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C: Surfaces, Interfaces, Porous Materials, and Catalysis

Electron Transfer Studies of Model Redox Active Species (Cationic, Anionic, Neutral) in Deep Eutectic Solvents (DES) Anu Renjith, and Lakshminarayanan Vedagiri J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b07749 • Publication Date (Web): 12 Oct 2018 Downloaded from http://pubs.acs.org on October 12, 2018

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The Journal of Physical Chemistry

Electron Transfer Studies Of Model Redox Active Species (Cationic, Anionic, Neutral) in Deep Eutectic Solvents (DES) Anu Renjith a and V. Lakshminarayanan *b a Indian Institute of Science, Bangalore b Raman Research Institute, Bangalore.

ACS Paragon Plus Environment

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ABSTRACT The redox potentials of electroactive species are significantly influenced by the solvation characteristics of the medium. This is manifested in the shift of half peak potentials with the change in the solvent medium. There have been many approaches till date, both experimental and theoretical to understand the role of molecular solvents in the peak potentials of redox species. The electrochemical studies reported here, are aimed at understanding the effect of Deep eutectic solvents (DES) which is distinct from conventional solvents in terms of highly concentrated ionic composition, on the half peak potentials of some standard redox reactions. The redox species selected for this study are distinct either in terms of their charge [Fe(CN)64-/3-, Ru(NH3)62+/3+, ferrocene methanol, FcMeOH0/+] or their hydrophilic / hydrophobic properties [methyl viologen, ferrocene]. The redox potentials are compared with the values obtained in the aqueous medium which is very well characterized in terms of solvent reorganization energy and free energy changes. The cyclic voltammetric behavior of the redox species is significantly different from that of aqueous medium. The diffusion coefficients of the redox species in DES measured by EIS and cyclic voltammetry showed significant deviations from that predicted by Stokes-Einstein equation, indicating the dominant effect of coulombic interactions within the components of DES.


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1. INTRODUCTION

Ionic liquids and deep eutectic solvents (DES) are recently being used as novel electrolytic media owing to their inherent ionic conductivity, wide electrochemical potential window and excellent solubilizing properties

1,2

. DES prepared from a homogeneous mixture of quaternary ammonium

salt and a hydrogen bond donor, has already received wide acceptance over ionic liquids in metal finishing/plating, nanoparticle synthesis, organic reactions, biocatalysis and enzymatic studies owing to their low cost and bulk production of components 3-6. Despite considerable progress in the applications of DES, there is still insufficient understanding of the electron transfer and transport properties of redox species in this medium. It is to be noted that the electrochemical behavior of DES as a solvent medium differs considerably from the conventional molecular solvents owing to their unique structure, high viscosity and ionic composition. While the viscosity affects the mass transport properties of the medium, the ionic composition can influence the redox potentials of electroactive species. A detailed study on the solvation effects on electron transfer in DES, therefore will be advantageous for applications involving electron transfer processes. In our work, the redox potentials and diffusion characteristics of a set of standard redox species in DES medium are studied by cyclic voltammetry and electrochemical impedance spectroscopy and their redox behaviour in DES wrt aqueous medium are demonstrated based on plausible solvation models. Among the few electrochemical studies carried out in DES, several of them were focused on understanding the process of diffusion of analyte species in the medium 7-9. The viscosity of DES varies considerably with the size and hydrogen bonding nature of its components which inturn significantly affects the diffusional properties of redox species in the medium. The diffusion coefficients of redox species in ionic liquids which is similar to DES in terms of solvent properties are typically in the range of 10-8 - 10-7 cm2/s, which is two to three orders lesser than that of 10-4-10-


3


5

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cm2/s in molecular solvents 9. The conventional theories on the diffusion of redox species in an

electrolytic medium are proposed for solvation in molecular solvents. Hence in the case of DES and ionic liquids, there is a significant deviation of the measured values of diffusion coefficients from that predicted theoretically 7,9. The observed deviations are due to the fact that the diffusing species are in concentrated ionic solutions, a situation vastly different from that of molecular solvents. For example, a molecular solvent usually contains supporting electrolyte concentrations of 0.1-1M while the ionic components in DES correspond to a concentration of about 5-9M. The high concentration makes the ion-ion interactions in these solutions quite significant which therefore cannot fit into a model of non-interacting solvents. Abbott et al. used the ‘hole theory’ to explain the viscosity and diffusion of ions in DES, according to which the ionic transport occurs only when the size of the diffusing ions matches the holes within the solvent structure

10

. Till date, only the

diffusion coefficient values of ferrocene methanol and Cu+/2+ in DES have been reported 7,8. The solvation mechanism of redox species in DES and ionic liquids are quite different from that normally observed in molecular solvents. The knowledge of the solvation mechanism in DES is essential to understand the redox behavior of electroactive compounds in the electrolytic medium. The Marcus theory on electron transfer processes in molecular solvents is based on the reorganization of solvent dipoles. However in the case of ionic liquids/DES, the solvent species is constituted mainly of anions and cations. Therefore, the solvent reorganization in ionic liquids occurs by ionic translation unlike the solvent dipole reorientation in molecular solvents

11

.

Therefore in the case of ionic liquids, the self-diffusion of ions/molecules plays a more significant role in the reorganization process. The differences in the solvent reorganization mechanism in DES are expected to cause drastic shifts in the redox potentials of electroactive species when compared


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to that observed in molecular solvents. A model of the solvation structure of a charged species in water, and in ionic liquids is depicted in scheme-1.

Scheme 1. Solvation structure of a charged species in (a) water (b) ionic liquids The redox potential is a direct measure of the thermodynamic feasibility of an electrochemical process. The free energy change (ΔG°) associated with the redox process is related to electrode potential (ΔE°) by the equation, ΔG° = –nF (ΔE°)

(1)

Where, n and F are respectively the number of electrons and Faraday constant. The change in the free energy associated with a redox reaction is obviously affected by the solvation structure of the redox species in the solvent. The reduction and oxidation processes are accompanied by solvent reorganizations causing entropy/enthalpy changes. The shift in redox potential depends on the extent of the solvent reorganization, which in turn depends on the ability of the oxidized/reduced complex to make/break the solvation structure. The redox processes in aqueous medium involve changes the charge of the redox species which in turn causes reorganization of highly ordered solvation structure of water molecules, leading to entropy changes. As compared to aqueous medium, organic solvents and ionic solutions have much less ordered structure and hence changes in the free energy associated with redox reactions occurring in these solvents are significantly lower

12

. The optimized structure of DES proposed by Wagle based on


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quantum chemical calculations is used as solvent model in our studies 13. According to the proposed structure, DES are associated extensively by hydrogen bonding between the components and forms six and seven membered cage like structures in the case of urea based reline and ethylene glycol based ethaline respectively 13. Abbott et al. established an electrochemical series for metallic redox couples in a model deep eutectic solvent of ethaline

14

. The proposed electrochemical series predicts the preferential

stabilization of one oxidation state of a redox species over another. The variations in the electrochemical behavior of redox species in DES wrt aqueous medium were attributed to the speciation effects in the high chloride medium of DES. We report here the results of our study on the nature and origin of the shift in the redox potentials observed in DES vis-à-vis molecular solvents on the basis of differences in the solvation structures in these two distinct media. 2. EXPERIMENTAL 2.1. Materials Potassium

ferrocyanide

(Loba),

potassium

ferricyanide

(Qualigens)

[Fe(CN)64-/3-],

hexaammineruthenium (II) and (III) chloride (Alfa Aesar) [Ru(NH3)62+/3+], ferrocene (Acros Organics) [Fc], ferrocenium tetrafluoroborate (Aldrich) [FcBF4], ferrocene methanol (Aldrich) [FcMeOH], methyl viologen (Aldrich) [MV0/2+] were used as received without further purification. The DES used in this study as electrolyte medium were prepared as two component mixtures of choline chloride and either urea or ethylene glycol. Choline chloride purchased from Aldrich was dried for one hour in a hot air oven at 50 0C prior to use. The hydrogen bond donors for the preparation of DES, viz., ethylene glycol and urea (Merck) were used as received. Choline chloride


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and the hydrogen bond donor component were magnetically stirred at 80 0C in a molar ratio of 1:2 until a homogeneous colorless liquid was formed. Millipore water having a resistivity of 18 MΩ cm was used for the preparation of aqueous solutions of redox species with potassium chloride (Merck) as supporting electrolyte. The electrochemical studies of hydrophobic redox species were carried out in propylene carbonate (Spectrochem) with Tetrabutylammonium tetrafluoroborate (Alfa Aesar) as supporting electrolyte. 2.2. Electrode pretreatment A gold disc electrode of diameter 1mm (CH instruments) was polished with 1µm, 0.3 µm, and finally 0.05 µm alumina slurries, followed by thorough cleaning with water. Further, the disc electrode was cleaned in Piranha solution (3:1 conc.H2SO4 : H2O2) and washed in water. 2.3. Electrochemical studies The cyclic voltammetric studies were performed with a three electrode system using gold disc electrode as the working electrode, Platinum foil as a counter electrode and Ag/AgCl/3M NaCl as the reference electrode using model 263A EG&G potentiostat. The electrochemical cell of capacity 10 ml was maintained at 25 0C using Julabo temperature control systems (F25) for all the electrochemical studies. All the cyclic voltammetric plots shown in this work correspond to the second cycle as it was close to the stabilized behavior. The impedance measurements were carried out over a wide frequency range from 100 kHz to 100 mHz and at an AC amplitude of 10 mV using a model 5210 lock-in amplifier (Perkin-Elmer Instruments) with Power Suite software (EG&G) interfaced with a PC. The ac signal was applied to the electrochemical cell containing redox species maintained at the half peak potential (dc) of the


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redox species measured from cyclic voltammetry. The electrochemical impedance spectroscopy data was analysed using ZSimpWin software (EG&G) software. FTIR studies were carried out using Shimadzu FTIR-8400. 3. RESULTS AND DISCUSSIONS 3.1.

Cyclic voltammetric studies

Cyclic voltammetric studies of the redox probes Fe(CN)64-/3-, Ru(NH3)62+/3+, Fc0/+, FcMeOH0/+ and methyl viologen MV0/2+ were carried out in DES media viz., ethaline and reline and also in appropriate molecular solvent viz., water or propylene carbonate for comparison. Prior to the addition of redox species, cyclic voltammetric scan in the DES medium was performed as blank experiments at different electrochemical potential ranges. The cyclic voltammetric scan of the blank exhibits the typical capacitive features confirming the absence of any redox reaction within the potential range. 3.1.1. Anionic complex [Ferrocyanide/Ferricyanide System] [Fe(CN)6]4-

Fe(CN)6]3-

The cyclic voltammogram of 5mM [Fe(CN)6]4-/3- in DES and aqueous medium displayed in the Fig. 1 corresponds to a scan rate of 20mV/sec. It can be observed from the figure that the peak currents are significantly lower in DES when compared to that in aqueous medium. Interestingly, both the anodic and cathodic peak potentials and therefore the half peak potential have shifted to more negative values in DES medium from that in aqueous electrolyte.


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6 Current (A)

40

4

20

c

0 -20 -40 -0.2

2

Current (A)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


0.0

0.2

0.4

Potential (V)

0

-2

a b

-4

-6 -0.3

-0.2

-0.1

0.0

0.1

0.2

0.3

0.4

0.5

Potential (V)

Figure 1. Cyclic voltammogram of 5 mM Fe[(CN) 6]4−/3− in (a)ethaline (b)reline (c)aqueous medium/0.1M KCl at a scan rate of 20 mV/s. (Dotted lines indicate the corresponding blank DES)

The large shift in the half peak potential of about -200mV in DES medium can be explained as follows : The redox couple [Fe(CN)6]4-/3- is highly stable in a chloride rich medium as that of DES. In fact, owing to this, Abbott et al. formulated an electrochemical series using [Fe(CN)6]4-/3- as an internal reference standard

14

. It is seen from Table 1 that the half peak

potentials of [Fe(CN)6]4-/3- in reline and ethaline differ by a mere 12 mV. However, the half peak potential (E1/2) in the aqueous medium is considerably more positive when compared to that in the DES medium by about 192mV. The positive shift of the redox potential in aqueous medium indicates that [Fe(CN)6]4- as a reduced species is stabilized more than its corresponding oxidized species. The significant shift in the peak potential of [Fe(CN)6]4-/3redox reaction can be understood from the nature of the two solvent media which differ markedly in their solvation properties. In order to understand the differences in the solvation features, separate FTIR studies of these two redox species [Fe(CN)6]4- and [Fe(CN)6]3-, were carried out in DES as well as in aqueous media. By this study, any shift in individual FTIR


9


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bands of [Fe(CN)6]4- and [Fe(CN)6]3- can be correlated to the preferential interaction of one oxidation state over the other with the solvent. Table 1. Anodic and cathodic peak potentials (EPa and EPc), peak separation (ΔEp), halfpeak potentials (E1/2) and peak currents (ipa and ipc) for the 5 mM [Fe(CN)6]4−/3− redox reaction

EPa mV

EPc mV

ΔEp m V

E1/2 mV

iap µA

-icp µA

18.4

-40.9

59.3

-11

2.76

2.72

Reline

33

-32

65

1

1.02

1.01

Aqueous

222

155

67

189

37.9

35.2

Solvent medium Ethaline

Earlier Yu et al. compared the solvation dynamics of cyanoferrate anions ([Fe(CN)6]4- and [Fe(CN)6]3-) in aqueous medium by ultrafast infrared spectroscopy 15. The results obtained from the frequency-frequency correlation function, extracted from the 2D IR spectra indicate that the hydration shell interacts more efficiently with the cyano group of ferro- species than with that of ferri-species and hence is more tightly bound to the -C≡N. The C≡N stretching frequencies of [Fe(CN)6]4- and [Fe(CN)6]3- in the linear IR spectra are at 2039 cm-1 and 2115 cm-1 respectively. The relative shift in the C≡N stretching band in these two oxidation states was attributed to the stronger hydrogen bonding interactions of C≡N……HOH in the [Fe(CN)6]4- species. The ionic hydrogen bond length of N

……

H was found to be shorter in the ferro species than in the ferri

species. The hydrogen bonding stabilization energy of [Fe(CN)6]4- and [Fe(CN)6]3- in water was reported to be -10.4 and -6.9 kcal/mol respectively 15. The higher hydrogen bonding stabilization of


10

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[Fe(CN)6]4- over [Fe(CN)6]3- also implies that the reduced form is preferred in aqueous medium. The electrochemical oxidation of the [Fe(CN)6]4- therefore requires more positive potential as can be seen from the E1/2 values shown in Table 1 for the redox reaction in aqueous as well as DES media.

(a) in H2O in DES

(b) in H2O

in DES

Figure 2. A comparison of FTIR spectra of (a) [Fe(CN) 6]4- (b) [Fe(CN)6]3- obtained in water and DES In order to compare the solvation features of [Fe(CN)6]4-/3- in DES and aqueous medium under similar conditions, we have carried out FTIR studies of the same in both the media. Reline medium prepared from choline chloride and urea was used as model DES system for FTIR studies. The linear IR spectra carried out in DES shows C≡N stretching frequencies of [Fe(CN)6]4- and [Fe(CN)6]3- to be very close, namely at 2152 cm-1 and 2156 cm-1 respectively (Fig. 2). This indicates that the hydrogen bonding interactions of the ferro-species and the ferri-species are not


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significantly different in DES unlike in aqueous medium where ν C≡N is observed at 2038 and 2116 cm-1 respectively. The solvation studies of a thiocyanate ion (-SCN) by FTIR and 2D IR in different DES (reline, ethaline and glyceline) carried out by Cui et al. have also confirmed that there are no hydrogen bonding interactions between the –CN group and components of various DES systems 16. Our observations from the FTIR studies carried out in DES and aqueous medium, clearly provide support to the shift in the electrochemical redox potentials measured by cyclic voltammetry. In the absence of preferential solvation of either of the two oxidation states, the electrochemical oxidation of [Fe(CN)6]4- occurs at relatively lower potential in DES than that observed in the aqueous medium where the observed half peak potential of [Fe(CN)6]4-/3 is ~190 mV lower in DES compared to that in aqueous medium. A pictorial representation of the probable solvation of the [Fe(CN)6]4-/3- in aqueous medium and DES is depicted in scheme-2. An organized solvation shell is shown in this scheme in the aqueous medium which is more tightly bound to [Fe(CN)6]4- when compared to [Fe(CN)6]3-. This is in contrast to that in DES, where the hydrogen bonded cage like structures are completely scattered around the redox species giving rise to a disorganized solvation pattern around both [Fe(CN)6]4- and [Fe(CN)6]3-. Previously, X-ray spectroscopic study of [Fe(CN)6]4-/3- in ethylene glycol and water by Penfold etal have shown that ethylene glycol does not provide any organized solvent structure around either [Fe(CN)6]4- nor [Fe(CN)6]3- owing to its high viscosity and slightly weaker interaction with the solute 17. Considering the high viscosity of DES compared to EG and water and also the weak interaction of DES-[Fe(CN)6]4-/3- as supported by FTIR data, the proposed model of a disorganized solvent shell of DES for [Fe(CN) 6]4- and [Fe(CN)6]3can be justified. This is also supported by a recent study on solvation dynamics of an ionic


12

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probe in DES by molecular dynamics simulations and ab initio calculations, where it is shown that DES acts as a disorganized solvent medium 16.

i(a)

i(b)

(ii)

Scheme 2. Proposed solvation structure of (a)[Fe(CN)6]4- (b)[Fe(CN)6]3- in (i) aqueous medium (ii) Fe(CN)6]4-/3- in reline (model not to scale) 3.1.2. Cationic complex [Hexaammineruthenium (II / III) system] [Ru (NH3)6]2+

[Ru (NH3)6]3+

Fig.3 shows the cyclic voltammograms of a positively charged complex [Ru (NH3)6]2+/3+ in ethaline and reline while that in aqueous medium is provided in the inset. It can be observed


13


that the change in the cyclic voltammetric features in DES and aqueous media follow the similar trend as observed for [Fe(CN)6]3-/4- system. The peak potentials shift towards negative by more than 100 mV in DES wrt aqueous medium. It may be noted that the half peak potentials for the redox process do not vary significantly in reline or ethaline (Table 2) while the peak current in reline is considerably lower compared to that in ethaline. 20

Current ()

2

c

0

-20

-40

-0.3

-0.2

-0.1

0.0

Potential (V)

1

Current ()

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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0

a b

-1

-0.5

-0.4

-0.3

-0.2

-0.1

0.0

Potential (V)

Figure 3. Cyclic voltammogram of 5 mM [Ru (NH3)6]2+/3+ in a)ethaline (b)reline (c) aqueous medium/0.1M KCl at a scan rate of 20 mV/s. (Dotted lines indicate the corresponding blank DES)

Table 2. Anodic and cathodic peak potentials (EPa and EPc ), peak separation ( ΔEp ), half peak potential (E1/2) and peak currents (ia p and ic p) of [Ru (NH3)6]2+/3+

EPa mV

EPC mV

ΔEp m V

E1/2 mV

iap µA

-icp µA

Ethaline

-262

-323

61

-292

0.74

0.87

Reline

-282

-344

62

-313

0.42

0.47

Aqueous

-149

-217

68

-183

20.7

28.7

Solvent medium


14

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The solvent reorganization energy λ plays a crucial role in a redox reaction as it is a measure of the structural changes that the redox species undergoes during the reaction. The reorganization energy is contributed by an inner sphere component, λi, and an outer sphere component, λo. While the inner sphere component is associated with the changing bond lengths and geometries within a molecule, the outer sphere component depend on the alignment of the dipole moment of a polar solvent as is the case in the [Ru(NH3)6]2+/3+ redox reaction. Yee et al. reported that in the case of [Ru(NH3)6]3+, the weakly acidic ammine protons form hydrogen bonds with the surrounding water molecules which brings about solvent ordering, with the entropy change associated with the process during the oxidation of [Ru(NH 3)6]2+ being 18 cal/ K/mole

18

. The solvent reorganization energy of ionic liquids is significantly

different than that of molecular solvents as reported in earlier works

11

. The large enthalpy

of formation of DES of about -45kcal/mol indicates that the weakly acidic nature of ammine protons of [Ru(NH3)6]2+ cannot interact strongly with the DES components so as to disrupt the hydrogen bond stabilized caged complexes of DES during oxidation

13

. Hence the caged

structure of the hydrogen bonded anionic complex and also the bulky choline cation is most likely to be preserved and cannot form an organized solvation shell around [Ru(NH 3)6]2+/3+. This is in contrast to the oriented molecular dipoles observed in molecular solvents for either oxidation states of Ru [scheme-3]. It can be noted from Table 2 that the electrochemical oxidation of [Ru(NH3)6]2+ in DES occurs at a more negative potential (-120 mV) compared to that in aqueous medium an observation which is quite similar to that of [Fe(CN)64-/3-]. The lowering of half peak potential can be explained by the fact that while oxidation/reduction causes substantial


15


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changes in the bulk organization of water molecule around the redox species, it hardly has any effect on the solvation structure of DES. The solvation patterns are almost similar in ethaline and reline with a very small difference (~ 20 mV) in the half wave potentials of [Ru(NH3)6]2+/3+, which again could be attributed to the more disorganized solvation shell of ethaline than reline 16. i(a)

i(b)

Scheme 3. A pictorial representation of the water solvation structure of (i) (a) [Ru(NH3)6]2+ and (b) [Ru(NH3)6]3+ in aqueous medium (model not to scale) 3.1.3. Neutral redox active species (a)

Ferrocene compounds Fc0

Fc+

Ferrocene (Fc), a neutral hydrophobic molecule undergoes a single electron oxidation of its central Fe in +2 oxidation state to ferrocenium ion in +3 state which is polar and hydrophilic. The cyclic voltammetric studies of certain water soluble ferrocene derivatives such as ferrocene methanol, FcMeOH and ferrocenium tetraflouroborate, FcBF4 are studied here as model neutral redox species in DES (Fig.4). These compounds are chosen as they are soluble and stable in both the DES as well as in aqueous media. The trend observed in cyclic


16

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voltammograms in terms of the shifts in redox peak potential of the neutral compounds of Fc, FcBF4 and FcMeOH in both the media is in contrast to that observed in the charged redox species [Ru(NH3)62+/3+] and [Fe(CN)64-/3-] discussed earlier. For example, the half peak potential of FcMeOH was observed to be more positive in DES by ~ 130 mV compared to that in aqueous medium (Table-3). However, the peak current which is related to the viscosity characteristics (Table-1) of the medium follows the same trend in all the ferrocene compounds similar to that of the charged complexes discussed in the earlier sections. Since the Fe2+ is sandwiched amidst the bulky cyclopentadienyl ligands, solvent effects are more pronounced around the Cp ligands than at the metal ion centre. On oxidation, the neutral ferrocene converts to cationic ferrocenium and consequently the electron density around the cyclopentadienyl ligands decreases causing drastic changes in its solvation properties. The solvation free energies of Fc and Fc + were computed in five polar solvents namely water, acetonitrile, methanol, acetone, and dimethylsulfoxide by Kuznetsov et al.

19

.

Due to the polar nature of ferrocenium species, its solvation in polar solvents was more favorable for ferrocenium (Fc+) compared to ferrocene (Fc) which is expected. For Fc+ and Fc, the solvation free energy in aqueous medium was found to be -46.1 kcal/mol and +0.1 kcal/mol respectively. The positive solvation free energy of Fc in water can be attributed to the non-polar nature of ferrocene molecule. Since the redox reaction causes a change in the nature of the molecule from non-polar to polar, the dielectric constant of the medium has a strong effect on the process 20. In the case of DES, there is a positive shift in the half peak potentials of ferrocenium tetraflouroborate (FcBF4) compared to that in aqueous medium (Table 3). This phenomenon can be attributed to the enhanced solvation of Fc by DES when compared to that in aqueous


17


electrolyte. Unlike DES molecules, which can solvate and stabilize the reduced product (Fc) which is non-polar and hydrophobic, water does not effectively solvate and stabilize the Fc species. The redox potentials of Fc in propylene carbonate (ε=64) can be seen to be shifted to more positive potential wrt to water (ε=80) due to the same reasons. A favorable stabilization of Fc drives the electrochemical oxidation of Fc to more positive potentials and hence in general, a decrease in the polarity of the solvent is accompanied by a positive shift in the E1/2.

Current ())

c

16

4

8

(ii)

Current ())

16

(i)

4

0 -8 0.4

otential (V)

P

0

2

Current ()

0.2

a b

-2

0.0

0.1

0.2

0.3

0.4

0.5

c

8 0 -8

-16

-16

2

Current ()

0.6

0.2

0

a b

-2

0.0

0.7

0.4

otential (V)

P

0.1

0.2

0.3

0.4

0.5

0.6

0.7

Potential (V)

Potential (V) 4 16

(iii)

8 Current ()

3

2

Current ()

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 18 of 34

c

0 -8 0.2

0.4

0.6

Potential (V)

1

a b

0

-1

-2 0.0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

Potential (V)

Figure 4. Cyclic voltammograms of (i)FcMeOH (ii)FcBF4 (iii)Fc in a)ethaline (b)reline (c) molecular solvent at a scan rate of 20 mV/s (Dotted lines indicate the corresponding blank DES)


18

Page 19 of 34 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


While the absolute polarity of DES has not been calculated till date, the most common approach is to describe solvent polarity of DES in Reichardt's scale where the values are reported wrt water in terms of ET N. According to the Reichardt's scale of polarity, ethaline (ChCl + EG) and reline (ChCl + Urea) are associated with ET N of 0.80 and 0.84 respectively wrt water, ET N = 1 and tetramethyl silane (ET N = 0)

21

. The shift in the redox potentials of FcBF4 in ethaline, reline and aqueous

medium can also be explained based on the polarity scales. Ethaline with lower ET N of 0.80 can solvate Fc better when compared to reline with ET N of 0.84. Table 3. Anodic and cathodic peak potentials (EPa and EPc ), peak separation ( ΔEp ), half peak potential (E1/2) and peak currents (ia and ic) of ferrocene compounds in different solvent media Redox

E Pa

EPc

ΔEp

E1/2

iap

- icp

mV

mV

mV

mV

µA

µA

Ethaline

381

318

63

350

2.77

1.78

Reline

330

260

70

295

1.35

1.13

Aqueous

261

196

65

229

8.08

5.88

Solvent

species

FcMeOH

FcBF4

Ethaline

370

300

70

335

3.19

2.66

Reline

303

242

61

273

1.03

0.94

Aqueous

306

206

100

256

15.77

18.48

Ethaline

378

302

76

336

2.38

1.54

Reline

312

246

66

276

0.4

0.29

PC (Vs Ag/Ag+)

505

418

87

484

11.2

5.40

Fc


19


Page 20 of 34

Even though there are no reports on the nature of solvation of Fc/Fc+ in a DES medium, certain implications can be derived from the solvation model of Fc in ionic liquid 1,3-dimethylimidazolium tetrafluoroborate (MMIMBF4) as reported by Yang et al 22. The study reported that the Fc+ cation formed on oxidation of Fc, experiences electrostatic repulsion with the Im+ and is stabilized only by the interaction of BF4-. However, in the case of DES, the chloride anions of DES do not interact and stabilize the oxidized Fc+ ion due to the weak ion-pair formation of Cl- - Fc+ unlike BF4- - Fc+ in the ionic liquid MMIMBF4

23

. It can therefore be concluded that the neutral Fc is stabilized better in

DES than Fc+ which makes ferrocene to undergo oxidation at higher positive potentials. The E1/2 of Fc in ethaline (ChCl + EG) is more positive by 50-70 mV than in reline (ChCl + urea) which suggests that the solvation characteristics of Fc are slightly different in these two DES media. Choline chloride being the common component in both these solvents, it is the characteristics of the HBD components (EG and urea) that needs to be considered. Urea with more polar components (-NH2 groups) cannot solvate the non-polar ferrocene while ethyl chain of EG can solvate and stabilize ferrocene better. This results in more positive E1/2 for the reaction in ethaline since the reactant ferrocene is stabilized by EG molecules. The Reichardt’s scale of polarity also supports the trend as discussed before. (b)

Methyl viologen

Methyl viologen (1,1’-Dimethyl-4,4’-bipyridinium dichloride) is an example for a water soluble organic redox active species exhibiting three oxidation states as given in scheme-4. Ionic liquids using viologens in combination with ferrocenes are promising candidates for development of durable, non-volatile electrochromic devices 24,25.


20

Page 21 of 34

Scheme 4. Redox reaction of methyl viologen

8

c

Current ()

40

6

0

-40 -1.0

4

Current ()

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-0.8

-0.6

-0.4

Potential (V)

2 0 -2

a b

-4 -6 -1.1

-1.0

-0.9

-0.8

-0.7

-0.6

-0.5

-0.4

-0.3

-0.2

Potential (V)

Figure 5. Cyclic voltammogram of 5 mM MV in a)ethaline (b)reline (c)aqueous medium/0.1M KCl at a scan rate of 20 mV/s (Dotted lines indicate the corresponding blank DES)

Table 4. Anodic and cathodic peak potentials (EPa and EPc ), peak separation ( ΔEp ), half peak potential (E1/2) and peak currents (ia and ic) of MV in different solvent media.

E2c mV

ΔEp2 m V

Ethaline -503

-558

55

-531

2.91

2.07

Reline

-536

-603

67

-570

0.262

0.227

Aqueous -645

-692

47

-668

20.37

17.76

E2a mV

E21/2 mV iap µA - icp µA


21


Page 22 of 34

Fig. 5 shows the cyclic voltammogram of methyl viologen in aqueous medium, ethaline and reline. The two redox peaks observed in the voltammogram correspond to the successive electron transfer processes shown in the above scheme. The peak separations observed in all the three solvent media studied here correspond to a typical reversible redox reaction. The half peak potentials were very close to the recently developed metal free redox active viologen based DES 26. The half peak potentials of MV have shifted considerably by +130mV in DES compared to aqueous medium a behaviour quite similar to that observed for the ferrocene compounds (Table 4). The cyclic voltammetric behaviour of MV in DES is quite similar to that observed in ferrocene compounds. The neutral MV0 with methyl and pyridine rings can have secondary interactions with the alkyl components of DES. However, on oxidation, the positively charged choline cation experiences electrostatic repulsion with the positively charged MV+ and moves away. On the other hand, in aqueous medium while the neutral methyl viologen cannot be adequately solvated, the oxidized hydrophilic species of MV+ are very well solvated by polar water molecules which facilitates the oxidation process. Therefore the positive shift in the peak potentials in DES compared to that in aqueous medium can be attributed to the preferential solvation of charged oxidized species MV+, MV2+ in water and repulsive interactions experienced by the charged species in DES. A similar observation was reported earlier for methyl viologens in lyotropic liquid crystalline phase 27. 3.2.

Scan rate dependence of the redox species

The peak currents of the redox species as measured in cyclic voltammograms are almost an order of magnitude lower in DES when compared to that in aqueous medium. Among the DES, the peak currents are almost halved in reline as compared to that in ethaline. This is attributed to the much lower viscosity of the ethaline compared to reline [ηethaline – 37 cP, ηreline – 750 cP]. It can be seen


22

Page 23 of 34

from Fig.6 and 7 for different redox reactions in reline and ethaline, the ν1/2 Vs ip plots exhibit a linear relationship following the Randles - Sevcik equation for a diffusion controlled process. This confirms that the redox reactions follow the planar semi- infinite linear diffusion in the DES media too.

12

6

(1a) Equation

9

Adj. R-Sq

ipa ipc

6

Current A

(1b)

4

3 0

F F G

-3

linear fit of ipa

-6

linear fit of ipc

G

-9

Equation

ipa

2

Current A

15

0

Adj. R-Squ

ipc

F

linear fit of ipa

G

F G

linear fit of ipc

-2

-4

-12 -15

-6

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.1

0.2

0.3

 5 4

0.4

0.5

0.6

0.7

0.8

 3

Equation

(2a)

y = a + b*x

Adj. R-Square B

3

ip

1

ip

0

0.99403 Intercept

B

Slope

C

Intercept

C

Slope

(2b)

2

a

1

Current 

2

Current 

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


c

-1

linear fit of ip

-2

linear fit of ip

a c

ip ip

0

a c Equation

linear fit of ip -1

linear fit of ip

Adj. R-Square

a

B B C

c

C

-3

-2

-4 -5

-3 0.1

0.2

0.3

0.4

0.5



0.6

0.7

0.8

0.1

0.2

0.3



0.4

0.5

0.6

0.7

0.8

 

Figure 6. Plot of average peak current vs. square root of scan rate (1) [Fe(CN)6]4−/3− (2) [Ru(NH3)6]2+/3+ in (a) ethaline (b) reline 3.3. Calculation of diffusion coefficients of the redox species 3.3.1. Calculation of diffusion coefficients from EIS Electrochemical impedance spectroscopy (EIS) studies of the redox probes [Fe(CN)6]4−/3−, Ru(NH3)6]2+/3+, Fc0/+, FcMeOH and MV0/+2 were carried out in the DES media as well as in aqueous medium. The EIS studies were carried out at the half-peak potentials of the redox probes obtained from their corresponding cyclic voltammograms and these results were analysed by fitting to a


23


Randles equivalent circuit consisting of an active electrolyte resistance Ru in series with the doublelayer capacitance Cdl which is in parallel with the charge transfer resistance Rct with Warburg impedance Zw. Fig. 8 shows both the experimental and fitted EIS data of all redox species. It can be seen from the Table 5 that the solution resistance Ru is higher in the DES media compared to the aqueous media for all redox systems studied. The Warburg admittance which is a measure of the diffusion of the redox species in a medium is considerably lower by almost an order of magnitude 20

6

(1a)

15

(1b)

4 Equation

ip

5

ip 0

c

linear fit of ip

-5

ip

2

a

Current 

Current 

10

linear fit of ip

a

ip

0

B B C

c

C

linear fit of ip

Equation

-2

c

Adj. R-Square

a

Adj. R-Square

linear fit of ip

B

a c

B C C

-10

-4

-15 0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.1



0.2

0.3

15

0.4

0.5

0.6

0.7

0.8



fcoh et

fcoh re 8

(2a)

(2b)

6 10

ipa ipc

5

0

Current A

Current A

4

linear fit of ipa

ipa ipc

2 0

linear fit of ipa

-2

linear fit of ipc

-5

linear fit of ipc

-4 -10

-6 0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.1

0.2

0.3

2.0

(3a)

0.5

0.6

0.7

0.8

(3b)

1.5

10

ip 5

ip

1.0

a c

linear fit of ip

0

Current 

15

0.4

 

 

Current 

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 34

linear fit of ip

a c

0.5

Equation Adj. R-Square

Equation

B B C C

ip ip

ia ia

c

ic ic

linear fit of ip

0.0

linear fit of ip

-0.5

-5

Adj. R-Square

a

a c

-1.0

-10 0.1

0.2

0.3

0.4



0.5

0.6

0.7

0.8

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8





Figure 7. Plot of average peak current vs. square root of scan rate (1) FcBF4 (2) FcMeOH (3) MV in (a) ethaline (b) reline


24

Page 25 of 34 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Figure 8. EIS studies of the redox systems (1) [Fe(CN)6]4-/3- (2) [Ru(NH3)5] 2+/3+(3) MV (4) FcBF4 (5) FcMeOH in (a) ethaline (b) reline. Red spot denotes experimental data and green circle denotes fitted data. Randles circuit used for equivalent circuit fitting of EIS data is also provided.


25


Page 26 of 34

in DES when compared to that in aqueous medium. The high solution resistance is attributed to lower ionic conductivity of DES while lower Warburg admittance is due to the higher viscosity of DES. The diffusion coefficient values of the redox species have been calculated from the Warburg impedance using the following equations 28, (2) Where,

(3)

where ZW is the Warburg impedance, ω is the angular frequency applied, n is the number of electrons involved in the redox reaction, F is the Faraday constant, R is the universal gas constant, T is the temperature (in K), D is the diffusion coefficient of the redox species, A is the area of the electrode, and CO and CR are the bulk concentrations of the oxidised and reduced species in the solution, respectively. The effect of viscosity on the diffusion coefficients is clear by comparing the √D values of redox species obtained in ethaline (η – 37 cP) and reline (η – 750 cP) given in the table. The diffusion coefficient D of a compound is inversely related to the viscosity (η) of the solvent by the Stokes Einstein equation given by, (4) Earlier Nkuku et al. observed that the diffusion coefficients of ferrocene (2.7x10-8 cm2s-1) measured from electrochemical experiments were found to deviate from the values predicted from the equation (6x10-8 cm2s-1) 7. This is because of the fact that the equation (3) assumes a non-interacting solvent which of course is not the case in DES. The coulombic interactions among the ions constituting DES / ionic liquids and the redox species cannot be neglected. The


26

Page 27 of 34 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


diffusion coefficients of ferrocene obtained in our studies agree well with the literature values of the species reported in DES (2 - 5 × 10−8 cm2 s−1) and ionic liquids of comparable viscosities 7,29,30

.

Table 5. Values of the solution resistance (Ru), Warburg admittance (Y0) and diffusion coefficients of various redox species from equivalent circuit fit of the impedance data

System

[Fe(CN)6]4-/3-

[Ru(NH3)6] 2+/3+

FcMeOH

FcBF4

MV

Y0 x10-5 (S.√sec) √DEIS (cm/√s)

Solvent

Ru (Ω)

Ethaline

283

7.39

6.268 x10-4

Reline

2255

2.41

2.042 x10-4

Aqueous

131

106

8.980 x10-3

Ethaline

289

2.205

1.867 x10-4

Reline

1387

1.205

1.021 x10-4

Aqueous

139

74.3

5.270 x10-3

Ethaline

343

7.356

6.234 x10-4

Reline

2289

3.031

2.568 x10-4

Aqueous

136

60.1

5.080 x10-3

Ethaline

305

8.411

7.128 x10-4

Reline

917

2.156

1.827 x10-4

Aqueous

289

17.67

1.520 x10-3

Ethaline

272

6.329

5.363 x10-4

Reline

2273

0.4656

3.940 x10-5

Aqueous

-

50.55

4.280 x10-3


27


Page 28 of 34

3.3.2. Calculation of diffusion coefficients from cyclic voltammetry The values of diffusion coefficients obtained for different redox species from the EIS studies are compared with the results obtained from cyclic voltammetry. The results are in reasonable agreement with each other as shown in Table 6 considering a simple Randles equivalent circuit model assumed for fitting EIS data. Table 6. Comparison of the diffusion coefficients of different redox species calculated from EIS and CV technique

System

[Fe(CN)6]4-/3-

[Ru(NH3)6] 2+/3+

FcMeOH

FcBF4

MV

√DEIS x10-4

√DCV x10-4

(cm/√s)

(cm/√s)

Ethaline

6.268

4.667

Reline

2.042

1.733

Ethaline

1.867

1.666

Reline

1.021

0.877

Ethaline

6.234

3.413

Reline

2.568

2.018

Ethaline

7.128

4.537

Reline

1.827

1.725

Ethaline

5.363

2.813

Reline

0.394

0.379

Solvent


28

Page 29 of 34 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


4. CONCLUSIONS Electron transfer studies of model anionic, cationic and neutral redox species in DES were systematically studied by cyclic voltammetry and electrochemical impedance spectroscopy. The shift in peak potentials of the redox species in DES wrt molecular solvents could be correlated with the nature of the solvation of the redox species in the medium. FTIR studies of ferrocyanide species which was carried out to interpret the electrochemical parameters, confirmed that the DES components remain largely unaffected by the oxidation-reduction processes of the charged species. The results obtained from the electrochemical studies reflects the lack of solvation order around redox species in DES as suggested earlier for charged species by molecular dynamics simulations and ab initio computations. The solvation behavior of DES around redox species studied here implies that DES has the potential to be used in the half peak potential measurements of novel charged redox probes with minimal solvent interaction. The half peak potentials of neutral/cationic (0/+) species have been found to be shifted more positive wrt molecular solvents. This was attributed to the stabilizing interactions of the nonpolar components of DES with the neutral redox species. While this work deals with solvation behavior

of

DES

for

neutral/cationic

redox

species,

it

is

possible

that a different effect might well be observed for the reduction of a neutral species

to

an

anionic

one

(-/0)

or

of,

say,

a

proton-coupled

oxidation

that

maintains a neutral molecule. The diffusion coefficient of redox species in DES was found to be an order of magnitude less than that observed in aqueous medium. The electron transfer studies carried out in DES medium with a set of standard redox species provide some valuable information on the solvation effects, diffusion coefficients and redox potentials which can be useful for further studies in biological and physiological media.


29


Page 30 of 34

ASSOCIATED CONTENT AUTHOR INFORMATION Corresponding Author *V.Lakshminarayanan [email protected] ACKNOWLEDGMENTS We would like to thank Mrs. K. N. Vasudha for her help in carrying out FTIR studies. REFERENCES 1. Abbott, A. P.; Boothby, D.; Capper, G.; Davies, D. L.; Rasheed, R. K.; Deep Eutectic Solvents Formed Between Choline Chloride and Carboxylic Acids: Versatile Alternatives to Ionic Liquids. J. Am. Chem. Soc., 2004, 126, 9142- 9147. 2. Abbott, A. P.; Capper, G.; Davies, D. L.; Rasheed, R. K.; Tambyrajah, V. Novel Solvent Properties of Choline Chloride/Urea Mixtures. Chem. Commun., 2003, 70-71. 3. Gu, C.; Tu, J. One-Step Fabrication of Nanostructured Ni Film with Lotus Effect from Deep Eutectic Solvent. Langmuir, 2011, 27, 10132-10140. 4. Wei, L.; Fan, Y.; Tian, N.; Zhou, Z.; Zhao, X.; Mao, B.; Sun, S. Electrochemically Shape-Controlled Synthesis In Deep Eutectic Solvents -A New Route to Prepare Pt Nanocrystals Enclosed by High-Index Facets with High Catalytic Activity. J. Phys. Chem. C, 2012, 116, 2040- 2044. 5. Pawar, P. M.; Jarag, K. J.; Sharkarling, G. S. Environmentally Benign and Energy Efficient Methodology for Condensation: An Interesting Facet to the Classical Perkin Reaction. Green Chem., 2011, 13, 2130-2134 6. Gorke, J. T.; Srienc, F.; Kazlauskas, R. J. Deep Eutectic Solvents for Candida Antarctica Lipase B - Catalyzed Reactions. ACS Symposium Series, Vol. 1038, Chapter 14, 169-180. 7. Nkuku, A.; Lesuer, R. J. Electrochemistry in Deep Eutectic Solvents. J. Phys. Chem. B, 2007, 111, 13271-13277. 8. Lloyd, D.; Vainikka, T.; Murtomäki, L.; Kontturi, K.; Ahlberg, E. The Kinetics of the Cu2+/Cu+ Redox Couple in Deep Eutectic Solvents. Electrochim.Acta., 2011, 56, 49424948.


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9. Silvester, S.; Compton, R. G. Electrochemistry in Room Temperature Ionic Liquids: A Review and some Possible Applications. Z. Phys. Chem., 2006, 220, 1247-1274. 10. Abbott, A. P.; Capper, G.; Gray, S. Design of Improved Deep Eutectic Solvents Using Hole Theory. Chem. Phys. Chem., 2006, 7, 803-806. 11. Nikitina, V. A.; Kislenko, S. A.; Nazmutdinov, R. R.; Bronshtein, M. D.; Tsirlina, G. A. Ferrocene / Ferrocenium Redox Couple At Au(111)/Ionic Liquid and Au(111)/Acetonitrile Interfaces: A Molecular-Level View at the Elementary Act. J. Phys. Chem. C, 2014, 118, 6151-6164. 12. Aurbach, D. Nonaqueous Electrochemistry, CRC Press, 1999. 13. Wagle, D.V.; Deakyne, C. A.; Baker, G. A. Quantum Chemical Insight into the Interactions and Thermodynamics Present in Choline Chloride Based Deep Eutectic Solvents. J. Phys. Chem. B, 2016, 120, 6739–6746. 14. Abbott, A. P.; Frisch, G.; Gurman, S. J.; Hillman, A. R.; Hartley, J.; Holyoak, F.; Ryder, K. S. Ionometallurgy: Designer Redox Properties for Metal Processing. Chem. Commun., 2011, 47, 10031-10033. 15. Yu, P.; Yang, F.; Zhao, J.; Wang, J. Hydration Dynamics Of Cyanoferrate Anions Examined by Ultrafast Infrared Spectroscopy. J. Phys. Chem. B, 2014, 118, 3104-3114. 16. Cui, Y.; Fulfer, K. D.; Ma, J.; Weldeghiorghisa, T. K.; Kuroda, D. G. Solvation Dynamics of an Ionic Probe in Choline Chloride-Based Deep Eutectic Solvents. Phys. Chem. Chem. Phys., 2016,18, 31471-31479. 17. Penfold, T. J.; Reinhard, M.; Rittmann-Frank, M. H.; Tavernelli, I.; Rothlisberger, U.; Milne, C. J.; Glatzel, P.; Chergui, M. X-Ray Spectroscopic Study of Solvent Effects on the Ferrous and Ferric Hexacyanide Anions. J. Phys. Chem. A, 2014, 118, 9411–9418. 18. Yee, E. L.; Cave, R. J.; Guyer, K. L.; Tyma, P. D.; Weaver, M. J. A Survey Of Ligand Effects Upon the Reaction Entropies of Some Transition Metal Redox Couples. J. Am. Chem. Soc., 1979, 101, 1131-1137. 19. Kuznetsov, M.; Maslii, A. N.; Krishtalik, L. I. Solvation Of Ferrocene, Cobaltocene, and their Ions by the Data of Quantum-Chemical Calculations. Russ. J. Electrochem., 2009, 45, 87-92. 20. Conte, V.; Floris, B.; Galloni. P.; Bizzarri, C. Solvent Effects of Ionic Liquids: Investigation Of Ferrocenes as Electrochemical Probes. J. Phys. Org. Chem., 2010, 24, 327-334. 21. Gorke, J. T.; Kazlauskas, R. J.; Srienc, F. US 20090117628 A1


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22. Yang, Y.; Yu, L. Theoretical Investigations of Ferrocene/Ferrocenium Solvation in Imidazolium-based Room-Temperature Ionic Liquids. Phys. Chem. Chem. Phys., 2013, 15, 2669-2683. 23. Ju, H.; Leech, D. Effect of Electrolytes on the Electrochemical Behaviour of 11(Ferrocenylcarbonyloxy) Undecanethiol SAMs on Gold Disk Electrodes. Phys. Chem. Chem. Phys., 1999, 1, 1549-1554. 24. Tahara, H.; Baba, R.; Iwanaga, K.; Sagara, T. ; Murakami, H. Electrochromism of a Bipolar Reversible Redox-Active Ferrocene–Viologen Linked Ionic Liquid. Chem. Commun., 2017, 53, 2455-2458 . 25. Gelinas, B.; Das, D.; Rochefort, R. Air-Stable, Self-Bleaching Electrochromic Device Based on Viologen and Ferrocene-Containing Triflimide Redox Ionic Liquids. ACS Appl. Mater. Interfaces, 2017, 9, 28726-28736. 26. Goeltz, J.C.; Matsushima, L. N.; Metal-free redox active deep eutectic solvents. Chem. Commun., 2017, 53, 9983-9985. 27. Kumar, P. S.; Lakshminarayanan, V.; Electrochemical Studies of Redox Probes in SelfOrganized Lyotropic Liquid Crystalline Systems. J. Chem. Sci., 2009, 121, 629–638. 28. Bard, A. J.; Faulkner, L.R. Electrochemical Methods Fundamentals and Applications, John Wiley and Sons, Noida, 2nd edn., 2004. 29. Lovelock, K. R. J.; Ejigu, A.; Loh, S. F.; Men, S.; Licence, P.; Walsh, D. A. On the Diffusion of Ferrocenemethanol in Room-Temperature Ionic Liquids: An Electrochemical Study. Phys. Chem. Chem. Phys., 2011, 13, 10155-10164. 30. Bahadori, L.; Manan, N.S.A.; Chakrabarti, M. H.; Hashim, M. A.; Mjalli, F.S.; AlNashef, I. M.; Hussain, M. A.; and Low C. T. J.; The Electrochemical Behaviour of Ferrocene in Deep Eutectic Solvents Based on Quaternary Ammonium and Phosphonium Salts. Phys. Chem. Chem. Phys., 2013, 15, 1707-1714.


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(Cationic, Anionic, Neutral) in Deep Eutectic ... - ACS Publications

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