Electrochemical and Engineering Approaches

Electrochemical and Engineering Approaches Toward Technological Advancement of Non-aqueous Redox Flow Batteries By Musbaudeen O. Bamgbopa A Dissertation Presented to the Masdar Institute of Science and Technology in Partial Fulfillment of the Requirements for the Degree of Doctor of Philosophy In Interdisciplinary Engineering © 2017 Masdar Institute of Science and Technology All rights reserved

Abstract Non-aqueous redox flow batteries (NARFBs) – developed with the aim of providing higher energy density than their aqueous redox flow batteries (RFBs), have presently not been able to fulfil the target. The setbacks including; low active species solubility, costly electrolytes, high internal resistance and incompatible membranes, still persist – contributing to low cycle efficiency. This dissertation addresses these issues towards technological

advancements

of NARFBs for possible

future

commercialization. Firstly, a systematic examination of the use of solvent mixtures for NARFBs was explored to increase energy density and efficiency, considering active species solubility, electrolyte conductivity, and redox reaction rates among a number of binary and ternary solvent mixtures. An optimum solvent mixture was identified – which also provided good capacity retention, following studies on the influence of the solvent on crossover through different membranes was done. This work also presents a developed inexpensive, fast-charging capable iron–chro mium NARFB – pointing out an additional potential advantage NARFBs can have over aqueous RFBs. Development and testing of a composite Nafion/SiO 2 membrane provided low active permeability which helped achieve long cycling of the battery, thereby opening up the potential of NARFBs with multiple active species. To address issues encountered with commercial ion exchange membranes in NARFB electrolyte environments, as well as a cost reduction measure, this work also explored the prospects of applying membraneless designs for macro-scale RFBs. A new design of a membraneless RFB based on immiscible electrolytes was presented, capable of recharging and recirculation of the same electrolyte streams for multiple cycles and maintains the advantages of the decoupled power and energy densities. ii

Publications Excerpts from this dissertation have been published in the following form at present: M.O. Bamgbopa, Y. Shao-Horn, S. Al Mheiri, The Potential of Non-Aqueous Redox Flow Batteries as Fast Charging Capable Energy Storage Solutions: Demonstratio n with an Iron-Chromium Acetylacetonate Chemistry. Journal of Materials Chemistry A 5 (2017) 13457–13468. M.O. Bamgbopa, S. Al Mheiri, Influence of Solvents on Species Crossover and Capacity Decay in Non-Aqueous Vanadium Redox Flow Batteries: Characterization of Acetonitrile and 1, 3 Dioxolane Solvent Mixture. Journal of Power Sources 342 (2017) 371–381. M.O. Bamgbopa, N. Pour, Y. Shao-Horn, S. Al Mheiri, Systematic selection of solvent mixtures for non-aqueous redox flow batteries – vanadium acetylacetonate as a model system. Electrochmica Acta 223 (2017) 115–123. M.O. Bamgbopa, S. Al Mheiri, H. Sun, Prospects of recently developed membrane less cell designs for redox flow batteries. Renewable and Sustainable Energy Reviews 70 (2017) 506–518.

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This research was supported by the Government of Abu Dhabi to help fulfill the vision of the late President Sheikh Zayed Bin Sultan Al Nahyan for sustainable development and empowerment of the UAE and humankind.

ii

Acknowledgments With all praises to Allah; the great orchestrator of the universe, this work is dedicated to all my teachers. My first teachers, my parents and siblings; who put me on the right path and gave me a solid foundation - dear mother who recognized my talents before I even realized, those who have taught me formally and everyone who I have learned one or two things from so far, in this school of life. This is a major output from all your efforts. Special appreciation to Dr. Saif Almheiri; who has been not just an academic advisor, but a mentor, adviser and a big brother. To my dear wife, Rayan, my pillar of support and every member of Electrochem family – Ibrahim, Ayoob, Ali, Rahmat, Saad, Zainab, Fatma, Tawadood, Eman, Sara. Thanks to you all, that boy from Iponri estate is on the path to being the great scientist he wanted to be. Musbaudeen O. Bamgbopa, Masdar City, October 31 2017

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Contents _____________________________________________________________________

1 General Introduction and Research Motivation.............................................................. 1

2

1.1

Literature review and Background ...................................................................... 3

1.2

Research Objectives......................................................................................... 11

Systematic Solvent Mixture Selection for NARFBs...................................................... 13 2.1.

Background ..................................................................................................... 13

2.2.

Experimental ................................................................................................... 16

2.3.

Review of Reported Solvent Properties ............................................................. 20

2.4.

Solubility and Electrolyte Properties of Solvent Mixtures Containing Active Species 23

2.5. Charge-Discharge Performance of S3 Solvent Mixture (84/16 vol% AC/1,3DO) vs. Pure AC (status-quo) ................................................................................................... 41 2.6. 3

Summary......................................................................................................... 45

Multi-metal Active Species for NARFB: Fe - Cr Chemistry with Fast-charging Capability 47 3.1.

Background ..................................................................................................... 47

3.2.

Experimental ................................................................................................... 56

3.3. Voltammetric Measurements of Iron(III) and Chromium(III) acetylacetonate Redox Reactions.................................................................................................................... 59 3.4.

Fast-charging Capability ................................................................................... 63

3.5.

Charge–discharge performance of the Fe–Cr NARFB in H-cell ............................. 66

3.6.

Comparison of the Fe–Cr NARFB with other reported NARFBs......................... 69

3.7.

Summary......................................................................................................... 74

4 Solvent Influence on Crossover Through Membranes and Development of a Low permeability Membrane ................................................................................................. 75 4.1.

Background ..................................................................................................... 75

4.2.

Introduction .................................................................................................... 79

4.3.

Charge-discharge performance and EIS for V(acac)3 NARFB............................. 81

4.4.

Solvent influence on crossover and capacity decay............................................. 86 iv

5

6

4.5.

Examination of cycled membranes .................................................................... 93

4.6.

Charge-discharge performance of the Fe-Cr NARFB using NafionSi in S3 .......... 96

4.7.

Summary........................................................................................................101

Prospects of Membraneless Cell Designs for Macro-scale RFBs .................................103 5.1.

Background ....................................................................................................103

5.2.

Performance metrics for cell designs................................................................106

5.3.

What membraneless cell designs bring to the table ..........................................111

5.4.

Application of performance metrics to evaluate flow battery configurations ......130

5.5.

Summary........................................................................................................136

A Cyclable All-Iron Membraneless RFB Based on Immiscible Liquid Electrolytes .........138 6.1.

Background ....................................................................................................138

6.2.

Experimental ..................................................................................................142

6.3.

Battery Design ................................................................................................145

6.4.

Electrolyte properties and voltammetric measurements ...................................149

6.5.

Battery performance.......................................................................................154

6.6.

Immiscible Liquid (Two-Phase) Flow Modelling.................................................162

6.6.2. SIMULATION RESULT.......................................................................................164 6.7. 7

Summary........................................................................................................166

Conclusions and Future Work..................................................................................168

B Abbreviations ...........................................................................................................174 Bibliography..................................................................................................................175

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List of Tables _____________________________________________________________________

Table 2.1: Critical solvation parameters and properties of selected pure solvents at room temperature.* .................................................................................................................. 14 Table 2.2: Regression constants for solubility limit of V(acac) 3 in binary solvent mixtures examined in this study, which are used to fit experimental data shown in Figure 2.3. .......... 25 Table 2.3: Measured V(acac)3 solubility (𝑥𝐴) in ternary solvent systems. ........................... 27 Table 2.4:Sample voltammetric data of V(acac) 3 redox reactions in pure AC (shown in Figure 6b & Figure 6c) and 80/20 vol % AC/1,3DO mixture (for demonstration) from CV results at different scan rates. ......................................................................................................... 32 Table 2.5: Standard heterogeneous rate constants (𝑘) for V(acac)3 redox reactions in BII ternary mixture. ......................................................................................................................... 36 Table 2.6: Observed Raman vibrations in pure AC and pure 1,3DO as assigned in [75] and [76], respectively. ................................................................................................................... 37 Table 2.7: Measured conductivities of 0.1 M supporting electrolyte solution in pure AC and in the S3 mixture (84/16 vol% AC/1,3DO). *......................................................................... 38 Table 2.8: Performances of V(acac)3 NARFB electrolyte systems of present study and literature – with applied electrolytes and solvents. * .......................................................................... 43 Table 3.1: Comparison of previously reported single-species NARFBs with a Neosepta anionexchange membrane or a similar anion-exchange membrane. a ........................................... 51 Table 3.2: Redox couples, reactions, and potentials for Cr(acac) 3 [32] and Fe(acac)3 [83] in acetonitrile. .................................................................................................................... 52 Table 3.3: Cyclic voltammetry data for Fe3–Fe2 and Cr3–Cr2 at various scan rates on platinum.a ....................................................................................................................... 61 Table 3.4: Diffusion coefficients 𝐷 and standard heterogeneous rate constants 𝑘 of Fe3–Fe2 and Cr3–Cr2 couples in 0.05 M TEABF4 in acetonitrile at a platinum electrode. a ................ 61 Table 3.5: Performance parameters of the present Fe–Cr NARFB with its aqueous counterpart. ...................................................................................................................................... 72 Table 4.1: Characteristics of membranes in pure acetonitrile. ............................................. 81 Table 4.2: EIS parameters obtained using the equivalent circuit model shown in Figure 4.2 for cells containing a Nafion or NafionSi membrane and S3 solvent. ....................................... 84 Table 4.3: Estimated Vacac3 + (V4) permeability (𝑃) through the membranes.................... 91 Table 4.4: Comparison of previously reported single-species NARFBs with a Neosepta anionexchange membrane or a similar anion-exchange membrane. a ........................................... 98 Table 5.1: Summary of well-known performance metrics. ................................................107 Table 5.2: Fuel, flowing separating electrolyte, and oxidant stream combinations. a ............120 Table 5.3: Performance of various membraneless cell designs for fuel cells and flow batteries. a .....................................................................................................................................128

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Table 5.4: Design parameters and testing conditions for comparison of three cell configurations.a .............................................................................................................132 Table 5.5: Summary of performance metrics. a .................................................................133 Table 6.1: Efficiency and capacity decay comparison of previously reported membraneless RFB cell designs with some membrane based cell designs. a ..............................................140 Table 6.2: Cyclic voltammetry data for [Fe(acac)3 ]-/[Fe(acac)3 ]0 redox at various scan rates on carbon paper for the EA/IL electrolyte system. a ...............................................................151 Table 6.3: Design and nominal operating parameters for model ........................................163 Table 6.4: Summary of performance parameters of the present membraneless RFB.a..........166

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List of Figures _____________________________________________________________________

Figure 1.1: (a) Ragone plots (showing power vs. energy density) [4], and (b) Capital cost (energy vs power) based on prices estimates from 2002 [5] – for different secondary batteries and other energy storage options. ....................................................................................... 1 Figure 1.2: Schematic of a redox flow battery ..................................................................... 3 Figure 1.3: Schematic of electrode reaction and ion transport in typical (aqueous) redox flow battery.............................................................................................................................. 6 Figure 1.4: Classification of redox flow batteries by electrolyte and redox couple. ................ 7 Figure 2.1: Picture of test setup and expanded schematic of cell assembly. ......................... 18 Figure 2.2: (a) Previously reported densities (𝜌), (b) Molar volumes (𝑉𝑚), and (c) Dynamic viscosities (𝜇) of binary mixtures of AC with PC, EC, DMSO, DMF, and 1,3DO................ 22 Figure 2.3: V(acac)3 solubility (𝑥𝐴) in AC binary solvent mixtures. Colored symbols show experimental data, and dashed lines show correlations using the regression constants listed in Table 2.2. The data points for AC/DMSO were obtained from ref. [37]. Errors were estimated as mentioned in Section 2.2.1 - page 16 (averaged over mole fractions) were ±1.52, ±1.32, ±1.45, and ±1.36 (×10-3 ) for the PC, EC, DMF, and 1,3DO mixtures, respectively. ............. 24 Figure 2.4: Three methods of combining binary ratios/portions to give ternary ratios. The size of the circles are not representative of the actual rations (i.e. not drawn to scale). ................ 26 Figure 2.5: V(acac)3 solubility (𝑥𝐴) as a function of solvent vol % for AC/1,3DO/DMSO (top left), AC/1,3DO/DMF (top right) and AC/DMSO/DMF (bottom) ternary solvent mixture. Smooth areas in the figures were simulated from a combination of respective binary and ternary experimental results presented in Table 2.3 and Figure 2.4. ............................................... 28 Figure 2.6: Measured ion conductivities in binary mixtures of AC and co-solvents: (a) 0.1 M TBAClO4 , (b) electrolyte containing 0.01 M V(acac)3 and 0.05 M TBAClO4 ...................... 29 Figure 2.7: Steady-state cyclic voltammograms for redox reactions of V(acac) 3 (0.01 M) in solution with 0.05 M TBAClO 4 in pure AC. (a) Full CV showing both positive and negative side redos couples at 49 mv s -1 , (b) Negative side redox at different scan rates, and (a) Positive side redox at different scan rates. Three electrode setup was used with platinum working electrode, graphite foil counter and Ag/Ag+ reference electrodes. All CVs were compensated for iR drop...................................................................................................................... 30 Figure 2.8: Cyclic voltamograms of positive and negative side reactions in all binary solvent mixtures. 0.01M V(acac)3 with 0.05 M TBAClO4 support at 100 mV s-1 ............................. 33 Figure 2.9: Calculated standard heterogenoues rate constants (a) 𝑘neg for negative redox Vacac30 ↔ Vacac3 −, (b) 𝑘pos for negative redox Vacac3+ ↔ Vacac30 and (c) V(acac)3 diffusivity. Pt working electrode (uncertainty, ±7 × 10-4 cm s-1 )............................ 34 Figure 2.10: Observed Raman spectra of pure AC, pure 1,3DO, and various proportions of AC/1,3DO mixtures; (a) without solute, and (b) with V(acac) 3 as the test solute. Where S3 solvent mixture is approximately 20/80 vol% 1,3DO/AC seen in both figures..................... 39 viii

Figure 2.11: Cyclic voltammetry at a Pt wire electrode. The solutions comprised 0.1 M TBABF4 in pure AC, S3, or pure 1,3DO (inset). Scan rate was 100 mV s -1 . ...................................... 40 Figure 2.12: Charge-discharge efficiencies; (a) voltage (b) columbic and (c) energy for flow cell with 0.05 M V(acac)3 + 0.2 M TEABF4 in AC or S3. Conditions: 10 mA cm-2 constant current charge and discharge, with 2.4 V and 1.0 V high and low cut-of potentials. ............. 41 Figure 3.1: Cyclic voltammograms of Fe(acac) 3 and Cr(acac)3 with 0.05 M TEABF4 as the support electrolyte in acetonitrile at a platinum electrode. Also shown are the battery potential window, the desired reactions (Fe3–Fe2 and Cr3–Cr2), the undesired reactions (other peaks/redox couples), the charging allowance window, and the solvent potential window. Note that both voltammograms do not have the same y-axis scale. ............................................. 53 Figure 3.2: Glass test H-cell. Magnetic stirrers present in each arm (not clearly seen).......... 57 Figure 3.3: Cyclic voltammograms with holds at the vertex potentials for the; (a) Fe3–Fe2 and (b) Cr3–Cr2 redox couples at a platinum electrode with 0.02 M Fe(acac) 3 , 0.05 M Cr(acac)3, and 0.05 M TEABF4 in acetonitrile. Plots of (c) 𝑗p vs. 𝑣0.5 and (d) ln𝑗p vs. (𝐸p − 𝐸0′) for reactions of both species at the platinum electrode. Sample cyclic voltammograms with holds at the vertex potentials for the; (e) Fe3–Fe2 and (f) Cr3–Cr2 redox couples at a carbon paper electrode with 0.05 M Fe(acac)3 , 0.05 M Cr(acac)3 and 0.25 M TEABF4 in acetonitrile. Currents were normalized by the geometrical area. ......................................................................... 60 Figure 3.4: Constant-voltage charging profiles measured using a Neosepta membrane with 0.01 M Fe(acac)3 and Cr(acac)3 and 0.05 M TEABF4 in acetonitrile at 1.5, 2.5, 3.5, and 4.0 V in an H-cell. Overall recorded voltage and current (I) are plotted against time. The negative side (Cr side) potential was tracked vs. Ag/Ag+, and the positive (Fe side) potential was calculated as; Cr side + overall. The dashed red and blue lines represent the equilibrium potentials of Fe3– Fe2 (–1 V) and Cr3–Cr2 (–2.2 V), respectively, vs. Ag/Ag+, as deduced from Figure 3.1; both of these are desired reactions. The dotted red and blue lines represent the potentials of further Feacac30 oxidation (~1.5 V) and Cracac3 − reduction (–2.4 V), respectively, vs. Ag/Ag+, both of which are undesired reactions. The experiments were stopped when charging stabilized. . 64 Figure 3.5: Charging rates in an H-cell at two different charging voltages: 0.01 M Fe(acac) 3 and Cr(acac)3 with 0.05 M TEABF4 in acetonitrile using a TEA+-conducting Nafion membrane. 65 Figure 3.6: Coulombic, voltage, and energy efficiencies (CE, VE, and EE) versus cycle number under constant-voltage charging at 2.5 V in H-cells with (a) a BF4 -conducting Neosepta membrane (the inset shows an image of the H-cell used for cycling) and (b) a Na +-conducting and TEA+-conducting Nafion 115 membranes. Fe(acac) 3 (0.01 M) and Cr(acac)3 (0.01 M) with 0.05 M supporting electrolyte were used for all the tests, except tests of cells with the Na +conducting Nafion 115 membrane; in these tests, NaClO 4 instead of TEABF4 was used as the supporting electrolyte...................................................................................................... 66 Figure 4.1: Charge-discharge efficiencies; (a) voltage (b) columbic (c) energy of 0.05 M V(acac)3 + 0.2 M TEABF4 in AC or S3, using either Nafion or NafionSi membranes. Conditions: 10 mA cm-2 constant current charge and discharge in a 5 cm2 flow cell, with 2.4 V and 1.0 V high and low potential limits. ........................................................................... 83 Figure 4.2: Electrochemical impedance spectra obtained at open-circuit voltage in a solution of 0.05 M V(acac)3 + 0.2 M TEABF4 in S3, using either a Nafion or NafionSi membrane. The spectra were fitted with the equivalent circuit model shown in the inset [58]....................... 84 Figure 4.3: Relative discharge capacity (capacity normalized to first cycle) for the experiment shown in Figure 4.1; 0.05 M V(acac) 3 + 0.2 M TEABF4 in AC or S3, using either Nafion or NafionSi membranes....................................................................................................... 86

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Figure 4.4: Measured open-circuit voltage ( 𝐸OCV ) during self-discharge using an H-cell (shown in the inset) with AC or S3 as the solvent and Nafion or NafionSi as the membrane. Notice all four scenarios start at near equal 𝐸OCV (or SOC). .............................................. 88 Figure 4.5: Concentration profiles calculated from the recorded 𝐸OCV (Figure 4.4) during selfdischarge using Equation (4.8) for: Nafion in AC, NafionSi in AC, Nafion in S3, and NafionSi in S3. ............................................................................................................................. 89 Figure 4.6: Optical and SEM images, with tabulated EDX analyses, for cycled and uncycled membranes..................................................................................................................... 92 Figure 4.7: Neosepta AHA membrane (a) pristine, (b) after cycling in AC, (c) after cycling in S3 and (d) close-up of torn grating. .................................................................................. 95 Figure 4.8: Flow cell cycling results for 0.1 M Fe(acac) 3 and Cr(acac)3 with 0.4 M TEABF4 in S3 (84/16 (volume %) acetonitrile/1,3-dioxolane) with a NafionSi membrane. (a) CE, VE, EE, and relative discharge capacity over 50 cycles, recorded using 5 and 1 mA cm–2 charging and discharging currents, respectively. High and low cutoff voltages were 1.6 and 0.3 V, respectively. (b) Voltage profile for cycles 1 and three consecutive median cycles; 24, 25, and 26. The flow cell is pictured in the inset. (c) Electrochemical impedance spectra obtained at open-circuit voltage in the flow cell before cycling. (d) iR - free discharge polarization curve obtained at 50 % SOC (using procedure from [112]). Inset – constant voltage charging profile at 3 V until 25% theoretical capacity accumulated. ........................................................... 96 Figure 5.1: Schematic of a single-phase co-laminar flow cell design. ................................112 Figure 5.2: Schematic of a T-shaped single-phase co-laminar flow fuel cell. Reprinted with permission from [138]. Copyright (2005) Elsevier B. V....................................................114 Figure 5.3: Radial reactant inflow cell in (a) isometric view and (b) cross-sectional view. Reprinted with permission from [146]. Copyright (2008) Elsevier B. V.............................115 Figure 5.4: (a) Flow battery cell with single-phase co-laminar flow fabricated by Kjeang et al. with three-dimensional porous electrodes. Images of operation of the cell in (b) discharge mode and (c) recharge mode. Reprinted with permission from [148]. Copyright (2008) American Chemical Society...........................................................................................................116 Figure 5.5: Membraneless redox flow battery reactor PTFE, poly(tetrafluoroethylene). Reprinted with permission from [11]. Copyright (2010) Elsevier B.V. ..............................117 Figure 5.6: Flow regime in a typical flowing separating electrolyte cell design. .................118 Figure 5.7: Improved flowing separating electrolyte fuel cell design with reactant flow-over conducting fluid. Reprinted with permission from [153]. Copyright (2008) Elsevier B.V....119 Figure 5.8: Multiphase co-laminar flow fuel cells (MLFFCs) using gaseous fuel and oxidant with acidic flowing separating electrolyte (left). Reprinted with permission from [154]. Copyright (2007) American Chemical Society. Alkaline (right) flowing separating electrolytes. Reprinted with permission from [155]. Copyright (2009) The Electrochemical Society. PMMA, poly(methyl methacrylate)..............................................................................................121 Figure 5.9: Planar fuel cell design with liquid fuel and electrolyte streams and an air-breathing cathode. PMMA, poly(methyl methacrylate). Reprinted with permission from [160]. Copyright (2009) Elsevier B.V. ......................................................................................................122 Figure 5.10: Membraneless hybrid flow Cu–Pb redox flow battery with metal (Cu) deposition and dissolution. Reprinted with permission from [124]. Copyright (2008) Elsevier B.V. ....124 Figure 5.11: Co-laminar flow fuel cell with graphite rods that act both as electrodes and as electrically inert spacers. Reprinted with permission from [177]. Copyright (2007) Elsevier B.V...............................................................................................................................126

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Figure 5.12: Two-inlet/two-outlet all-vanadium redox flow battery with single-phase colaminar flow. Reprinted with permission from [13]. Copyright (2013) The Royal Society of Chemistry. ....................................................................................................................130 Figure 5.13: (a) Charge–discharge experiment results for D1. Reprinted with permission from [13]. Copyright (2013) The Royal Society of Chemistry. (b) & (c) Charge–discharge experiment results for D2n.aq & D2aq respectively. Reprinted with permission from [33]. Copyright (2010) Taylor & Francis Group. ......................................................................131 Figure 6.1: (a) Schematic voltammogram of simple ion transfer across an interface between an organic phase containing a strongly hydrophobic supporting electrolyte and an immiscible aqueous phase containing NaCl or K2SO 4 . The potential window of the interface is dictated by the aqueous anion/cation transfer potential ( ∆oraq ); ‘aq’ – aqueous, ‘or’ – organic. (b) Schematic of battery operation, showing redox reactions and accompanying ion transport. .146 Figure 6.2: (a) Measured ionic conductivity (κ), and estimated density (𝝆) of, EA/IL binary mixture at room temperature with varying volume percentages of IL. (b) Measured ionic conductivity of aqueous Fe 2+ solutions from FeCl2 and FeSO4 precursors at room temperature ‘aq’ – aqueous, ‘or’ – organic. ........................................................................................149 Figure 6.3: 𝑖𝑅 compensated cyclic voltammograms of 0.05 M Fe(acac)3 at a carbon paper electrode in 75/25 (v/v) EA/IL, with holds at the vertex potentials. The potential windows of pure IL and 75/25 (v/v) EA/IL on a platinum electrode are also shown. The y-axis of the voltammogram is normalized by the geometric area of the electrode. ................................151 Figure 6.4: (a–c) Schematic illustrations of the tested cell types: (a) static, gravity-controlled cylindrical cell with either graphite foil or carbon felt electrodes (Cell 1; d = approximate distance from liquid–liquid interface to electrode), (b) flowing-electrolyte cylindrical cell with graphite foil electrodes (Cell 2), and (c) conventional planar flow cell (Fuel Cell Technologies) with carbon paper electrodes (Cell 3). (d) Image of 0.05 M Fe(acac) 3 in 75/25 (v/v) EA/IL being placed in contact with 0.05 M FeSO 4 + 0.25 M K2SO4 aqueous (left), and image after 48 h with angular stirring at 100 rpm (right). ..................................................................................153 Figure 6.5: (a) Constant-current charge–discharge profiles (at 0.43 and 0.08 mA cm–2 , respectively) and (b) self-discharge profile. The profiles were measured using Cell 1 with graphite foil electrodes and 0.05 M FeSO 4 + 0.25 M K2 SO4 in water and 0.05 M Fe(acac)3 in 75/25 (v/v) EA/IL as the aqueous and organic phases, respectively. The dashed red and blue lines represent the equilibrium potentials of Fe 2+/Fe3+ (0.5 V vs Ag/AgCl) and [Fe(acac) 3]– /[Fe(acac)3 ]0 (–1.0 V vs. Ag/Ag+), as deduced from Equations (3.1) and (3.2), respectively. Cut-off potentials of 1.45 and 0.4 V were used during charging and discharging, respectively. .....................................................................................................................................155 Figure 6.6: (a) 𝑖𝑅 compensated polarization curves obtained with Cell 1 (static cell) using graphite foil electrodes (d = 0.5 cm) and carbon felt electrodes (d = 1.5 cm). 0.1 M FeSO4 + 0.5 M K2 SO4 was used in the aqueous electrolyte, and 0.1 M Fe(acac) 3 in the organic electrolyte at the indicated EA/IL volume ratio. (b) 𝑖𝑅-compensated polarization curves obtained with Cell 2 (flow cell) using carbon felt electrodes (d = 0.5 cm) and with Cell 3 (flow cell) using carbon paper electrodes (zero gap). ............................................................................................157 Figure 6.7: Coulombic efficiency (CE), voltage efficiency (VE), energy efficiency (EE), and relative discharge capacity over 25 cycles for Cell 3 cycling with constant-current charging– discharging at 1 and 0.2 mA cm–2 , respectively. Currents were normalized by the geometric area. For the flow cell experiments, 0.1 M FeSO 4 + 0.5 M NaCl in water was used as the aqueous phase, and 0.1 M Fe(acac)3 in 60/40 (v/v) EA/IL was used as the organic phase. 𝑅Ω,

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average area ohmic resistance across cell, obtained by means of electrochemical impedance spectroscopy..................................................................................................................158 Figure 6.8: Considered flow channels for the membraneless battery (used with Cell 3) and the prevalent flow regimes. ..................................................................................................160 Figure 6.9: 3D modelling domains for the membraneless battery under the considered flow channel setups. ..............................................................................................................162 Figure 6.10: (a) Predicted phase separation in the modelled free (left) and parallel (right) channel flow cell domains, as indicated by aqueous phase saturation (𝑠𝑎𝑞), recall 𝑠𝑜𝑟 = 1 − 𝑠𝑎𝑞. (b) Velocity magnitude in the porous electrodes of both arrangements. ....................165

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1

CHAPTER 1 1 General Introduction and Research Motivation

Energy storage is one of the most pressing challenges for renewable energy power plants, particularly solar photovoltaic facilities and wind farms, where the applicatio n of electrochemical energy storage demonstrates high response times and round-trip efficiencies [1]. Redox flow batteries (RFBs) are considered to be a viable energystorage option that should be thoroughly evaluated for utility-scale use [2]. Although commercial RFBs currently have relatively lower energy density (amongst other electrochemical

energy storage options - Figure

1.1a), their most attractive

characteristics are the decoupled nature of their volumetric power and energy densities [3]. This decoupling economizes capacity expansion costs (Figure 1.1b). RFBs also have low self-discharge rate; thus making them suitable for long-duration large scale storage.

Figure 1.1: (a) Ragone plots (showing power vs. energy density) [4], and (b) Capital cost (energy vs power) based on prices estimates from 2002 [5] – for different secondary batteries and other energy storage options.

2

Therefore, wider adoption of RFBs is, subject to their ability to compete with other available options (such as lithium ion batteries) in terms of cost, charge–discharge cycle efficiency, reliability, and so on – under a range of operating conditions.

3

1.1 Literature review and Background

Figure 1.2: Schematic of a redox flow battery Figure 1.2 depicts a schematic of a typical RFB, which is in principle akin to a regenerative fuel cell operating with liquid electrolytes. As suggested earlier, the RFB electrolyte volume capacity of the cell stack is decoupled from the external electrolyte storage capacity, whereas a conventional (solid-state) battery has a fixed total amount of electrolyte dictated by the cell volume. During discharging of a RFB, the anolyte and catholyte containing the electroactive species are supplied from their respective external storage tanks and are pumped through the cell to supply current to a load. Active species are oxidized at the negative electrode (anode) and reduced at the positive electrode (cathode) respectively, and the spent electrolyte returns to the external storage tank. During charging, oxidation and reduction occur at the positive and negative electrodes respectively, (which is the opposite of what occurs during discharging), while current is supplied by a charging power source (Figure 1.2). As seen in Figure 1.2, the main components of a redox flow cell are the electrode (electrode material), the membrane or separator, the electrolyte, and the cell casing, similar to other electrochemical devices [7].

4

1.1.1. ELECTRODE MATERIAL It is worthy of note to distinguish between the term ‘electrode’, which describes a halfcell, and the term ‘electrode material’, which refers to the material used for the electrically active area of the half-cell. Although the electrode material does not typically participate directly in the electrochemical reaction, it acts as or carries the reaction sites (as seen in Figure 1.3) and thus must have a high active surface area. Generally, the requirements for a suitable electrode material include; A. Compatibility with redox couple chemistry The electrode material must be compatible with the physical chemistry of the redox couple. For example, in aqueous hybrid chemistries such as zinc–bromine [2] and zinc– cerium [8], where the electrode is plated with an active species (zinc), the electrode must be planar and have a high surface area. Electrode materials made of bare or coated metals and meshes of metals such as titanium, platinum, gold, and palladium have been proposed for use in such RFBs [9], [10]. For redox chemistries that do not involve plating of the active species, nonmetal porous electrode materials including graphite or carbon felt [11] [12], foams [13], fibers [14], and papers have been widely reported. B. Compatibility with electrolyte chemical composition Selection of a suitable electrode material depends on the working environment in which it will be used. For example, if a corrosive electrolyte will be used, the electrode material must be resistant to corrosion [15]. Furthermore, the material should be stable in its working environment, which is one major reason that carbon-based electrodes are more widely used than metal based electrodes. The operating temperature and pressure range of the RFB and the possibility of gas evolution during operation are also aspects of the working environment that must be considered.

5

C. Compatibility with flow architecture RFB designs can be grouped into three categories on the basis of their flow architecture : flow over, flow through, and air breathing [16]. As the name suggests, in the flow-over architecture, the electrolyte flows over a planar electrode material, as is the case on the lithium side of the RFB reported by Hamelet et al. [17] and most of the laminar flow fuel cells reviewed by Shaegh et al. [16]. Porous electrode materials are used in flowthrough RFBs, as well as in air-breathing RFBs. The porosity of the electrode material is usually higher for air-breathing designs than for flow-through designs. The most critical properties that must be considered in investigation of the suitability of an electrode material for a particular flow architecture are the material’s directiona l electrical conductivity and its mass transfer characteristics [18]. The electrode material’s electrochemical activity is also important [19]. 1.1.2. ION-EXCHANGE MEMBRANE The IEM physically separates the two half-cells of a RFB (Figure 1.3) so that reactants do not mix, and is responsible for selective ion transfer. The choice of membrane is largely dictated by the working electrolyte, where the ions to be conducted emanate from. Consequently, anion or cation-exchange membranes are used to transport anions and cations, respectively, and to limit reactant crossover. Figure 1.3 depicts typical ion transport (aqueous system) that consists of active species A and B. Hydrogen ions from sulfuric acid dissolved in the electrolyte (as supporting electrolyte) are exchanged between the positive and negative sides of the cell through the membrane. Like the electrode, the IEM must be stable in its chemical environment and mechanically capable of withstanding the pressures exerted on its opposite sides.

6

e-

Charge Discharge

eSO42-

A(n+x)+

SO42-

H+

-xe-

+

B(m-y)+ -ye-

H+

An+

SO42-

Bm+ SO4

2-

-

Figure 1.3: Schematic of electrode reaction and ion transport in typical (aqueous) redox flow battery. The currently available membranes exhibit various drawbacks. For example, the Nafion membrane that is widely used suffers from reactant crossover, although it possesses adequate ion conductivity

[7]. Commercially

available

membranes

for other

applications such as desalination may be modified to try meet the ion conductivity and selectivity requirements for use in RFBs in a chemica lly and mechanically stable manner. Ion-exchange membranes (IEMs), which were adopted from the conventional fuel ce ll architecture can account for 20–40% of the cost of a RFB [20]. Evaluation of cost reduction opportunities is therefore seen as a critical area for RFB research. 1.1.3. ELECTROLYTE The bulk of the research on RFBs in recent years has focused on the electrolyte, and various combinations of active species have been investigated. These investigatio ns typically involve the following steps: a. A redox couple is selected on the basis of its standard redox potential.

7

b. The precursors of the active species of the redox couple are dissolved in a solvent, and a supporting electrolyte is added to impart conductivity. c. Electrochemical characterization is performed on the electrolyte mixture to determine the electrochemical window, the open-circuit voltage (OCV), and the reversibility of the redox couple. d. Charge and discharge experiments are performed to evaluate the cyclability and the performance of active species over an extend period of time; hence simulating the actual conditions of a working battery. e. Subsequently, the performance of a novel cell design is then evaluated using the new electrolyte mixture. In Figure 1.4, a classification scheme for RFB electrolytes on the basis of solvent or redox couple is presented, using information presented in a detailed review by Wang et al. [21] of all electrolytes reported up to late 2012. RFB electrolytes, which are mixtures of active species, solvents, and supporting electrolytes, are usually classified by solvent because their complexity makes classification by redox couple more difficult.

RFB Classification

By Solvent

BY Redox Couple

Aqueous

Non-Aqueous

Metallic

Organic Solvents

Ionic Liquids and Deep Eutectic Solvents

Hybrid Solvent Systems

Organic

Figure 1.4: Classification of redox flow batteries by electrolyte and redox couple.

8

A. Aqueous redox flow batteries The most common aqueous RFBs use all-vanadium, bromide/polysulfide, iron/chromium electrolyte systems. The iron/chromium technology,

and

which was

developed first, is based on an Fe2+/Fe3+ redox couple on the positive side and a Cr2+/Cr3+ redox couple on the negative side, with hydrochloric acid as the supporting electrolyte. The major drawbacks of this technology are low OCV (~1.2 V) [7], poor reversibility on the chromium side, and cross-contamination from reactant crossover [22]. The all-vanadium chemistry developed by Skyllas-Kazacos et al. [10] is used in most of the commercially available RFB systems and is the most studied. In this system, sulfuric acid is the supporting electrolyte, and the redox couples on the negative and positive sides are V2+/V3+ and VO 2+/VO 2 +, respectively. The theoretical standard cell voltage for this couple is 1.26 V, and the energy density ranges from 25 to 35 Wh/kg [6]. These aqueous RFBs are generally limited to relatively low OCVs and energy densities, owing to the narrow electrochemical window of water. As a result, achieving long term cycling of the aqueous electrolytes is challenging due to the possibility of electrolyzing the solvent. Therefore, alternative solvents have been considered to overcome this limitation. B. Non-aqueous redox flow batteries as present research focus This doctoral research majorly focuses on technological advancement of a non-aqueous redox flow battery (NARFB) as an alternative to the available aqueous systems. This follows recent research endeavors investigating various non-aqueous solvents for RFBs, the most common being organic solvents. In addition to having a wider electrochemical window than water, some organic solvents also have higher and lower boiling and melting points respectively [23]. Such solvents eradicate precipitatio n

9

problem of active species as with water [24]–[28], permitting operation in wider temperature ranges. The most frequently used non-aqueous electrolyte systems consist of metal coordination complexes as active species dissolved in organic solvents. In a typical non-aqueous electrolyte system with a transition metal (M)–ligand (L) complex active species, the electrochemical reaction equations are as follows: Cathode

[MLn ](𝑧+𝑥 )+ ↔ [MLn ]𝑧+ + 𝑥e−

(1.1)

Anode

[MLn ](𝑧−𝑦)+ + 𝑦e− ↔ [MLn ]𝑧+

(1.2)

Matsuda et al. [29] were among the first to report such a system; they used a single ruthenium coordination complex, tris(2,2-bipyridine)ruthenium(II) tetrafluorobor ate, [Ru(bpy)3 ](BF4 )2 , as the active species in acetonitrile with tetraethylammo nium tetrafluoroborate as the supporting electrolyte. They obtained an OCV of 2.6 V (twice that of an aqueous all-vanadium system) in a RFB with an anion-exchange membrane and carbon fiber cloth as the electrode material. Unlike many aqueous RFBs, systems based on a single- metal active species are less subjected to the detrimental effects of reactants crossover. Liu et al. [30]–[32] have investigated

the potential

of

acetylacetonate (acac) complexes of manganese, chromium, or vanadium dissolved in acetonitrile with tetraethylammonium tetrafluoroborate as the supporting electrolyte. Cyclic voltammetry results showed OCVs of 3.6, 2.2, and 1.1 V for Cr(acac)3 , V(acac)3 , and Mn(acac)3 , respectively. Chakrabarti et al. reported a notable membrane-free cell design, which they used to investigate

few metal-complex

active

species such as Ru(acac)3

[11] and

[Fe(bpy)3 ](ClO 4 )2 [33]. These investigators also evaluated the physical properties and electrochemical activities of metal complex electrolytes brewed in deep eutectic solvents [34]. Deep eutectic solvents and ionic liquids are seen as cheaper, more

10

environmentally friendly alternatives to widely used organic solvents. Chakrabarti et. al. have reviewed the prospects of these alternative solvent options [35], but, only very few such electrolytes have been tested in a real flow battery cell. Although non-aqueous solvent based RFBs have higher OCVs than their aqueous counterparts (indicating that they have the potential to have high energy densities), a major challenge for these systems is the low solubility of the active metal complexes in the solvents; thus limiting their commercial availability. Furthermore, some of the investigated systems showed slow electrochemical kinetics, which leads to a low recharge capacity. Note that these are chemical problems rather than design problems, which has led to the suggestion that “hybrid” electrolyte systems—that is, systems involving a mixture of different types of electrolytes, such as multiple organic solvents [36], [37] or combinations of organic solvents and ionic liquids [38]—be used to impart the desired individual properties. The proposed doctoral research therefore focuses on improvement in these various aspects (discussed in Section 1.1.1 to 1.1.3) to develop new and improved non-aqueous redox flow battery (NARFB) design approaches and systems.

11

1.2 Research Objectives The objective of this doctoral research is to investigate various improvements as technological advancements on the NARFB status-quo and consequently develop novel NARFB system/approaches with adequate scientific/theoretical background. The main objectives that will be focused on as performance metrics are; energy density, energy efficiency and capacity retention. Given the end goal would be to pave the way for commercialization of NARFBs, cost implications are also considered at each developmental/investigation stage. Recall the difference between a NARFB and its aqueous counterpart is the solvent which is the bulk carrier of the electrolyte system. Subsequently, the first point of focus of this study is to achieve a new electrolyte system. Specifically focus on investiga ting improvements to solvent and active species – with improved energy density and efficiency as targets, in addition to electrode kinetics between the resulting electrolyte system and applicable electrode materials. Secondly, In the light of the compatibility issues highlighted in Sections 1.1.1 (page 4), where energy efficiency is dependent on the interaction of other cell components with the electrolyte system, membrane compatibility and improvement studies are also undertaken herein. Lastly, the viability of a developed membraneless cell design will be investigated to address certain issues (such as cost, stability, etc.) that arise when commercia l membranes are adopted for NARFBs. Each chapter of this dissertation (after this introduction), will be initiated with a chapter specific background to reinforce the motivation/objectives of the investigatio ns presented in the chapter and the goals to be achieved. The chapter specific introductio ns will also be followed by an ‘experimental’ subsection - highlighting the experimenta l

11

12

methods used for each chapter. The dissertation takes this compilation approach to aid reader assimilation because of the multitude of experiments to be reported.

12

13

CHAPTER 2 2

Systematic Solvent Mixture Selection for NARFBs

2.1.Background An upper bound in the energy density (𝐸max in J L-1 ) for a RFB can be estimated by the following equation: 𝐸max =

1 𝑉 𝑛 𝐹 𝐶A 2 oc

(2.1)

where 𝑛 is the number of electrons transferred per molecule of active species during the redox reaction, 𝑉oc is the open circuit voltage (V), 𝐶A is the solubility of the active species (mol L-1 ) and 𝐹 is Faraday’s constant (96487 C mol-1 ). Much research has been focused on the electrolyte chemistry to increase the solubility of the active species for enhanced energy density and raise the conductivity for enhanced power densities [39]. The state-of-the-art RFBs are based on all-vanad ium chemistry in aqueous electrolytes developed by Skyllas-Kazacos et al. [10]. The vanadium RFB (VRFB) is comprised of the VIVO2+/VVO 2 + redox couple for the positive electrode and the VII/VIII redox couple for the negative electrode, where the reactants are typically dissolved

in aqueous sulfuric

acid solution.

Commercial VRFB

electrolytes contain between 1.6 to 2 M vanadium ions (𝐶A in Equation (2.1)) and the major challenge to further increase their energy density is the limited solubility of the different vanadium ions outside the temperature range of 10 – 40o C [40]. The 𝑉oc values of such systems are also constrained by the electrochemical window of water.

13

14

The intention of the present study to seek an alternative approach to aqueous RFBs by developing non-aqueous RFBs (NARFBs) - thereby replacing water with non-aqueous solvents with a wider electrochemical window [29], [41]–[43] such as acetonitrile (AC, ~6 V) or propylene carbonate (PC, 6.5 V) [23]. However, the most used active species in non-aqueous electrolytes (i.e., metal coordination complexes) have lower solubility [44] than those used in aqueous electrolytes, which results in a trade-off between 𝑉oc and 𝐶A . Although the metal coordination complexes used as active species in nonaqueous RFBs [11], [30]–[32] typically undergo single-electron redox reactions [45], recent studies show that new metal complexes can provide multi-electron redox reactions [46], [47]. Besides, complexes with non-innocent ligands have been investigated [48]. Table 2.1: Critical solvation parameters and properties of selected pure solvents at room temperature.* 𝑽𝒎

Solubility

[49]

𝑴𝒎

μ

κ

AN

Pure

(mol

(g

[49]

(μS

[23],

DN

[50],

ε

[37], [51]

δp [50]

Solvent

cm-3 )

mol -1 )

(cP)

cm-1 )

[50]

[51]

[52]

[49]

(M)

(MPa0.5 )

1,3DO

69.89

74.08

0.60

0.66

18.3

15.1

43.1

7.34

0.80

6.6

AC

52.23

41.05

0.34

0.77

18.9

14.1

45.6

35.90

0.60

18.0

DMF

77.10

73.09

0.80

0.35

16.0

26.6

43.8

36.70

0.51

13.7

DMSO

71.00

78.13

1.99

0.92

19.3

29.8

45.1

46.50

0.30

16.4

EC†

66.66

88.06

2.56

1.56

—

16.4

48.6

89.60

—

21.7

PC

84.72

102.09

2.53

0.43

18.3

15.1

46.0

64.92

0.07

18.0

ET(30)

of V(acac)3

* Symbols

and abbreviations: ε, dielectric constant; δp , Hansen polarity parameter (polar contribution to Hildebrand solubility parameter, δ); κ, conductivity; μ, dynamic viscosity; Vm , molar volume; Mm , molar mass; AN, acceptor number; DN, donor number; ET(30), solvent polarity parameter. †EC is a solid at room temperature, so the listed values for κ and ET(30) are for melted EC at ~40 °C.

Most NARFB chemistries have used AC as the solvent [29] because it can provide high conductivity resulting from its high polarity (see Table 2.1). Unfortunately, a number

14

15

of active species developed have shown low solubility in AC [53], and research efforts have been focused on searching for alternative solvents to enhance solubility. For example, Herr et al. [36], [37] replaced AC or mixed it with other organic solvents, and Zhang et al. [38] used combinations of organic solvents and ionic liquids. However, these studies only focused on enhancing solubility without attention to possible compromises in reaction rates and consequently, energy efficiency. Thus, this calls for a systematic approach of solvent mixtures selection. Herein, binary solvent mixtures for NARFBs are evaluated by considering solubility, conductivity and reaction rates to increase: (a) volumetric energy density (via addressing the trade-off between 𝑉oc and 𝐶A - Equation (2.1)), and (b) reaction kinetics towards improved energy efficiencies of electrolytes from select solvent mixtures. This work differs from previous studies [36], [37] in that it systematically addresses the trade-offs (or absolute improvements) that affect NARFB energy density and efficie nc y when binary solvents mixtures are used. This approach allows reasonable extrapolatio n of the results to predict the performance of NARFB with ternary solvents mixtures. Vanadium acetylacetonate V(acac)3 is selected as a model active species to demonstrate the systematic approach. The results presented herein can potentially encourage the NARFB community

to better address organic solvent selection issues before

developing new active species while maintaining the same solvent. Such issues will have to be addressed if non-aqueous rivals to the existing commercial all-vanadium and other aqueous RFBs are to be developed.

15

16

2.2.Experimental Acetonitrile (AC, anhydrous 99.8+%, Alfa-Aesar), ethylene carbonate (EC) and propylene carbonate (PC, BASF), 1,3-dioxolane (1,3DO, anhydrous 99.8% with ~75 ppm inhibitor, Sigma-Aldrich), dimethyl sulfoxide (DMSO, ≥99.9%, Sigma-Aldric h), and dimethyl formamide (DMF, anhydrous 99.8%, Sigma-Aldrich) were dried over activated 4 Å molecular sieves (Alfa-Aesar) to a moisture content of ~6 ppm as verified with a Karl Fischer titration setup (899 Coulometer, Metrohm). The supporting electrolytes;

Tetrabutyl

tetraethylammonium tetrafluoroborate

ammonium

tetrafluoroborate

perchlorate (TEABF4 ,

(TBAClO 4 , 99%),

≥99.0%),

tetrabuylammo nium

(TBABF4 , 99%) and tetraethylammonium

hexaflorophosp hate

(TEAPF6 , 98%), and the active species vanadium III acetylacetonate (V(acav)3 , 97%), were obtained from Sigma-Aldrich. Sodium perchlorate (NaClO 4 , 98.0-102.0%) and tetraethylammonium perchlorate (TBAClO 4 , 98%) were obtained from Alfa-Aesar. 2.2.1. ACTIVE SPECIES SOLUBILITY MEASUREMENTS The classic gravimetric method was used to determine solubility [36], [37], [54]–[56]. The weighed solute was gradually dissolved in the solvent or solvent mixture with vigorous stirring. At overnight intervals, more solute (in 0.02–0.04 g increments) was added until precipitation occurred. Therefore, the solubility limits determined in this way were accurate to approximately ±0.04 g (the maximum increment). The method used to gauge maximum solubility in the present study produced results consistent with literature results [37] when checked for the value for V(acac)3 in pure PC (0.07 M, Table 2.1). 2.2.2. CONDUCTIVITY MEASURMENTS The conductivities of solutions in the dried solvents were tested with a VWR Bench/Portable Conductivity

Meter (VWR International)

probe with platinum

16

17

electrodes or with a Metrohm 912 conductometer. Readings were taken after meter stabilization while solutions were magnetically stirred at approximately 600 rpm. 2.2.3. CYCLIC

VOLTAMMETRY

AND

ELECTROCHEMICAL

IMPEDANCE SPECTROSCOPY (EIS) Electrochemical measurements were conducted with a Biologic SP 300 potentiostat or a with a Parstat MC potentiostat (Princeton Applied Research) with supplied softwares. All experiments were performed in an argon-filled glove box at room temperature (25 ± 3 °C) as gauged by an in-box mercury-in- glass thermometer. Cyclic voltammetry was performed with a 3-electrode setup, with platinum wire (3 mm diameter, 99.9%, VWR) as the working electrode. An Ag/Ag+ reference electrode was prepared with 0.01 M AgNO 3 and 0.1 M TBAClO 4 in AC. The reference electrode was tested and found to be stable when used in electrolytes based on solvent mixtures as the case with pure AC. A comparison of voltammograms indicated little or no shift in peak potentials in most cases, and even when a shift did occur, both the positive and the negative side peak shifts were precisely consistent. The area of the working electrode exposed to the electrolyte was 2 mm2 . A graphite foil counter electrode was used. Test electrolyte solutions consisted of 2.5 mL of solvent or solvent mixture containing 0.01 M V(acac)3 and 0.05 M TBAClO 4 . Each electrolyte mix was tested 3 times, and the values were averaged for final reporting. New platinum electrodes were polished with 2 µm grit polishing paper (3M) and rinsed in the test solvent before and after test repetitio ns, whereas the graphite foil counter electrode was discarded and replaced for each test. Potential window CVs were performed in a three-electrode setup, with a platinum wire (3 mm diameter, 99.9%, Sigma Aldrich) as the working electrode. Lithium foil (0.2 mm thick, 99.9%, Sigma Aldrich) and a graphite foil (Alfa-Aesar) were used as the reference and counter electrodes, respectively. All cyclic voltammograms were

17

18

compensated for the accompanying iR drop (Ohmic losses). For electrochemica l impedance spectroscopy, the acquisition of the impedance spectra was performed at open-circuit voltage with frequencies between 100 kHz and 0.1 Hz and at an amplitude of 10 mV RMS.

Figure 2.1: Picture of test setup and expanded schematic of cell assembly. 2.2.4. CHARGE-DISCHARGE CYCLING Charge-discharge testing was conducted using the test setup and flow cell assembly (Fuel Cell Technologies) shown in Figure 2.1 with 0.05 M active species and 0.2 M TEABF4 in the solvent. The flow cell assembly had a 5 cm2 single serpentine flow channel cross-section. The flow cell tanks on each side of the assembly contained 30 mL of electrolyte. The electrodes consisted of carbon paper (SGL 10 BA, SGL Carbon) positioned on each side of a two-electrode setup for the flow cell. Peristaltic pumps

18

19

(Masterflex) were used to circulate the electrolyte through the flow cell at a rate of 35 mL min–1 . 2.2.5. ELECTRODE AND MEMBRANE PREPARATION The as-received carbon paper electrodes and the membranes were cleaned by sonicatio n in HNO 3 solution, rinsed thoroughly in deionized water, and dried overnight under vacuum. Nafion 115 (Dupont) membranes conductive to tetraethylammonium cation (TEA+) were prepared as described by Park et al. [57]. After cleaning and drying the Nafion 115 membranes, they were heated in 0.5 M aqueous sulfuric acid at 70°C for 2 h to convert them to the H+ form. The membranes were rinsed in deionized water and then soaked in deionized water for 1 h at 70°C to remove excess acid. The H+-forms of the membranes were neutralized with a 1 M methanolic solution of tetrabutylammonium hydroxide at room temperature for 18 h. The membranes were again rinsed with deionized water and vacuum dried at 80 o C. Following preparation of the membranes, the electrodes and membranes were transferred to the glove box. Prior to use, the electrodes and membranes were conditioned as described previously ([31], [58]) by soaking in the 0.1 M support solution for at least 4 h. 2.2.6. RAMAN SPECTROSCOPY Spectra from molecular vibrations of the solvent were evaluated with a Raman microscope (WiTech Alpha 300) using a 532 nm green laser.

19

20

2.3.Review of Reported Solvent Properties 2.3.1. PURE SOLVENTS A wide range of fundamental physical and chemical parameters have been used to describe and model single solvent, solvent–solvent, and solvent–solute interactions, as well as the physical properties of the resulting solutions. These physical and chemica l parameters can be used in conjunction with experimental data (as is done in herein) to narrow down the choice of a solvent for a given purpose. As far as NARFBs are concerned, the solubility of the active species and the conductivity of the resulting electrolyte solution are highly important, as mentioned earlier. To search for solvent mixtures to replace AC, the properties of select pure solvents (PC, EC, DMSO, and DMF, and 1,3DO) that govern ionic solubility and ionic conductivity were reviewed; these properties are listed in Table 2.1. To describe the trends in the contents of Table 2.1, Shinkle [51] had shown that the solubility of V(acac)3 and ionic conductivity is enhanced in solvents with increasing solvent polarity parameter (ET (30)), Hildebrand solubility parameter (δ), and solvent molar volume (𝑉𝑚 ), in an order of less correlation. The influence of acceptor number and donor number on solubility and ionic conductivity came in a distant fourth and fifth, respectively [51]. It should be noted that ionic conductivity is proportional to the concentration of solvated ions and their mobility [59], which is governed largely by the solvent viscosity [60]. 2.3.2. BINARY SOLVENT MIXTURES The standard approach to describe the density (𝜌𝑚𝑖𝑥 ) or molar volume (𝑉𝑚 ) of a mixture of two miscible liquids is the additive volume approach, shown in Equation (2.2) for density [61]: 𝜌𝑚𝑖𝑥 = 𝑥1 𝜌1 + (1 − 𝑥1)𝜌2

(2.2)

20

21

Where 𝑥 1 is the mole fraction of solvent 1, and 𝜌1 and 𝜌2 are the densities of solvents 1 and 2, respectively. However, experimental observations have shown that this approach often results in underestimation of density or molar volume (𝑉𝑚 = 𝜌⁄𝑀 ). 𝑚 Previously reported experimental results of densities and molar volumes of binary mixtures of AC and PC [62], EC [63], DMSO [64], DMF [65], or 1,3DO [66] are plotted as a function of (1 – 𝑥𝐴𝐶 ) - the mole fraction of the co-solvent with AC, as shown in Figure 2.2a and Figure 2.2b, respectively. Measured viscosities of mixtures of AC and PC [62], EC [63], DMSO [64], DMF [65], or 1,3DO [66] as a function of 1 – 𝑥𝐴𝐶 are seen in Figure 2.2c. Jouyban et al. [67] modeled the dynamic viscosity (𝜇𝑚𝑖𝑥,𝑇 ) of binary solvent mixtures and found good correlation with experimental data using Equation (2.3) instead of Equation (2.2). ln 𝜇𝑚𝑖𝑥,𝑇 = 𝑥 1 ln 𝜇1,𝑇 + 𝑥 1 ln 𝜇1,𝑇 𝑥 𝑥 + 1 2 [−61.784 + 54.566 (𝐸𝑚1 − 𝐸𝑚2 )2 𝑇 − 129.759 (𝑆1 − 𝑆2 )2 − 1987.988 (𝐴1 − 𝐴2 )2 + 331.691 (𝐵1 − 𝐵2 )2 + 190.370 (𝑉1 − 𝑉2 )2 ] 𝑥 𝑥 (𝑥 − 𝑥 2 ) [706.352 (𝐴1 − 𝐴2 )2 + 1 2 1 𝑇 + 65.119 (𝑉1 − 𝑉2 )2 ]

(2.3)

Where the subscripts “1” and “2” designate the two solvents, 𝑥 is mole fraction, 𝑇 is temperature, 𝐸𝑚 is excess molar fraction, 𝑆 is dipolarity/polarizability, 𝐴 is hydrogenbond acidity, 𝐵 is hydrogen-bond basicity, and 𝑉 is McGowan volume [67]. As seen in Figure 2.2, pure AC has the lowest molar volume and viscosity. Both molar volume and viscosity increases with the addition of the co-solvents. Densities as predicted by Equation (2.2) are depicted as straight lines in Figure 2.2. Notice that the difference (excess) between predicted values (using Equation (2.2)) and experimenta l results for density (𝑉𝑚 similarly) increases with the severity of the difference between the pure co-solvent density and that of pure AC.

21

22

Figure 2.2: (a) Previously reported densities (𝝆), (b) Molar volumes (𝑽𝒎 ), and (c) Dynamic viscosities (𝝁) of binary mixtures of AC with PC, EC, DMSO, DMF, and 1,3DO. Therefore, this suggests that Equation (2.2) can be adequate to model mixtures in which the density of the individual components making the mixture is similar; for example, mixtures of AC with DMF, 1,3DO and DMSO as seen Figure 2.2a. For those mixtures of AC with EC and PC also seen in Figure 2.2a, a higher order (non-linear) correlation is needed for a more accurate fit with experimental data. Evaluation of differe nt correlations that fit the experimental density data in Figure 2.2a is not the focus of the present study. Further correlation of observed trends to electrolyte properties are established in the following sections.

22

23

2.4.Solubility and Electrolyte Properties of Solvent Mixtures Containing Active Species 2.4.1. ACTIVE SPECIES SOLUBILITY IN SOLVENT MIXTURES The solubility 𝑥𝐴 (mol mol-1 ) in this study was defined by Equation (2.4) [68], where subscripts 𝐴 and 𝑖 designate the active species and solvent 𝑖 in the multi-solve nt system, 𝑚 is mass, and 𝑀 is molar mass. We used 𝑥𝐴 rather than molarity, to compare among solvent mixtures consisting of various amounts of different solvents. 𝑚𝐴 ⁄𝑀 𝐴 𝑥𝐴 = 𝑚 𝑚 𝐴 ⁄𝑀 + ∑𝑖 𝑖⁄𝑀 𝐴

(2.4)

𝑖

The measured solubility of V(acac)3 in a number of binary solvent mixtures are plotted as a function of 1 – 𝑥𝐴𝐶 , as shown in Figure 2.3. Solubility was found to decrease as greater mole fractions of PC and EC. Increasing the mole fraction of DMF, DMSO, or 1,3DO first increased the solubility then followed by a decline. This increase in the solubility by introducing DMF, DMSO, or 1,3DO as co-solvents to AC can be attributed to preferential solvation of V(acac)3 [40], where one solvent aggregates around the solute to a greater extent than its co-solvent in a solvent mixture solution.

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Figure 2.3: V(acac)3 solubility (𝒙𝑨 ) in AC binary solvent mixtures. Colored symbols show experimental data, and dashed lines show correlations using the regression constants listed in Table 2.2. The data points for AC/DMSO were obtained from ref. [37]. Errors were estimated as mentioned in Section 2.2.1 - page 16 (averaged over mole fractions) were ±1.52, ±1.32, ±1.45, and ±1.36 (×10 -3 ) for the PC, EC, DMF, and 1,3DO mixtures, respectively. V(acac)3 solubility was next correlated with the measured values of solvatio n parameters as listed in Table 2.1. For binary mixtures of AC with EC and PC, V(acac)3 solubility (Figure 2.3) is inversely correlated with molar volume (𝑉𝑚 , Figure 2.2b). For the binary mixtures of AC with 1,3DO, DMF, and DMSO where preferential solvatio n occurs, the strong correlation between solubility and ET(30) earlier concluded in [51] is suspected. This is due to the fact that ET(30) has been shown to be capable of describing preferential solvation in binary solvent mixtures [69]. The V(acac)3 solubility was further modelled using the Jouyban–Acree-type model [Equation (2.5)], as applied by D’Aprano [63] in a binary solvent system. The model was used in early studies of ET(30) in binary solvent mixtures [70]. 24

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ln 𝑥𝐴 = 𝐵0 + 𝐵1 𝑥 1 1 + 𝐵2 𝑥 1 2 + 𝐵3 𝑥 1 3 + 𝐵4 𝑥 1 4

(2.5)

Where 𝑥 1 is the mole fraction of one of the solvents in the binary mixture, and 𝐵0 –𝐵4 are regression constants. The obtained regression constants from the experimenta l results are presented in Table 2.2. Table 2.2: Regression constants for solubility limit of V(acac)3 in binary solvent mixtures examined in this study, which are used to fit experimental data shown in Figure 2.3. Binary Mixture

B0

B1

B2

B3

B4

AC/PC

-3.2681

-1.1565

3.8401

-7.3900

2.9155

AC/EC

-3.2659

-0.5093

1.4351

-3.9568

0.9851

AC/DMF

-3.2686

1.9339

-3.5282

1.7456

0.0623

AC/DMSO

-3.2765

6.3517

-21.9758

28.7069

-14.1790

AC/1,3DO

-3.2439

-1.9654

10.9980

-12.8671

4.4175

Using the solubility results of the binary mixtures, ternary solvent mixtures were explored to take further advantage preferential solvation to enhance the solubility. The following procedure was used to determine how to combine these binary results to select and test ternary solvent mixtures. The peaks in solubility (Figure 2.3) were observed at the following vol % values: 25/75 AC/1,3DO, 83/17 AC/DMSO, and 68/32 AC/DMF, which suggested the following 3 ternary mixtures: a) AC/1,3DO + AC/DMSO to make an AC/1,3DO/DMSO ternary mixture. b) AC/1,3DO + AC/DMF to make an AC/1,3DO/DMF ternary mixture. c) AC/DMSO + AC/DMF to make an AC/DMSO/DMF ternary mixture. As illustrated in Figure 2.4, 3 possible ways to combine the binary mixtures were considered, which are: 

Ternary ratio I: add up the binary ratios – binary A and binary B.

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

Ternary ratio II: replace the AC portion of the first binary (say A) mixture with the second binary mixture ratio (say B) so that (AC B + nonACB) : nonACA = ACA : nonACA.



Ternary ratio III: replace the AC portion of the second binary mixture (B) with the first binary mixture (A) ratio so that (AC A + nonACA) : nonACB = ACB : nonACB.

nonACB nonACB

nonACA

ACB nonACA

nonACB

ACB

nonACA

ACA

ACA

Ternary ratio I

Ternary ratio II

Ternary ratio III

Figure 2.4: Three methods of combining binary ratios/portions to give ternary ratios. The size of the circles are not representative of the actual rations (i.e. not drawn to scale). Therefore, 3 ternary mixtures combined in 3 ways each gives 9 possible combinatio ns. V(acac)3 solubility in these nine ternary solvent mixtures were measured and are listed in Table 2.3. Combining experimental data obtained from these ternary mixture combinations with their related binary experimental data, complete ternary system plots were simulated using MATLAB and presented in Figure 2.5; the smooth areas in the figure were extrapolated from the experimental data. As shown in Figure 2.5 and Table 2.3, ternary ratio II gave higher solubility than ternary ratio I or III for the A and B combinations because the solubility of V(acac) 3 was higher in 1,3DO than in the other 2 solvents. In method II, the 1,3DO percentage was kept constant for the A and B combinations. In addition, B combinations gave higher

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solubility than the A and C combinations. Comparing with those reported by Herr et al. [37], the systematic approach in this study produced a ternary solvent combination with higher V(acac)3 solubility. Table 2.3: Measured V(acac)3 solubility (𝒙𝑨 ) in ternary solvent systems. Vol % 1,3DO

Vol % DMSO

Vol % AC

Vol % DMF

𝒙 𝑨 × (10 -2 )*

AI

37.50

08.50

54.00

0.00

7.13

AII

75.00

04.25

20.75

0.00

8.64

AIII

62.25

17.00

20.75

0.00

6.99

BI

37.50

0.00

46.50

16.00

7.42

BII

75.00

0.00

17.00

08.00

9.04

BIII

51.00

0.00

17.00

32.00

8.57

CI

0.00

08.50

75.50

16.00