gloveox (Innovative Technology, Pure LabHE, USA) purged with argonoxygen-free). The water mass fraction of all the chemicals as perhe manufacturer’s guide is <0.05%. The DESs are formed by mixinghe two components together at 120 ◦C for a minimum period of
h until a homogeneous, hazy light greenish liquid forms [33]. Theynthesis experiments are conducted in an inert glove box with lesshan 5 ppm moisture. The physical properties are typically mea-ured in a glove box, which provides an atmosphere low in oxygennd moisture.
The viscosities of the DESs were measured using a Brookfieldnstrument. The pH was determined by means of a MetrohmH meter, which was calibrated by means of standard pH bufferolutions. The conductivities were measured using a DZS-708ulti-parameter analyser, which was calibrated using a 0.001 M
tandard solution of KCl (Merck). Temperatures were controlledsing a water bath.
Densities of the synthesized DESs were determined using aMA 4100 Density Meter (Anton Paar, Austria) with three repli-ates for each reading with an uncertainty of ±0.00007 g cm−3.o adjust the density meter for accuracy, the density of dis-illed water was measured at 25 ◦C and compared with theorresponding value in density tables. The results showed
difference of ±0.00004 g cm−3 which confirmed goodccuracy.
.2. Electrochemical measurements
Electrochemical measurements were undertaken with aomputer-controlled í-Autolab potentiostat (PGSTAT302 N) using
conventional three-electrode cell [5,33]. The electrochemical cellas located within a Faraday cage to minimize electrochemicaloise, which in turn was situated inside the dry argon-filled gloveox. This ensured that the moisture content in all samples wereelow 5 ppm in all cases.
A platinum microelectrode (Pt., 20 �m diameter) or a 3 mmiameter Glassy Carbon (GC) were used as the working elec-rodes in alternative experiments, while Pt. wire and silver wirepre-treated as described earlier [33]) were used as the counternd quasi-reference electrodes (AgQRE), respectively. Workinglectrodes were carefully polished before each voltammetry exper-ment with 0.25 �m alumina suspensions and ultrasonically rinsedn acetone. The limiting potentials for the electrochemical win-ow (defined as the difference between the cathodic and anodic
imits) of each DES were arbitrarily defined as the potentialt which the current density reached higher than 0.2 mA cm−2
hat was selected as the cut-off current density for comparison35].
T / K
Fig. 1. Densities (�) of ammonium based DESs as a function of temperature.
3. Results and discussion
3.1. Physical properties of DESs
Table 1 summarizes the chemical formulae of the DESs preparedand studied in the present work and their fundamental physico-chemical properties for a better understanding of the interactionsin these new compounds. The chemical structures of some of theDESs employed in this work are given in our previous publications[24,33].
Being a mixture, DES (a tri-component system with cations,anions, and neutral ligands) is more complicated than IL (a bi-component system with anions and cations) [36]. In addition tocations and anions, DES also contains neutral ligand molecules suchas urea that can exert structure-directing effects. In this sense, DESpossesses features of both IL and molecular solvents. Each of thethree individual components (cations, anions, and neutral ligands)has the potential to participate in the self-assembly process, eitherindividually or in combination with another component, making itpossible to create new types of chemistry that are not accessible ineither IL or molecular solvents.
3.1.1. DensityThe densities of several DESs (ranging from 1.07 to 1.47 g cm−3)
were measured at a temperature range of 30–90 ◦C. Fig. 1 shows thetemperature dependence of density for all the DESs. The behav-ior here was linear as expected. The results for the temperaturedependent densities (�) of these DESs could be well fitted by thefollowing equation:
� = AT + B (1)
where T is the absolute temperature and B is a constant thatdepends on the type of DES being investigated. The adjustableparameters of Eq. (1) for the density of the tested DESs are shown inTable 2. As indicated in Table 1, the DES 4 has higher density thanits other counter parts due to a higher intermolecular packing ofthe compound’s denser structure.
3.1.2. ViscosityThe viscosity of DESs is influenced by the salt to hydrogen bond
donor (HBD) interaction and their coordinating ability. It is an
important property for electrochemical studies because it stronglyinfluences the rate of mass transport within a solution. DESs withthe amide HBD exhibits lower viscosities than other DESs. More-over, nitrate based DESs show lower viscosities in comparison to
L. Bahadori et al. / Electrochimica Acta 113 (2013) 205– 211 207
Table 1Physicochemical properties of ammonium-based DESs.
DESs Formulae Salt HBDa Molar ratio Mw (g mol−1) � (g cm−3) � (mPa s) � (mS cm−1) pH
defined as � = Ve�. The equivalent volume (Ve) of the DES is calcu-lated from the experimental density using the equation Ve = M/�,where M is the molar mass and � is the density. Fig. 4 is a graph ofthe equivalent conductivity against the inverse of viscosity (Walden
6
7(b)
T / K
290 300 31 0 320 330 34 0 350 360 370
/ mP
a s-1
0
200
400
600
800
1000
DES1DES2DES3DES4DES5DES6DES7
(a)
ammonium chloride3 2
HBD = hydrogen bond donor.
cidic and amine based DESs thus resulting in higher conductivitieshan other DESs.
As expected, Fig. 2a shows that viscosity values of the mixtureecrease as the temperature increases from 25 to 90 ◦C. This effect
s due to the higher mobility of ions. The temperature dependentiscosity plot for DESs is represented by the Arrhenius plot (Fig. 2b)s shown in Eq. (2), which describes the temperature dependenceor non-associating electrolytes [37].
n� = ln�o + E�RT
(2)
here T is the absolute temperature, � is the viscosity, E� is thectivation energy, �o is a constant and R is the gas constant. Thealue of E�, �o and sums of square errors are depicted in Table 3.he regression correlation coefficients have values higher than 0.96howing a reasonably good fit.
.1.3. ConductivityIonic conductivity is one of the most important properties of
ESs as electrolyte materials. Fig. 3 shows the temperature depend-nce of conductivity for the DESs; the ions move faster at higheremperatures due to the relatively lower viscosity of the neat DESs.he electrical conductivity exhibits exponential behavior with tem-erature for all measured DESs. The conductivity of DESs (�) variesith temperature according to Eq. (3):
n� = ln�o + E�RT
(3)
here E� is the activation energy for conduction, T is the absolute
emperature and �o is a constant. Consequently, from Eq. (3), E� , �o
nd sums of square errors are depicted in Table 3. The regressionorrelation coefficients have values higher than 0.97 indicating aeasonable fit of the data.
able 2he adjustable parameters for density and pH of DESs used in Eq. (2).
3.1.4. Molar conductivityThe molar conductivity � (m2 S mol−1) of the electrolyte is
1000 / T (K-1)
2.7 2.8 2.9 3.0 3.1 3.2 3.3 3.4
ln (
/ m
Pa
s-1)
1
2
3
4
5
DES1DES2DES3DES4DES5DES6DES7
Fig. 2. (a) Dynamic viscosity (�) of ammonium based DESs as a function of temper-ature; (b) Arrhenius plot of viscosity (�) for the ammonium based DESs.
208 L. Bahadori et al. / Electrochimica Acta 113 (2013) 205– 211
Table 3Regression Parameters for viscosity and conductivity of DESs.
lot) over a temperature range of 25–80 ◦C. This scheme is par-icularly applicable in ILs [38,39], because it provides the basisor understanding the relationship between conductivity and lowapor pressures and is a useful measure for examining the ion pair-ng problem in electrolytes. If the conductivity and viscosity of thelectrolyte comply with Walden’s rule, the conductivity is corre-ated to the temperature-dependent viscosity using a qualitativepproach [40]:
�˛ = C (4)
In the log–log plot of Fig. 4, the slope of the Walden line is givens ˛, while C is a temperature-dependent constant. In comparisono the ideal line it can be seen that DES 5 and DES 6 lie below theideal” Walden line. The other DESs examined from Walden plots
T / K
280 30 0 320 340 360 380
/ mS
cm
-1
0
10
20
30
40DES1DES2DES3DES4DES5DES6DES7
(a)
1000/T( K-1)
2.6 2. 8 3. 0 3. 2 3.4
ln (m
S c
m-1
)
-1
0
1
2
3
4
DES1DES2DES3DES4DES5DES6DES7
(b)
ig. 3. (a) Dependence of specific conductivity (�) on temperature for the DESs; (b)rrhenius plot of specific conductivity (�) for the ammonium based DESs.
log( P-1)
-3 -2 -1 0 1 2 3
log
-3
-2
-1
Non-ionic Liquid s
DES5DES6DES7
Fig. 4. Walden plot for the various DESs applied in this work (the dotted line indi-cates the data for a dilute aqueous KCl solution to fix the position of the “ideal”Walden line).
lie on or above the ideal line which indicates that they are good ILs.To fix the position of the ideal Walden line in Fig. 4, the dotted idealline depicts the data for dilute aqueous KCl solution [40] comprisingof equal mobilities of fully dissociated ions [19].
3.1.5. pH
In this study, the pH of the synthesized DESs is measured as a
function of temperature as shown in Fig. 5. The pH of DES1is around7 and does not change much with temperature. On the other hand,
T / K
300 320 340 360
pH
0
2
4
6
8
10
12
DES1DES3DES4DES 5DES 6DES7
Fig. 5. A linearized plot of pH variation of different DESs as a function of tempera-ture.
L. Bahadori et al. / Electrochimica Acta 113 (2013) 205– 211 209
Cur
rent
den
sity
/ m
A cm
-2
-1. 5 -1.0 -0.5 0. 0 0.5 1.0 1.5 2.0
Cur
rent
den
sity
/ m
A c
m-2
Poten tial ag ainst A g/AgQRE / V
Fe
ttotn
3
tmoeb0FaapAapcit
Table 4Electrochemical potential windows for the DESs obtained by means of GC and Pt.electrodes.
DESs Electrodematerial
Anodiclimit (V)
Cathodiclimit (V)
Potentialwindows
DES1 GC 1.15 −1.49 2.64Pt 1.11 −0.51 1.62
DES2 GC 1.11 −0.99 2.10Pt 1.14 −0.29 1.43
DES3 GC 0.61 −1.91 2.52Pt 0.65 −1.11 1.76
DES4 GC 0.93 −1.09 2.02Pt 1.31 −0.55 1.86
DES5 GC 1.31 −2.01 3.32Pt 1.13 −0.81 1.94
DES6 GC 1.31 −1.51 2.82Pt 1.21 −0.39 1.60
DES7 GC 1.51 −0.91 2.42
ig. 6. (a) Electrochemical potential windows of all seven DESs using GC workinglectrode; (b) potential windows obtained using Pt. microelectrode.
he pH of DES 2 is very low and acidic and tends to increase withemperature. It seems that the type of the HBD has a strong effectn determining the acidity and alkalinity of the DESs. The tempera-ure dependence of pH has been fitted linearly and the parameterseeded for fitting are shown in Table 2.
.2. Electrochemical potential window
The electrochemical potential window is an important indica-or of electrochemical stability, when DESs are used as electrolytic
edia for electrochemical devices. This is because the redox couplef an analyte must fall within the potential window of the DESs. Thelectrochemical windows of the DESs prepared in this work haveeen determined using cyclic voltammetry (CV) at a scan rate of.1 V s−1 (recorded at GC and Pt. microelectrode) and illustrated inig. 6(a) and (b). The magnitudes of the electrochemical windowsre defined by a cutoff current density of 0.2 mA cm−2. The limitingnodic and cathodic potentials of the DESs are listed in Table 4. Theotential window lies approximately between 1.78 and 2.90 V vs.g/AgQRE for the GC electrode and between approximately 1.18nd 1.84 V vs. Ag/AgQRE for the Pt. microelectrode. The largest
otential window is obtained using the GC working electrode. Theathodic limit for the Pt. working electrode is smaller than that ofts GC counterpart. The anodic limit obtained at a Pt. microelec-rode is similar to those obtained by means of the GC electrode.
Pt 1.47 −0.97 2.44
It is discovered that the tested DESs have similar potential rangesas compared to some typical ILs [16] and this is consistent withreported electrochemical windows of DESs [33]. However, someILs have wider electrochemical windows [17,41]. The screened EWsmay assist in selection of proper DESs as electrolytes and solventsin various spectroscopic and electrochemical applications. It needsto be noted that some tested DESs, e.g., DES 3 and DES 5 have arelatively high reduction limit. This allows the use of these DESs inmany important electrochemical processes like the generation of astable superoxide ion [6].
3.3. Determination of diffusion coefficients of ferrocene andcobaltocenium in DESs
For the accurate determination of the diffusion coefficients oftwo metallocene derivatives, two different electrochemical meth-ods, CV and chronoamperometry (CA), have been conducted in theDESs. The electrochemical oxidation of Fc and reduction of Cc+
occur as shown in Eqs. (5) and (6):
Fc(C5H5)2 = [Fc(C5H5)2]+ + e− (5)
[Co(C5H5)2]+ + e− = [Co(C5H5)2] (6)
Ferrocene/ferrocenium (Fc/Fc+) and cobaltocenium/cobaltocene (Cc+/Cc) are conventional reference electrodes inelectrochemistry because the redox reactions represent standardreversible one-electron transfer processes. However, the Cc+/Ccprocess provides a more useful reference scale for DESs since ithas a higher thermal stability and solubility than Fc [42]. TheCV of 10 mM Fc and Cc+ at different scan rates (0.01–1 V s−1) inseven different DESs have been measured. At each scan rate, �, themagnitudes of the ratio of the anodic to the cathodic peak current,ipa/ipc, for Fc or cathodic to anodic peak current, ipc/ipa, for Cc+ areclose to 1. The anodic and cathodic peak potentials, Epa and Epc, areindependent of � and ip increases with �1/2. These observations areconsistent with the fact that mass transfer of DESs to the electrodesurface is diffusion-controlled in the solutions. The peak-to-peakpotential separation (�Ep) in the cyclic voltammograms, obtainedat slow scan rates for Cc+/Cc converges to the ideal value of57 mV predicted for an electrochemically reversible one-electrontransfer process at 298 K. The magnitudes of �Ep for Fc/Fc+ show
discrepancy from this ideal value at higher scan rates in DESs,which is attributed to the presence of uncompensated solutionresistance.
210 L. Bahadori et al. / Electrochimica Acta 113 (2013) 205– 211
Table 5The kinetic parameters and the diffusion coefficients for ferrocene and cobaltocenium redox couples in DESs at 298 K.
DES D (cm2 s−1) �D T−1 (g cm s−2 K−1) ks (cm s−1) E1/2(V) vs. Ag/Ag+
V = cyclic voltammetry; CA = chronoamperometry; E1/2 = half wave potential as defi
The diffusion coefficient (D) of Fc and Cc+ in DESs were estimatedrom CV using the Randles–Sevcik equation (Eq. (7)) using the GClectrode.
p = 0.4463(nF)3/2(RT)−1/2AD1/2C0�1/2 (7)
here ip is the peak current (A), n is the number of electronquivalents exchanged, A is the active surface area of the workinglectrode (cm2), D is the diffusion coefficient (cm2 s−1), C is the bulkoncentration of the diffusing species (mol cm−3), � is the potentialcan rate (V s−1), F is Faraday’s constant and R is the universal gasonstant.
The experimental diffusion current was further studied usingA by means of the Pt. micro working electrode in terms of thehoup–Szabo correlation as shown in Eqs. (8)–(10) [42]:
here D was the diffusion coefficient, n was the number of elec-rons transferred, F was Faraday’s constant, rd was the radius of the
icrodisk electrode, c was the bulk concentration of the electro-ctive species and t was the time. CA was measured by meansf a potential step from 0 where no Faradaic reaction occurredo 0.7 V vs. Fc/Fc+ and from 0 to −0.9 V vs. Cc+/Cc. The nonlinearurve fitting function in the software package Origin 6.0, follow-ng the approximation outlined in Eqs. (8)–(10), was used to fithe experimental data. The software was instructed to performne hundred iterations on the data, stopping when the exper-mental data had been optimized. The diffusion coefficients forc/Fc+ and Cc+/Cc in DESs (Table 5) were found to follow the order:ES5 > DES4 > DES7 > DES2 > DES6 > DES1 > DES3.
.4. Stokes–Einstein behavior of metallocene derivatives
The mass transport properties of Fc and Cc+ in DESs can be pre-icted on the basis of the Stokes-Einstein expression (Eq. (11)),hich predicts the linear dependence of diffusion coefficient of the
pecies, D on the reciprocal of the dynamic viscosity of the medium,. kB is the Boltzmann constant, ̨ is the hydrodynamic radius of the
pecies and T is the absolute temperature.
D�
T= kB
6a(11)
2.75 × 10−11 1.59 × 10−4 −0.856
arlier [33].
Table 2 summarizes the calculated Stokes–Einstein products, D�/T(� from Table 1, D from CA method, T = 298 K) of Fc and Cc+ in DESs.D�/T of Fc in DESs is larger than that of Cc+ in the same solvent. SinceD�/T of metallocene couples is dependent on the hydrodynamicradii of the diffusing species, their comparable values reveal thatthe radius of the diffusing entity is independent of the DESs. Con-sequently, one can estimate the viscosity of the new DES from theexperimentally determined D as well as from D�/T values of Fc andCc+. D of Fc and Cc+ in the seven DESs (Table 5) decrease as the vis-cosity of the solvent increases, as expected from the slower rate ofmass transport to the electrode in the more viscous medium.
3.5. Heterogeneous rate constant for the oxidation of Fc andreduction of Cc+
The apparent heterogeneous rate constant (k0) was deter-mined using the classical method of Nicholson [43]. Anodic andcathodic peak separations from a background were subtracted froma voltammogram for a simple one electron transfer reaction andfurther used to determine from which k0 was estimated (Eq.(12)):
= k0
(aD)1/2(12)
where a = nF�/RT, D was the diffusion coefficient, � was the scan rateand all other symbols had their usual meanings. For this experimentthe data was acquired at 25 ◦C, C = 5 mM and � = 0.1 V s−1. Once Dwas determined, k0 was estimated by performing CV at variousscan rates and fitting the observed variation in peak separation totabulated values. k0 of Fc and Cc+ in all DESs were summarized inTable 5.
The k0 calculated for Fc was between 5.36 × 10−5 and5.87 × 10−4 cm s−1. Other work in the literature [44–46] showedthat k0 for Fc/Fc+ was lower in DESs as compared to those in ILs andcommon organic electrolyte solutions [44]. It was reported that thek0could be correlated to the solution’s viscosity. The k0calculatedfor Cc+ was between 4.99 × 10−5 and 4.11 × 10−4 cm s−1. As notedby Tsierkezos [47], studies of the heterogeneous rate constant, k0,for the Cc/Cc+ reaction in organic solvents produced values about100× higher than those determined in DESs.
4. Conclusions
Several investigations on the physicochemical and electrochem-ical properties of seven DESs are presented in this study. Thecathodic and anodic limits for GC and Pt. electrodes in the DESs
imica
aDoctcSl
A
IfSNL
R
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
[
L. Bahadori et al. / Electroch
re also determined and DES 5 (representing an amide basedES) shows a wider electrochemical potential window than allther counterparts investigated here. D and k0 of ferrocene andobaltocenium in the DESs are found to be an order of magni-ude lower than those determined in ionic liquids under similaronditions reported in the literature. The values of the calculatedtokes–Einstein product (�D T−1) of Fc and Cc+ in these DESs areess than those determined in other ILs.
cknowledgements
The authors are grateful for financial support from the Highmpact Research Grant (UM.C/HIR/MOHE/ENG/18) in Malaysia androm the National Plan for Science, Technology and Innovation,audi Arabia (10-ENV1315-02). MHC shows gratitude to Professor.P. Brandon for allowing the use of full facilities in Imperial Collegeondon.
eferences
[1] C. Zhao, G. Burrell, A.A.J. Torriero, F. Separovic, N.F. Dunlop, D.R. MacFarlane,A.M. Bond, Electrochemistry of room temperature protic ionic liquids, The Jour-nal of Physical Chemistry B 112 (2008) 6923.
[2] L. Waligora, A. Lewandowski, G. Gritzner, Electrochemical studies of fourorganometallic redox couples as possible reference redox systems in 1-ethyl-3-methylimidazolium tetrafluoroborate, Electrochimica Acta 54 (2009) 1414.
[4] M. Galinski, A. Lewandowski, I. Stepniak, Ionic liquids as electrolytes, Elec-trochimica Acta 51 (2006) 5567.
[5] M.H. Chakrabarti, R.A.W. Dryfe, E.P.L. Roberts, Organic electrolytes for redoxflow batteries, Journal of the Chemical Society of Pakistan 29 (2007) 294.
[6] M. Hayyan, F.S. Mjalli, I.M. AlNashef, M.A. Hashim, Chemical and Electro-chemical Generation of Superoxide Ion in 1-butyl-1-methylpyrrolidiniumbis(trifluoromethylsulfonyl) imide, International Journal of ElectrochemicalScience 7 (2012) 8116.
[7] A.P. Doherty, L. Diaconu, E. Marley, P.L. Spedding, R. Barhdadi, M. Troupel, Appli-cation of clean technologies using electrochemistry in ionic liquids, Asia-PacificJournal of Chemical Engineering 7 (2011) 14.
[8] C. Fu, L. Aldous, E.J.F. Dickinson, N.S.A. Manan, R.G. Compton, Volatilisation ofsubstituted ferrocene compounds of different sizes from room temperatureionic liquids: a kinetic and mechanistic study, New Journal of Chemistry 36(2012) 774.
[9] M.J.A. Shiddiky, A.A.J. Torriero, Application of ionic liquids in electrochemicalsensing systems, Biosensors and Bioelectronics 26 (2011) 1775.
10] B. Li, L. Wang, B. Kang, P. Wang, Y. Qiu, Review of recent progress in solid-statedye-sensitized solar cells, Solar Energy Materials and Solar Cells 90 (2006) 549.
11] T. Sato, G. Masuda, K. Takagi, Electrochemical properties of novel ionic liquidsfor electric double layer capacitor applications, Electrochimica Acta 49 (2004)3603.
13] S. Menne, J. Pires, M. Anouti, A. Balducci, Protic ionic liquids as electrolytes forlithium-ion batteries, Electrochemistry Communications 31 (2013) 39.
14] S. Pandey, Analytical applications of room-temperature ionic liquids: a reviewof recent efforts, Analytica Chimica Acta 556 (2006) 38.
15] R. Barhdadi, C. Courtinard, J.Y. Nedelec, M. Troupel, Room-temperature ionicliquids as new solvents for organic electrosynthesis. The first examples ofdirect or nickel-catalysed electroreductive coupling involving organic halides,Chemical Communications 12 (2003) 1434.
16] X. Lu, G. Burrell, F. Separovic, C. Zhao, Electrochemistry of room temperatureprotic ionic liquids: a critical assessment for use as electrolytes in electrochem-ical applications, The Journal of Physical Chemistry B116 (2012) 9160.
17] T.-Y. Wu, S.-G. Su, H.P. Wang, Y.-C. Lin, S.-T. Gung, M.-W. Lin, I.W. Sun, Elec-trochemical studies and self diffusion coefficients in cyclic ammonium basedionic liquids with allyl substituents, Electrochimica Acta 56 (2011) 3209.
18] L.E. Barrosse-Antle, A.M. Bond, R.G. Compton, A.M. O’Mahony, E.I. Rogers, D.S.Silvester, Voltammetry in room temperature ionic liquids: comparisons andcontrasts with conventional electrochemical solvents, Chemistry: An AsianJournal 5 (2010) 202.
trochemical studies and self-diffusion coefficients in cyclic ammonium basedionic liquids with allyl substituents, Electrochimica Acta 56 (2010) 853.
20] A.P. Abbott, G. Capper, D.L. Davies, R.K. Rasheed, V. Tambyrajah, Novel sol-vent properties of choline chloride/urea mixtures, Chemical Communications1 (2003) 70.
[
Acta 113 (2013) 205– 211 211
21] Q. Zhang, K.D.O. Vigier, S. Royer, F. Jerome, Deep eutectic solvents: syntheses,properties and applications, Chemical Society Reviews 41 (2012) 7108.
22] D. Carriazo, M.C. Serrano, M.C. Gutierrez, M.L. Ferrer, F. del Monte, Deep-eutectic solvents playing multiple roles in the synthesis of polymers and relatedmaterials, Chemical Society Reviews 41 (2012) 4996.
23] M.A. Kareem, F.S. Mjalli, M.A. Hashim, I.M. AlNashef, Phosphonium-based ionicliquids analogues and their physical properties, Journal of Chemical & Engi-neering Data 55 (2010) 4632.
24] M.H. Chakrabarti, N.P. Brandon, M.A. Hashim, F.S. Mjalli, I.M. AlNashef, L.Bahadori, N.S. Abdul Manan, M.A. Hussain, V. Yufit, Cyclic voltammetry of iron(III) acetylacetonate in quaternary ammonium and phosphonium based deepeutectic solvents, International Journal of Electrochemical Science 8 (2013)9652.
25] A.P. Abbott, D. Boothby, G. Capper, D.L. Davies, R.K. Rasheed, Deep eutec-tic solvents formed between choline chloride and carboxylic acids: versatilealternatives to ionic liquids, Journal of the American Chemical Society 126(2004) 9142.
26] S.B. Phadtare, G.S. Shankarling, Halogenation reactions in biodegradable sol-vent: efficient bromination of substituted 1-aminoanthra-9,10-quinone indeep eutectic solvent (choline chloride: urea), Green Chemistry 12 (2010) 458.
27] E.I. Rogers, D.S. Silvester, D.L. Poole, L. Aldous, C. Hardacre, R.G. Comp-ton, Voltammetric characterization of the Ferrocene|Ferrocenium andCobaltocenium|Cobaltocene redox couples in RTILs, The Journal of PhysicalChemistry C112 (2008) 2729.
28] S. Eisele, M. Schwarz, B. Speiser, C. Tittel, Diffusion coefficient of ferrocene in1-butyl-3-methylimidazolium tetrafluoroborate – concentration dependenceand solvent purity, Electrochimica Acta 51 (2006) 5304.
29] V.M. Hultgren, A.W.A. Mariotti, A.M. Bond, A.G. Wedd, Reference potential cal-ibration and voltammetry at macrodisk electrodes of metallocene derivativesin the ionic liquid [bmim][PF6], Analytical Chemistry 74 (2002) 3151.
30] A. Lewandowski, L. Waligora, M. Galinski, Electrochemical behaviour of cobal-tocene in ionic liquids, Journal of Solution Chemistry 42 (2013) 251.
31] J. Zhang, A.M. Bond, Conditions required to achieve the apparent equivalence ofadhered solid- and solution-phase voltammetry for ferrocene and other redox-active solids in ionic liquids, Analytical Chemistry 75 (2003) 2694.
32] A.A.J. Torriero, J. Sunarso, P.C. Howlett, Critical evaluation of reference systemsfor voltammetric measurements in ionic liquids, Electrochimica Acta 82 (2012)60.
33] L. Bahadori, N.S. Abdul Manan, M.H. Chakrabarti, M.A. Hashim, F.S. Mjalli, I.M.AlNashef, M.A. Hussain, C.T.J. Low, The electrochemical behaviour of ferrocenein deep eutectic solvents based on quaternary ammonium and phosphoniumsalts, Physical Chemistry Chemical Physics 15 (2013) 1707.
34] P. De Vreese, K. Haerens, E. Matthijs, K. Binnemans, Redox reference systemsin ionic liquids, Electrochimica Acta 76 (2012) 242–248.
35] T.-Y. Wu, S.-G. Su, H.P. Wang, I.W. Sun, Glycine-based ionic liquids as potentialelectrolyte for electrochemical studies of organometallic and organic redoxcouples, Electrochemistry Communications 13 (2011) 237.
36] J. Zhang, T. Wu, S. Chen, P. Feng, X. Bu, Versatile structure-directing roles of deepeutectic solvents and their implication in generation of porosity and open metalsites for gas storage, Angewandte Chemie International Edition 48 (2009) 3486.
37] P. Bonhote, A. Dias, N. Papageorgiou, K. Kalyanasundaran, M. Gratzel,Hydrophobic, highly conductive ambient-temperature molten salts, InorganicChemistry 35 (1996) 1168.
38] D.R. MacFarlane, M. Forsyth, E.I. Izgorodina, A.P. Abbott, G. Annata, K. Frasera,On the concept of ionicity in ionic liquids, Physical Chemistry Chemical Physics11 (2009) 4962.
39] C.A. Angell, N. Byrne, J.-P. Belieres, Parallel developments in aprotic and pro-tic ionic liquids: physical chemistry and applications, Accounts of ChemicalResearch 40 (2007) 1228.
40] M. Yoshizawa, W. Xu, C.A. Angell, Ionic liquids by proton transfer: vapor pres-sure, conductivity, and the relevance of �pka from aqueous solutions, Journalof the American Chemical Society 125 (2003) 15411.
41] P.A.Z. Suarez, Vn.M. Selbach, J.E.L. Dullius, S. Einloft, C.M.S. Piatnicki, D.S.Azambuja, R.F. de Souza, J. Dupont, Enlarged electrochemical window indialkyl-imidazoliumcation based room-temperature air and water-stablemolten salts, Electrochimica Acta 42 (1997) 2533.
42] D. Shoup, A. Szabo, Chronoamperometric current at finite disk electrodes, Jour-nal of Electroanalytical Chemistry and Interfacial Electrochemistry 140 (1982)237.
43] R.S. Nicholson, Theory and application of cyclic voltammetry for measurementof electrode reaction kinetics, Analytical Chemistry 37 (1965) 1351.
44] D.Y. Kim, J.C. Yang, H.W. Kim, G.M. Swain, Heterogeneous electron-transfer rateconstants for ferrocene and ferrocene carboxylic acid at boron-doped diamondelectrodes in a room temperature ionic liquid, ElectrochimicaActa 94 (2013)49.
45] M. Matsumiya, M. Terazono, K. Tokuraku, Temperature dependence of kineticsand diffusion coefficients for ferrocene/ferricenium in ammonium-imide ionicliquids, Electrochimica Acta 51 (2006) 1178.
46] Y. Pan, W.E. Cleland Jr., C.L. Hussey, Heterogeneous electron transfer kinetics
and diffusion of ferrocene/ferrocenium in bis(trifluoromethylsulfonyl) imide-based ionic liquids, Journal of The Electrochemical Society 159 (2012) F125.
47] N.G. Tsierkezos, Electron transfer kinetics for the cobaltocene (+1/0) couple atplatinum disk electrode in acetonitrile/dichloromethane binary solvent sys-tem, Journal of Molecular Liquids 138 (2008) 1.
