CHARACTERIZATION OF SOLVENTS FOR
ELECTROCHEMICAL ENERGY STORAGE:
DEEP EUTECTIC SOLVENTS AND IONIC LIQUIDS
By
HENRY JOHN SQUIRE
Submitted in partial fulfillment of the requirements for the degree of
Master of Science
Thesis Advisor: Professor Burcu Gurkan
Department of Chemical and Biomolecular Engineering
CASE WESTERN RESERVE UNIVERSITY
May, 2020
1 Case Western Reserve University
School of Graduate Studies
We hereby approve the thesis of Henry John Squire
Candidate for the degree of Master of Science
Committee Chair Burcu Gurkan
Committee Member Uziel Landau
Committee Member Robert Savinell
Committee Member Daniel Lacks
Date of Defense April 27, 2020
*We also certify that written approval has been obtained for any proprietary material contained therein.
2 Copyright
Portions of this work has been/will be published in the following papers:
Gurkan, B., Squire, H. & Pentzer, E. Metal-Free Deep Eutectic Solvents:
Preparation, Physical Properties, and Significance. J. Phys. Chem. Lett. 10, 7956–
7964 (2019). Chapter 3
Spittle, S. et al. Glyceline Dynamics and Structure. J. Am. Chem. Soc. to be submitted. Chapter 4
Zhang, Y. et al. Liquid Structure and Transport Properties of the Eutectic Mixture of Choline Chloride and Ethylene Glycol (Ethaline). J. Phys. Chem. B to be submitted. Chapter 4 and Chapter 5
Klein, J., Squire, H., Dean, W. & Gurkan, B. From salt in solution to soley ions - solvation of methyl viologen in deep eutectic solvents and ionic liquids. J. Phys.
Chem. B under review. Chapter 5
Klein, J. M., Squire, H. & Gurkan, B. Electroanalytical Investigation of the
Electrode–Electrolyte Interface of Quaternary Ammonium Ionic Liquids: Impact of Alkyl Chain Length and Ether Functionality. J. Phys. Chem. C 124, 5613–5623
(2020). Chapter 6
3 Table of Contents
List of Tables ...... 6 List of Figures ...... 7 Acknowledgements ...... 10 Chapter 1: Introduction ...... 12 Chapter 2: Background ...... 16 2.1 Definition of DESs and ILs ...... 16 2.2 Brief History of DESs and ILs ...... 18 2.3 Application of DESs for Energy Storage ...... 20 2.4 Applications of ILs for Energy Storage ...... 21 2.5 Structure-Function Relation ...... 22 Chapter 3: Methods ...... 24 3.1 Preparation and Handling of DESs ...... 24 3.2 Characterization Considerations ...... 27 3.3 Physical Property Characterization Techniques ...... 29 Chapter 4: Properties of Glyceline and Ethaline Deep Eutectic Solvents ...... 38 4.1 Introduction ...... 39 4.2 Methods and Materials ...... 41 4.3 Results and Discussion ...... 42 4.4 Conclusion ...... 50 Chapter 5: Solvation of Methylviologen Dichloride in an IL, DES, and DES Analogues ...... 52 5.1 Introduction ...... 52 5.2 Methods ...... 56 5.3 Results and Discussion ...... 60 5.4 Conclusions ...... 71 Chapter 6: Redox Behavior of Quinones in Ethaline...... 73 6.1 Introduction ...... 73 6.2 Materials and Methods ...... 77 6.2 Results and Discussion ...... 78 6.3 Conclusions ...... 87 Chapter 7: Ionic Liquids: Bulk Properties and Interface Characteristics ...... 88
4 7.1 Introduction ...... 89 7.2 Methods ...... 90 7.3 Results and Discussion ...... 94 7.4 Conclusion ...... 101 Chapter 8: Conclusions and Future Work ...... 102 References ...... 106
5 List of Tables
Table 2.1 Viscosity, conductivity, and electrochemical window (EW) of representative solvents at 25ºC; comparing organic and aqueous with ILs and DESs...... 18 Table 4.1 Water contents of studied DESs (33 mol% ChCl) and DES analogues, and frequency range of EIS used for the calculation of ionic conductivities...... 41 Table 4.2 Fitted Arrhenius parameters and for the mixtures of ChCl and G...... 43 𝐴𝐴 𝐸𝐸𝐸𝐸 Table 4.3 Fitting parameters and for temperature-dependent viscosity of ChCl in EG...... 44 𝐴𝐴 𝐸𝐸𝐸𝐸 Table 4.4 Fitting parameters and for temperature-dependent conductivity of ChCl in G and ChCl in EG solutions...... 47 𝐴𝐴 𝐸𝐸𝐸𝐸 Table 4.5 Fitted parameters of and for temperature-dependent densities of ChCl in G and ChCl in EG solutions...... 50 𝑎𝑎 𝑏𝑏 Table 5.1 Measured viscosity ( ) and surface tension ( ) at 25 °C and the estimated average hole size (r) using equation 1...... 62 𝜂𝜂 𝛾𝛾 Table 5.2 Peak assignments and peak locations for MV+2 and EG vibrational modes in solvents studied. 1:2, 1:4, and 1:6 represent the molar ratios of the choline-based salt to EG. Peaks were extracted by Gaussian function fitting. The symbol ⁑ represents a mode that comes up at two different frequencies. MVCl2 concentration was 10 mM in the corresponding mixtures. The table is color coded to indicate blue shifts and red shifts of MV2+ and EG modes in solutions with respect to the pure compounds...... 68 Table 6.1 Results of linear fit to concentration versus absorbance for different modes of H2BQ. Parameters m and b are the slope and intercept of the fit respectively. R2 is the goodness of fit measure...... 80 Table 6.2 Solubility limit of various quinones in ethaline listed according to the functional groups associated with each quinone species. Decreasing polarity and increasing size of functional groups decreases solubility of quinone in ethaline. 82 Table 7.1 Water contents of dried ILs and frequency range of linear fit to EIS data for determining conductivities...... 93 Table 7.2 Measured viscosity conductivity and density of studied ILs, each with [TFSI] anion, at 298.15 K ...... 95
6 List of Figures
Figure 1.1 Set-up for a redox flow battery in which pumps supply electrolyte to the electrodes (A) and a lithium-ion battery in which electrolyte is stagnant, stored between the electrodes (B)...... 13 Figure 2.1 Structures of representative DESs and ILs ...... 17 Figure 3.1 Preparation of ethaline: choline chloride and ethylene glycol in a 1:2 molar ratio. The sample on the left was heated improperly or exposed to water leading to crystal formation. The homogenous sample on the right was heated at 80 ºC for 2 hours and stored in an argon atmosphere glovebox...... 26 Figure 3.2 Viscosity as a function of temperature for ethaline. Our measurements are compared with other reports in literature. Viscosity values at 313 K are reported in the legend following the sample name and the water content in parenthesis...... 27 Figure 3.3 Schematic of the setup of a vibrating U-tube density meter...... 31 Figure 3.4 Example Nyquist plot with linear fit. Intercept with the real axis (Z’) represents solution resistance...... 34 Figure 3.5 Schematic of a T-Cell setup (A) and a traditional three electrode setup (B). The reference electrode is insulated to prevent accidental contact with the coiled counter wire electrode in the T-cell setup...... 36 Figure 4.1 Structures of investigated hydrogen bond donors (HBDs), ethylene glycol (EG) and glycerol (G), and hydrogen bond acceptor (HBA), choline chloride (ChCl)...... 40 Figure 4.2 Temperature-dependence of viscosity for mixtures of ChCl and G (A). Lines are Arrhenius fits. Viscosity as a function of mol% ChCl in G at 298.15 K (B)...... 42 Figure 4.3 Temperature-dependent viscosity of mixtures of ChCl and EG (A). Lines are fits to Arrhenius model. Viscosity as a function of mol% ChCl in EG at 298.15 K (B)...... 44 Figure 4.4 Temperature-dependent conductivities of mixtures of ChCl and G (A) and ChCl in EG (B). Lines indicate Arrhenius fits...... 46 Figure 4.5 Temperature-dependent density of various mol% ChCl in G (A). Density as a function of ChCl mol% in G at 25 ºC (B). Temperature-dependent density of various mol% ChCl in EG (C). Density as a function of ChCl mol% in EG at 25 ºC (D). Lines are linear fits...... 49
Figure 5.1 Structures of the three solvent systems studied: [PYR13][TFSI] – the IL(left), ChCl:EG – components of a DES (middle), ChTFSI:EG – a DES analogue (right)...... 55
7 Figure 5.2 Two-step reduction mechanism of methylviologen, MV2+, in acetonitrile as reported by Kohler et al...... 56 Figure 5.3 Temperature-dependent viscosity (A) and conductivity (B) of the solvent systems from 298 to 328 K. The legend in the inset of panel A applies to both panels A and B...... 61 Figure 5.4 Walden plot for the solvents investigated (A). Calculated diffusivities as a function of inverse viscosity. The dotted lines represent fits to the Stokes- Einstein relation in equation 3 where two average sizes of radii are observed: one for concentrated systems (mixtures with 1:2 molar ratio of salt-to-solvent and the IL) and one for dilute systems (mixtures with salt-to-solvent molar ratios of 1:4 and 1:6) (B). Representative CVs of the systems at 15 mV s-1, the dotted line indicates the half wave potential of the first reduction of MV2+ in 1:2 ChCl:EG. The shift in reduction potential of MV2+ is due to the choice of reference electrode and should not be used to link to hydrogen bonding affects (C)...... 64 Figure 5.5 Relationship between the diffusion coefficient of MV2+ determined from CV using eqn. 2 and the ratio of the solvation radii (eqn. 3) to the hole size (eqn. 1)...... 66 Figure 5.6 Normalized Raman spectra for representative systems as indicated in labels. Green highlights are used to identify the regions of the spectra discussed in the text. Spectra were normalized to a maximum intensity of one so all systems could be compared on a common scale. It should be noted that non-normalized spectra were used for the analysis...... 67 Figure 5.7 Change in Raman shift of MV rings modes at 1302 cm-1 (A) and 1554 -1 cm (B) in [PYR13][TFSI], 1:2 ChCl:EG, 1:2 ChTFSI:EG, and EG. Ring modes are red-shifted in the IL and DES but blue-shifted in the DES analogue and EG. 70 Figure 5.8 Illustration of the change in bonding environment for EG upon the addition of MVCl2 (10 mM), followed by ChCl or ChTFSI. Blue indicates vibrational modes that were blue-shifted and red indicates modes that were red- shifted with respect to pure compounds. Green dashed lines represent hydrogen bonds that yield peak splitting and broadening in the experimental Raman spectra 2+ for the corresponding modes. Pink dashed lines indicate MV -EG interactions. 71 Figure 6.1 Redox pathway of BQ in protic (A) and aprotic media (B)...... 75 Figure 6.2 Structures of quinones investigated in this study, from left to right: hydroquinone (H2BQ), benzoquinone (BQ), naphthoquinone (NQ), and anthraquinone (AQ)...... 76
Figure 6.3 IR spectra for neat ethaline and various concentrations of H2BQ in ethaline. Vertical black, blue, and red lines indicating modes of interest for H2BQ, ethylene glycol, and choline chloride, respectively (A). Linear fit to concentration -1 versus normalized absorbance for the 1205 cm v CC mode of H2BQ (B)...... 80
8 Figure 6.4 IR spectra for neat ethaline and various concentrations of BQ in ethaline. Vertical black, blue, and red lines indicating modes of interest for BQ, ethylene glycol, and choline chloride respectively (A). Linear fit to concentration versus normalized absorbance for the 1653 cm-1 v CO mode of BQ (B)...... 81 Figure 6.5 EW of ethaline on glassy carbon and platinum working electrodes. ...84
Figure 6.6 Voltammograms of 0.5 M H2BQ in ethaline (A) and 20 mM BQ in ethaline (B). Anodic and cathodic peaks are marked Ea and Ec respectively. The anodic limit of ethaline is indicated by a black line...... 85 Figure 6.7 Potential-restricted voltammograms, -0.75 to 0.75 V v Ag/AgCl (A) and 0.125 V to 1.25 V vs Ag/AgCl (B)...... 86 Figure 7.1 Walden plot of investigated ILs (A) and cyclic voltammograms of neat ILs on glassy carbon at 298 K, scale bar represents the current density of 200 µA/cm2 (B)...... 96 Figure 7.2 Differential capacitance evaluated with EIS on glassy carbon electrode. Error bars represent a 90% confidence interval calculated from the statistical average of capacitance values. Red lines represent fits to the Goodwin- Kornyshev model. Potential was swept from open circuit voltage (OCV) to negative potentials, returned to OCV then scanned to positive potentials...... 99
9 Acknowledgements
Thank you to Professor Burcu Gurkan for three years of mentorship and support throughout my undergraduate and graduate research careers. Thank you to Professor Dan
Lacks for providing numerous opportunities and unique ideas. Thank you to Professor
Rohan Akolkar for constant advice and unmatched dedication to teaching.
Thank you to Professor Uzi Landau, Professor Bob Savinell, and Professor Dan Lacks for serving as committee members.
Thank you to Jeff Klein for challenging me to “figure it out” and being a great friend.
Thank you to Andrew Wang, Adriaan Reit, and Joe Toth for giving me a home and providing advice, friendship, and entertainment over the last two years.
Thanks to my “ChemE best friend” Sidd Rajupet.
Finally, much love and thanks to mom, dad, Claire, and Jay.
10 Characterization of Solvents for Electrochemical Energy Storage:
Deep Eutectic Solvents and Ionic Liquids
By
HENRY JOHN SQUIRE
Abstract
Deep eutectic solvents (DESs) and ionic liquids (ILs) are emerging classes of tunable
solvents with potential applications in electrochemical energy storage systems. These
solvents exhibit favorable properties including high solvent strength, wide electrochemical
windows, and lack of flammability. However, high viscosity and low conductivity has thus
far limited their practical application. Additionally, their heterogenous liquid structure and
lack of basic structure-function design rules presents challenges in understanding fundamental properties. Here, standards for consistent preparation, handling, and characterization of DESs are proposed. The impact of chemical structure on DES physical, solvation, and transport properties is evaluated demonstrating anion substitution as a future pathway for solvent tuning. A preliminary study investigating the feasibility of utilizing quinones in DESs demonstrates anomalous previously unreported redox behavior. Finally, the electrode-electrolyte interface of ammonium ILs is probed. Long alkyl chains are demonstrated to stabilize accumulation of anions at the electrode surface.
11 Chapter 1: Introduction
Grid-scale energy storage systems are becoming increasingly important to foster a
sustainable, modern electric grid. In response to the threat of global warming, zero- emission renewables like solar and wind represent an increasing proportion of US energy sources every year. However, these sources are intermittent; peak supply is often mismatched with peak demand. Grid-scale energy storage systems are able to integrate renewables into the grid and bridge supply and demand gaps thereby improving grid sustainability and reliability. Pumped hydro, electrochemical capacitors, compressed air, flywheels, and batteries are all currently implemented as grid-scale energy storage solutions with varying degrees of success.1,2 Among these solutions, batteries are
particularly attractive due to their high overall efficiency, high energy density, and
moderate power density.1 State of the art rechargeable batteries currently used to power
portable electronic devices are based lithium-ion battery chemistry. However, ongoing
safety concerns with lithium-ion batteries and incompatibility of its design with flow
systems such as redox-flow batteries necessities the development of new materials,
electrolytes, and battery architectures for large scale energy storage.3,4 Redox flow batteries
have attracted attention for grid-scale application due to their scalability and decoupling of
energy and power density,1,5 defined as the amount of stored energy per weight or volume
of battery and the speed at which energy can be supplied per weight or volume of the
battery.
Lithium-ion and redox flow batteries are composed of the same components:
electrodes, separator, and electrolyte, but with differences in functionality and material
identities. Figure 1.1 shows the design of a lithium-ion and a redox flow battery.6,7 Both
12 architectures rely on reversible coupled reduction and oxidation (redox) reactions at the electrode surfaces to either store or supply energy. To balance charge across the system and complete the circuit for the flow of ions, an ion-conductive electrolyte is placed in between the electrodes. Finally, a separator between the electrodes prevents short circuiting of the system.
Figure 1.1 Set-up for a redox flow battery in which pumps supply electrolyte to the electrodes (A) and a lithium-ion battery in which electrolyte is stagnant, stored between the electrodes (B).
Lithium-ion batteries typically utilize a porous graphite electrode coupled with a lithium doped metal oxide counter electrode such as cobalt oxide. While charging a lithium-ion battery, lithium ions deintercalate from the metal oxide electrode and move through the electrolyte eventually intercalating into the graphite electrode.7 The reverse process occurs upon discharging. In this way, the electrolyte simply functions as a
13 conductive pathway in lithium-ion batteries. Because species intercalate into the
electrodes, the total energy stored by the battery depends on the electrode mass and volume.
The speed at which lithium-ion batteries can supply energy is also dependent on the electrode architecture. In contrast, redox species in flow batteries typically remain in solution and do not intercalate into the electrode.6 Pumps maintain electrolyte flow across
the electrodes, which are typically a carbon felt, to charge or discharge the battery. Thus, the electrolyte is not only a conductive medium but also an energy storage medium in a redox flow battery. Furthermore, this means the energy storage capacity of a redox flow
battery is proportional to the quantity of stored electrolyte. The speed at which energy can
be supplied is then a function of the electrode area, independent of the amount of energy
stored.
To enable redox flow batteries and safe lithium-ion batteries with higher energy
and power densities, novel materials are needed. Among these materials, electrolytes need
particular improvement as traditional aqueous and organic electrolytes are limited in
voltage window and solvent strength. Currently, the most common industrially relevant
redox flow batteries utilize vanadium metal redox species in aqueous electrolytes with low
electrochemical stabilities which limit the energy and power density of the battery.8
Lithium-ion batteries utilize highly flammable organic carbonate electrolytes which
present serious safety concerns, especially as larger battery architectures are built.4 Two
novel classes of electrolytes, ionic liquids (ILs) and deep eutectic solvents (DESs), with
wide electrochemical stability and lack of flammability could solve these ongoing
problems and enable a new generation of safer grid-scale energy storage devices with
greater energy storage capacity.9 ILs are salts which are liquid below 100ºC while DESs
14 are mixtures of acids and bases with a significantly depressed freezing point compared to the parent compounds.10,11 In this study, we develop structure-function relation of these electrolytes by characterizing and relating macroscopic properties like viscosity to molecular level interactions such as solvation with the aim of enabling ILs for use in lithium-ion batteries and DESs for application in redox flow batteries.
The following thesis is organized as follows. Chapter 2 introduces ILs and DESs relating specific electrolyte properties to their potential applications. Chapter 3 reviews the methods used to characterize ILs and DESs which was published as a Perspective in the Journal of Physical Chemistry Letters.12 In Chapter 4, the physical properties of glycerol and ethylene glycol DESs are investigated. The results are under review in the
Journal of Physical Chemistry B and the Journal of the American Chemical Society.14,15
Chapter 5 presents an investigation of solvation of redox active species in DESs and ILs, published in an article under review in the Journal of Physical Chemistry B.16 Chapter 6 introduces a screening study investigating the solubility and redox activity of quinones in
DESs. Chapter 7 presents work related to understanding the IL-electrode interface, the results of which were published in the Journal of Physical Chemistry C.13 Finally, Chapter
8 summarizes conclusions and future work.
15 Chapter 2: Background
Deep eutectic solvents (DESs) and ionic liquids (ILs) are distinct classes of tunable
solvents with potential for application in energy storage devices. This chapter provides
information on the background and relevant properties of these systems for application in electrochemical devices such as lithium ion, lithium metal, and redox flow batteries. In
particular, the challenges and opportunities in the field will be reviewed. Focus is given to
ILs and DESs due to their tunable properties and potential designation as “designer
solvents”.
2.1 Definition of DESs and ILs
DESs and ILs possess a number of similar properties yet are distinct classes of solvents;
although DESs are sometimes classified as a subclass of ILs. DESs are mixtures of Lewis
and/or Bronsted acids and bases with significantly depressed freezing points as compared
to the parent compounds.10 The constituent species of DESs engage in a “hydrogen bonding network” that is believed to govern the liquid structure of DESs.17 DES parent compounds
are classified as hydrogen bond donors (HBDs) or hydrogen bond acceptors (HBAs). In
contrast to DESs, ILs are salts composed of discrete cations and anions, not combinations
of HBAs and HBDs. Figure 2.1 illustrates the structure of representative DESs and ILs.
ILs are arbitrarily defined to have a melting point below 100ºC to distinguish them from
high temperature molten salts.18 While charge delocalization over large cationic and
anionic structures weakens the ionic bond between ions resulting in a low melting point in
ILs, the origin of freezing point depression in DESs is not yet completely understood.
16
Figure 2.1 Structures of representative DESs and ILs
DESs and ILs share a number of favorable properties that make them attractive for application in electrochemical devices. Traditional solvents such as organic and aqueous solvents have found extensive use in electrochemical energy storage due to their high conductivity, low viscosity, and relatively low-cost.19 Table 2.1 summarizes properties of
DESs, ILs, and traditional solvents. Differences in these properties can be seen across solvent classes. Aqueous solvents are not electrochemically stable; they can only be used in narrow a voltage window of less than 1.23 V and are not suitable for devices such as lithium ion batteries where open circuit voltages approach 4 V.11 Organic solvents by contrast have wider electrochemical windows and have found use in lithium-ion batteries.
However, organic solvents are volatile, toxic, and flammable which limit the operating temperature ranges and present a fire or explosion risk. DESs and ILs have become attractive due to lack of flammability, little to no volatility, and wide electrochemical
17 windows.19-9 Due to their properties, these tunable solvents could pave the way for a new
generation of safer electrochemical storage devices with greater energy densities.
Table 2.1 Viscosity, conductivity, and electrochemical window (EW) of representative
solvents at 25ºC; comparing organic and aqueous with ILs and DESs.
aethaline: a mixture of choline chloride and ethylene glycol in a 1:2 molar ratio10 bTFA-TPB: a mixture of methyltriphenylphosphonium bromide and 2,2,2-triflouracetamide in a 1:8 molar ratio20 c[EMIM][TFSI]: 1-Ethyl-3-methylimidazolium tetrafluoroborate21 d[N1114][TFSI]: butyltrimethylammonium bis(trifluoromethylsulfonyl)imide13 e1 M KCl in water22 f1 M LiBF4 in DMF: 1 molar lithium tetrafluoroborate in dimethylformamide5
2.2 Brief History of DESs and ILs
In 2001, Abbott et al. reported the first DES system, a mixture of zinc chloride and choline
chloride which was liquid at room temperature despite the parent species each melting over
150ºC.23 The term “DES” was later coined by Abbott et al. in 2004.24 Since then, a vast
array of DESs have been developed utilizing HBDs such as metal halides,24 amides,25
18 sugars,26 polyols,27 quinones,28 and carboxylic acids24 and HBAs such as ammonium
salts,24 phosphonium salts,29 amino acids,30 and viologens.31 Given the array of parents
compounds, the design space is seemingly infinitely wide allowing for tuning of DES
properties for any application. DESs have shown promise in areas including catalysis,32
electrodeposition,33 separations,34 and energy storage.9
ILs have a longer history then DESs and have therefore been more thoroughly investigated. ILs were first reported in 1914 when Walden et al. found the salt ethylammonium nitrate had a melting point of 12ºC.35 However, there were few
investigations into ILs until the 1980s when sustained interest in ILs generated the birth of
the modern “IL field”.36 Like DESs, ILs are tunable with a wide variety of cation and anion
structures. Common cation structures include phosphoniums, ammoniums, sulfoniums,
pyridiniums, and imidazoliums while common anion structures include imides, sulfonates,
sulfates, phosphates, and borates.36 ILs have been more extensively studied then DESs
37 38 finding application in biomaterial processing, CO2 conversion or separation, and
electrochemical devices such as super capacitors,39 sensors,40 and dye sensitized solar
cells.41
Despite the potential of DESs and ILs, only a few niche applications utilize these
systems in practice. Currently DESs and ILs are not widely used in electrochemical
applications. This could be attributed to their high viscosity and low conductivity. Aqueous
and organic electrolytes are typically two orders of magnitude less viscous and one order
of magnitude more conductive.42,43 Additionally, unlike traditional solvents such as
aqueous or organic solvents, ILs and DESs typically display dynamic and/or spatial
heterogeneities. These heterogeneities present challenges in understanding
19 electrochemical, transport, rheological, and phase properties.12 Traditional solvents by contrast are typically dynamically and spatially homogenous, composed of simple structures with relatively simple molecular interactions. Thus, classical liquid theories are applicable to traditional solvents which are not applicable to ILs or DESs. 13,12 This presents a significant challenge in applying DESs and ILs and therefore more fundamental studies are needed to understand their physical properties, liquid structure, solvation, and dynamics.
2.3 Application of DESs for Energy Storage
DESs have a number of favorable properties that make them attractive for application in electrochemical devices, in particular redox flow batteries. As mentioned before DESs are non-flammable and exhibit little to no volatility. Additionally, most are inherently conductive as DESs typically incorporate a salt HBA. DESs, as seen in Table 2.1, also have wider electrochemical stability windows then aqueous solvents.9 DESs have a number
of attractive properties that ILs do not such as low-cost, ease of preparation, low-toxicity, and solvation strength towards redox active solutes.12 DESs can be prepared from materials produced on a mega-ton scale. Additionally, since DESs are prepared and not synthesized, extensive purification is not typically required other than drying.12 Finally, redox active
species such as metal halides have been shown to have high solubilities in DESs potentially
allowing for a high energy density electrolyte.44
The properties of DESs make them ideal candidates for electrolytes in redox flow
batteries. Replacing the aqueous electrolyte currently used in flow batteries would allow
20 the operating open circuit voltage to double due to the wider EW of DESs. This could
correspondingly double the power and energy density of the system. Furthermore, since
DESs can be prepared cheaply, a DES based electrolyte could be cost competitive with
aqueous electrolytes.9 While it has previously been proposed that organic solvents could
be used as an electrolyte in flow batteries due to their wide EW, organic solvents are highly
flammable and toxic which presents a major safety concern for flow batteries due to their
large size. Even though the EW of some of the common DESs such as ethaline (1:2 molar
ratio of choline chloride to ethylene glycol) are smaller than organic solvents, DESs are
non-flammable and non-toxic so a redox flow battery based on DES would have fewer
safety and environmental concerns. Furthermore, the EW of DESs might be widened
through structural design. Finally, due to the solvation strength of DESs, the concentration
of certain redox active species can be increased in DESs as compared to aqueous or other
organic electrolytes. Since the concentration of redox active species is proportional to the
energy density of the battery, utilizing a DES based electrolyte could greatly increase the
energy density of a redox flow battery. While DESs are ideal for increasing the energy
density of redox flow batteries, most of the currently known DESs could not be utilized in
high energy density devices like lithium-ion batteries as their EW is not wide enough to operate under high potentials.
2.4 Applications of ILs for Energy Storage
ILs, similar to DESs, are non-flammable with no volatility. However, in contrast to
DESs, ILs have much wider EWs approaching 5.5 V in some cases (Table 2.1). Due to their wide EW, ILs can be used in high voltage, high energy density applications like
21 lithium-ion batteries. Additionally, due to their lack of flammability, ILs could potentially
enable lithium metal batteries which are not adapted in practice due to a variety of problems
including safety issues.45 Lithium metal batteries replace the graphite electrode of lithium
ion batteries with a lithium metal electrode. Lithium metal is extremely reactive and cannot
be safely implemented with traditional organic solvents which are flammable. Because ILs
are non-flammable and have wide EWs, they are potential electrolytes for lithium-ion or lithium metal batteries.46 Compared to DESs, ILs are more expensive and intensive to
synthesize. Additionally, the synthesized ILs often require extensive purification, adding
to the expense and difficulty of producing ILs.
2.5 Structure-Function Relation
As illustrated above, solvent properties determine the feasibility of DESs and ILs for
various applications. Ultimately, the structure of the molecules that comprise a solvent
determines its properties or function. Thus, understanding the relationship between the
structure and the function of a solvent is the key to developing new generations of high-
performance solvents. Despite the potential of complex solvents, thus far they have not
been widely adopted in practical applications, as mentioned earlier. The design space for
both ILs and DESs is seemingly infinite with a vast array of possible structures and
combinations of cation and anions or HBDs and HBAs. As such, in theory, there could be
an ideal DES or IL for any imaginable application. However, currently the process of
tuning these systems is haphazard with few underlying theories and guiding principles for
developing a solvent with a desirable set of properties.12 Through developing structure- function relation, underlying principles can be established such that ILs and DESs could
22 be crafted to fit particular applications. Much of the work in the field of complex solvents
seeks to develop this understanding so ILs and DESs can truly be utilized as tunable
designer solvents. This work seeks to further develop structure-function understanding through a physical chemistry approach. Specifically, through evaluation of the physical, electrochemical, and spectroscopic properties of complex solvents, trends in molecular level interactions are related to macroscopic properties like viscosity, solvation, and transport, tying structure to function. In the consecutive chapters, these aspects will be discussed in the light of performed experimental studies.
23 Chapter 3: Methods
A variety of methods are used to investigate ILs and DESs including the physical property
characterization, cyclic voltammetry (CV), and Fourier-transform infrared spectroscopy
(FTIR). These techniques are well established for ILs and more information on
characterization of ILs can be found in the works of Wilkes et al.,18 Angell et al.,47 and
Welton et al.36. Since the DES field remains in its infancy, characterization considerations
are not unified. Additionally, DES preparation and handling procedures lack uniformity.
In many cases, this has led to discrepancies in reported physical properties complicating
the establishment of structure-function relations. This section recommends standards for
preparation, handling, and characterization of DESs. While the focus here is DESs, these
techniques can be applied to study ILs as well. Much of this chapter was published as a
perspective article in The Journal of Physical Chemistry Letters.12
3.1 Preparation and Handling of DESs
DESs are prepared rather than synthesized typically only requiring mixing combined with heating. While DESs are simple to prepare, DESs and their constituent components are often hygroscopic. Reports investigating DESs frequently provide no information regarding what measures, if any, were used to remove or prevent water uptake. In fact,
many reports simply state constituent components were mixed in a particular ratio until a
homogenous mixture was obtained. In order to consistently obtain reliable properties from
DESs, factors such as component purity, water content, preparation environment, and
preparation time should be carefully considered.
24 Many of the components used in DESs are hygroscopic. Hydrogen bond acceptors
(HBAs) tend to be more hygroscopic then hydrogen bond donors (HBAs) since HBAs are
typically salts. Drying of HBAs can be accomplished by using a vacuum oven. For example, choline chloride, which is utilized in DESs such as ethaline, glyceline, and reline, if used as received could introduce a significant amount of water to the final DES. This could significantly impact physical properties. Thermally stable solids such as choline chloride or urea can be dried under vacuum at <120 ºC. Occasionally, if pure species cannot be obtained, it may be necessary to recrystallize samples from an appropriate solvent prior to the drying procedure. Volatile components such as ethylene glycol can be dried with molecular sieves if necessary. Once components are dried, the DES can be prepared. This should be done under a dry inert environment such as an argon or nitrogen atmosphere.
The exact preparation procedure depends on the DES of interest. Ethaline which is a mixture of choline chloride and ethylene glycol in a 1 to 2 molar ratio, can be prepared on a hotplate using a stir bar. Ethaline should be mixed at approximately 80 ºC for to 2 hours.
In contrast, the DES reline, a mixture of choline chloride and urea in a 1 to 2 molar ratio, cannot be mixed on a hotplate initially. Since both constituent materials start as solids, stir bars cannot initially mix components. Thus, the mixture should initially be stirred by hand with a spatula under a dry environment such as a glovebox to prevent water uptake. When the choline chloride/urea mixture begins liquifying, it can be transferred to a hotplate and mixed at 80 ºC for 2 hours.
Once a dry DES is prepared, precautions must be taken to prevent water uptake.
DESs are hygroscopic and must be stored in a dry environment such as a glovebox.
Occasionally DESs such as ethaline will form crystals. These crystals were confirmed by
25 X-ray diffraction to be choline chloride leaving solution.12 This is more frequently seen if
samples are improperly heated or improperly dried. Figure 3.1 shows an example of a
properly prepared ethaline sample as well an improperly prepared ethaline sample with
crystal growth.
Figure 3.1 Preparation of ethaline: choline chloride and ethylene glycol in a 1:2 molar ratio. The sample on the left was heated improperly or exposed to water leading to crystal formation. The homogenous sample on the right was heated at 80 ºC for 2 hours and stored in an argon atmosphere glovebox. From ref 12.
If crystal growth is seen in an ethaline sample, simply reheating to 80 ºC and slowly cooling for 2 hours dissolves crystals and prevents further crystal growth. Ibrahim et al. reported formation of crystals in 1:1 molar ratio mixtures of choline chloride and ethylene glycol but other reports disclosing crystal formation have not thus far been published.42 As
such, the cause of crystal growth remains ambiguous. Crystal formation could be due to
formation of choline chloride rich domains during cooling which expel ethylene glycol
destroying the local hydrogen bonding network. Given a homogenous mixture could not
be obtained with a 1:1 molar ratio of choline chloride and ethylene glycol,42 the hydrogen
26 bonding network of ethaline could be saturated at a 1:2 molar ratio and cannot
accommodate any additional choline chloride without phase separation. Regardless of the
cause of crystal formation in ethaline, crystal formation can be prevented through proper
preparation and handling as explained above.
3.2 Characterization Considerations
Maintaining consistent preparation and handling techniques is important as, for example,
even a small amount of water can drastically affect properties such as viscosity and
conductivity. To further illustrate this point, the impact of water on the viscosity of ethaline
was evaluated. Figure 3.2 shows viscosity as a function of temperature for a variety of
ethaline samples including samples from literature.48,49,50,51,52
Figure 3.2 Viscosity as a function of temperature for ethaline. Our measurements are compared with other reports in literature. Viscosity values at 313 K are reported in the legend following the sample name and the water content in parenthesis. From ref 12.
27
Samples had different reported values of water content ranging from 100 ppm up to 1000 ppm. Some literature samples had no reported water contents. In general, as the water content of the sample increases, the viscosity decreases. This is due to the difference in viscosity between water and ethaline, 1 cP for water and 48 cP for dry ethaline at 25ºC.
The viscosity of our sample with a water content of 100 ppm matches the viscosity reported by Barrdo et al.48 who had a similarly low water content of 145 ppm. Samples with higher water contents, such as Gajardo-Parra et al.49 (410 ppm) and Harafi-Mood et al.50 (600 ppm), consistently had lower viscosities as compared to the 100 ppm and 145 ppm samples.
As the temperature increases, differences in viscosity between samples of varying water content become less pronounced due to the dominance of molecular motion over intermolecular interactions between the HBA, HBD, and water at higher temperatures.
The biggest discrepancies in reported viscosity occurs in samples with no reported water content. For example, Ozturk et al.51 and D’Agostino et al.52 do not report the water content of their samples. The viscosities reported by Ozturk et al.51 are consistently lower than all other samples which suggest their water content was the highest but no water content is reported. In general, all samples follow the same trend; exponential decrease in viscosity with increase in temperature as predicted by the Arrhenius or Vogel-Fulcher-
Tammann53 (VFT) models.54 A notable exception is the viscosities reported by D’Agostino et al.52 which follow an inverse linear relationship with temperature. While the exact reason for this discrepancy is unknown, it should be noted that the authors do not report the water content of the sample, the preparation/handling procedure, or the method used to determine viscosity. Thus, the importance of consistent preparation, handling, and characterization
28 cannot be overstated. To obtain reliable results with utility to the rest of the community,
stringent precautions need to be taken to avoid water uptake and all water contents should
be reported.
3.3 Physical Property Characterization Techniques
A few common DES properties are typically characterized when attempting to develop
structure-property relation and evaluating the applicability of a DES to a practical system.
These properties include density, viscosity, conductivity, thermal stability, electrochemical
stability, and polarity. A variety of setups or techniques can be used to evaluate properties
but certain techniques are preferred for DES systems. These considerations are reviewed
here with particular attention to electrochemical, spectroscopic, and rheological properties.
3.3.1 Water Content Determination
Before any physical characterization, the water content of the DES needs to be evaluated
so that it can be reported. This is typically accomplished through Karl Fischer Titration
(KFT) in which water is consumed through the KF reaction expressed below:
+ + [ ] + 2 [ ] + 2[ ] + − + − + − 𝐻𝐻2𝑂𝑂 𝐼𝐼2 𝑅𝑅𝑅𝑅𝑅𝑅 𝑆𝑆𝑆𝑆3𝐶𝐶𝐻𝐻3 𝑅𝑅𝑅𝑅 → 𝑅𝑅𝑅𝑅𝑅𝑅 𝑆𝑆𝑆𝑆4𝐶𝐶𝐻𝐻3 𝑅𝑅𝑅𝑅𝑅𝑅 𝐼𝐼 where represents a base.55 For every mole of iodine consumed, one mole of water is
consumed𝑅𝑅𝑅𝑅 allowing for water content determination. Iodine is added until excess is present
at which point all water has been consumed. Two methods for adding iodine are utilized,
coulometric titration were iodine is electrochemically supplied or volumetric titration
29 where iodine is directly supplied. Coulometric titration is more precise than volumetric
titration with up to 1 ppm resolution, especially for low water content samples, so all water
contents reported here utilized coulometric KFT.56
A typical coulometric KFT setup consists of an anolyte compartment and a
catholyte compartment separated by a diaphragm. The anolyte solution consists of a
combination of an alcohol, a base, sulfur dioxide, and potassium iodide.55,56 A platinum
anode oxidizes potassium iodide to elemental iodine which is available to react with sulfur
dioxide and water. Once water is consumed, iodine rapidly accumulates in the anolyte
available to react at the electrodes. Current flow between the cathode and anode which was
initially zero increases marking the equivalence point. The amount of water present in the
sample is then proportional to the amount of charge released forming iodine. In this work,
water contents were measured with a Metrohm Coulometric KF 889 D with a resolution of
0.1 µg water.
3.3.2 Density
Density can be accurately measured with vibrating-tube density meters. While other methods exist, few are as accurate as vibrating-tube density meters which can achieve accuracies of ± 0.000007 g/cm3. Figure 3.3 illustrates the setup of a vibrating tube density meter.
30
Figure 3.3 Schematic of the setup of a vibrating U-tube density meter.
A U-shaped tube housed inside a temperature control unit is filled with a small
volume of sample on the order of 1 mL. To ensure the accuracy of the measurement, it is important that the tube is completely filled with sample and that the sample is free of air bubbles. The tube is fixed at both ends of the U and vibrated with a piezo electric element.
The instrument measures the frequency of the vibrations of the tube. This frequency is related to the density of the fluid by modeling the tube sample system like a mass and
spring. In this work, density was measured with an Anton Paar DMA 4500M with an
accuracy of ± 0.03 K for temperature and ± 0.00005 g/cm3 for density. Solvent density has
implications on the eventual energy/power density of devices. Additionally, density is
commonly measured to provide a first principles glimpse into molecular level interactions
between solvents and solutes by investigating excess molar properties.
3.3.3 Viscosity
Viscosity can be measured with a wide variety of rheometers. Regardless of the type of instrument, viscometers/ rheometers rely on a well-defined fluid flow profile. Fluid is
31 forced to flow in a known profile and the instrument measures a resulting shear stress or
pressure drop.57 The measured shear stress or pressure drop is related to viscosity through
a number of empirical relations depending on the exact flow conditions. Temperature
control is typically achieved through use of a control unit designed to work with the particular viscometer.
Temperature-dependent viscosity can be fit to a number of different models.
Arrhenius and VFT models are among the most commonly used models to predict
viscosity.54, The Arrhenius expression is given in equation 1:
= equation 1 𝐸𝐸𝑎𝑎 𝜂𝜂 𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴 �𝑅𝑅𝑅𝑅� where is temperature-dependent viscosity (cP), is the pre-exponential constant (cP),
is the𝜂𝜂 activation energy for viscous flow (J mol𝐴𝐴-1), is the gas constant (J mol-1 K-1) ,
𝑎𝑎 and𝐸𝐸 is temperature (K).58 While the Arrhenius 𝑅𝑅relation is sufficient for narrow
temperature𝑇𝑇 ranges, the VFT model is preferred for wider temperature ranges54 and is
known to be more accurate for glass-formers like ILs.59,60 The VFT model utilizes three-
parameters as given in equation 2:
= equation 2 𝐵𝐵 𝜂𝜂 𝜂𝜂0𝑒𝑒𝑒𝑒𝑒𝑒 �𝑇𝑇−𝑇𝑇𝑜𝑜� where is temperature-dependent viscosity (cP), is the pre-exponential constant (cP),
0 is the𝜂𝜂 activation temperature (K), and is the ideal𝜂𝜂 glass-transition temperature (K).
𝐵𝐵 𝑇𝑇𝑜𝑜 The viscosity of an electrolyte has direct impacts on device performance. More
viscous electrolytes are more resistant to mass and ion transport, increasing the
overpotential. Overpotentials are simply inefficiencies associated with processes such as
32 ion movement and electron transfer that result in greater charging voltages and lower discharging voltage then theoretically predicted. Ideally, electrolytes are minimally resistant to mass and ion transport such that little energy is spent pushing reactants through the electrolyte to the electrode surface during charge and discharge. Viscosity also has important implications in the case of redox flow batteries where electrolyte is continuously pumped through the electrochemical cell. As electrolyte becomes more viscous and
resistant to flow, more energy is needed to pump the electrolyte. Although pumping energy
is expected to be minimal,61 greater pumping energy results in lower overall efficiencies
for the redox flow battery.
3.3.4 Conductivity
Conductivity is commonly measured by electrochemical impedance spectroscopy (EIS)
using a potentiostat equipped with a frequency response analyzer and a two-electrode cell
that are separated by a known distance. To determine conductivity as a function of temperature, the cell can be placed in an oil bath or on an aluminum block. Impedance is
determined by applying a small potential wave of varying frequency to the system and measuring the current response. When performing EIS, it is important to ensure a linear response between potential and current.62 For DESs, this typically means the frequency of
the applied potential wave must remain low (kHz range or lower). The lower the
conductivity of the DES, the lower the maximum allowable frequency of the potential
wave. Plots of the imaginary component of impedance (Z’’) versus the real component of impedance (Z’), Nyquist plots, can be generated from EIS spectra as illustrated in Figure
3.4
33
Figure 3.4 Example Nyquist plot with linear fit. Intercept with the real axis (Z’) represents solution resistance.
Solution resistance is the intersection of complex impedance data with the real impedance axis in Figure 3.4. A linear fit to a Nyquist plot can be used to extrapolate the intersection as show by the red line. Conductivity is calculated by dividing the cell constant of the two-electrode cell by the solution resistance. The cell constant is determined by measuring the resistance of a standard solution with a well-established conductivity.22
Electrolyte conductivity is especially important to consider when designing an electrochemical device as it has a direct impact on the ohmic overpotential. Conductivity is inversely proportional to ohmic overpotential so the more conductive the electrolyte the lower the ohmic overpotential and associated inefficiencies. As previously stated, DESs and ILs are typically orders of magnitude less conductive then traditional solvents due to
34 the low mobility of ions. Maximizing ion mobility and therefore electrolyte conductivity
is important for the eventual application of DESs and ILs.
3.3.5 Electrochemical Window
The electrochemical window (EW) of a DES is typically determined with slow scan cyclic
voltammetry (CV). CV is performed with a potentiostat and three electrode
electrochemical cell consisting of a working, counter, and reference electrode. Figure 3.5
illustrates two common three-electrode electrochemical cell setups. The T-Cell setup,
Figure 3.5 A, is commonly used when electrolyte volume needs to be minimized for fast
screening.63 T-cell setup requires about 20 to 600 µL of electrolyte. A traditional three- electrode setup, Figure 3.5 B, typically utilizes more electrolyte. Reference electrodes
should be carefully chosen to provide reproducible results. Silver/silver chloride reference
electrode was chosen for DESs containing choline chloride but must be checked against an internal reference for potential drift over time.64 Ferrocene/ferrocenium redox couple can
be used as an internal reference.63
When conducting a CV experimenting, a potential is applied to the working
electrode with respect to the reference electrode. The potential is scanned between a
minimum and maximum at a constant scan rate; 5 mV/s was used in this study for
electrochemical window determination. The current response between the working and
counter electrode is measured as a function of applied potential. To determine the potentials
at which the electrolyte begins breaking down, a cut off current density is arbitrarily
chosen, typically between 0.01 and 1 mA/cm2. The electrochemical window of DESs
35 changes depending on the material of the working electrode. For example, platinum is
catalytic to many different redox reactions so EWs are typically smaller when evaluated
with platinum working electrode as compared to other materials like glassy carbon.65
Figure 3.5 Schematic of a T-Cell setup (A) and a traditional three electrode setup (B). The reference electrode is insulated to prevent accidental contact with the coiled counter wire electrode in the T-cell setup.
The EW of an electrolyte is critically important to the design of an electrochemical
device. In many cases, the operating voltage of the device is limited by the electrolyte. For example, none of the known DESs can be used in a lithium-ion or a lithium metal battery because their EW is too narrow, typically 2 to 3 V. Lithium-ion batteries operate over 3.5
V which is outside the stability window of many electrolytes including DESs and aqueous systems. The EW of ILs by contrast is wide enough for lithium-ion batteries. Thus,
determining EWs is important for establishing potential applications for electrolytes.
36 3.3.6 FTIR Spectroscopy
FTIR is an infrared (IR) spectroscopy technique used to identify species and molecular
level interactions such as hydrogen bonding. Compared to other IR techniques such as
dispersive IR spectroscopy, FTIR is faster with better accuracy and resolution. FTIR relies
on determining the amount of infrared light absorbed by a sample at different frequencies.66
As incoming light interacts with the sample, bonds are stretched and compressed depending
on the energy of the light. If the stretching or compressing vibration of the bond results in
a change in the dipole moment of the molecule then energy is absorbed and the vibration
is visible in IR spectra. The wavenumber location of a particular vibrational mode can be
roughly predicted by approximating a bond like a ball and spring system where the atoms
represent balls and the bond represents a spring. IR has previously been used to investigate
molecular interactions in DESs especially hydrogen bonding.67
When investigating DESs with FTIR, it is especially important to consider the water
content of the sample. Hydrogen bonds between water molecules appear as a strong, broad
signal in the 3500-3000 cm-1 region.68 This signal can be so strong and broad that it washes out signal from other vibrational modes in the region. This can be avoided by properly preparing and storing DES samples as discussed previously such that water contents are minimized. FTIR requires a spectrometer and can be operated in variety of modes including transmission and attenuated total reflectance (ATR). FTIR-ATR is more suited to investigating liquids and thus it is the most commonly used technique for DESs.69 FTIR-
ATR only requires a single drop of sample placed on an ATR crystal. All FTIR spectra
reported here were collected with a Nicolet iS50 FTIR (Thermo Scientific, USA) with a
diamond ATR crystal.
37 Chapter 4: Properties of Glyceline and Ethaline Deep Eutectic Solvents
Due to the diversity of species which can be used to form DESs, understanding of the impact chemical structure and concentration on properties is limited. Here, fundamental physical properties of two DESs are investigated: (1) ethaline which is choline chloride
(ChCl) and ethylene glycol (EG) in a 1 to 2 molar ratio and (2) glyceline which is choline chloride and glycerol (G) in a 1 to 2 molar ratio. In addition, solutions away from the DES composition are investigated to determine the impact of concentration on physical properties and to gain insight into DES “formation”. We find that addition of ChCl to G decreases viscosity and increases conductivity. Addition of ChCl to EG results in more complex behavior where viscosity increases as ChCl added but conductivity increases as
ChCl is added until a critical concentration after which addition of more ChCl results in a decrease in conductivity.
This study was a part of a collaboration with the groups of Profs. Ed Maginn (U of
Notre Dame), and Joshua Sangoro (U of Tennessee, Knoxville) and led to two manuscripts that are currently under review in the Journal of Physical Chemistry B and the Journal of the American Chemical Society.14,15 Specifically, measured densities and viscosities were used in the development and validation of the force field specific to these DESs in molecular dynamic simulations performed by Maginn group. The development of force field was critical in enabling predictions for other properties such as liquid structure, diffusivities, and coordination in the bulk liquid. Furthermore, the measured conductivities were used by Sangoro group to confirm the dynamical properties measured by broadband dielectric spectroscopy (BDS). BDS measurements were essential to understand the
38 dynamics of the liquid and specifically the dynamics of the hydrogen bonds in DESs that correlate with macroscopic viscosities and conductivities.
4.1 Introduction
In designing a DES, the two most basic aspects to be considered are the species to be used
(ethylene glycol, glycerol, urea, etc.) and the mole fraction of each species (33 mol% HBA,
25 mol% HBA, 20 mol % HBA, etc.).70 Reline, which is a mixture of ChCl and urea at a molar ratio of 1 to 2 respectively, demonstrates a remarkable depression in freezing point
(FP) (below room temperature) compared to ChCl and urea which are solids with FPs above 150 ºC.71 Away from the molar ratio of 1 to 2 (33 mol% ChCl, 77 mol% urea), the mixture is no longer a DES and does not necessarily even form a liquid at room temperature. Mixtures of DES components which are away from the eutectic composition are referred to as DES analogues from here on.
As previously discussed, understanding of structure-function relation is lacking in
DESs.12 Additionally, the distinction between interactions found in DESs and DES analogues is not understood. While many studies have hypothesized the origin of FP depression, there is no conclusive evidence demonstrating the origin of FP depression.72
Typically, FP depression in DESs is attributed to the hydrogen bonding network, which simply refers to the complex inter- and intramolecular forces present in DESs. To gain insight into molecular interactions in DESs, we investigated DES analogues in parallel.
Previous work by Abbott et al. has established FP as a function of HBD or HBA mole fraction, generating phase diagrams for mixtures of the HBA choline chloride and
39 HBDs such as oxalic acid, malonic acid, and succinic acid.24 Studies have also investigated variations in physical properties such as conductivity, density, and viscosity as a function of DES composition for ammonium and phosphonium based DESs.29 Another approach to
elucidating fundamental DES interactions involves addition of a component such as water
or alcohol to the DES and studying corresponding excess property changes.73,49
Here, the physical properties, density, viscosity, and conductivity of glyceline and
ethaline are investigated. Glyceline and ethaline are widely studied DESs. ChCl based
DESs are among the least viscous and are therefore of particular interest for application in
electrochemical devices.9 In addition to ethaline and glyceline (both with 33 mol% ChCl),
DES analogues composed of 20, 10, 5, and 0 mol% ChCl are also characterized. By
maintaining identical HBA, slightly varying the HBD structure through the addition of a
hydroxyl group, and varying solution compositions, we aim to develop understanding of
structure-function relation and fundamental molecular interactions found in DESs and DES
analogues. Previous work has investigated the viscosity of various mixtures of ChCl and
G but without proper preparation precautions to prevent water uptake.70
Figure 4.1 Structures of investigated hydrogen bond donors (HBDs), ethylene glycol (EG)
and glycerol (G), and hydrogen bond acceptor (HBA), choline chloride (ChCl).
40 4.2 Methods and Materials
Choline chloride (99%), anhydrous ethylene glycol (99.8%), and glycerol (99%) were
obtained from Acros Organics. All mixtures were prepared and characterized according
the standard procedure outlined in Chapter 3 subsection 1.12 The water contents of each
sample were measured by coulometric Karl Fisher titration (Metrohm Coulometric KF
889D) and are reported in Table 4.1. Physical property data was collected as previously discussed. 12,13 The range of frequencies used in EIS to determine solution conductivities are reported in Table 4.1. Density, viscosity, and conductivity experiments were replicated
three times at a given temperature for each solvent system and the average is reported.
Table 4.1 Water contents of studied DESs (33 mol% ChCl) and DES analogues, and frequency range of EIS used for the calculation of ionic conductivities.
41 4.3 Results and Discussion
Temperature-dependent viscosity for different ChCl in G solutions are reported in Figure
4.2. As the mol% of ChCl increases, the viscosity of ChCl:G solutions decrease. Addition of just 5 mol% ChCl drops the viscosity of G from 921 cP to 715 cP at 25 ºC. The 33%
ChCl in G solution has the lowest viscosity at 371 cP at 25 ºC. Addition of more ChCl has a less profound impact on viscosity at higher mol% of ChCl as seen in Figure 4.2 B. All solutions exhibit an Arrhenius like relationship between viscosity and temperature. Fitted parameters of the Arrhenius equation, and , the pre-exponential constant and
𝑎𝑎 activation energy, respectively, are reporte𝐴𝐴d for each𝐸𝐸 solution in Table 4.2.
Figure 4.2 Temperature-dependence of viscosity for mixtures of ChCl and G (A). Lines are Arrhenius fits. Viscosity as a function of mol% ChCl in G at 298.15 K (B). Results from ref 15.
42 Table 4.2 Fitted Arrhenius parameters and for the mixtures of ChCl and G.
𝐴𝐴 𝐸𝐸𝑎𝑎
The fitted Arrhenius parameter is a representation of the energy required to force
𝑎𝑎 flow and as expected the solution with𝐸𝐸 the greatest viscosity, 0% ChCl:G, has the greatest
while the solution with the lowest viscosity, 33% ChCl:G, has the lowest . These
𝑎𝑎 𝑎𝑎 𝐸𝐸activation energies are on the same order of magnitude as those previously reported𝐸𝐸 for
DESs and ILs.49 At higher temperatures the difference in viscosity of solutions is less
pronounced, Figure 4.2 A, due to increased molecular motion which dominates
intermolecular interactions between ChCl and G.
Temperature-dependent viscosity for solutions of ChCl in EG are reported in
Figure 4.3. ChCl in EG solutions are about an order of magnitude less viscous then ChCl in G solutions. In contrast to ChCl in G solutions, the viscosity of ChCl in EG solutions increase as the mol% of ChCl increases. The viscosity of neat EG is 19 cP at 25ºC while the viscosity of 33% ChCl in EG is 48 cP at 25ºC. Addition of 5% and 10% of ChCl has little impact on viscosity but addition of 20% and 33% greatly reduces fluidity, Figure 4.3
B. Similar to ChCl in G solutions, the viscosity of ChCl in EG solutions shows an
Arrhenius like dependence on temperature, Figure 4.3 A. The fitting parameters and
𝐴𝐴 𝐸𝐸𝑎𝑎 43 are reported in Table 4.3. As expected, the activation energy for viscous flow is less for
ChCl in EG solutions then ChCl in G solutions.
Figure 4.3 Temperature-dependent viscosity of mixtures of ChCl and EG (A). Lines are fits to Arrhenius model. Viscosity as a function of mol% ChCl in EG at 298.15 K (B).
Results from ref 14.
Table 4.3 Fitting parameters and for temperature-dependent viscosity of ChCl in EG.
𝐴𝐴 𝐸𝐸𝑎𝑎
44 ChCl in G and ChCl in EG display different behaviors as the DES composition, 33 mol% ChCl, is approached. Whereas the viscosity of ChCl in G plummets with addition of
ChCl, the viscosity of ChCl in EG rapidly increases with the addition of the salt. This is likely explained by the hydrogen bonding behaviors of the neat liquids, EG and G. EG is a diol (recall Figure 4.1) taking either the trans or gauche conformation.74 In the trans position, EG forms a linear hydrogen bonding network between distinct EG molecules. In the gauche conformation, the hydroxyl groups of an EG molecule primarily interact with each other, minimally interacting with separate EG molecules. In contrast, G is a triol which forms a highly ordered, three-dimensional hydrogen bonding network.70 This network is observable in the high viscosity of G, 921 cP at 25ºC. Adding ChCl disrupts this long-range structure as hydroxyl groups complex chloride anions decreasing viscosity.
Addition of ChCl to EG however results in greater long-range order as the chloride anion complexes with hydroxyl groups thus increasing viscosity.
Interactions between ChCl and the HBD, EG or G, are further probed through temperature-dependent conductivity measurements, Figure 4.4. In general, ChCl in EG solutions are more conductive then ChCl in G solutions. This is expected due to the high viscosity of ChCl in G solutions as compared to ChCl in EG solutions.
45
Figure 4.4 Temperature-dependent conductivity of mixtures of ChCl in G (A) and ChCl
in EG (B). Lines indicate Arrhenius fits. Results from ref 14 and 15.
Fitted Arrhenius parameters for conductivities are reported in Table 4.4. Across all
solutions, conductivity increases with increasing temperature. The ionic conductivity of a
solution is dependent on the number of mobile ions and the mobility of these ions. The
number of ions in solution increases with increasing ChCl content whereas the mobility of ions in solution is best represented by viscosity.
46 Table 4.4 Fitting parameters and for temperature-dependent conductivity of ChCl in
𝑎𝑎 G and ChCl in EG solutions. 𝐴𝐴 𝐸𝐸
As previously discussed, the relationship between ChCl content and viscosity
depends on the HBD species. For ChCl in G solutions, as ChCl content increases,
conductivity increases as seen in Figure 4.4 A. This is expected as both the number of ions and the mobility of ions, recall Figure 4.2, increases with increasing ChCl mol%. Solutions of ChCl in EG show a more complex behavior, conductivity increases with increasing ChCl mol% until a critical composition of 20% ChCl in EG. Beyond this point increasing ChCl content results in a decrease in conductivity; the DES composition, 33% ChCl in EG, is less conductive then 20% ChCl in EG. This is explained by a trade-off in number of ions and mobility of ions. At 25ºC, the 33% ChCl in EG solution is 20 cP more viscous then the
20% ChCl in EG solution. While 33% ChCl in EG contains more ions then 20% ChCl in
EG, ions in the DES composition are much less mobile then in the 20% composition resulting in the former being less conductive.
47 Finally, temperature-dependent density data is used to further evaluate the impact of the HBD and ChCl content. Figure 4.5 reports densities of ChCl in G and ChCl in EG as a function of temperature. Increased molecular motion at higher temperatures results in molecules occupying a greater volume and correspondingly lower density. In general, ChCl in G solutions are denser than ChCl in EG solutions across all compositions. At 25ºC, neat
G has a density of 1.25795 g/cm3 while neat EG has a density of 1.10988 g/cm3. The greater density of G as compared to EG is attributed to highly ordered hydrogen bonding network of G. As ChCl is added to G, density decreases linearly with increasing ChCl mol%, Figure
4.5 B. This is in agreement with the conclusions drawn from viscosity data. Addition of
ChCl disrupts the highly ordered hydrogen bonding network of G as chloride anions coordinate with hydroxyl groups, resulting in decreased viscosity. Additionally, this coordination complex does not have the volume economy of neat hydrogen bonded G molecules resulting in lower density with greater ChCl content.
Addition of ChCl to EG results in the opposite trend, density increases with increased ChCl mol%, Figure 4.5 D. As discussed previously, EG has intermolecular bonds between hydroxyl groups or intramolecular bonds between separate molecules forming linear aggregates. Similar to ChCl in G, chloride anions complex with the hydroxyl groups of EG. However, in contrast to ChCl in G, this coordination complex is denser than the neat hydrogen bonded EG. Additionally, recall the viscosity of ChCl in EG solutions increased with greater ChCl content. Taken together, it is probable ChCl results in a greater order in EG. In contrast, ChCl decreases order in G.
48
Figure 4.5 Temperature-dependent density of various mol% ChCl in G (A). Density as a function of ChCl mol% in G at 25 ºC (B). Temperature-dependent density of various mol%
ChCl in EG (C). Density as a function of ChCl mol% in EG at 25 ºC (D). Lines are linear fits. Results from ref 14 and 15.
Temperature-dependent density is well predicted by a linear fit: = + where
is temperature-dependent density in g cm-3, is the temperature in K,𝜌𝜌 and𝑎𝑎 𝑏𝑏and𝑏𝑏 are
𝜌𝜌fitting parameters. As temperature increases, density𝑇𝑇 decreases for both ChCl𝑎𝑎 in G𝑏𝑏 and
ChCl in EG for all compositions; the term is always negative. The fitting parameters
and are reported in Table 4.5. 𝑏𝑏𝑏𝑏 𝑎𝑎
𝑏𝑏
49 Table 4.5 Fitted parameters of and for temperature-dependent densities of ChCl in G and ChCl in EG solutions. 𝑎𝑎 𝑏𝑏
4.4 Conclusion
Temperature-dependent viscosity, conductivity, and density have been measured for choline chloride-based DESs ethaline and glyceline. The solutions away from the DES composition were also studied. Solutions of ChCl in G were found to be an order of magnitude more viscous and less conductive then ChCl in EG solutions. Addition of ChCl to G resulted in decreased viscosity and increased conductivity. In contrast, increasing
ChCl content in EG resulted in increased viscosity. Conductivity increased up until 20%
ChCl in EG after which decreasing ion mobility resulted in decreased conductivity with greater ChCl content. Finally, ChCl in G solutions were denser than ChCl EG solutions.
Taken together, addition of ChCl results in greater order in EG but decreased order in G.
Simply by varying the number of hydroxyl groups, significant changes in physical
50 properties are observed. HBDs which display highly ordered intermolecular hydrogen bonds such as triols are likely to display improved conductivity and viscosity with addition of HBAs. In contrast, less ordered systems like diols are likely to display more complicated behavior where a tradeoff between ion mobility, limited by formation of coordination complexes, and number of ions must be balanced.
51 Chapter 5: Solvation of Methylviologen Dichloride in an IL, DES, and DES Analogues
Fundamental solvation and transport properties of DESs are not understood due to the
complexity of molecular interactions. Here, solvation and transport of methlyviologen
dichloride (MVCl2) in 1:2, 1:4, and 1:6 molar mixtures of choline chloride (ChCl) and
ethylene glycol (EG), including the DES ethaline (1:2 mixture), were investigated with
physical property determination, cyclic voltammetry, and Raman spectroscopy. Solvation
and transport in ChCl:EG mixtures was compared to ChTFSI:EG mixtures in order to
evaluate the impact of the anion and hydrogen bond donor interactions. As a reference, the
IL, 1-methyl-1-propylpyrrolidinium bis(trifluoromethylsulfonyl) imide, [PYR13][TFSI]
was also investigated.
Exchanging the chloride anion of ethaline with TFSI is found to increase the fluidity
of the solvent and promote stronger solute-solvent interactions. Raman spectroscopy suggests MVCl2 is strongly solvated by EG in ChTFSI:EG solutions but interstitially
accommodated in holes in ChCl:EG and [PYR13][TFSI]. Anion exchange is demonstrated
to be a potential pathway for tuning the solvation environment of DESs. This work was
completed in collaboration with Jeffrey M Klein and William Dean and is under review in
the Journal of Physical Chemistry B.16
5.1 Introduction
Unlike classical solvents, molecular transport in ILs and DESs is hypothesized to occur
through a hopping mechanism where species move through empty holes in the solvent.52,75
Hole theory describes the formation of holes in liquid structures as a result of thermally
52 driven fluctuations in local density.76,77 Transport of solutes in ILs and DESs has been suggested to occur by the same mechanism.52 Studies probing the bulk liquid structure and structure-property relations, though limited in DESs as compared to ILs, support the notion that these solvents are functionally similar.
ILs are heterogeneous liquids with implications on solute diffusion mechanisms.52-
78,79 Hayes, Warr, and Atkin, summarize the heterogeneous nature of ILs as structured
solvents with the bulk liquid containing ion pairs, ion clusters, H-bond networks, and
micellular domains.80 As a consequence of heterogeneity, microviscosities orders of
magnitude lower than the bulk viscosity have been reported in tetraalkylammonium and
pyrrolidinium based ILs.81 Nanoscale heterogeneity in ILs was first reported by Triolo et al.82 through the application of XRD; it was suggested that diffusion processes significantly
affect the morphology of the nanodomains in the IL.82 The complexity of IL solvation
manifests spectroscopically where multiple time scales of solvent relaxation are observed,
associated with internal ion motions (fast) and solvent diffusive reorganization
(slow).79,83,84 Computational and experimental investigations using techniques such as
cyclic voltammetry75 and pulsed-field gradient NMR85 have shown self-diffusion and solute diffusion do not follow the Stokes-Einstein relation, an additional consequence of
IL heterogeneity.86,87 Hu and Margulis were able to relate the red-edge effect to the time a solvent molecule is trapped in local “quasistatic solvent cages”, further suggesting the solvation structure of ILs play a significant role in solute diffusion.88
In contrast to ILs, generalizations about the liquid structure of DESs cannot be
made, certain DESs display heterogeneities while others do not due to the diversity of the
species which form DESs.89,90,91 For example ethaline, a low-viscosity DES composed of
53 choline chloride (ChCl) and ethylene glycol (EG) in a 1:2 molar ratio respectively, has been shown to be dynamically heterogeneous but spatially homogenous.92 In contrast to
ethaline, Kaur et al. have shown nanoscale segregation (spatial heterogeneities) in DESs
composed of lithium perchlorate and alkyamides.17 Similar to ILs, solvation, and transport
has been shown to be complex in certain DESs.70 Physical property measurements suggest solutes can be interstitially accommodated in holes or actively participate in local bonding networks depending on the nature of the DES and solute.49,73 Pulsed-field gradient NMR
studies have shown that self-diffusion in ethaline does not follow classical Stokes-Einstein
behavior, similar to ILs.52 It is clear that the DES intermolecular hydrogen bonding network, plays a significant role in solvation and transport properties of DESs.
While a plethora of studies have investigated DES formation and properties, the influence of the hydrogen bond donor/acceptor structure and concentration on solvation and transport properties is not understood. This study investigates the properties of
ChCl:EG mixtures (1:2, 1:4 and 1:6 molar ratios) and the impact of replacing the Cl anion with TFSI. As reference, the solvation properties of the ChCl:EG and ChTFSI:EG mixtures are compared to that of n-methyl-n-propylpyrrolidinium bis(trifluoromethylsulfonyl)imide, [PYR13][TFSI], an IL as an extreme case of a liquid
lacking neutral solvent molecules. The molecular structures of the studied solvents are
shown in Figure 5.1.
54
Figure 5.1 Structures of the solvent systems studied: [PYR13][TFSI] – an IL(left),
ChCl:EG – DES components (middle), ChTFSI:EG – DES analogue (right). From ref 16.
[PYR13][TFSI] is a widely studied IL especially in electrochemical systems due to
its wide electrochemical window and relatively low viscosity compared to other ILs.93
Ethaline is among the least viscous and most conductive of commonly studied DESs.24,94,95
By replacing the Cl anion with TFSI, the impact of anion structure on physical properties
and solvation environments is investigated. In ethaline, the primary intermolecular
interaction occurs between EG, the hydrogen bond donor, and Cl anions.94 Thus, both the
nature and the solvation environment of the DES should be fundamentally changed through
TFSI substitution of the Cl anion. It should be noted that only the 1:2 ChCl:EG mixture is
formally defined as a DES. Since the 1:4 and 1:6 ChCl:EG mixtures and all ChTFSI:EG
mixtures are not confirmed to be DESs, the term “DES analogue” is used to describe these
systems instead.
2+ Methylviologen dichloride (MVCl2 or for simplicity MV ) is used as an
electrochemical probe in this study. MV2+ has been extensively investigated as a redox
active species in flow batteries96, 97 and as an excited state fluorescent probe.98 MV2+
55 undergoes a two-stage reduction and oxidation process following the sequence given in
Figure 5.1.98,99
Figure 5.2 Two-step reduction mechanism of methylviologen, MV2+, in acetonitrile as
reported by Kohler et al.98 From ref 16.
Here, the measured bulk properties for the three electrolyte systems in Figure 5.1
including viscosity, conductivity, density and surface tension are reported. Bulk properties
are used to estimate the mobility of ions within each solution while cyclic voltammetry is
used to estimate diffusion coefficients and solvated radii of MV2+. Finally, Raman
spectroscopy is used to estimate the strength and nature of interactions between the MV2+
solute and solvent system.
5.2 Methods
Choline chloride, ChCl, (99%), anhydrous ethylene glycol, EG, (99.8%), and
methylviologen dichloride hydrate, MVCl2, (98%) were purchased from Acros Organics.
N-methyl-n-propylpyrrolidinium bis(trifluoromethylsulfonyl)imide, [PYR13][TFSI],
(99%) was purchased from Iolitec. Nitrogen gas (99.998%) was purchased from Airgas.
56 Silver foil (0.25 mm thick, both 99.998%) and silver wire (0.5 mm and 1 mm in diameter,
99.9%, and 99.999% respectively) were purchased from Alfa Aesar. Potassium chloride
(99%) was purchased from Fisher Scientific.
5.2.1 Preparation
Preparation of ethaline DES and DES analogues, and the characterization of the solvents
were performed according to the standard techniques previously described12 and detailed
in Chapter 3. The water contents of the samples were measured before each experiment and are reported in Table 5.1.
5.2.2 Surface Tension
Surface tension experiments were conducted by William Dean. A Krüss drop shape analyzer was used to perform pendant drop measurements to determine the surface tensions of all the mixtures of ChCl:EG and ChTFSI, as well as the IL, [PYR13][TFSI]. A one mL
syringe with a 1.8 mm diameter blunt tip needle was used to ensure a large drop size. An
initial drop was used to determine the break-off point for each liquid. Measurements were
then performed on a second drop formed to just before the break-off point. The average radius of holes in the DESs and IL are calculated from measured surface tension ( ) using equation 1: 𝛾𝛾
4 ( ) = 3.5 equation 1 2 𝑘𝑘𝐵𝐵𝑇𝑇 𝜋𝜋 𝑟𝑟 𝛾𝛾
57 where r is the average hole size, is the Boltzmann constant, and T is the absolute
𝐵𝐵 temperature.76 𝑘𝑘
5.2.3 Cyclic Voltammetry
Cyclic voltammetry (CV) experiments were conducted by Jeffrey Klein. Experiments were
performed with a T-cell100 consisting of a glassy carbon working electrode (3 mm dia,
BASi), silver wire counter electrode (1 mm dia), and silver wire quasi-reference electrode
(1 mm dia). Silver wire quasi-reference was used as silver-silver chloride reference is not suitable for ILs. All measurements were conducted in a positive pressure nitrogen glove box (Terra Universal). The electrodes were polished using the procedure previously described.101 All CVs began at open-circuit voltage which was near 0 V vs Ag wire for all
solutions. Scan rates of 1, 5, 10, 15, 20, and 50 mV/s were used for all systems starting at
1 mV/s for 5 cycles and moving up sequentially performing 5 cycles at each scan rate. The
CV’s were iR corrected using the resistance determined from a single frequency (100 kHz)
impedance measurement.
The peak current of each cyclic voltammogram was used to estimate the diffusion
coefficients of MV2+ using the Randles-Sevcik (RS) equation, equation 2:
= 0.4463 1 1 1 equation 2 2 𝑛𝑛𝑛𝑛 2 2 𝑖𝑖𝑝𝑝 𝑛𝑛𝑛𝑛 �𝑅𝑅𝑅𝑅� 𝐴𝐴𝐷𝐷 𝐶𝐶𝜈𝜈 where is the peak current of the first reduction measured in CV, is the number of
𝑝𝑝 electrons𝑖𝑖 transferred in the redox event, is Faraday’s constant, is𝑛𝑛 the electrode area
(cm2), is the diffusion coefficient of the𝐹𝐹 redox species (cm2 s-1),𝐴𝐴 is the concentration
𝐷𝐷 𝐶𝐶
58 of redox species (mol cm-3), is the scan rate (V s-1), is the gas constant (J K-1 mol-1),
and is the temperature (K). 𝑣𝑣 𝑅𝑅
𝑇𝑇 Equation 2 assumes that the concentration of the solute at the electrode surface is
equal to the bulk concentration and that the species are freely diffusing. These assumptions
may not be necessarily true if the solubilities of reduced and oxidized species are
dramatically different. Fitting a line to vs / for the entire sequence of scan rates 1 2 𝑝𝑝 resulted in a near-perfect linear fit. However,𝑖𝑖 during𝑣𝑣 experiments, a blue color developed
near the working electrode which is believed to be due to the accumulation of MV+, the reduced radical cation. Therefore, diffusion coefficients for MV+2 were calculated at 1
mV/s (the first scan rate of the applied sequence). The radius of the solvation structure was estimated from the Stokes-Einstein relationship in equation 3 using the obtained diffusion
coefficients, , and the measured viscosities, .
𝐷𝐷 𝜂𝜂 = equation 3 𝑘𝑘𝐵𝐵𝑇𝑇 𝐷𝐷 6𝜋𝜋𝜋𝜋𝜋𝜋 where is Boltzmann's constant (m2 kg s-2 K-1), is temperature (K), is the solvated
𝐵𝐵 radius, 𝑘𝑘and is viscosity (kg m-1 s-1). 𝑇𝑇 𝑟𝑟
𝜂𝜂
5.2.4 Raman Spectroscopy
Raman spectra were collected using a Horiba Xplora One Raman system with a 785 nm
excitation laser and a spectral range of 100 to 3500 cm-1. The spectra were collected using
a 20 x magnification objective, 1000 um hole, 200 um slit, 10 s exposure time, and 10
accumulations. In order to amplify the Raman signals surface enhancement was employed.
59 An Ag substrate (0.25 mm thick Ag foil), was electrochemically roughened by adapting
the procedure by Tian et al.102 Briefly, a 3 potential step sequence was used starting at -
0.25 V (vs Ag wire quasi reference) for 15 s followed by 0.25 V for 8 s, and finally -0.25
V for 60 s in 0.1 M KCl aqueous electrolyte. The electrochemical setup for the substrate
fabrication and Raman measurements are described elsewhere.13
5.3 Results and Discussion
The temperature dependent viscosity and conductivity of each solvent system are shown
in Figure 5.3. [PYR13][TFSI], the IL, is the most viscous solvent (59 cP at 25 °C) followed by 1:2 ChCl:EG (49 cP at 25 °C) while 1:6 ChTFSI:EG (22 cP at 25 °C) was the least
viscous as seen in Figure 5.3 A. Substituting Cl with TFSI results in a significant decrease in viscosity for solutions of identical molar ratios. For example, at 25°C the viscosity of
1:2 ChTFSI:EG was measured to be 31 cP compared to 49 cP for 1:2 ChCl:EG. According
to hole theory, solvents composed of smaller ions should have lower viscosities. Viscosity
tracks with the size of ions because as ion size shrinks the probability of an ion moving
through an appropriately sized hole increases.76 Thus, it might be surprising that
ChTFSI:EG solutions are less viscous then ChCl:EG solutions; however, due to strong
intermolecular interactions between EG and Cl, molecular motion is limited in ethaline.
Charge delocalization in the TFSI anion weakens interactions with both the Ch cation and
EG, promoting decreased viscosity. As the EG content increases in ChTFSI:EG mixtures,
fluidity approaches the viscosity of pure EG (16.1 cP at 25 °C).103 Introduction of 10 mM
of MVCl2 slightly increases the viscosity of both [PYR13][TFSI] and 1:2 ChTFSI:EG by 5
cP and 3 cP at 25ºC respectively, while the viscosity of 1:2 ChCl:EG remains unchanged.
60 Although the concentration of MVCl2 is very small, solute-solvent interactions in
[PYR13][TFSI] and 1:2 ChTFSI:EG might result in larger solvation complexes leading to
increased viscosity. In contrast, the general solvation complex in 1:2 ChCl:EG is not
impacted by MVCl2 perhaps due to the strength of Cl and EG interactions and thus viscosity
is relatively unchanged.
Figure 5.3 Temperature-dependent viscosity (A) and conductivity (B) of the solvent
systems from 298 to 328 K. The legend in the inset of panel A applies to both panels A and
B. Results from ref 14 and 16.
Conductivity, shown in Figure 5.3 B, demonstrates a noticeable tradeoff between
concentration and mobility of ions in solution resulting in the maximum value of
conductivity at 1:4 ChCl:EG. In this case, ChCl:EG systems are more conductive then both
[PYR13][TFSI] and ChTFSI:EG systems. Although ChTFSI:EG mixtures are less viscous
then ChCl:EG mixtures, the conductivity of ChTFSI:EG mixtures is about half of ChCl:EG
mixtures demonstrating that despite the greater fluidity of ChTFSI:EG, the sum of all ions
mobilities in ChCl:EG are on average greater.
61 Surface tension measurements were used to explore the relationship between hole size and viscosity in the reported systems. Surface tension and the calculated hole size for the solvent systems are shown in Table 5.1.
Table 5.1 Measured viscosity ( ) and surface tension ( ) at 25 °C and the estimated average hole size (r) using equation𝜂𝜂 1. 𝛾𝛾
Solvent (cP) (mN/m) r (Å)
104 105 EG 𝜼𝜼15 𝜸𝜸47.7 1.55 1:2 ChCl:EG 48 52.0 1.48 1:4 ChCl:EG 27 46.9 1.56 1:6 ChCl:EG 23 46.9 1.56 1:2 ChTFSI:EG 30 38.0 1.73 1:4 ChTFSI:EG 29 38.4 1.73 1:6 ChTFSI:EG 22 38.4 1.73
[PYR13][TFSI] 59 31.9 1.9
According to hole theory, the average hole size of a liquid should correlate with macroscopic viscosity. The larger the average hole size, the lower the viscosity of the liquid. An increase in the average hole size, relative to 1:2 ChCl:EG, is observed in the 1:4
ChCl:EG mixture, as expected this correlates with a significant decrease in viscosity. The
1:4 ChCl:EG and 1:6 ChCl:EG had identical surface tensions and hole size, essentially equal to that of pure EG. Correspondingly, the magnitude of decrease in viscosity from 1:4 to 1:6 ChCl:EG was less as compared to the viscosity drop from 1:2 to 1:4 ChCl:EG. In general, the average hole sizes of ChTFSI:EG mixtures were larger than the average hole sizes of ChCl:EG mixtures, but ChTFSI:EG mixtures did not have significantly lower
62 viscosities than ChCl:EG mixtures with the exception of the 1:2 molar ratios. This is likely
due to the drastic increase in anion size between Cl and TFSI. While the average hole size
is larger in ChTFSI:EG then ChCl:EG, the anion size of TFSI offsets the effect of the larger
average hole size. A similar phenomena is observed in [PYR13][TFSI] which exhibited the
largest average hole size but also the highest viscosity.
To further investigate the mobility of ions, Walden analysis was performed, shown
in Figure 5.4 A. This technique has previously been applied to ILs106,107 and DESs,108,109
to understand the relative extent of the ion associations. Figure 5.4 shows all systems fall
below the KCl reference (diagonal solid line representing an ideal non-interacting electrolyte). Data for 1:2 ChCl:EG and [PYR13][TFSI] are consistent with prior
literature.8,110 Diluting ChCl and ChTFSI with EG pushes respective lines further below
the non-interacting reference line due to strong intermolecular hydrogen bonding within
EG. ChCl in a 1:2 molar ratio disrupts this inherent EG-EG bonding and yields improved
mobility of ions. ChTFSI has a similar, but less dramatic, effect in the 1:2 molar ratio
shown by the line still falling far below the reference.
63
Figure 5.4 Walden plot for investigated solvents (A). Calculated diffusion coefficients versus inverse viscosity. Dotted lines represent fits to Stokes-Einstein relation, eqn 3. Two average sizes of radii are observed: one for concentrated systems (mixtures with 1:2 molar ratio of salt-to-solvent and the IL) and one for dilute systems (mixtures with salt-to-solvent molar ratios of 1:4 and 1:6) (B). Representative CVs of the systems at 15 mV s-1. Dotted line indicates the half wave potential of the first reduction of MV2+ in 1:2 ChCl:EG. The shift in reduction potential of MV2+ is due to the choice of reference electrode and should not be used to link to hydrogen bonding affects (C). Results from ref 16.
64 Transport of MV2+ (10 mM) in the electrolytes of interest was probed by CV and
the trends in the estimated diffusivity of MV2+ in the studied solvents is interpreted as the
consequence of the differences in the local solvation environment. The obtained
diffusivities from the Randles-Sevcik equation (equation 2) are plotted in Figure 5.4 B as
a function of inverse viscosity. Representative CVs at 15 mV/s for all systems are shown
in Figure 5.4 C. If all diffusion coefficients fell on a line from Stokes-Einstein analysis
then a single radius of solvation could be used to describe each system, however, it is
obvious that one line does not sufficiently fit the data in Figure 5.4 B. The two linear fits
(dashed lines in Figure 5.4 B) are for the concentrated systems (1:2 mol ratio and the IL)
and the dilute systems (1:4 and 1:6 mol ratios). The smaller slope of the fit to the
concentrated systems suggests MV2+ in these systems has a larger solvation radius than the dilute systems. The MV2+ solvation radius is nominally the same in 1:2 ChCl:EG, 1:2
ChTFSI:EG, and [PYR13][TFSI]. Estimated radii for the individual systems in descending order follow [PYR13][TFSI] (6.7 Å) > 1:2 ChTFSI:EG (6.4 Å) > 1:2 ChCl:EG (5.8 Å) >
1:4 ChCl:EG (4.4 Å) > 1:6 ChTFSI:EG (3.4 Å) > 1:6 ChCl:EG ~ 1:4 ChTFSI:EG (3.1 Å).
If MV2+ is indeed transported through the holes in the liquid, accompanied by its solvation sphere, then it would be expected that the system with the smallest solvation radius to hole size ratio would have the fastest diffusion of MV2. Indeed, an inverse relationship between
the ratio of solvation radius to hole size and diffusivity of MV2+ is seen in Figure 5.5. Even
in Figure 5.5 there appears to be an unaccounted factor such as the extent of dissociation
between MV2+ and Cl that is impacting both the size of holes and the diffusion of MV2+.
65
Figure 5.5 Relationship between the diffusion coefficient of MV2+ determined from eqn.
2 and the ratio of solvation radii (eqn. 3) to solvent hole size (eqn. 1). Results from ref 16.
In an attempt to bridge the gap between macroscopic properties and molecular level
interactions, Raman spectroscopy was used to evaluate solvent-solute interactions. Figure
5.6 shows the Raman spectra of 1:2 molar mixtures of ChCl:EG and ChTFSI:EG in
comparison to [PYR13][TFSI], pure EG, pure MVCl2 and MVCl2 in EG. Table 5.2 lists the identified vibrational modes for EG and MV2+ and the corresponding wavenumbers at
which those modes were observed in the studied solvents.
66
Figure 5.6 Normalized Raman spectra for representative systems as indicated in labels.
Green highlights are used to identify the regions of the spectra discussed in the text. Spectra were normalized to a maximum intensity of one so all systems could be compared on a common scale. Non-normalized spectra were used for analysis. Results from ref 16.
67 Table 5.2 Peak assignments and peak locations for MV+2 and EG vibrational modes in solvents studied. 1:2, 1:4, and 1:6 represent the molar ratios of the choline-based salt to
EG. Peaks were extracted by Gaussian function fitting. The symbol ⁑ represents a mode that comes up at two different frequencies. MVCl2 concentration was 10 mM in the corresponding mixtures. The table is color coded to indicate blue shifts and red shifts of
MV2+ and EG modes in solutions with respect to the pure compounds. Results from ref 16.
[PYR13] MVCl2 MVCl2/ ChCl:EG MVCl2/ChTFSI:EG [TFSI] / EG
This study Lit. 1:2 1:4 1:6 1:2 1:4 1:6
843 (ν C-C) 84547 838 842 840 839 838 839 838 839
47 1200 (δ CH2) 1200 --
1286 (ν C-C) 128748 -- 1283 1287 1288 1290 1291 1290 1291 MV+2 1302 (ν Ring) 130247 1297 1298 1301 1302 1303 1305 1305 1306
1554 (ν Ring) 154847 1537 1534 1540* 1540 1561 1561 1561 1560
1657⁑ (ν Ring) 163547 1645 1646 1647 1649 1649 1649 1650 1647
865 (ν C-C) 86349 863 866 858 867 865 867 865
1042⁑ (ν CH2- 1031 104449 1054 1052 1051 1052 1053 1050 CH2-O) 1051 EG 106250 1064 1064 1060 1066 1069⁑ (δ C-OH) N/A 1068 1065 1079 106949 1081 1076 1080 1078 50 1092 (δ,τ CH2) 1089 1088 1084 1090 -- 1086 1090 1091
50 1463 (νs CH2) 1460 1464 1466 1461 ------1456
Shifts in the vibrational frequency of specific modes between solvents indicate differences in local solvent environments. An increase in Raman shift (blue shift) of a peak indicates that the bond responsible for the peak is becoming shorter (more constrained) whereas a decrease in Raman shift (red shift) of a peak indicates the opposite effect – bond lengthening. The modes listed in Table 5.2 are color coded to represent changes in Raman
68 shift with respect to MV2+ or EG in their pure state, blue to represent a blue shift, red to
represent a red shift.
The first section of Table 5.2 contains the peaks attributed to MV2+ and the location
of these peaks in different solvent systems. Changes in peak locations suggest that the ring
2+ structures of MV are strongly influenced by the solvent system. For example, in solid
2+ -1 -1 MV a ring vibration at 1302 cm is red-shifted to 1297 cm in [PYR13][TFSI] but blue-
shifted to 1306 cm-1 in EG. In 1:2 ChCl:EG, this ring mode is red-shifted to 1298 cm-1,
similar to the peak location in [PYR13][TFSI], as illustrated in Figure 5.7 A. In 1:4 and 1:6
ChCl:EG and 1:2 ChTFSI:EG solutions, this ring mode matches its original frequency in solid MV2+. In 1:4 and 1:6 ChTFSI, the frequency of this particular ring mode is blue- shifted, similar to neat EG. Figure 5.7 B illustrates the shifts observed for the MV2+ ring
mode at 1554 cm-1. The same changes in Raman shift seen in the 1302 cm-1 mode is seen
in the 1554 cm-1 mode. Therefore, the concentration and identity of the salt influences the energy of the ring mode – it is at a higher energy (faster frequency) when interacting with
EG and a lower energy state when contacting [PYR13][TFSI] or when surrounded by the
bonding network of 1:2 ChCl:EG.
69
Figure 5.7 Change in Raman shift of MV rings modes at 1302 cm-1 (A) and 1554 cm-1 (B) in [PYR13][TFSI], 1:2 ChCl:EG, 1:2 ChTFSI:EG, and EG. Ring modes are red-shifted in the IL and DES but blue-shifted in the DES analogue and EG. Results from ref 16.
The observed red and blue shifts suggests that solute-solvent interactions are similar in [PYR13][TFSI] and 1:2 ChCl:EG. Both [PYR13][TFSI] and 1:2 ChCl:EG might accommodate MV2+ in holes. Inside these holes, solutes are not strongly confined and are allowed to expand. The red shifts (bond lengthening) of MV2+ ring modes in both
2+ [PYR13][TFSI] and 1:2 ChCl:EG support the idea that MV is accommodated in the holes of these solvents. Additionally, the magnitude of the red shifts observed in the ring modes
2+ are greatest in [PYR13][TFSI] suggesting MV is less confined in [PYR13][TFSI] as compared to 1:2 ChCl:EG. This suggests 1:2 ChCl:EG contains smaller holes then the IL, which is in agreement with surface-tension measurement results in Table 5.1. In contrast, blue shifts in all ring modes of MV2+ in ChTFSI:EG solutions, with the exception of the
1657 cm-1 mode, suggest molecular interactions are resulting in bond shortening of MV2+.
Given these ring modes are blue-shifted similarly in MVCl2 in EG and ChTFSI:EG solutions, Figure 5.7, interactions with EG might explain the observed blue shifts of MV2+
70 in ChTFSI:EG. Furthermore, this suggests that in ChCl:EG, interactions between EG and
MV2+ are less favorable or do not occur. Finally, the complex bonding environment,
focusing on hydrogen boding and specific solute-solvent interactions, analyzed via Raman is illustrated in Figure 5.8.
Figure 5.8 Illustration of changes in bonding environment of EG upon addition of MVCl2
(10 mM), followed by ChCl or ChTFSI. Blue indicates vibrational modes that were blue-
shifted and red indicates modes that were red-shifted with respect to pure compounds.
Dashed lines represent hydrogen bonding between species. Pink dashed lines indicate
2+ MV -EG interactions. Results from ref 16.
5.4 Conclusions
MVCl2 has been investigated as an electrochemical probe and redox active solute in three
classes of electrolytes: IL ([PYR13][TFSI]), ChCl:EG (1:2, 1:4, 1:6), and ChTFSI:EG (1:2,
1:4, 1:6). The viscosity of ChTFSI:EG systems were significantly lower than ChCl:EG
71 mixtures, which results in faster diffusion of MV2+ in the former as observed by
voltammetry. Using the Stokes-Einstein relation, the solvation radius for 1:2 ChTFSI:EG,
1:2 ChCl:EG, and [PYR13][TFSI] were estimated to be similar. Raman spectroscopy
suggests MV2+ is most likely accommodated in the ‘holes’ of the IL and DES – the free
2+ volume in the bulk of [PYR13][TFSI] and 1:2 ChCl:EG. In contrast, MV in ChTFSI:EG
interacts strongly with EG actively participating in the local bonding network. Thus, the
anion of the DES plays an important role in developing the local solvation environment
and associated dynamics. Weakly coordinating anions with charge delocalization, such as
TFSI, allows for reduction in viscosity, increase in diffusion coefficients, and increased
solute-solvent interactions. The complex bonding environment that becomes apparent in
the studied electrolyte systems directly influences how solute molecules move through
solution by means of modifying the hole size and solvation radius. Thus, the extent of both
intramolecular and intermolecular hydrogen bonding can be tuned through anion exchange
and through variation of hydrogen bond donor concentration, providing a potential
pathway for future investigations in tuning DESs and DES analogues for redox flow
batteries.
72 Chapter 6: Redox Behavior of Quinones in Ethaline
For DESs to be utilized in redox flow batteries, reversible, kinetically facile redox couples
must be identified. Small organic redox active molecules have attracted attention as
potential candidates. Here, a variety of quinone species are investigated for their solubility
and redox activity in ethaline as a screening study to evaluate the prospects of utilizing
quinones in ethaline as an electrolyte for redox flow batteries.
A wide variety of quinones are investigated here based on variations of the basic
benzoquinone structure. Hydroxyl groups are shown to greatly improve solubility in
ethaline while large aromatic side groups limit solubility. The observed redox behavior in
cyclic voltammetry measurements demonstrated significant deviation from typical redox
behavior of quinones reported in other non-aqueous solvents. Specifically, this intial study suggests certain quinones may not be redox active and while others display irreversible redox activity. This chapter provides a preliminary assessment of the redox activity of a
few common quinones but further research is needed to fully evaluate the potential of
quinone containing DESs for redox flow batteries.
6.1 Introduction
As previously discussed, DESs typically have electrochemical windows which are not wide
enough for high voltage applications such as lithium ion or lithium metal batteries.
However, their electrochemical windows are wider than aqueous electrolytes once again
presenting an opportunity to increase the energy density of redox flow batteries.
Furthermore, since DESs are composed of species produced on a megaton scale, their use
73 is economically feasible, potentially competitive with classical aqueous and organic
solvents.
In order for DESs to find application in redox flow batteries, a number of challenges
must be overcome. Perhaps the most important step is development of a reversible redox
couple as the energy storage medium. As such, identifying promising redox active species
has become an area of vigorous research. Desirable properties of redox active species
include high solubility in DES, reversibility of redox reaction, fast redox kinetics, and low
cost.9,115 Previous work has investigated the redox behavior of species such as iron,8
copper,116,64 ferrocene,117 and viologens31 in DESs. Thus far, the majority of reports
investigating redox active species in DESs have studied metal salts. Small organic
molecules represent an understudied and potentially favorable class of redox active species
due to their low cost, molecular diversity, and tailorable properties.115 Quinones are a
particularly interesting class of small redox active organic molecules. Quinones are
aromatic compounds which has been extensively studied in aqueous118 and organic119
solvents in biological,120 electrochemical,38 and chemical applications.121 The redox
behavior of quinones is known to depend on the solvent, in particular the protic or aprotic
nature of the solvent.118,122 While previous reports have investigated the redox behavior of
quinones in ILs, to date, the author is not aware of any investigations of quinone redox
behavior in DESs.38,123
Figure 6.1 illustrates the structure of archetypical quinone, 1,4-benznoquinone
(BQ) and potential redox pathways for BQ.118 In protic media, BQ undergoes a single two electron transfer reduction coupled with double proton transfer to form hydroquinone
(H2BQ), Figure 6.1 A. This reaction appears as one irreversible waveform in cyclic
74 voltammetry with typical cathodic and anodic peak separations of ~0.5 V.124 In aprotic
media, BQ undergoes two separate single electron reduction steps, first to a radical anion
then to a dianion with no proton transfer steps, Figure 6.1 B. This reaction appears as two
separate waveforms with the first step occurring reversibly and the second step slightly
irreversibly depending on the solvent.122 The reaction scheme illustrated in Figure 6.1 is
representative of the general redox pathways which most quinone structures undergo.
Figure 6.1 Redox pathway of BQ in protic (A) and aprotic media (B).
To evaluate the feasibility of utilizing quinones as redox couples in DES based flow batteries, the redox activity of benzoquinone (BQ), hydroquinone (H2BQ), naphthoquinone
75 (NQ), and anthraquinone (AQ) was investigated in ethaline. The structures of each quinone
is reported in Figure 6.2. Differences in functional groups allow for probing of the impact
of small polar groups like hydroxyl and bulky nonpolar groups like carbon rings. As
previously discussed, ethaline is a low viscosity DES composed of choline chloride and
ethylene glycol in a 1 to 2 molar ratio respectively. The solubility of the quinone species
was investigated and the redox behavior probed with cyclic voltammetry. In a 2014 review,
Abbott et al. suggested DESs such as ethaline were protic.10 Few other studies have
investigated or reported on the protic nature of ethaline. If indeed ethaline behaves like a
protic media, the redox mechanism of quinones would be expected to follow the path given in Figure 6.1 A.
Figure 6.2 Structures of quinones investigated in this study, from left to right: hydroquinone (H2BQ), benzoquinone (BQ), naphthoquinone (NQ), and anthraquinone
(AQ).
76 6.2 Materials and Methods
Choline chloride (99%) and anhydrous ethylene glycol (99.8%) was obtained from Acros
Organics. Hydroquinone (99.5%), benzoquinone (99%), naphthoquinone (99%), and
anthraquinone (99%) were obtained from Acros Organics. Silver wire (99.9%) and
platinum wire (99.9%) was obtained from Alfa Aesar. A glassy carbon working electrode
(diameter, 3 mm) and platinum working electrode (diameter, 1.6 mm) was obtained from
BASi.
Ethaline was prepared as previously discussed.12 Ethaline water contents were measured with Karl Fischer titration and found to be < 100 ppm. Quinones are expected to contain water yet are heat sensitive prone to decomposing or subliming, as such quinones were used as received and are expected to increase the water content of ethaline samples.
Water contents of quinone and ethaline mixtures cannot be evaluated by coulometric titration as quinones participate in side reactions with the anolyte solution.
Cyclic voltammetry experiments were conducted in a T-cell setup as previously described in Chapter 3. Experiments were conducted with either platinum or glassy carbon working electrodes. Coiled platinum wire and silver wire plated with silver chloride were used as counter and reference electrodes respectively. The electrode assembly was immersed in a liquid well of approximately 300 µL filled with quinone ethaline solution.
Experiments were conducted in a positive pressure nitrogen glove box to prevent additional water uptake. All voltammograms were iR corrected using a single 100 kHz impedance measurement. Experiments were conducted beginning at open circuit voltage (OCV) and swept anodically to a maximum potential before reverse sweeping cathodically to a minimum potential.
77 Fourier Transform Infrared Spectroscopy (FTIR) was used to probe the solubility of quinones. A Nicolet iS50 FTIR (Thermo Scientific) was utilized with attenuated total reflectance (ATR) capability (diamond crystal). All spectra were collected using the ATR attachment in the range of 4000-400 cm-1 wavenumber. Before an experiment, the ATR crystal was cleaned with methanol and a background ATR spectrum collected. This background spectrum was subtracted from experimental spectra. Spectra were collected with a 0.5 cm-1 resolution in wavenumber and normalized to a maximum absorbance of one.
To evaluate quinone solubilities, ethaline solutions with different concentrations of quinone were prepared. Saturated solutions were prepared by adding excess quinone to ethaline at room temperature. The supersaturated solution was heated to 60 ºC and stirred vigorously until all quinone dissolved. The solution was then slowly cooled to 25 ºC resulting in precipitation of some quinone. The solution was allowed to equilibrate for 24 hours at 25 ºC. Ethaline saturated with quinone was then slowly pipetted from the vial for
FTIR analysis. Solutions of known quinone concentration were prepared and analyzed with
FTIR to generate calibration curves and determine the concentration of quinone in the saturated solution based on the Beer-Lambert law.
6.2 Results and Discussion
Figure 6.3 A shows the IR spectra for different concentrations of H2BQ in ethaline as well as a neat ethaline spectrum. Five different concentrations of H2BQ were analyzed, 0.1 M,
0.5 M, 1 M, 2 M, and a saturated solution of unknown concentration. Regions of the
78 spectrum in Figure 6.3 A marked with lines represent vibrational modes of interest. Black
lines at 1511 cm-1, 1205 cm-1, and 760 cm-1 indicate v CC, v CC, and v CO vibrational
125 -1 modes of H2BQ respectively. The blue and red lines indicate the 1036 cm v CO mode of ethylene glycol and 953 cm-1 v CC mode of choline chloride.74,67 The absorbance of all
H2BQ modes at all investigated concentrations reported above were normalized to the
absorbance of the v CO mode of ethylene glycol. Plots of H2BQ concentration versus
normalized absorbance revealed a linear relationship for all modes probed, Figure 6.3 B.
Using a linear fit, the approximate concentration of the saturated solution could be estimated. The fitting parameters m (slope) and b (intercept) as well as goodness of fit (R2) are reported in Table 6.1. It should be noted the intercept b is not zero due to the background. Based on this analysis, the solubility limit of H2BQ in ethaline at 25 ºC is
estimated to be 6.0 ± 0.30 M. A solvent capable of holding redox active species up to 6 M
is unprecedented. To ensure choline chloride was not expelled from the ethaline mixture
-1 to accommodate H2BQ, the absorbance of the 953 cm v CC mode of choline chloride was investigated. In neat ethaline and saturated H2BQ in ethaline, the normalized absorbance
of this mode was found to be 0.55 AU and 0.53 AU respectively. This indicates the relative
amount of choline chloride changed negligibly between the neat ethaline and saturated
H2BQ solution. Similar results were observed in the 0.1 M, 0.5, 1 M, and 2.5 M solutions
of H2BQ.
79
Figure 6.3 IR spectra for neat ethaline and various concentrations of H2BQ in ethaline.
Vertical black, blue, and red lines indicating modes of interest for H2BQ, ethylene glycol, and choline chloride, respectively (A). Linear fit to concentration versus normalized
-1 absorbance for the 1205 cm v CC mode of H2BQ (B).
Table 6.1 Results of linear fit to concentration versus absorbance for different modes of
2 H2BQ. Parameters m and b are the slope and intercept of the fit respectively. R is the goodness of fit measure.
Similar analysis was performed for BQ in ethaline. The IR spectra of neat ethaline and BQ in ethaline are reported in Figure 6.4 A. Solutions with known concentrations of
80 BQ in ethaline, 0.03 M, 0.06 M, 0.12 M, and 0.24 M, were used as a calibration to
determine the concentration of BQ in the saturated sample. Due to the low concentration
of BQ in ethaline, only one mode at 1653 cm-1 (v CO) can be attributed to BQ. This mode is marked by a vertical black line in Figure 6.4 A. Blue and red vertical lines once again indicate the 1036 cm-1 v CO mode of ethylene glycol and 953 cm-1 v CC mode of choline
chloride. Figure 6.4 B shows a plot of BQ concentration versus absorbance normalized to
the 1036 cm-1 ethylene glycol mode. From a linear fit to data in Figure 6.4 B, the solubility of BQ in ethaline at 25ºC is estimated to be 0.29 ± 0.03 M. The normalized absorbance of the 953 cm-1 v CC mode of choline chloride was found to be identical for all solutions suggesting the composition of the base solvent was not altered.
Figure 6.4 IR spectra for neat ethaline and various concentrations of BQ in ethaline.
Vertical black, blue, and red lines indicating modes of interest for BQ, ethylene glycol, and
choline chloride respectively (A). Linear fit to concentration versus normalized absorbance for the 1653 cm-1 v CO mode of BQ (B).
81 As compared to H2BQ and BQ, NQ and AQ exhibited significantly lower
solubilities in ethaline. A 10 mM solution of NQ could not be successfully prepared
although a small portion of NQ was successfully dissolved. AQ exhibited essentially no
solubility in ethaline remaining a suspended powder even under heating at 60ºC and
vigorous mixing. The solubility limit of each quinone in ethaline is summarized in Table
6.2. Simply by substituting the hydroxyl groups of H2BQ for the carbonyl groups of BQ,
the solubility of the quinone is greatly reduced. Addition of bulky ring groups further
reduces quinone solubility. This trend suggests the greater the polarity of a species, the
greater the solubility of the species in DESs. Additionally, it might be excepted that large
molecules which are initially valent and highly soluble, might be less soluble in the DES
upon reduction or oxidation to the ground state. This is an important consideration as redox
active species must remain soluble in the electrolyte across oxidation states to have
practical application in flow batteries.
Table 6.2 Solubility limit of various quinones in ethaline listed according to the functional
groups associated with each quinone species. Decreasing polarity and increasing size of
functional groups decreases solubility of quinone in ethaline.
82 The difference in solubility of H2BQ and BQ in ethaline might be surprising given
the similarities of their structures. While the hydroxyl groups of H2BQ are slightly more
polar then the carbonyl groups of BQ, variation in polarity might not fully explain the
result. The discrepancy between the solubility of H2BQ and BQ in ethaline was further
probed by attempting to make a DES using the quinones as hydrogen bond donors and
choline chloride as the hydrogen bond acceptor. A previous study demonstrated a DES
could be made from choline chloride and the quinone catechol.126 Two separate mixtures
were prepared, choline chloride and H2BQ in a 1:2 molar ratio respectively and choline
chloride and BQ also in a 1:2 molar ratio. Both mixtures were manually mixed before
heating to 80 ºC. Interestingly, the solution of choline chloride and H2BQ formed a liquid
while the choline chloride and BQ mixture did not form a liquid. The solution was slowly
cooled and remained liquid above 30 ºC. Given hydroquinone and choline chloride have
melting points of 172 ºC and 302 ºC respectively, this solution displays significant melting
point depression characteristic of a DES. H2BQ, in a manner similar to ethylene glycol, is able to donate hydrogen bonds to choline chloride forming a coordination complex which results in freezing point depression. This could explain the significant solubility of H2BQ
in ethaline and the much lower solubility of BQ. A 6 M concentration of H2BQ in ethaline can be achieved because H2BQ and ChCl can form a DES. A DES composed of redox
active species, like the DES formed from choline chloride and H2BQ, represents the most energy dense of electrolytes and thus is an exciting opportunity for potential application in redox flow batteries.
While H2BQ displays high solubility in ethaline, to be utilized as a redox couple,
its redox behavior in ethaline must be understood. The electrochemical window (EW) of
83 ethaline on platinum and glassy carbon was evaluated with cyclic voltammetry, Figure
6.5. Ethaline exhibited greater electrochemical stability on glassy carbon as compared to platinum. At a cut-off current density of 0.05 mA/cm2, the EW of ethaline was found to be
2.98 V and 1.92 V on glassy carbon and platinum respectively. Given the greater stability of ethaline on glassy carbon, the redox activity of H2BQ and BQ were evaluated on glassy carbon rather than platinum.
Figure 6.5 EW of ethaline on glassy carbon and platinum working electrodes.
Figure 6.6 shows the redox behavior of H2BQ and BQ as evaluated through cyclic voltammetry. In Figure 6.6 A, the anodic peak associated with oxidation of H2BQ is seen to be outside the anodic window of ethaline, as indicated by the vertical black line.
Significant bubble formation was observed on the counter and working electrodes. As such,
H2BQ is not a suitable redox active solvent in ethaline and its redox activity cannot be probed further. BQ in contrast displays its full redox behavior inside the EW of ethaline as seen in Figure 6.6 B. However, it does not exhibit behavior expected from previous
84 investigations of BQ in protic solvents. Four clear peaks are observed, two anodic at 0.31
and 1 V vs Ag/AgCl and two cathodic at 0.33 and -0.11 V vs Ag/AgCl respectively. While
not particularly similar to the expected protic or aprotic redox behavior, this behavior is
most similar to aprotic redox behavior where two redox events are observed.
Figure 6.6 Voltammograms of 0.5 M H2BQ in ethaline (A) and 20 mM BQ in ethaline (B).
Anodic and cathodic peaks are marked Ea and Ec respectively. The anodic limit of ethaline is indicated by a black line.
To determine the coupling of anodic and cathodic peaks in Figure 6.6 B, potential- restricted cycling was employed. Based on the results displayed in Figure 6.7, the redox
events labeled Ea,1 and Ec,1 are coupled while Ea,2 and Ec,2 must also be coupled. Peak Ea,1
is not observed without peak Ec,1 (Figure 6.7 B) while peak Ea,2 is not observed without
peak Ec,2 (Figure 6.7 A). While this coupling is similar to aprotic redox behavior, large
separation of cathodic and anodic peaks suggests the kinetics is sluggish. Additionally, the
ratio of peak currents for coupled anodic and cathodic events do not appear close to one.
These factors point to the irreversibility of BQ redox behavior in ethaline. If the BQ redox
85 reactions were reversible, peak separation between cathodic and anodic events would ≤ 59
mV and the ratio of peak currents would ~1. However, peak separation approaches 0.75 V.
As such, BQ is likely not a promising couple in ethaline.
Figure 6.7 Potential-restricted voltammograms, -0.75 to 0.75 V v Ag/AgCl (A) and 0.125
V to 1.25 V vs Ag/AgCl (B).
The question remains, why does BQ exhibit anomalous behavior in ethaline? While
ethaline has been stated to be protic,10 the author is not aware of definitive evidence
demonstrating the protic nature of ethaline. BQ certainly does not exhibit redox behavior
expected within protic solvents.124 In aprotic solutions, BQ redox behavior has shown
dependency on hydrogen bonding, where formation of the quinone diradical is stabilized
by hydrogen bonding. In fact up to 0.7 V shifts in the redox potential of the quinone
diradical has been reported in aprotic solvents simply by doping the solvent with hydrogen
bond donors like ethylene glycol.127 The extensive hydrogen bonding network of ethaline could be stabilizing the formation of dianion species resulting in the observed behavior.
Additional, investigation of this behavior is needed to fully understand its origin.
86 6.3 Conclusions
The solubility and redox activity of four different quinone structures in ethaline has been investigated. To date no studies have investigated electrochemical properties of quinones in DESs. H2BQ was found to be the most soluble at 6.0 M followed by BQ at 0.29 M while
AQ was essentially insoluble. Discrepancies in solubility are attributed to differences in the polarity of functional groups. Importantly, a DES composed of choline chloride and
H2BQ was observed to form at a 1:2 molar ratio respectively. This is an interesting result as DESs composed of redox active species offer unparalleled energy density potentials.
Future studies should investigate the redox behavior of this new DES. While H2BQ displayed high solubilities in ethaline, it was not redox active within the EW of ethaline.
BQ in contrast displayed redox activity within the window of ethaline but displayed previously unreported behavior. The hydrogen bonding network of ethaline is suggested to stabilize the formation of the dianion quinone species perhaps contributing to the anomalous behavior. However, the origin of the redox behavior anomaly is not well understood and requires further investigation.
87 Chapter 7: Ionic Liquids: Bulk Properties and Interface Characteristics
ILs lack neutral solvent molecules so classical models governing ion transport and interfacial charging established for traditional solvents must be reevaluated for ILs. In this study, we characterized a series of quaternary ammonium ILs in terms of their bulk macroscopic properties such as density, viscosity, conductivity and interfacial charge density and structure near a glassy carbon electrode. For the interfacial characterization, electrochemical impendence spectroscopy (EIS) was employed. Specifically, the impact of cation structure on interfacial structures is investigated by varying alkyl chain length of the ammonium cation.
Walden analysis, a measure of ionicity, revealed that capacitance, a measure of charge density, at the point of zero charge (PZC) is independent of the strength of cation- anion associations for the ILs investigated: varied cations with the same anion (Table 4.1).
Long alkyl chains are found to stabilize the formation of dense anion layers at positive electrode potentials by screening ion-ion repulsion. This work was done in collaboration with Jeffrey M. Klein, who performed differential capacitance fits for the measured EIS and surface enhanced Raman spectroscopy experiments. It was published in The Journal of Physical Chemistry C titled: Electroanalytical Investigation of the Electrode–Electrolyte
Interface of Quaternary Ammonium Ionic Liquids: Impact of Alkyl Chain Length and
Ether Functionality.13 This chapter is organized to provide background on electrode- electrolyte interfaces of ILs followed by the details of the specific methods used and the results of physical property measurements and EIS analysis.
88 7.1 Introduction
The formation of the electrical double layer, an ion structure formed on electrode surfaces,
which can be described by the Gouy-Chapman-Stern (GCS)128 model for traditional
solvents must be reevaluated for ILs. For traditional solvents, the GCS model describes the
position of ions and electric potential profile from the surface of the electrode into the bulk
electrolyte and is given by equation 1:
= equation 1 𝑑𝑑𝑑𝑑 𝐶𝐶 𝑑𝑑𝑑𝑑 where is differential capacitance, is surface charge density, and is the electrode
charge.𝐶𝐶 ILs stray from this behavior 𝛿𝛿typically displaying a camel shaped𝑞𝑞 dependence of
differential capacitance on electrode potential.129 A local minimum in differential
capacitance is located at the PZC while “camel” local maximums are found at positive and
negative electrode potentials. This behavior can be modeled by the relation developed by
Kornyshev et al.130 given by equation 2:
= 𝛼𝛼𝛼𝛼 equation 2 2 𝛼𝛼𝛼𝛼 1 2𝛾𝛾 𝑠𝑠𝑠𝑠𝑠𝑠 ℎ � 2 � 𝛼𝛼𝛼𝛼 0 2 2 𝛼𝛼𝛼𝛼 𝐶𝐶 𝐶𝐶̌ 𝑐𝑐𝑐𝑐𝑐𝑐ℎ � 2 � 1+2𝛾𝛾 𝑠𝑠𝑠𝑠𝑠𝑠ℎ � 2 � �𝑙𝑙𝑙𝑙�1+2𝛾𝛾 𝑠𝑠𝑠𝑠𝑠𝑠ℎ � 2 ��
where is the Debye capacitance, is thermal potential, is the ion interaction
0 parameter,𝐶𝐶̌ and is the compressibility parameter.𝑢𝑢 The compressibility𝛼𝛼 parameter has also
been defined as 𝛾𝛾the fraction of unpaired ions at the interface. Given ILs are salts, composed
exclusively of ions, the parameter might depend on the ionicity of the solvent. Ionicity is a measure of ion interactions to qualitatively𝛾𝛾 distinguish between paired ions and fully dissociated ions. The greater the degree of dissociation between ions, perhaps the greater
89 the ability of ions to pack at the interface. The ionicity of ILs can be evaluated through
Walden analysis, plots of molar conductivity versus inverse viscosity.106 Walden analysis is accomplished through measurements of viscosity, conductivity, and density.
The IL electrode-electrolyte interface has been previously studied using methods such as electrochemical impedance spectroscopy (EIS),131 X-ray scattering,132 surface- enhanced Raman spectroscopy (SERS),133 chronoamperometry,134 and molecular simulations.135 These techniques have suggested the IL ion structure takes on alternating layers of cations or anions.131-135 This structure can extend nanometers into the electrolyte and contrasts the structure formed by traditional solvents where an inner layer of hydrated ions followed by a diffuse layer of ions and solvent dipoles develops. Previous work has suggested cation rigidity can prevent accumulation of IL ions at the electrode interface.101
That is, ILs composed of less rigid cations like quaternary ammoniums display greater local maximums in differential capacitance as compared to ILs composed of rigid cations like pyrrolidinium. As such, this work further investigates ammonium ILs to understand the relationship between structure and the IL-electrode interface. In particular the impact of ionicity and alkyl chain length is investigated through Walden analysis and EIS.
7.2 Methods
The following ILs were purchased from Iolitec (Alabama): methyl trioctylammonium bis
- (trifluoromethylsulfonyl)imide [N1888][TFSI] (99%), octyltriethylammonium bis(trifluoromethylsulfonyl)imide [N2228][TFSI] (98%), butyltrime thylammonium bis -
(trifluoromethylsulfonyl)imide [N1114][TFSI] (99.5%), N,N- diethyl-N-methyl-N-(2-
90 methoxyethyl)ammonium bis-(trifluoromethylsulfonyl)imide [N122(2O1)][TFSI] (99%),
and hexadecytrimethylammonium bis(trifluoromethylsulfonyl)imide [N111(16)][TFSI]
(99%). Prior to characterization, each IL was dried under vacuum for 12 h at 80°C and
stored in a nitrogen glovebox to prevent water uptake. Water contents were measured by
Karl Fischer Titration (Metrohm Coulometric KF 889D) and are reported in Table 7.1.
Nitrogen gas (99.998%) was purchased from Airgas. Mineral oil N100 and S20 viscosity
standards were purchased from Cannon Instrument Company (Pennsylvania). Polystyrene
sheets, 1 mm thick, were purchased from McMaster Carr. Silver foil (0.1 and 0.05 mm
thick) and silver wire (0.5 and 1 mm in diameter) were purchased from Alfa Aesar.
7.2.1 Density
Densities of all ILs were measured in the temperature range of 298.15 K to 323.15 K with
an Anton Paar DMA 4500M vibrating tube density meter (accuracy of ± 0.03 K and ±
0.00005 g/cm3). Prior to an experiment, the density of air and water were measured at room temperature to confirm the accuracy of the equipment. The vibrating tube was rinsed between experiments with methanol followed by deionized water and dried with an onboard air pump. Each experiment required approximately 1 mL of sample.
7.2.2 Viscosity
Viscosities were measured with a RheoSense MicroVISC microchannel viscometer inside a RheoSense MicroVISC Temperature Control unit (± 0.10 K). Before testing each IL, the microchannel chip was rinsed with methanol followed by deionized water. The viscosity
91 of N100 and S20 Cannon viscosity standards were measured at room temperature to
confirm the reliability of measurements and found to be within 0.2 %. Viscosity
measurements were repeated three times at each temperature for each IL and the average
is reported.
7.2.3 Conductivity
Conductivity was measured with a dual platinum electrode cell with a cell constant of 1.41
cm-1 (MMA 500, Materials Mates Italia, Italy) by electrochemical impedance spectroscopy
(EIS) with a BioLogic SP240 potentiostat equipped with frequency response analyzer (7
MHz - 10 µHz). Temperature was controlled for conductivity measurements by placing the
cell inside a RheoSense MicroVISC Temperature Control unit (± 0.10 K). The conductivity
is reported as the average value calculated from three repeats. Resistance was determined
from Nyquist plots where complex impedance data intersected the real impedance axis.
From resistance, conductivity was calculated as the cell constant divided by solution
resistance. As the ILs had a wide range of conductivities, the frequency range which was fit to a line varied between ILs and is reported in Table 7.1.
92 Table 7.1 Water contents of dried ILs and frequency range of linear fit to EIS data for
determining conductivities. Results from ref 13.
7.2.4 Electrochemical Window
Electrochemical windows (EWs) of ILs were determined by cyclic voltammetry (CV) on a 3 mm diameter glassy carbon working electrode (BASi) at a scan rate of 20 mV/s. A T-
cell was used with silver wire counter and reference electrodes (1 mm diameter, Alfa
Aesar) contained in a positive pressure nitrogen glove box. A cutoff current density of 200
µA/cm2 was used to determine the cathodic and anodic limits. The working electrode was
polished by sliding the disk surface in a figure eight motion over a microdisk cloth covered
with the alumina polish for 2 minutes. The electrode was then rinsed with distilled water
(DI), and sonicated for 120 seconds in DI water. Finally, the electrode was rinsed with DI water followed by methanol and dried.
93 7.2.5 Differential Capacitance
Differential capacitance data was collected and analyzed by Jeffrey Klein. Here the method
used to collect the data is briefly explained. Differential capacitance of the IL-electrode
interface was measured by EIS in a nitrogen glovebox using a 3 mm diameter glassy carbon
working electrode. The counter electrode was Ag plate (2.5 cm2, 0.05 mm thick from Alfa
Aesar) and the reference electrode was Ag wire (1 mm diameter, Alfa Aesar). Before an
experiment, ILs were degassed by bubbling nitrogen gas for least 45 minutes. A 5 mL
three-electrode cell was used with the counter electrode placed at the bottom of the cell,
the working electrode dipped to form a hanging meniscus, and the reference electrode
placed in the bulk electrolyte near the hanging meniscus. EIS measurements were
performed at potentials inside the measured EW of each IL with a frequency range of 500
kHz - 100 mHz. A Biologic potentiostat equipped with a frequency response analyzer
applied a 10 mV amplitude sine wave perturbation. Prior to applying the sinus perturbation
either the base potential was held for 1 minute. EIS data was fit to an equivalent circuit
consisting of a resistor in series with a constant phase element using EC-Lab (Bio-Logic
science instruments software).
7.3 Results and Discussion
The room temperature density, viscosity and conductivity of each IL along with the cation
structure is reported in Table 7.2. It is seen that as cation alkyl chain lengths decrease, viscosity decreases while density and conductivity increase. [N1888][TFSI] is the most viscous and least conductive of the investigated ILs. This might be expected due to the size
94 of the cation which contains three separate octyl alkyl chains. The size and highly branched
nature of the cation limits it mobility with detrimental effects on viscosity and conductivity.
Furthermore, [N1888][TFSI] is the least dense of the reported ILs. This could also be
attributed to the size of the cation which prevents efficient ion packing resulting in low
density. The investigated ILs are likely too viscous to be utilized in lithium-ion or lithium
metal batteries, even [N122(2O1)][TFSI] which is the least viscous of the investigated ILs
has a viscosity of 75 cP which is an order of magnitude greater than organic solvents.
However, these ILs provide a case study for IL-electrode structuring which has a direct impact on electron transfer kinetics.
Table 7.2 Measured viscosity conductivity and density of studied ILs, each with [TFSI] anion, at 298.15 K. Results from ref 13.
95 Physical property measurements were used to construct a Walden plot (log molar
conductivity vs log inverse viscosity) shown in Figure 7.1 A. The Walden plot is
interpreted as a measure of the strength of cation-anion interactions in each IL. The dashed
line in Figure 7.1 A represents an aqueous solution of 0.01 M KCl, a completely
dissociated ion pair system and is used as a reference of a non-interacting ionic solution.
ILs that lie close to the reference line are suggested to contain non-interacting, or weakly
interacting, ions. Deviations below and to the right of the KCl line indicates increased ion associations which can lead to neutral aggregate ion pairs. Accordingly, the strength of ion
interactions for the respective ILs follows [N1888][TFSI] > [N2228][TFSI] >
[N122(2O1)][TFSI] ~ [N1114][TFSI].
Figure 7.1 Walden plot of studied ILs (A) and voltammograms of neat ILs on glassy carbon
at 298 K, scale bar represents a current density of 200 µA/cm2 (B). Results from ref 13.
EWs of the neat ILs are shown in Figure 7.1 B. All ILs have nearly the same anodic
and cathodic limits which allows the interfacial capacitances to be determined across the
same range of potentials. Additionally, the voltammograms of each IL show broad low
96 current peaks likely attributed to impurities. While the identity of the impurities is unknown, previous work suggested dissolved gasses or halide intermediates leftover from
IL synthesis could cause broad peaks in voltammograms.136
Figure 7.2 shows the measured differential capacitance of the four ILs. All ILs display a camel shape differential capacitance curve. Three key features of the camel shape differential capacitance curves should be noted. First, the minimum capacitance at the PZC between the anodic and cathodic peaks. The capacitance at PZC was expected to correlate to the position of the ILs on the Walden plot in Figure 7.1 A; recall the parameter from the Goodwin-Kornyshev model, equation 2, which is suggested to depend on the fraction𝛾𝛾 of unpaired ions. ILs with weakly interacting ions were expected to have greater capacitance with no applied electrode potential. Weakly interacting ions should more easily dissociate and accumulate at the electrode interface as unpaired ions without applied potential. As stated earlier, [N122(2O1] and [N1114] fall close to the KCl line on a Walden plot and can be interpreted as more dissociated then [N1888] and [N2228]. Thus,
[N122(2O1] and [N1114] might have the greatest differential capacitance at PZC.
Interestingly, the differential capacitance curves indicate no clear trend between degree of ion interaction and the capacitance at PZC. Contrary to the expectation, capacitance near
PZC for [N1888][TFSI], [N2228][TFSI], and [N1114][TFSI] are all nominally 2 - 3
µF/cm2. Free ion density at the interface at PZC depends on more than just the number of participating ions. Perhaps interactions between like ions, as well as steric effects and electrostatics, must be considered to understand differential capacitance at PZC.
The second feature in Figure 7.2 is related to the two maxima; one at a positive electrode potential and one at a negative electrode potential. The peak at positive potentials
97 corresponds to excess [TFSI] anion density while peaks at negative potentials correspond
to excess cation density. The increase in capacitance from PZC to these maxima are related to increasing ion density as the magnitude of potential increases which is made possible by rotations and reorientations of ions at the interface. The wings of the capacitance curves
(maxima to limiting potentials) reflect decrease in ion density due to repulsive forces between like ions. Figure 7.2 demonstrates that cation structure modulates excess anion density at the interface. For instance; [TFSI] density at the interface increases to a greater extent than the cation [N1888] density for the IL [N1888][TFSI], Figure 7.2 A. The octyl groups present in [N1888][TFSI] are expected to hinder anion – anion repulsion resulting in large differential capacitance maxima at positive potentials. The octyl group present in
[N2228][TFSI] has the same effect on the accumulation of anions and shows similar peak asymmetry to [N1888][TFSI], Figure 7.2 B. A maxima of 19.9 μF/cm2 is observed for
[N1888][TFSI] while a maxima of 14.4 μF/cm2 is observed for [N2228][TFSI]. Each octyl
group is capable of screening repulsion forces between anions leading to more effective
screening in [N1888][TFSI] compared to [N2228][TFSI]. More effective screening of like
charges results in larger capacitance maxima. The ILs [N1114] and [N11(22O1)] lack long octyl alkyl groups to screen anions at positive potentials. Consequently, the value of capacitance at positive electrode potentials is lower and capacitance peaks are more symmetric compared to the [N1888] and [N2228] cations, Figure 7.2 C and D.
98
Figure 7.2 Differential capacitance evaluated with EIS on glassy carbon electrode. Error bars represent a 90% confidence interval calculated from the statistical average of capacitance values. Red lines represent fits to the Goodwin-Kornyshev model. Potential
was swept from open circuit voltage (OCV) to negative potentials, returned to OCV then
scanned to positive potentials. Results from ref 13.
The final feature of interest in Figure 7.2 are the tails of the curves where differential capacitance decreases from the local maximum as the electrodes are further polarized. The tails reflect decreases in ion density due to repulsive forces between like ions. Decreasing capacitance is associated with either steric forces or ion-ion repulsion, both of which result in an increase in the interface thickness. [N1888] and [N2228] exhibit slow capacitive decay as seen in the tails at negative potentials in Figure 7.2 A and B,
99 respectively. This is due to steric cation-cation repulsions which allow [TFSI] with its
comparatively smaller size to access the void space between the cations leading to an
overscreened interface. Overscreened systems contain a monolayer of same charged ions
followed by alternating oppositely charged layers as previously described by Bazant and
Kornyshev.137 In contrast to the slow capacitive decay seen at negative potential tails, the
sharp decrease in capacitance for [N1888][TFSI] and [N2228][TFSI] at positive potential
tails can be explained by interface crowding. As more [TFSI] ions pack into the interface, the proximity of like charges creates a large degree of ion-ion repulsion which eventually overcomes the screening introduced by the C8. The leads to rapid increases in interfacial thickness resulting in a tail with steep slope.
Red lines in Figure 7.2 represent fits to differential capacitance data using the
Goodwin-Kornyshev model as expressed in equation 2. The model is unable to predict the
tail behavior of differential capacitance curves for certain ILs. While the model is able to
account for ILs with slow decay in capacitance, [N2228] at negative potentials, it does not
fit ILs exhibiting quick decay in capacitance, [N1888] at positive potentials for example.
Differences between the experimental data and the fit may arise from reorientations of ions
at the interface. While the model takes into account short range ion associations, ion
reorientations directly impact crowding and overscreening in the interface resulting in the
observed differences between experiments and model.
100 7.4 Conclusion
The ammonium ILs investigated here display “camel” shaped differential capacitance curves. As cation structure became bulkier with longer alkyl chains, viscosity increased while density and conductivity decreased. Ether functionality on the
[N122(2O1)] resulted in the highest capacitance at PZC, 6.19 μF/cm2. However, no trend in IL ionicity, as evaluated by Walden analysis, and differential capacitance at PZC was observed. Any relationship between ionicity and differential capacitance is likely hidden by more dominating effects such as steric repulsion between alkyl chains. Incorporating octyl groups on the cation lead to the highest capacitances at positive potentials: 19.9
μF/cm2 for [N1888][TFSI] and 14.4 μF/cm2 for [N2228][TFSI]. Long carbon chain substituents of the ammonium cation demonstrate the ability to increase interfacial ion density of anions by screening neighboring ions from like charges. SERS analysis of
[N1888][TFSI], described elsewhere, supports the conclusion that large alkyl chains can penetrate the anion structure.13 SERS also suggests octyl alkyl chains buckle forming ring- like structures at greater electrode polarizations reducing anion-anion repulsions, leading to increased capacitances at positive polarization. This study demonstrates the ability to control the charge density of electrode-electrolyte interfaces with the choice of IL structure and applied potential.
101 Chapter 8: Conclusions and Future Work
With this study, standards for preparation, handling, and characterization of DESs were
developed. Even small quantities of water can have a dramatic impact on DES properties.
Therefore, DES components must be thoroughly dried, and prepared and stored under a
dry, inert atmosphere. Although DESs are relatively simple to prepare, preparation,
handling, and characterization techniques and the water content of samples should always
be reported. By using consistent techniques detailed here, trends between studies can be
identified and structure-function understanding can be established.
With the developed techniques, DESs and DES analogues based on choline
chloride, ethylene glycol, and glycerol were then characterized. Choline chloride was
found to disrupt the highly ordered three-dimensional hydrogen bonding network of neat glycerol. Correspondingly, the viscosity and density of glycerol decreased with increased choline chloride content as chloride complexed with hydroxyl groups reducing glycerol- glycerol interactions. In contrast, choline chloride increased the viscosity and density of ethylene glycol. The chloride-hydroxyl complex formed in choline chloride and ethylene glycol solutions increased long ranger order. Interestingly, conductivity increased as choline chloride content increased until a critical composition after which conductivity decreased with increased choline chloride content. Triols like glycerol or other molecular solvents with highly ordered hydrogen bonding structures may be expected to exhibit improved conductivity and viscosity with addition of HBAs. In contrast, diols like ethylene glycol may be expected to exhibit more complicated behavior where tradeoffs between ion mobility, number of ions, and strength of ion coordination must be balanced.
102 The solvation environment of an IL, DES, and DES analogues were probed with the redox active species methyl viologen. Replacing the chloride anion of choline chloride with the TFSI anion in ethylene glycol solutions was found to significantly increase fluidity but decrease conductivity. Average hole sizes were estimated for each solvent and found to correlate with viscosity and molecular size. Diffusion of methyl viologen was greatest in solvents with large holes sizes and small solvated radii. Raman spectroscopy coupled with physical property data suggested methyl viologen largely weakly coordinates with ethylene glycol-choline chloride solutions and the IL, likely accommodated in the holes of the solvents. In contrast, methyl viologen coordinates more strongly with ethylene glycol when choline TFSI salt was introduced. Thus, anion substitution provides a pathway for tuning solvent environments. Weakly coordinating ions can improve fluidity and diffusion.
Various quinones were studied in terms of their solubility and redox reversibility in ethaline. The solubility of quinones was found to depend on the polarity and size of functional groups. Small, polar quinones were found to have greater solubility in ethaline.
Hydroquinone was found to have a solubility of 6 M in ethaline but was not redox active in the EW of ethaline. A DES based on choline chloride and hydroquinone was demonstrated to be liquid at 30 ºC, an important realization as DESs composed of redox active species offer the greatest concentration of redox active solutes and therefore energy density. Finally, benzoquinone was show to display redox behavior unlike previously established protic or aprotic behaviors. This is tentatively attributed to the hydrogen bonding network of ethaline which might stabilize dianion quinone formation.
Finally, the electrode-electrolyte interface of ammonium ILs was studied. The developed research method can be adapted for the study the DES-electrode interfaces in
103 the future. The capacitance at the point of zero charge (PZC) was determined to be
independent of IL ionicity as evaluated through physical property measurements by
Walden analysis. Increasing the alkyl chain length of cations was found to support the
formation of dense anion stern layers at positive electrode polarizations. Complimentary
surface enhanced Raman studies suggest long alkyl chains shield anions, allowing
increased anion density and increased differential capacitance. At negative electrode
polarization, these alkyl chains buckled forming ring structures. This study demonstrates a
pathway for tuning the charge density at the electrode-IL interface by varying cation alkyl chain lengths. Additionally, it furthers understanding of the complex interfacial capacitance behaviors displayed by ILs.
In order for ILs and DESs to achieve the potential as designer solvents utilized in electrochemical systems, a variety of additional studies building off the work discussed here are required. For ammonium ILs, no relationship between PZC capacitance and ionicity was observed likely due to a number of competing forces such as steric repulsion.
The nature of these interactions should be further probed to understand this discrepancy and fully understand the nature of the electrode-electrolyte interface. It is expected that
DESs display a similarly complicated electrode-electrolyte interface to ILs. While
preliminary work has probed this interface, the breadth of studies which have investigated
IL interfaces with methods like SERS, EIS, x-ray reflectometry, and simulations need replicating with DES systems. Additional electrochemical properties taken for granted in aqueous or organic solvents must be reevaluated in DESs and ILs. For example, while electrochemical windows have been characterized here, the breakdown species due to anodic oxidation or cathodic reduction are unknown.
104 For application of DESs in flow batteries, additional promising redox couples must be identified. Methyl viologen, while suitable for probing solvation dynamics, is not soluble enough in ethaline for practical use as a redox couple in energy storage. The quinones evaluated here did not demonstrate favorable properties for application as a redox couple in ethaline. However, additional studies should evaluate small redox active species in DESs as few currently exist. Instead of relying on dissolution, DESs composed of redox active species offer the potential to vastly improve the energy density of flow battery systems. In particular, the redox activity of the hydroquinone/choline chloride DES reported here should be evaluated in future studies. While significant progress has been made in understanding the tuning of DES and IL physical properties, the low conductivity and high viscosity of these systems still limit their application. While choline chloride based DESs exhibit the best performance thus far, novel DES species should be considered to further improve properties.
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