Applied Science, Faculty of
Chemical and Biological Engineering, Department of
DSpace
UBCV
Jameson, Alexander
Doctor of Philosophy - PhD
University of British Columbia
Redox flow batteries are considered to be one of the most realistic candidates for stationary energy storage applications in the range of kW/kWh up to tens of MW/MWh. The recently proposed zinc-iodine flow battery (ZIFB) is among the most promising flow batteries for such applications, having achieved the highest reported energy densities of any aqueous flow batteries to date. However, the energy efficiency of the ZIFB during cycling is still very low due to the thick iodine film that forms during charge and activation overpotentials associated with the iodine electrode reactions. In this thesis, I first investigated the effect of organic cosolvent additives and operating conditions on the ZIFB performance. The acetonitrile (5 vol%) cosolvent showed the most significant improvement to charge-discharge performance, increasing the energy efficiency from 59.4% to 70.0%. Cosolvents that were strongly protophilic preferentially interacted with Zn²⁺ over I₂ leading to the formation of insulative iodine deposits on the positive electrode during charge, but constructing a ZIFB with a posolyte absent of Zn²⁺ with 5 vol% DMF was able to increase the energy efficiency from 59% to 65.4% while maintaining the iodine dissolved in solution. I then studied the kinetics of the iodide redox reactions on various carbonaceous materials (carbon blacks and carbon nanotubes) for the purpose of being incorporated as a catalyst layer (CL) into the graphite felt (GF) positive electrode of the ZIFB. The Ketjenblack EC-300J carbon black performed the best out of the electrode materials studied due to a pore network capable of encapsulating iodine and promoting its adsorption on the pore walls. A ZIFB constructed with a Ketjenblack EC-300J CL on the GF surface adjacent to the battery separator increased the energy efficiency (at 40 mA/cm⁻²) from 63% to 76% and the discharge power density from 39.1 mW/cm⁻² to 44.1 mW/cm⁻² compared to the pristine GF without any CL. Addition of 5 vol% cosolvent further increased the operating efficiency of the ZIFB to 80.3% and demonstrates the synergic effect of the improved electrolyte and electrode compositions.
Development of a PolyiodideAqueous/Non-Aqueous Flow BatterybyAlexander JamesonB. Sc., Queen’s University, 2016MASc, Queen’s University, 2018A THESIS SUBMITTED IN PARTIAL FULFILLMENT OFTHE REQUIREMENTS FOR THE DEGREE OFDOCTOR OF PHILOSOPHYinThe Faculty of Graduate and Postdoctoral Studies(Chemical and Biological Engineering)THE UNIVERSITY OF BRITISH COLUMBIA(Vancouver)May 2024© Alexander Jameson 2024The following individuals certify that they have read, and recommend to the Faculty of Graduate and Post-doctoral Studies for acceptance, the dissertation entitled:Development of a Polyiodide Aqueous/Non-Aqueous Flow BatterySubmitted by Alexander Jameson in partial fulfillment of the requirements for the degree of Doctor of Philosophyin Chemical and Biological Engineering.Examining Committee:Dr. Elod Gyenge, Professor, Chemical & Biological Engineering, UBCSupervisorDr. Pierre Kennepohl, Professor, Chemistry, University of CalgarySupervisory Committee MemberDr. Kevin Smith, Professor, Chemical & Biological Engineering, UBCSupervisory Committee MemberDr. Dan Bizzotto, Professor, Chemistry, UBCUniversity ExaminerDr. David Dreisinger, Professor, Chemical & Biological Engineering, UBCUniversity ExamineriiAbstractRedox flow batteries are considered to be one of the most realistic candidates for stationary energy storageapplications in the range of kW/kWh up to tens of MW/MWh. The recently proposed zinc-iodine flowbattery (ZIFB) is among the most promising flow batteries for such applications, having achieved the highestreported energy densities of any aqueous flow batteries to date. However, the energy efficiency of theZIFB during cycling is still very low due to the thick iodine film that forms during charge and activationoverpotentials associated with the iodine electrode reactions.In this thesis, I first investigated the effect of organic cosolvent additives and operating conditions onthe ZIFB performance. The acetonitrile (5 vol%) cosolvent showed the most significant improvement tocharge-discharge performance, increasing the energy efficiency from 59.4% to 70.0%. Cosolvents that werestrongly protophilic preferentially interacted with Zn2+ over I2 leading to the formation of insulative iodinedeposits on the positive electrode during charge, but constructing a ZIFB with a posolyte absent of Zn2+with 5 vol% DMF was able to increase the energy efficiency from 59% to 65.4% while maintaining theiodine dissolved in solution.I then studied the kinetics of the iodide redox reactions on various carbonaceous materials (carbon blacksand carbon nanotubes) for the purpose of being incorporated as a catalyst layer (CL) into the graphitefelt (GF) positive electrode of the ZIFB. The Ketjenblack EC-300J carbon black performed the best outof the electrode materials studied due to a pore network capable of encapsulating iodine and promotingits adsorption on the pore walls. A ZIFB constructed with a Ketjenblack EC-300J CL on the GF surfaceadjacent to the battery separator increased the energy efficiency (at 40 mA cm−2) from 63% to 76% andthe discharge power density from 39.1 mW cm−2 to 44.1 mW cm−2 compared to the pristine GF withoutany CL. Addition of 5 vol% cosolvent further increased the operating efficiency of the ZIFB to 80.3% anddemonstrates the synergic effect of the improved electrolyte and electrode compositions.iiiLay SummaryFlow batteries have been successfully implemented for several applications involving stationary large scaleenergy storage over the last two decades. The zinc-iodine flow battery in particular has been successfulin achieving some of the highest energy densities of all flow batteries. Like all electrochemical cells, theseflow batteries store and liberate energy through electrochemical reactions involving the electro-active speciesdissolved in the electrolyte at the battery electrode. This study aims to improve the performance of the zinc-iodine flow battery by systematically varying the composition of the battery electrolyte and electrodes andother operating parameters of the system. The addition of low amounts of non-aqueous additives to theprimarily aqueous electrolyte and incorporating low cost carbon catalysts onto the electrode were able toimprove the energy and power capabilities of the zinc-iodine flow battery, further increasing the viability ofthe technology for stationary energy storage applications.ivPrefaceThe work presented in this thesis, including literature review, research questions, experimental design andexecution, data analysis, and dissertation preparation were performed by Alexander Jameson under thesupervision of Dr. Elod Gyenge. All the experimental works were performed at the Chemical and BiologicalEngineering Department at the University of British Columbia.Chapter 1 of this thesis has been published as:• A. Jameson and E. Gyenge. “Halogens as Positive Electrode Active Species for Flow Batteries andRegenerative Fuel Cells”. Electrochemical Energy Reviews. 2020, 3, 431–465. This review paperwas written by me under the supervision and academic guidance of my supervisor Elod Gyenge.Much of the works of Chapter 3 was presented at the following conferences by the author:• A. Jameson and E. Gyenge. “Comparative Electrocatalysis of Bromine and Iodine Redox Species”,PRiME 2020 Meeting, 2020.• Alexander Jameson and Elod Gyenge. “Resolving the Iodine Film Formed on Charge Through Iodine-Ion Interactions”, 241st ECS Meeting, 2022.Part of Chapter 4 is currently in preparation for submission as:• A. Jameson and E. Gyenge. “Selecting a Hybrid Aqueous/Non-Aqueous Electrolyte for Iodine RedoxFlow Batteries” (in preparation). I conducted all design, testing and analysis of the results and wrotethe manuscript under the guidance of E. Gyenge.Part of Chapter 5 is currently in preparation for submission as:• A. Jameson and E. Gyenge. “High Activity Carbonaceous Electrodes for Zinc-Iodine Flow Batteries”vPreface(in preparation). I conducted all design, testing and analysis of the results and wrote the manuscriptunder the guidance of E. Gyenge. Scanning Electron Microscopy images were taken at the BioimagingFacility (BIF), UBC.viTable of ContentsAbstract . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . iiiLay Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . ivPreface . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . vTable of Contents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . viiList of Tables . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xiList of Figures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xiiList of Symbols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xvList of Abbreviations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .xviiiAcknowledgements . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . xxi1 Chapter 1: Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11.1 The Need for Medium/Large-Scale Stationary Energy Storage . . . . . . . . . . . . . . . . 11.2 Redox Flow Batteries for Energy Storage . . . . . . . . . . . . . . . . . . . . . . . . . . . 4viiTable of Contents1.2.1 Electrochemical Cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 41.2.2 Flow Battery Design . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61.2.3 Figures of Merit for Flow Battery Technology . . . . . . . . . . . . . . . . . . . . 71.2.4 Minimizing Electrode Polarization in Flow Batteries . . . . . . . . . . . . . . . . . 121.3 Chemical and Electrochemical Properties of Iodine . . . . . . . . . . . . . . . . . . . . . . 151.3.1 Iodine/Triiodide Equilibria . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 151.3.2 Iodine Electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 171.4 Halogen/Halide Based Flow Batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191.4.1 The Zinc Anode . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191.4.2 The Zinc-Bromine Flow Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . 211.4.3 The Zinc Iodine Flow Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 231.5 Research Approach . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 252 Chapter 2: Experimental Methodology . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 272.1 Electrochemical Half-Cell Design and Procedures . . . . . . . . . . . . . . . . . . . . . . 272.2 Flow Battery Apparatus . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 292.3 Spectroscopy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 312.4 Catalyst Preparation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 322.5 Determination of the Electrochemically Active Surface Area . . . . . . . . . . . . . . . . . 332.6 Calculation of Kinetic Parameters . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 342.7 Materials and Reagents . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 36viiiTable of Contents3 Chapter 3:Interplay of I2-Ion, I2-Solvent, and Temperature Effects on I2/I−3 /I− Redox Couples 373.1 The Iodine Electrode . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 373.2 Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Film . . . . . . . . . . . . . . . 413.3 Effects of Temperature on the Zinc Iodide Flow Battery . . . . . . . . . . . . . . . . . . . 463.4 Iodine-Ion and Iodine-Solvent Interactions in the Electrolyte . . . . . . . . . . . . . . . . . 493.5 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 534 Chapter 4: Organic Cosolvents for the Zinc Iodine Flow Battery . . . . . . . . . . . . . . . . 554.1 Iodine Electrochemistry in Non-Aqueous Solvents . . . . . . . . . . . . . . . . . . . . . . 554.1.1 Solvent and Mixed Solvent Properties . . . . . . . . . . . . . . . . . . . . . . . . . 574.1.2 Interactions Influencing Iodine Species Solubility . . . . . . . . . . . . . . . . . . 594.2 Iodine Solubility & Spectroscopy in Non-Aqueous Cosolvents . . . . . . . . . . . . . . . . 634.3 Non-Aqueous Iodine Electrochemistry . . . . . . . . . . . . . . . . . . . . . . . . . . . . 674.4 Hybrid Aqueous-Nonaqueous ZIFB . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 724.5 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 745 Chapter 5: High Surface Area & Activity Carbon Materials for the I2/I− Redox Couple in theZIFB . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 765.1 Electrode Materials for the Iodine Redox Reactions . . . . . . . . . . . . . . . . . . . . . . 765.2 Iodine Kinetics on Carbonaceous Materials . . . . . . . . . . . . . . . . . . . . . . . . . . 795.3 Comparing the Kinetics of the Bromine and Iodine Electrodes . . . . . . . . . . . . . . . . 845.4 Electrode Impact on Flow Battery Performance . . . . . . . . . . . . . . . . . . . . . . . . 87ixTable of Contents5.5 Summary and Conclusions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 926 Chapter 6: Conclusions and Recommendations . . . . . . . . . . . . . . . . . . . . . . . . . 946.1 Summary of Key Findings . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 946.2 Future Recommendations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 95Bibliography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 97xList of Tables1.1 Advantages and disadvantages of common energy storage sources and technologies. . . . . . 41.2 Advantages and disadvantages associated with common positive half-cell redox species inflow batteries. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 81.3 Review of the most significant recent advances in the development of the ZIFB. . . . . . . . 242.1 Overview of carbon catalysts used in this study. . . . . . . . . . . . . . . . . . . . . . . . . 332.2 Values of ECSA and RF determined by calculating CDL and Cs. . . . . . . . . . . . . . . . . 342.3 List of reagants and solvents used in this study. . . . . . . . . . . . . . . . . . . . . . . . . 364.1 Properties of solvents investigated throughout this study . . . . . . . . . . . . . . . . . . . . 585.1 Kinetics parameters of I2/I− on various electrodes . . . . . . . . . . . . . . . . . . . . . . . 835.2 Kinetics parameters of Br2/Br− on various electrodes . . . . . . . . . . . . . . . . . . . . . 865.3 ZIFB performance as a function of positive electrode, with and without 5 vol% acetonitrilecosolvent. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 92xiList of Figures1.1 Composition of the gross global electricity generation by source based on percent share. . . 21.2 Comparison of power and energy capabilities for different energy storage systems. . . . . . 31.3 Schematic of a redox flow battery. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61.4 Schematic of the iodine film. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 191.5 Research approach for the dissertation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 262.1 Schematic and dimensions of the RRDE used in this study. . . . . . . . . . . . . . . . . . . 282.2 Schematic of the flow battery apparatus used for ZIFB tests . . . . . . . . . . . . . . . . . . 292.3 Comparison of the charge-discharge behaviour of a ZIFB with a microporous separator,ion-exchange membrane, and no separator/membrane. . . . . . . . . . . . . . . . . . . . . 302.4 UV-Vis spectra of 0.1 M I2 in water, with and without I−. . . . . . . . . . . . . . . . . . . . 322.5 Cyclic voltammograms of the capacitative current on various different electrode materialsfor calculation of the electrochemically active surface area. . . . . . . . . . . . . . . . . . . 353.1 Schematic of iodine film. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 383.2 Cyclic voltammograms showing the effect of scan rate on the electrochemical iodine reactions. 41xiiList of Figures3.3 Rotating ring disk experiments investigating the formation of the iodine film during iodideoxidation. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 433.4 Linear sweep voltammetry experiments on an RDE for iodide oxidation and iodine/triiodidereduction, with increasing rotation rate. . . . . . . . . . . . . . . . . . . . . . . . . . . . . 443.5 Chronoamperometry experiments with increasing I2 content, at 1 V (iodide oxidation) and-0.5 V (iodine/triiodide reduction). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 453.6 Cyclic voltammograms showing the effect of temperature on the iodine electrochemicalreactions. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 463.7 RDE showing the effect of temperature on the iodine electrochemical reactions. . . . . . . . 473.8 Chronoamperometry at 1 V and chronopotentiometry at ± 50 mA cm−2 showing the effectof temperature on the iodine electrochemical reactions. . . . . . . . . . . . . . . . . . . . . 483.9 Charge-discharge curves for a ZIFB as a function of temperature. . . . . . . . . . . . . . . . 493.10 Cyclic voltammograms comparing I−3 , I2Br−, and I2Cl− electrochemical reactions at 50 mVs−1, 10 mV s−1, and 5 mV s−1. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 503.11 CVs with varying Br− Cl− NO−3 ClO−4 SO2−4 . . . . . . . . . . . . . . . . . . . . . . . . . 513.12 Cyclic voltammograms showing electrochemical behaviour of the iodide electrolyte in thepresence of acetonitrile. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 533.13 Zinc-iodide flow battery charge-discharge tests with increasing vol% of acetonitrile. . . . . . 544.1 Schematic of hybrid aqueous/non-aqueous zinc iodine flow battery. . . . . . . . . . . . . . . 574.2 EC-model graphical representation of the strength of interaction between Lewis bases andiodine. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 624.3 UV-Vis analysis for I2 solubility in dimethylformamide, N-methyl-2-pyrrolidone, dimethylsulfoxide, dimethoxyethane, pyridine, and diethylamine. . . . . . . . . . . . . . . . . . . . 63xiiiList of Figures4.4 UV-Vis analysis for I2 solubility in acetonitrile and propylene carbonate. . . . . . . . . . . . 644.5 UV-Vis analysis for I2 solubility in various organic cosolvents with addition of KI or ZnI2. . 664.6 Photos showing decrease in I2 solubility in the presence of Zn2+. . . . . . . . . . . . . . . . 674.7 Cyclic voltammograms comparing the iodide electrochemical behaviour with addition ofvarious cosolvents along with KI or ZnI2. . . . . . . . . . . . . . . . . . . . . . . . . . . . 684.8 Chronoamperometry analysis at 1 V investigating cosolvent effect on charge reaction in KIand ZnI2 electrolytes. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 694.9 RRDE for KI and ZnI2 electrolytes containing different cosolvents. . . . . . . . . . . . . . . 704.10 Nyquist plots for KI and ZnI2 electrolytes containg different cosolvents. . . . . . . . . . . . 714.11 CVs and RDE analysis of zinc electrochemistry in the presence of various organic cosolvents. 724.12 ZIFB performance with various different organic solvents. . . . . . . . . . . . . . . . . . . 724.13 ZIFB performance with DMF with and without Zn2+ in the iodide half-cell. . . . . . . . . . 734.14 ZIFB performance with DMF over 15 cycles. . . . . . . . . . . . . . . . . . . . . . . . . . 745.1 Cyclic voltammograms comparing the iodide redox reactions on various carbonaceous cat-alysts. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 805.2 Cyclic voltammograms demonstrating the durability of various carbonaceous catalysts. . . . 815.3 Linear sweep voltammograms at 1000 rpm comparing the cathodic and anodic polarizationbehaviour on glassy carbon and Ketjen-300. . . . . . . . . . . . . . . . . . . . . . . . . . . 825.4 A comparison of the low-field polarization data on the RDE for GCE and other carboncatalyst materials. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 835.5 Cyclic voltammograms of the Br2/Br− redox couple on various electrode materials. . . . . . 84xivList of Figures5.6 Linear sweep voltammograms at 1000 rpm comparing the cathodic and anodic polarizationbehaviour of the bromine/tribromide/bromide redox reactions on glassy carbon (GCE) andKetjen-300. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 855.7 Comparison of the Br2/Br− and I2/I− redox reactions on various electrodes. . . . . . . . . . 875.8 SEM images of the carbon catalyst treated graphite felts. . . . . . . . . . . . . . . . . . . . 885.9 Charge–discharge curves comparing zinc-iodine flow batteries with different positive elec-trodes. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 895.10 Nyquist plots of the ZIFB with and without Ketjen-300 on the graphite felt electrode. . . . . 905.11 Cycling performance for the ZIFB with Ketjen-300 sprayed on the surface of the iodine-sideelectrode. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 905.12 Charge–discharge curves comparing zinc-iodine flow batteries with electrode and electrolytemodifications. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91xvList of Symbolsα Transfer coefficientδ Thickness of the Nernst diffusion layer (µm)∆ΩΦ j Ohmic potential losses in the cell (V)∆Ep Peak potential separation (V)∆G0f Standard Gibbs free energy of formation (J mol−1)∆Gr Non-standard Gibbs free energy (J mol−1)∆G0r Standard Gibbs free energy (J mol−1)∆Sr Entropy change of reaction (J mol−1 K−1)η Applied overpotential (V)ηd Concentration overpotential (V)ηs Activation (surface) overpotential (V)µ Dipole moment (D)ν Scan rate (V s−1)νk Kinematic viscosity (m2 s−1)ω Rotation rate (rad s−1 or rpm)ρ Resistivity (Ω m)ρe Electrical resistivity (Ω m)xviList of Symbolsρi Ionic resistivity (Ω m)εc Round-trip coulombic efficiencyεr Relative permittivity of the solventεt Thermal efficiencyεv Voltage efficiencyA Cross-sectional area (m2)a j Activity of species jC Concentration (mol L−1)Cb Concentration in the bulk electrolyte (mol L−1)Cs Concentration at the electrode surface (mol L−1)Cs Specific capacitance (F cm−2)CDL Double layer capacitance (F or C V−1)D j Diffusivity of species j (m2 s−1)E Non-standard cell potential under an applied load (V)Ee Equilibrium (non-standard) cell potential (V)E0e Equilibrium (standard) cell potential (V)EA Electron affinity (kJ mol−1)ED Energy density (J m−2)F Faraday’s constant (96485 C mol−1)H Hardness of a chemical speciesi Current density (mA cm−2)xviiList of SymbolsI Current (A)i0 Exchange current density (mA cm−2)il Mass transfer limiting current density (mA cm−2)icap Capacitance charging current (mA)iD Disk current density of RRDE (mA cm−2)ipa Peak anodic current density of cyclic voltammogram (mA cm−2)ipc Peak cathodic current density of cyclic voltammogram (mA cm−2)iR Ring current density of RRDE (mA cm−2)iss,CV Cyclic voltammogram steady state current density (mA cm−2)IE Ionization energy (kJ mol−1)Km Mass transfer coefficient (m s−1)l Length travelled by current (m)n Electrons transferred in balanced cell reactionPD Power density (W m−2)Qs Specific discharge capacity (Ah g−1)Qv Volumetric discharge capacity (Ah L−1)r Ionic radius (nm)R Resistance (Ω)s j Stoichiometric coefficientt Time (s)Wout Useful work produced by the electrochemical cell (J)z Charge of ionic speciesxviiiList of AbbreviationsACN AcetonitrileCA ChronoamperometryCAE Compressed Air Energy StorageCL Catalyst LayerCP ChronopotentiometryCVs Cyclic VoltammogramsDME MonoglymeDMF DimethylformamideDMSO Dimethyl sulfoxideECSA Electrochemically Active Surface AreaEN EthylenediamineESS Energy Storage SystemFW FlywheelsHFB Hybrid Flow BatteryHOMO Highest Occupied Molecular OrbitalIPA 2-PropanolLIB Li-ion BatteryxixList of AbbreviationsLSVs Linear Sweep VoltammogramsNMP N-Methyl-2-pyrrolidinonePC Propylene carbonatePHS Pumped-Hydro StoragePYR PyridineRDE Rotating Disk ElectrodeRFB Redox Flow BatteryRRDE Rotating Ring Disk ElectrodeRTILs Room Temperature Ionic LiquidSCs SupercapacitorsSMES Superconducting Magnetic Energy StorageSoC State of ChargeZBFB Zinc Bromine Flow BatteryZIFB Zinc Iodine Flow BatteryxxAcknowledgementsForemost I would like to express my sincere gratitude to my supervisor Professor Elod L. Gyenge for hissupport and encouragement throughout my studies at UBC. When initially applying for my PhD I ultimatelywanted to engage in a project that would allow me to apply my chemistry and chemical engineering back-grounds and I am grateful to have found such. In order to become an accredited engineer, I also had totake sixteen additional courses on top of my PhD workload and if it was not for the help and support fromthe CHBE community I would not have been able to make it through. I would like to thank my researchstudents throughout the years, including Michael Kuenzler and Elena Stark from the Technical Universityof Munich (Germany) and Alex Whitbeck (MASc student in Gyenge group). I would also like to thank therest of my labmates, including Ray, Yu Pei, Philip Wijaya, Shahid and Wendy Wu, as well as Dr PooyaHosseini-Benhangi (Agora Energy Technologies Ltd.).I thank my family for their love and support. I dedicate this thesis to my grandparents May and DavidJameson.xxiChapter 1Introduction1.1 The Need for Medium/Large-Scale Stationary Energy StorageThe depletion of fossil fuel stores along with the raising global population has caused energy supplyand demand to be forefront issues in engineering, the economy, and politics. This energy supply must alwaysbe accompanied by energy storage, whether it is an intrinsic property of the system itself or integrated withan auxiliary system capable of doing so. Energy storage systems (ESS) are capable of storing surplus energygenerated during off-peak hours that can then be discharged at a later time which is necessary for meetingfluctuations in energy demand throughout the day. Developing systems capable of efficiently storing anddelivering large quantities of energy (kW/kWh to MW/MWh) can provide advantages to new and existingelectrical power transmission systems and in other applications such as emergency backup systems. Oneof the seventeen Sustainability Developmental Goals established by the United Nations is to ensure accessto affordable, reliable, sustainable, and modern energy for all [1]. This includes increasing the proportionof the world population with access to energy and substantially increasing the share of renewable energy inthe global energy mix. Many sources of renewable energy (i.e. solar, wind) are by nature intermittent andunreliable as they can only be harvested during times where the source is presently available. Coupling re-newables to some form of ESS allows excess energy generated during periods of low customer demand to bestored over long periods and later liberated when needed, significantly improving the viability of widespreadimplementation of renewables. In 2022, 60% of the world’s electricity (16000 TWh) was obtained from coaland gas fossil fuels compared to approximately 11% (3000 TWh) from solar and wind (Figure 1.1) [2], butintegration of renewables has been on a steady increase over the past decade due in large to improvementsin the accompanying energy storage systems [3, 4].11.1. The Need for Medium/Large-Scale Stationary Energy Storage0102030402000 2005 2010 2015 2020Electricity Share by Source / %           Figure 1.1: Composition of the gross global electricity generation by source based on percent share. Datawas retrieved from Ember’s Yearly Electricity Data & Energy Institute Statistical Review of World Energy[2].Energy storage systems can be classified based on the physical form of energy stored (electrical, chem-ical, mechanical, and thermal) and a comparison of typical power and energy ratings for some of thesetechnologies is given in Figure 1.2 and Table 1.1. Electrical energy storage includes supercapacitors (SCs)and superconducting magnetic energy storage (SMES) which typically find use in the electricity sector forapplications that require large amounts of power in a short period of time, but suffer from very low energydensities (see bottom left quadrant of Figure 1.2) [5]. Mechanical energy storage systems such as pumped-hydro storage (PHS), compressed air energy storage (CAE), and flywheels (FW) have been commonplacein grid-scale energy management. Canada in particular relies heavily on PHS, accounting for over 60%(10200 kWh) per capita of electricity generated in 2022 compared to 6% (750 kWh) in the United States.Mechanical energy storage systems generally have large storage capacities and long cycle lives, but typicallyinvolve very high capital costs and suffer from being geographically dependent [6] (e.g. PHS requires twowater reservoirs separated by a large elevation).Chemical energy storage relies on the internal energy stored in the bonds of molecules which can be lib-erated (and/or stored) through chemical reactions, including the combustion of fossil fuels or carbon-neutral21.1. The Need for Medium/Large-Scale Stationary Energy StorageFigure 1.2: Comparison of power and energy capabilities for different energy storage systems (Image re-trieved from [7]).alternatives such as bio-ethanol. One category of chemical energy storage is electrochemical energy stor-age, where energy is stored and liberated via redox (reduction-oxidation) reactions within an electrochemicalcell. Electrochemical energy storage is advantageous for stationary large-scale energy applications as thistechnology has no geographical requirements, much lower capital costs compared to PHS and CAE, and theelectrochemical cells are stackable/modular in design (cells can be combined in series or parallel to meetvarious different requirements) [8]. Li-ion batteries (LIB) have found widespread success in both portableconsumer electronics and electrical vehicles because of the high energy densities (≈ 100-265 Wh/kg or250-690 Wh/L [9]) that they have been able to achieve. Redox flow batteries (RFBs) are another subclass ofelectrochemical energy storage and offer an alternative to the LIB (and other ESS) for reliably storing andsupplying energy with less environmental impacts over the course of its lifetime [10]. The energy densityof aqueous RFBs is lower than the LIB but for stationary energy applications this is largely eclipsed bythe availability/costs of lithium and safety concerns associated with the flammable LIB electrolyte that canresult in thermal runaway [11]. The commercialization and installation of redox flow batteries for variousapplications around the globe over the past two decades has demonstrated their ability to provide reliable31.2. Redox Flow Batteries for Energy Storageenergy storage supply over tens of thousands of charge-discharge cycles [12], and the general constructionand operation of RFBs is described in the proceeding Section 1.2.Table 1.1: Advantages and disadvantages of common energy storage sources and technologies.Source/Technology Advantages DisadvantagesFossil fuels Fully-developed, established infrastructure Limited resources and environmental risksBiofuels Uses carbon neutral biomass Limited sustainability, large land useSupercapacitors High power densities Very low energy densityFast response times High costsExcellent cyclability Prone to self-dischargeFlywheels High power density Energy densities unsuitable for long storageHigh cycle and calendar life Complications in construction/scaling from fa-tigue placed on rotating componentsPumped-Hydro Storage Low operating and maintenance costs High capital costs with long construction timesVery high storage capacity Much lower energy density than FW and CAETheoretically unlimited cycle lifetimes Geographically dependent, disrupts ecosystemsLi-ion Battery Very high energy and power densities Poor safety (e.g. thermal runaway)High round trip efficiency Limited lithium reserves and need for lithiumdisposal proceduresFlow Batteries Decoupled flexible design of energy & power Lower energy density than LIBSafer than LIB Pump power and maintenance costsCan be charged and discharged completelywithout damaging electrodes1.2 Redox Flow Batteries for Energy Storage1.2.1 Electrochemical CellsElectrochemical cells are devices capable of storing and liberating energy via electrochemical redoxreactions that convert electrical energy to chemical energy (energy storage) or chemical energy to electricalenergy (energy liberation). A battery (static or flow) may be composed of one or more individual cellsconnected in series or parallel, and the major components of an individual electrochemical cell includes:1. The negative electrode, which functions as the anode during discharge and as the cathode duringcharge2. The positive electrode, which functions as the cathode during discharge and as the anode during charge41.2. Redox Flow Batteries for Energy Storage3. The electrolyte, which is an ionically conductive phase that contains the electrochemical species in-volved in the redox reactions of the cell4. A separator, which physically separates the two half-cell compartments but allows them to remain inionic contact via diffusion of the charged species through the mediumOn battery discharge a spontaneous oxidation reaction at the anode liberates electrons which then travelthrough an external circuit connecting the electrodes towards the cathode where they are consumed in anelectrochemical reduction reaction. During charge, the reverse (non-spontaneous) electrochemical reactionsare driven by the application of an external voltage, and the negative electrode species is reduced while thepositive electrode species is oxidized. The electrodes serve the dual role of facilitating the redox reactionsof the cell as well as conducting electrons to/from the external circuit, resulting in the flow of current [9].For a majority of traditional static batteries, the electrodes directly participate in the cell reactions (e.g. thePb electrode of lead-acid batteries being converted to PbSO4 on discharge, and back to Pb on charge). Intraditional RFBs and fuel cells the electrodes are not consumed or produced during the electrochemical re-actions, but rather act as heterogeneous catalysts that facilitate the redox reactions in the electrolyte withouthaving their mass or chemical structure significantly changed. This feature of RFBs allows them to theoret-ically achieve 100% depth of discharge without altering or deteriorating the electrode over repeated cycling.Furthermore, this decouples the power and energy ratings of the flow battery providing a degree of systemdesign flexibility and scalability that static batteries like the LIB are not capable of.The electrolyte contains both the electroactive species involved in the electrode reactions along with asolvent to dissolve and disperse the redox species. Ideally, the electrolyte must have a high ionic conductivityin order to provide a pathway for the reacting species while blocking electronic leakage, and should be stablein the operating potential window of the battery. The storage capacity of a RFB increases with the volumeof the electrolyte tanks, and because of this the electrolyte tanks of commercial flow batteries can occupyover half the total volume of the entire RFB module [12]. The electrolyte of an RFB should ideally thereforebe non-toxic and inexpensive in order to better compete with other ESS technology. Battery separators arenecessary to prevent direct contact of the two electrodes and mixing of the half-cell electrolytes while stillpermitting certain ions and/or chemical species to permeate through and maintain ionic contact betweenthe electrode reactions. Separators are able to achieve selective permeability via the size of their pores,while ion-selective membranes are able to further differentiate between chemical species by their charge.51.2. Redox Flow Batteries for Energy StorageSeparator and membrane design is crucial to battery performance, where the porosity, pore size, thickness,and species permeability have to be properly selected. The separator/membrane is generally one of themostly costly components of the battery, totaling 20-30% of the total cost when expensive ion-selectivemembranes such as Nafion® are employed [12].1.2.2 Flow Battery DesignFigure 1.3: Schematic of a redox flow battery.Redox flow batteries (RFBs) are similar in principle to traditional static batteries, except the electrolyteis stored in tanks external to the battery cell stack that is delivered to and removed from the electrodes bypumps (Figure 1.3). In a RFB, the redox species involved in the electrochemical reactions remain dissolvedin the electrolyte phase during both charge and discharge and the electrodes serve as heterogeneous catalystsfor the redox reactions. Since the electrodes are not being consumed during the electrochemical reactions,RFB are able to be charged to a state of charge (SoC) of 100% and discharged to 0% repeatedly withoutloss of performance or capacity, and some flow batteries are advertised as having lifetimes of 10-20 years[13, 14]. For static batteries the repeated consumption and generation of the electrode accelerates its degra-dation which diminishes the performance and cyclability of the battery over repeated prolonged operation,61.2. Redox Flow Batteries for Energy Storagehighlighting how RFBs are ideally suited for stationary applications that require thousands of deep dischargecycles [15].A distinguishing feature of flow batteries is their power and energy capacities are uncoupled and ableto be designed separately [16]. Since the redox species remain in the electrolyte phase, the energy density(i.e. capacity to do work) of a RFB is primarily determined by the volume of electrolyte and concentrationof the relevant redox species. On the other hand, the power density (i.e. rate at which work can be done)depends heavily on the rate of the redox reactions at the electrode and the interfacial surface area betweenthe electrolyte and electrode where these redox reactions take place. Hybrid flow batteries (HFBs) containat least one electrode that is not completely uncoupled from the redox reactions, such as in the zinc-bromineand zinc-iodine flow batteries where the Br2/Br− and I2/I− charge/discharge species remain in the electrolytephase throughout operation but solid zinc is deposited onto the negative electrode during charge and thissolid zinc is electrochemically dissolved back into the electrolyte phase during discharge. For these HFBsthe power and energy density may not be completely uncoupled from each other depending on the battery,but it has been shown that the Br2/Br− and I2/I− species do limit the overall capacity of these zinc-halogenflow batteries and thus their power and energy capacities are still uncoupled [17].Based on this, the ideal electroactive species (redox pair) for a redox flow battery would be highlysoluble in the electrolyte while possessing facile and reversible (i.e. both the oxidation and reduction)kinetics on the electrode. The electroactive species and solvent should be inexpensive to support large-scaleapplications, and the kinetics of the redox reactions should ideally be sufficiently high on low cost electrodematerials to avoid expensive precious metal catalysts. The most common battery solvent is water which isthermodynamically susceptible to degradation by electrolysis at potentials above 1.23 V (vs. SHE), and sothe relevant redox reactions must also occur at potentials within the stability window of water. The Br2/Br−and I2/I− redox couples fit the above criteria well (Table 1.2), making them promising candidates for thepositive half-cell of flow batteries.1.2.3 Figures of Merit for Flow Battery TechnologyAn electrochemical cell is a device capable of converting the potential energy stored in the bonds of anelectroactive species into useful electrical energy via electrochemical reactions at the positive and negative71.2. Redox Flow Batteries for Energy StorageTable 1.2: Advantages and disadvantages associated with common positive half-cell redox species in flowbatteries.ElectroactiveCoupleE0 vs. SHEat 25 ◦CPros ConsI2/I−0.54 High volumetric energy densities Iodine cost is increasingHigh solubility due to formation of I−3 I−3 formation limited by iodine filmBr2/Br−1.09 Inexpensive and abundant Extremely toxicHigh solubility due to formation ofBr−3Electrode and membrane stability incorrosive electrolyteV5+/V4+0.99 Minimal crossover issues when cou-pled to V3+/V2+Low solubility/energy densities com-pared iodine and bromineCyclability (VRB has achieved 20+years of continual operation)Concentrated acid electrolytes degradecommon electrode and separator mate-rialsCe4+/Ce3+1.61 High standard reduction potential Slow electrode kineticsRelatively high solubility in MSA High costs of precious metal catalystsFe3+/Fe2+0.77 Highly reversible Very sensitive to pHInexpensive Precipitation issueselectrodes. The theoretical maximum electrical energy that can be delivered by the cell reaction is equal tothe Gibbs free energy (∆Gr) of the reaction, which can be expressed in terms of the Gibbs free energy atstandard state ∆G0r and the activities a j of all species involved in the cell reaction:∆Gr = ∆G0r +RTlnK∗r (1.1)K∗r =∏jas jj (1.2)where s j is the stoichiometric coefficient in front of species j in the balanced cell reaction. The value of∆Gr (or ∆G0r if all reactants and products are in their standard states) represents the maximum amount ofwork that can be performed during a chemical process. For a galvanic cell (i.e. a battery discharging) thisvalue represents the maximum amount of energy liberated from the cell reactions that can be converted intoelectrical energy, and for an electrolytic cell (i.e. a battery being recharged) represents the minimum amountof energy required to drive the cell reaction. For an electrochemical cell at equilibrium (i.e. when no netcurrent flows through the cell) the corresponding cell potential can be calculated from the change in Gibbs81.2. Redox Flow Batteries for Energy Storageenergy at standard (E0e ) and non-standard (Ee) conditions by Equation 1.3 and Equation 1.4, respectively:∆G0r =−nFE0e (1.3)∆Gr =−nFEe (1.4)where n is the number of electrons transferred per mole of reacting redox species and F is Faraday’s constant(96485 C/mol) which represents the amount of charge carried by one mole of electrons. The dependence of∆Gr on the activities a j of the species involved in the cell reaction is provided by the Nernst equation:Ee = E0− RTnF ln∏jas jj (1.5)The thermal efficiency εt of an isothermal electrochemical cell operating reversibly is defined as the amountof useful work Wout produced (represented by the Gibbs free energy of reaction ∆Gr) relative to enthalpychange of the reaction (∆Hr):εt =Wout−∆Hr =−∆Hr−T∆Sr−∆Hr =∆Gr∆Hr(1.6)where ∆Sr is the change in entropy of the reaction. This “first-law efficiency” is used to describe mostelectrochemical cells, but for exothermic cell reactions (negative ∆Hr) with positive values of ∆Sr the tab-ulated efficiency from Equation 1.6 will be over 100% and alternative “second-law efficiency” calculationsutilizing exergy are used in place [18, 19].The amount of charge (Q, units C) liberated from a battery is the total amount of electricity that canbe generated from the electrochemical reactions in a period of time, typically expressed for electrochemicalcells in units of ampere hours (Ah). The capacity of an electrochemical cell is determined by the type andamount of active materials in the cell, and the theoretical discharge capacity available can be normalizedfor electrode mass (Qs, common for static batteries) or the electrolyte volume (Qv, common for RFBs). Forflow batteries, the volumetric discharge capacity Qv is dictated by the concentration C of the electroactivespecies in the electrolyte and number of electrons transferred per mole of oxidized/reduced species n:MaximumVolumetric Discharge Capacity (AhL−1) =nFC3600(1.7)91.2. Redox Flow Batteries for Energy StorageThe round-trip coulombic/faradaic efficiency εc of an electrochemical cell is calculated byεc =QdischargeQcharge(1.8)and represents the efficiency of the charge transferred during a charge-discharge cycle. Electrochemicalreactions other than the desired cell reactions cause εc to decrease, such as when a redox species from onehalf-cell crosses through the separator and directly reacts with the opposite electrode (i.e. a short circuit).The volumetric energy density of a flow battery (in Wh L−1) can be calculated by combining Equation 1.7with the cell potential under operation (E):Volumetric Energy Density (WhL−1) =nFCE3600(1.9)The value of E when current passes through a cell is always different than the value of Ee calculated fromEquation 1.5. The cell potential calculated from the Nernst equation is based purely on a thermodynamicanalysis, and therefore this value can only be approximately achieved when the battery is at open-circuit (noapplied load, I = 0). The round-trip voltage efficiency εv is calculated asεV =EdischargeEcharge(1.10)where |E|> |Ee| during battery charge and |E|< |Ee| during discharge. The value of εv is always less thanunity due to non-equilibrium effects that are inevitable but can be minimized through careful design. Thesenon-equilibrium effects can be represented by a voltage balance on the electrochemical cell:E = Ee−∑jηa, j+∑jηc, j−∑j∆ΩΦ j (1.11)∑jηa, j = ηa,s+ηa,d (> 0) (1.12)∑jηc, j = ηc,s+ηc,d (< 0) (1.13)where Ee is the open circuit/equilibrium voltage of the cell, ∆ΩΦ j represents the ohmic potential losses inthe cell, ηs is the surface activation overpotential at the electrodes and ηd is the diffusion overpotential,with the a and c subscripts denoting the anode and cathode, respectively. That is, the actual operating101.2. Redox Flow Batteries for Energy Storagevoltage delivered by an electrochemical cell is reduced by losses associated with overcoming the activationenergy barrier to reaction (ηs), losses from sluggish mass transfer to/from the electrode surface (ηd), andlosses dictated by Ohm’s law that arise from the non-zero resistance (Ri) of the cell. The overall operatingefficiency of an electrochemical cell is calculated as the product of the thermodynamic (Equation 1.6),coulombic (Equation 1.8), and voltage (Equation 1.10) efficiencies:ε = εt × εc× ε f = ∆Gr∆Hr ×QdischargeQcharge× EdischargeEcharge(1.14)The value of εt depends on the temperature of the system, but the values of εc and εv can be improvedthrough appropriate design of the electrochemical cell.The power density (PD, W m−2) of a battery is determined by the rate at which the electrochemicalreactions at the electrode take place. The power density is calculated as the product of the current density (i,A cm−2) and the average potential (E, in V) that the cell is operating at:PD= i×E (1.15)The energy density ED of a battery is based on the identity and amount of electroactive species in the cell,calculated as the product of the power density and length of time (t) over which the energy is delivered:ED= i×E× t = PD× t (1.16)The power and energy efficiency of an electrochemical cell over a charge-discharge cycle can be calculatedby Equation 1.17 and Equation 1.18, respectively:εP =PDdischargePDcharge=(iE)discharge(iE)charge(1.17)εE =EDdischargeEDcharge=(iEt)discharge(iEt)charge(1.18)From Equations 1.15 and 1.16, a high power and energy density means maximizing both the operatingcurrent and potential of the electrochemical cell, but the three loss terms in Equation 1.11 increase with thecurrent density. Understanding and minimizing the overpotentials and internal resistance are vital in the111.2. Redox Flow Batteries for Energy Storagedesign of electrochemical cells, and are discussed in detail in the following Subsection 1.2.4.1.2.4 Minimizing Electrode Polarization in Flow BatteriesImproving the performance of any electrochemical cell relies on minimizing the factors that reduce thepower, energy, and overall operating efficiencies of the cell. However, the current and voltage of the batteryare coupled; they cannot be simultaneously controlled. When the rate of the redox reaction is slow at theelectrode in comparison to the mass transfer of reactants/products, the concentrations in the bulk electrolyteand electrode surface do not differ appreciably from each other and the relationship between the currentdensity i and electrode potential E for a half-cell reaction is given by [20]:i= i0[eαa fη − e−αc fη ]. (1.19)whereη = E−Ee (1.20)f =FRT= 38.9 at 298 K (1.21)The overpotential (η) is the deviation of the actual potential at the electrode surface E from the equilibriumvalue Ee (i.e. when i = 0), and is defined as positive for anodic potentials (E > Ee) and negative for cathodicpotentials (E < Ee). The first term in Equation 1.19 is the contribution from the anodic current while thesecond term is the contribution from the cathodic current to the total net current at any applied overpotential.At equilibrium, E = Ee and η = 0, which from Equation 1.19 causes the net current density i to be zeroas well. The i0 term is the exchange current density and is the absolute magnitude of either ic or ia atequilibrium where the rate of the cathodic and anodic reactions are equal to each other but opposite insign such that the net current density i is zero. The exchange current density is an extremely importantparameter in electrochemical kinetics, representing the rate at which an electrochemical reaction will reachequilibrium on a given electrode material. Another important kinetic parameter is the transfer coefficientTransfer coefficient (α) which is a measure of the symmetry of the energy barriers of reaction, with 0 < α <1 where α = 12 indicating symmetrical intersection of the free energy curves for the oxidation and reductionreactions of a redox pair.121.2. Redox Flow Batteries for Energy StorageIn Equation 1.19 mass transfer effects are negligible compared to the slow electrode reactions, andso the overpotential at a given current density i serves solely to provide the activation energy required todrive the electrode reaction and η = ηs, where ηs is the surface activation overpotential that arises fromovercoming the activation energy barrier to the electrode reaction (introduced previously in Equation 1.11).Rearranging Equation 1.19 for ηs, the surface overpotential associated with overcoming the activation barrierto reaction for the anodic and cathodic electrode reactions is respectively:ηs = balnii0(1.22)ηs =−bcln ii0 (1.23)where b is the Tafel slope (for either the anodic or cathodic branch)b=RTαF. (1.24)Equation 1.19 demonstrates that the lower the value of i0, the larger the (activation) overpotential requiredto produce a desired current density, while an electrochemical system with high i0 would require only smalloverpotentials to achieve the same current density (i.e. rate of reaction). To minimize surface overpotentialsthe electroactive species and electrode materials should be selected such that the desired redox reactions(both charge and discharge reactions) are sufficiently fast. Carbon based electrode materials (ex. felts, cloths,papers, pastes) are very common electrode materials for flow batteries (including the all-vanadium andzinc-halogen flow batteries) because they are inexpensive, electrically conductive and have highly tuneablestructures while also being sufficiently active towards many potential redox couples [21, 22].If the kinetics of the electrode reaction are very fast in comparison to the rate of mass transfer ofreacting species to/from the electrode surface, activation energy effects are not observed but concentrationpolarization that arises from the transport of reactants and/or products to/from the electrode surface becomesmore significant. The current at these conditions is dictated by the magnitude of the concentration gradientat the electrode surface [20]:i= nFKm[Cb−Cs] (1.25)Km =D jδ(1.26)131.2. Redox Flow Batteries for Energy Storagewhere Cb and Cs are the concentration of reacting species in the bulk electrolyte and electrode surface,respectively, Km (Equation 1.26) is the mass transfer coefficient for the system, D j is the diffusivity ofthe reacting species, and δ is the Nernst diffusion layer thickness. The concentration gradient grows as thecurrent is increased since the value ofCs will continuously decrease, which increases flux of reactant towardsthe electrode surface. In the limiting case of the electrode reaction being so fast that the electroactive speciesis immediately consumed upon reaching the electrode, the surface concentrationCs becomes 0 and the masstransfer limiting current density of the system is reached:il = nFKmCb (1.27)This limiting current density il is the highest possible current density for a system of electrolyte concentra-tion Cb and mass transfer coefficient Km; it is impossible to operate the cell at higher current densities untileither the concentration of the electrolyte is increased or mass transfer conditions are enhanced (or anothersecondary reaction takes place). The values of the diffusion overpotentials at the anode ηda or cathode ηdcare a function of the limiting current density:ηd,a = balog(1− iail ) (1.28)−ηd,c = bclog(1− icil ) (1.29)From Equation 1.28 and Equation 1.29, operating far below the limiting current density il ensures that thediffusion overpotential is small in Equation 1.11. The Thickness of the Nernst diffusion layer (µm) (δ )(Equation 1.25 and Equation 1.26) is a function of the hydrodynamic parameters of the system; comparedto traditional static batteries, the constantly circulating electrolyte in a flow battery decreases the value of δwhich in turn decreases ηd .The internal resistance in Equation 1.11 arises as an inevitability of physics, where the voltage E willdrop along any path of current I flowing through a circuit of finite resistance R (in units Ω) according toOhm’s law:E = I×R (1.30)The overall internal resistance includes contributions from both the electronic current flowing through the141.3. Chemical and Electrochemical Properties of Iodineexternal circuit and the ionic current travelling through the electrolyte and separator between the electrodes.For both electronic and ionic current, the resistance to the flow of current R is calculated byR= ρlA(1.31)where ρ is the resistivity of the material, l is the distance travelled, and A is the cross-sectional area. Theelectrical resistance arises from poor electrical conductivity of electrode, wires, and other cell components,as well as poor contact between mating cell components. In an electrochemical cell the electrical resistivityρe is almost always much lower than the ionic resistivity ρi and minimizing the ionic resistance is a commonissue in electrochemical cell design. Reducing the ionic contribution to the iR drop in an electrochemical cellrelies on using an electrolyte with a high ionic conductivity (decreasing ρi), using a separator material thatis highly permeable to the desired charge carriers, and keeping the spacing between the anode and cathodeas small as possible (to minimize l).1.3 Chemical and Electrochemical Properties of Iodine1.3.1 Iodine/Triiodide EquilibriaThe solubility of molecular iodine is very low in water, with 1 g dissolving in 3450 mL H2O at 20◦C[23]. However, in the presence of iodide I− the complexation reaction between electrophilic I2 and nucle-ophilic I− occurs:I2(s)+ I−(aq) ⇀↽ I−3(aq) (R1)The rate of both the forward and reverse reactions are high [24–26] and the equilibrium lies far to the right(Kstab = 723 at 20◦C in water solvent) [27–29], ensuring that there is very little free I2 in solution when I−is present in excess. Similar to iodide, the solubility of the triiodide ion I−3 in water is very high which cancircumvent the low solubility of molecular I2. There is proof of higher order polyiodide species (I−5 , I−7 ) inthe electrolyte formed from triiodide building blocks [30], and there have also been reports of polyiodinespecies stabilized by protons or other cations in the solution forming complexes such as I4H+ and I6H+[31]. For iodine flow batteries, I−3 is generally considered to be the most relevant polyiodide species, with151.3. Chemical and Electrochemical Properties of Iodineformation of I−5 or I−7 possible at high SoCs but readily dissociating to I−3 [32]. Due to its ns2np5 electronconfiguration iodine has a tendency to accept an additional electron (but is the weakest oxidizing agent outof the stable halogens). The triiodide ion has 22 valence electrons: six in the low energy s-orbitals, 12 inorbitals of π-symmetry, and 4 delocalized p-electrons that form the 3-center, 4-electron bond [33]. The I2highest occupied molecular orbital (Highest Occupied Molecular Orbital (HOMO)) is the major contributorto the 2σu orbital of I−3 , and the I− HOMO (5pz) is the major contributor to the 3σg orbital [23]. The stabilityof the I−3 species is due to (and can be enhanced by) the energy gap between the σ2 and σ2∗ orbitals of I2;when the gap is small, the stability of the complex increases due to a more favorable match with the 5pzorbital of I−. The small gap between the LUMO and HOMO of molecular I2 also causes the species to behighly polarizable, and will form strong interactions with nucleophiles that are likewise strongly polarizablewith low-energy orbitals.The electrophilic nature of iodine causes it to form charge-transfer adducts with nucleophiles other thaniodide. Iodine is capable of undergoing complexation reactions analogous to Reaction 1 in the presenceof other halide ions (Reaction 2 and Reaction 3) and most anion (or neutral charge nucleophile) species(Reaction 4):I2+Br− ⇀↽ I2Br− (R2)I2+Cl− ⇀↽ I2Cl− (R3)I2+X− ⇀↽ I2X− (R4)The stability of I2X− complexes (where X− is any electron rich species that is not I−) is generally lowerthan that of I−3 , with Kstab for Reactions Reaction 2 and Reaction 3 being 16 and 1.5, respectively [34–36]and the value of Kstab for Reaction 4 when X− = SCN− being 53.3 [37]. When dissolved in nucelophilicsolvents (including H2O) iodine forms adducts with the solvent molecules and the strength of the interactiondepends on the (Lewis) basicity of the solvent. The stability of the I−3 species is significantly higher inmost organic solvents than water, with values of Kstab for Reaction 1 being 105 in acetonitrile and 107 inpropylene carbonate [38] compared to 102.86 in water. Compared to water, less polar solvents do not solvatethe I− species as strongly which shifts the equilibrium further towards I−3 . A more detailed description ofthe solubility of I2 and stability of I−3 in different solvents is provided in Section 4.1.161.3. Chemical and Electrochemical Properties of Iodine1.3.2 Iodine ElectrochemistryThe identity of the iodine electrochemical reaction at the positive electrode is dependent on the amountof iodide and iodine relative to one another and the stability of the triiodide complex (Reaction 1) in thesolvent. When iodide is not available to complex with iodine, the main electrochemical reaction isI2(s)+2e− ⇀↽ 2I−(aq) E0 = 0.621VSHE · (R5)The ZIFB employs high concentrations of iodide salt to achieve high volumetric energy densities, and atthese conditions where triiodide formation is favourable the electrochemical reaction is more accuratelyrepresented byI−3 +2e− ⇀↽ 3I− E0 = 0.536VSHE (R6)At high anodic potentials, the triiodide species can be further oxidized to molecular iodine according to3I2+2e− ⇀↽ 2I3− E0 = 0.789VSHE · (R7)During charge (i.e. iodide oxidation) Reaction 5 and Reaction 6 compete with each other at the electrodesurface. At lower oxidation potentials (or low states of charge of the ZIFB) there is an abundance of I− atthe vicinity of the electrode surface and Reaction 6 is favoured. Reaction 1 and the oxidation reactions at theelectrode compete over the local supply of I−, and so at higher anodic potentials where there is limited I−at the vicinity of the electrode Reaction 5 and/or Reaction 7 become more significant and iodine is the mainoxidation product instead of triiodide. At even higher anodic potentials iodide and/or iodine can be furtheroxidized to monopositive iodine I+:2 [I(H2O)x]++2e− ⇀↽ I2+2xH2O E0 = 1.34VSHE (R8)with I+ being stabilized as 2[I(H2O)x]+ [39]. The I+ species can also be stabilized by the presence of othernucleophilic species, with the reported value of E0 for Reaction 8 being 0.99 and 0.741 V when carriedout in hydrochloric acid and hydrogen cyanide, respectively, due to the I+ species being stabilized throughformation of complexes such as ICl, 2ICl−2 and ICN [39, 40]. Depending on the electrolyte pH, iodine171.3. Chemical and Electrochemical Properties of Iodinecan also be oxidized to hypoiodite (IO−) and/or iodate (IO−3 ) [41, 42]. For the aqueous ZIFB where theelectrolyte is highly concentrated (1-5 M ZnI2) and neutral/mildly acidic, the most relevant iodine electrodereactions are Reaction 5, Reaction 6, and Reaction 7, but in non-aqueous solvents or in the presence ofstrongly nucleophilic additives (e.g. pyridine) the formation of I+ can become significant as well [39, 43].Comparing the stoichiometries of Reaction 5 and Reaction 6, 1 e− is transferred per I− in the former,but only 2/3rd an e− per I− for the latter, and so the theoretical capacity of the I−3 /I− redox pair will be a1/3rd less than that theoretically available from the I2/I− couple:I2/I−Capacity=ChargeTrans f erred f rom 2mol e−Masso f 1moleo f I2=53.6Ah126.9gcons= 0.422 Ah/g (1.32)I−3 /I−Capacity=ChargeTrans f erred f rom 2mol e−Masso f 1moleo f I−3=53.6Ah190.35gcons= 0.281 Ah/g (1.33)That is, although the formation of the I−3 species is advantageous in that it significantly increases the solubil-ity of iodine in aqueous solvents, a third of the I− in the electrolyte will have to bear the brunt of acting as acomplexing agent and its associated capacity will be unavailable. However, the electrochemical generationof iodine (as opposed to triiodide) during charge is still undesirable for flow batteries because the weaklysoluble iodine will precipitate out from the electrode surface and form what is commonly referred to in lit-erature as the iodine film. When the rate at which iodide is oxidized to iodine becomes greater than the rateat which iodine can be removed by dissolution, complexation as triiodide, or by diffusion, the concentrationof iodine adjacent to the electrode surface becomes saturated (solubility of I2 in water at 22 ◦C is 0.29 g/L)and solid iodine precipitates from the surface which impedes further oxidation (Figure 1.4). The iodine filmand minimizing its impact on iodine based flow batteries is discussed in greater detail in Chapter 3.Analogous to Reaction 6, combining Reaction 5 with Reaction 2 or Reaction 3 results in the electrodereactions for the I2Br− and I2Cl− species, respectively:I2Br−+2e− ⇀↽ 2I−+Br− (R9)I2Cl−+2e− ⇀↽ 2I−+Cl− (R10)Using Equation 1.3 with values of ∆Go for Reaction 6, Reaction 9, Reaction 10 (calculated from the standard181.4. Halogen/Halide Based Flow Batteries   Iodine Film 2I- I2(s) + 2e- 3I- I3- + 2e- I3- I2(s) + e- I2(s) I2 I2 + 3I- 2I- + I3-     (solv)(solv)(solv) (solv)(solv)(solv)(film)I2 + I- I3- 32(solv)(solv)   ElectrodeFigure 1.4: Schematic of the iodine film.Gibbs free energy of formation ∆G0f of the participating species [34]) leads to the values of E0 for the I−3 ,I2Br− and I2Cl− electrode reactions being 0.536 V, 0.594 V, and 0.614 V (versus the SHE) respectively. Theamount of charge transferred per I− ion in Reaction 9 and Reaction 10 is identical to that for the I2/I− pair,and so the theoretical capacity of these heteropolyhalide redox reactions is the same. Recent zinc-iodideflow battery designs have incorporated these I2Br− and I2Cl− complexes for the reasons outlined above, andthe analogous Br2Cl− heteropolyhalide species has been utilized in some bromine based flow batteries aswell [44, 45].1.4 Halogen/Halide Based Flow Batteries1.4.1 The Zinc AnodeThe advantages of the Br2/Br− and I2/I− redox pairs for the positive half-cell of flow batteries wasdiscussed previously in Section 1.1. The promise of these halogen/halide half-cells have made them verycommon for numerous different flow battery chemistries and have been coupled to different negative half-191.4. Halogen/Halide Based Flow Batteriescells such as zinc [46, 47], hydrogen [48], vanadium [49], polysulfide [50], organic species derived fromquinone [51, 52], and lithium [53]. The zinc anode has been employed in batteries since the first battery200+ years ago, and to date remains one of the most common electrodes used for the negative half-cellbecause of its reliability, stability to water electrolysis, and low cost [54, 55].In neutral and acidic electrolytes, the Zn2+/Zn redox reactions occur byZn2++2e− ⇀↽ Zn E0 =−0.76VSHE (R11)and in alkaline media occurs byZnOH2+2e− ⇀↽ Zn+2OH− E0 =−1.26VSHE (R12)The value of E0 for the ZnOH2/Zn redox couple is much higher than that of Zn2+/Zn, theoretically leadingto much higher energy densities because of the higher voltage of the cell. However, the electrochemicalbehaviour of iodine is poor in alkaline electrolytes and employing a ZnOH2/Zn negative electrode requiresimplementing additional measures to ensure that the iodine half-cell remains neutral/acidic [56]. Therefore,only the Zn2+/Zn pair will be discussed further in this study. Among the abundant low-cost metals (zinc,lead, chromium, cadmium, iron, manganese) zinc has the highest energy content due to its negative standardpotential of -0.76 VSHE and high gravimetric capacity:Zn2+/ZnCapacity=ChargeTrans f erred f rom 2mol e−Masso f 1moleo f Zn=53.6Ah65.38gcons= 0.8198 Ah/g (1.34)which is approximately double that of I2/I− (Equation 1.32) and triple that of I−3 /I− (Equation 1.33). Withthis in mind, it is the capacity of the halogen/halide half-cell that will typically limit the overall capacityof zinc-halogen flow batteries. Improvements to the iodine (or bromine) electrode allow these batteries tooperate closer to their maximum capacity and accordingly the majority of this thesis will be dedicated to thehalogen side of the flow battery instead of the negative zinc electrode.The zinc anode has been successful in primary (unrechargeable) batteries, where the solid zinc electrodeis electrochemically dissolved into Zn2+ ions (Reaction 11) until the end of its lifetime. Historically, it hasbeen much more difficult to implement the zinc electrode in secondary (rechargeable) batteries because the201.4. Halogen/Halide Based Flow Batteriesreduction of Zn2+ to Zn occurs unevenly over the electrode surface resulting in long dendritic deposits [57,58]. These zinc dendrites elongate over prolonged charge and can eventually pierce the battery separator,resulting in an electrical short-circuit of the cell. The dendrite morphology can also increase the specificsurface area of the zinc electrode which facilitates zinc corrosion, and the poor adhesion of the dendritedeposits causes them to detach easily which decreases the capacity of the zinc electrode [59, 60]. Theproblems associated with the reduction of zinc on charge has been a long known issue with over a centuryof research available on the topic, and there are now various rechargeable zinc batteries (static and flow)commercially available [61, 62] including the zinc bromine flow battery [63]. The circulating electrolytein flow batteries helps alleviate (but not completely eliminate) the dendritic behavior of the zinc anode [64]making the Zn2+/Zn promising for aqueous flow battery technology.1.4.2 The Zinc-Bromine Flow BatteryA patent for the zinc-bromine flow battery (ZBFB) was filed in 1978 by the Exxon company [65]. Overthe decades since, the ZBFB has become the second most researched and developed flow battery system(behind the all-vanadium RFB). Similar to iodine, the solubility of molecular bromine is increased in thepresence of Br− through the formation of tribromide (Br−3 ):Br2(s)+Br− ⇀↽ Br3− (R13)The stability of the tribromide species (Kstab = 17 at 20◦C for Reaction 13 [66–68]) is much lower than thatof triiodide, but the contribution of the Br−3 species at the positive electrode of the ZBFB is still significantenough to necessitate taking it into consideration and the bromine electrochemical reactions are therefore acombination of [69, 70]:Br2+2e− ⇀↽ 2Br− E0 = 1.09VSHE (R14)Br−3 +2e− ⇀↽ 3Br− E0 = 1.05VSHE (R15)211.4. Halogen/Halide Based Flow BatteriesCompared to Equation 1.32 and Equation 1.33, the gravimetric capacity offered by the bromine redox reac-tions is higher than that for iodine:Br2/Br−Capacity=ChargeTrans f erred f rom 2mol e−Masso f 1moleo f Br2=53.6Ah79.904gcons= 0.67 Ah/g (1.35)Br−3 /Br−Capacity=ChargeTrans f erred f rom 2mol e−Masso f 1moleo f Br−3=53.6Ah119.85gcons= 0.447 Ah/g (1.36)but on a volumetric basis the energy density offered by the iodine electrolyte is able to far surpass that ofbromine while also avoiding the use of the highly concentrated, toxic and corrosive Br2 electrolytes.Combining Reaction 11 with Reaction 14 or Reaction 15, the cell reaction describing the zinc-bromideflow battery (ZBFB) is:Zn+Br2 ⇀↽ Zn2++2Br−E0 = 1.85V (R16)Zn+Br3− ⇀↽ Zn2++3Br−E0 = 1.81V (R17)In practical ZBFBs, the evolution of bromine during charge (and preventing the release of its toxic vapoursinto the surrounding environment) must be carefully controlled. When the ZBFB was introduced in the1980s, unsymmetrical substituted quaternary ammonium bromides such as N-ethyl-N-methyl-morpholinium(MEM) and N-ethyl-N-methyl-pyrrolidinium (MEP) were employed as complexing agents for Br2. Thecomplex that results between these large organic cations and bromine (QBrn) exists as a dense hydrophobicorganic phase that separates to the bottom of the storage tank, thereby alleviating self-discharge and brominegas evolution [71, 72]. However, the inclusion of these organic bromine complexing agents add costs to thesystem, reduces the kinetics of the Br2/Br− redox reactions, and decreases the overall energy density andefficiency of the ZBFB [73, 74]. The slow kinetics of the Br2/Br− reaction (particularly the reduction ofbromine to bromide) is another major obstacle in improving the performance of current ZBFBs [75, 76] andlimits the current densities that they can feasibly operate at. The adsorption of bromide and bromine are therate determining steps for the oxidation and reduction reactions, respectively, and ZBFB research has pri-marily focused on identifying electrode materials that lower the activation energy barrier for the adsorptionprocesses [77]. Carbon based electrode materials have been extensively studied for the ZBFB because oftheir low cost, tunable properties, and electronic conductivity [78, 79], including carbon blacks [80], carbonnanotubes [81, 82] hetero-atom (fluorine, boron, nitrogen) doped carbons [83–86], and several other novel221.4. Halogen/Halide Based Flow Batteriesmaterials [87, 88].1.4.3 The Zinc Iodine Flow BatteryThe zinc iodine flow battery (ZIFB) offers an alternative to the commercialized ZBFB which avoids thehazardous nature of bromine [89, 90]. Iodide salts are highly soluble in aqueous solvents (solutions of ZnI2are able to reach concentrations of 12 M in water at room temperature [91]), and as a result lab-scale ZIFBshave been reported to achieve volumetric energy densities of 200 Wh L−1 [45]. In comparison, the prevalentall-vanadium redox flow battery operates at energy densities around 25-40 Wh L−1 [44] as a consequenceof the low solubility of V(II) and V(III) ions in the acidified battery electrolyte [92, 93] while the ZBFBtypically operates at 50-70 Wh L−1 [75]. The electrochemical reactions in the ZIFB are represented byZn+ I2 ⇀↽ Zn2++2I− E0 = 1.38V (R18)Zn+ I3− ⇀↽ Zn2++3I− E0 = 1.30V (R19)where I− is oxidized to I2 and/or I−3 on charge depending on the availability of iodide, and then reducedback to I− on discharge.The first ZIFB was reported in 2015, capitalizing on the high solubility and relatively benign nature ofthe I−3 /I− redox species [17]. Some notable advances in the ZIFB over the past decade are highlighted inTable 1.3. The first lab-scale ZIFB was able to achieve a discharge energy density of 166.7 Wh L−1 at 5.0M ZnI2, far surpassing that of the VRB and ZBFB despite only operating at an energy efficiency of 63%[17]. The low operating efficiency of the ZIFB was ascribed primarily to the I−3 /I− reaction dominating overthat of I2/I− and so a later design used an electrolyte containing ZnBr2 ([I−]:[Br−] of 2:1) in order to formI2Br− over I−3 and free up iodide in solution to participate in the electrode reactions instead of solubilizingI2 as I−3 . This flow battery was able to achieve an energy density of 202 Wh L−1, and at the time was thehighest energy density reported for any aqueous redox flow battery [45]. The chloride anion Cl− and I2Cl−species has similarly been used for this purpose [94]. Building on this method of solubilizing the I2 species,polyvinylpyrrolidone (PVP) has been employed in ZIFB electrolytes to form a I2-PVP complex akin to I−3[95]. Further modifications to the ZIFB electrolyte include opting to use an alkaline zinc half-cell, leading231.4. Halogen/Halide Based Flow Batteriesto an increase of 0.497 V in battery voltage and a higher energy density of 330.5 Wh L−1 [56] but is stillonly 57% of the theoretical energy density (577.6 Wh L−1). To avoid the cost of ZnI2, salts of ammonium(NH4I, NH4Br, NH4Cl) have been used in its place [96, 97]. The NH+4 salts were also found to mitigatedendrite formation at the zinc electrode by acting as a weakly acidic buffer that inhibits hydrolysis of Zn2+and promoting uniform deposition of zinc, while still allowing for the I2Cl− reaction to occur over I−3 in thepositive half-cell to minimize capacity loss from I−3 formation. While a majority of research on the ZBFBhas focused on the electrode composition, research on the ZIFB has primarily focused on the electrolytecomposition and the solution-phase iodine interactions that affect battery performance. In this thesis, theelectrolyte composition of the ZIFB is investigated in Chapter 3 and Chapter 4 and electrodes for the ZIFBare studied in Chapter 5.Table 1.3: Review of the most significant recent advances in the development of the ZIFB. CD = currentdensity, CE = current efficiency, EE = energy efficiency.Year Contribution/Contribution Electrolyte Performance2015 [17] First ZIFB 5 M M ZnI2 At 5 mA cm−2, CE = 96.3 %, VE =70.4 %, EE = 67.8 %, achieved 167Wh L−1 on discharge2017 [45] Iodide/bromide mixedelectrolyte I2Br−5 M ZnI2 + 2.5 MZnBr2At 10 mA cm−2, achieved 202 WhL−1 on discharge2018 [94] kW scale design 5 M KI + 2.5 M ZnBr2+ 1 M KClAt 80 mA cm−2, EE = 80 % andachieved 69 Wh L−1, 1000 charge-discharge cycles2018 [56] Alkaline electrolyte 6 M KI + 6 M KOH At 20 mA cm−2, CE = 100 %, VE(and EE) = 70 %, achieved 200 WhL−12019 [98] Single Flow Battery 7.5 M KI + 3.75 MZnBr2At 40 mA cm−2, EE = 81 % andachieved 205 Wh L−12021 [95] I2-PVP electrolyte 1 M KI + 5 vol% PVP At 20 mA cm−2, EE = 70 % andachieved 115 Ah L−1241.5. Research Approach1.5 Research ApproachThe preceding literature review has highlighted the advantages of the zinc-iodine flow battery for sta-tionary kWh and MWh energy storage. The ZIFB is able to achieve the highest energy density among allaqueous flow batteries, yet operates at a very low energy efficiency due to the iodine film that forms duringcharge and the large activation overpotentials associated with the iodine electrode reactions. The goals of thisresearch were to investigate the impacts of the operating conditions, electrolyte composition, and electrodematerial on the electrochemical behaviour of the iodine redox reactions and performance of the zinc-iodineflow battery Figure 1.5. The iodine and iodine charge-transfer redox reactions were studied under potentio-static and galvanostatic conditions in electrochemical half-cell experiments prior to being implemented intothe full zinc-iodine flow battery. The influence of various organic cosolvents on the solubility and electro-chemical behaviour of the iodine redox species in the zinc-iodine flow battery was extensively studied andcharacterized. Furthermore, the kinetics of the iodine redox reactions on different commercially availablehigh-activity carbon catalyst materials was studied and the composition of the positive electrode was variedto lower the large activation overpotentials of the ZIFB. The outline of the dissertation is as follows:• Chapter 1 provided a background on the relevant chemical properties of the iodide/iodine/triiodidespecies and the electrode reactions involved in the ZIFB.• Chapter 2 details the electrochemical half-cell and flow battery apparatus used throughout the study,the relevant electrochemical experiments, the non-electrochemical physical characterization methods,and calculations for commonly used metrics used throughout the dissertation.• Chapter 3 presents an introductory look at the electrochemical behaviour of the iodide/iodine/triiodideredox reactions, which includes a detailed look at the iodine film that forms on charge and its impactson the ZIFB. The effects of temperature, iodine-ion interactions, and interactions between iodine withacetonitrile cosolvent on the solubility of iodine, the iodine film, and the charge/discharge reactionswere investigated parametrically prior to constructing the full zinc-iodide flow battery.• Chapter 4 elaborates on the iodine-cosolvent interactions introduced in Chapter 3, providing a muchmore detailed investigation on the effects of different organic cosolvents on the solubility of iodineand overall flow battery performance.• Chapter 5 provides a detailed review on prior research of electrode materials used for the iodine redox251.5. Research Approachreactions in both static and flow batteries. The kinetics of the iodine electrode reactions are studiedon various comercially available carbon catalyst materials and incorporated into the positive electrodeof the ZIFB. The electrochemical kinetics of the iodide/iodine redox reactions were also directlycompared to that of bromide/bromine.• Chapter 6 gives a brief summary of the key findings from the electrolyte and electrode studies of theZIFB described in this research. Potential future research directions are listed as well.Figure 1.5: Research approach for the dissertation.26Chapter 2Experimental Methodology2.1 Electrochemical Half-Cell Design and ProceduresElectrochemical experiments were performed in a glass cell with a water jacket for temperature control.A typical three-electrode set up was used consisting of a circular glassy carbon electrode (GCE) (geometricarea = 0.2 cm2) for the working electrode (WE), a Pt sheet (geometric area = 1 cm2) counter electrode(CE), and a Ag/AgCl (sat’d KCl) (0.199 V vs. SHE) reference electrode (RE) unless otherwise stated.The temperature coefficient of this reference electrode is 0.2 mV/◦ and so changes in the RE potential withtemperature during the study were negligible. The circular glassy carbon WE was used for experiments instagnant electrolytes as well as for rotating disk electrode (RDE) experiments under forced convection atdifferent rotation rates. The working electrode was vertically mounted in a Modulated Speed Rotator for allhalf-cell experiments, which could also apply a constant rotational velocity to the WE during experimentsinvolving forced-convection. The electrode used for rotating ring disk experiments (RRDE) is similar tothat of the RDE, except that the glassy carbon disk electrode of the RRDE is encircled by a platinumconcentric ring, with a concentric ring of unreactive insulative material (made of polyether ether ketone,PEEK) separating the disk and ring electrodes (Figure 2.1). The RRDE was used extensively in this workfor studying the iodine film that forms on oxidation, where the potential at the disk electrode was linearlyswept to increasingly oxidative potentials while the platinum ring was held at a constant reduction potential(0 V Ag/AgCl [sat’d KCl]). The E6R2 Fixed-Disk RRDE had a geometric area of 0.275 cm2 and theconcentric Pt ring had a geometric area of 0.238 cm2. The collection efficiency Ne f f (percentage of materialthat can be captured by the ring based on radius of the disk and ring relative to each other) for the dimensionsof this RRDE is 38%. The glassy carbon electrode (both RDE and RRDE) was mechanically polished with272.1. Electrochemical Half-Cell Design and Procedures0.05 µm alumina and rinsed in IPA and DI water between every experiment due to the tendency of iodideand iodine to be strongly adsorbed onto carbon and platinum surfaces. The WE and CE were also routinelypolished electrochemically in 0.5 M H2SO4 by scanning several times at 50 mV s−1 from 1 V to -1 V(Ag/AgCl, sat’d KCl) to remove any residual adsorbed/deposited iodine or zinc species.                         ­ Figure 2.1: Schematic and dimensions of the RRDE used in this study. The disk is swept linearly to increas-ingly anodic potentials to oxidize iodide while the ring is held at a constant reduction potential of 0 V toreduce iodine/triiodide generated from the disk back to iodide. Collection efficiency of this RRDE geometryis 38%.All electrochemical measurements (cyclic voltammetry, linear sweep voltammetry with RDE or RRDE,chronoamperometry & chronopotentiometry, electrochemical impedance spectroscopy, flow battery charge& discharge) were performed using a Metrohm multi-channel bipotentiostat system on Nova software andexperiments using the RDE or RRDE were performed with a speed control module from Pine InstrumentCompany. All experiments were performed multiple times to ensure reproducibility. Impedance measure-ments were made by applying a perturbation voltage of 10 mV vs. the open-circuit voltage between afrequency range of 100 kHz to 0.1 Hz. Kinetic parameters were determined from iR-corrected RDE datataken at 1000 rpm. Several half-cell experiments used an inert supporting electrolyte of either 0.5 M NaClO4for zinc salt electrolytes or 0.5 M KNO3 for potassium salt electrolytes (due to formation of the insolubleKClO4 species) to decrease the internal resistance of the cell. For chronoamperometry and chronopotentiom-etry half-cell experiments, supporting electrolyte was not used because the concentration of the electrolytes282.2. Flow Battery Apparatusused in these experiments (typically 0.5 M ZnI2 or 1.0 M KI) are high enough to provide adequate ionicconductivity.2.2 Flow Battery Apparatus              Figure 2.2: Schematic of the flow battery apparatus used for ZIFB tests (ElectroCell Technologies). AFumapem F-1050 membrane is sandwiched by two gaskets, followed by a zinc-metal electrode on the nega-tive side and a graphite felt (with or without catalyst treatment) on the iodine side, with geometric electrodeareas of 10 cm2. The electrodes are in contact with separate graphite current collectors which are clampedtogether by a PTFE end-frame and stainless steel end-plates. The volume of posolyte and negolyte were 50mL each, and circulated through the cell by flow-through electrode mode at a flow rate of 20 mL/min.Zinc-iodine flow battery experiments were performed with the Micro Flow Cell purchased from Elec-troCell Technologies Figure 2.2. On the iodine side, graphite felts (AvCarb G280A, 2.8 mm thick and bulkdensity 0.08 g cm−3) with or without catalyst ink treatments were used as the positive electrodes, whereas apiece of commercially pure zinc metal (0.2 mm thick, McMaster-Carr) was used as the negative electrode.The geometric surface area of both electrodes was 10 cm2. The electrodes were sandwiched together be-tween graphite plate current collectors that were in turn supported by stainless steel end-plates, and the flow292.2. Flow Battery Apparatuscell was tightened to 30 lb in−1 for all tests to ensure the inter-electrode spacing was consistent betweenexperiments. The positive and negative electrolytes typically consisted of 50 mL of 1 M ZnI2 or 2 M KI(unless otherwise stated) circulating at a flow rate of 20 mL min−1, operated galvanostatically at appliedcurrent densities of 10, 20, or 40 mA cm−2.0.60.91.21.50 1000 2000 3000 4000Potential / VTime / s         20 mA cm-2(a)00.40.81.21.62 3 4 5 6 7 8Z'' / ohm ∙ cm2Z' / ohm ∙ cm2            ••••(b)0.20.50.81.11.41.720 2000 4000 6000 8000Potential / VTime / s  20 mA cm-2                    (c)Figure 2.3: Charge–discharge performance of a ZIFB constructued with a Fumapem F-1050 and Nafion®cation exchange membranes (a), Nyquist plots for the Fumapem F-1050 and Nafion® ZIFBs (b), and charge-discharge performance of a ZIFB constructued with a Fumapem F-1050 CEM, a Daramic microporousseparator, and no separator/membrane at all (c). Flow battery tests were performed at 20 mA cm−2geom with50 mL of 1 M ZnI2 electrolyte circulating at 20 mL/min through both half-cells (T = 22 ◦C).The separator employed for a majority of flow battery experiments throughout this study was FumapemF-1050 (abbreviated as F-1050), a commercially available cation exchange membrane that provides a cheaperalternative to the Nafion® membranes commonly employed in other ZIFBs. Nafion® membranes also re-quire time intensive pre-activation methods and require stringent storage conditions, while the F-1050 mem-302.3. Spectroscopybrane is advertised as being pre-activated on arrival. The performance of a ZIFB constructed from F-1050and Nafion® are compared in Figure 2.3a and Figure 2.3b, where it can be seen that the internal resistanceof the F-1050 membrane in the ZIFB environment is higher than that of Nafion®. However, F-1050 was stillable to prevent any noticeable crossover after several days of cycling, and was selected as the membrane for amajority of flow battery experiments herein because of its lower costs. Modifying the composition/thicknessof the ZIFB membrane (which falls out of the scope of this study) can further improve the performance ofthe flow battery. Figure 2.3c compares the performance of a ZIFB constructed from F-1050 to an identicalZIFB that had been constructed using a polyethylene-based microporous separator (Daramic), and a ZIFBthat was constructed with no membrane or separator at all. With no separator, the coulombic efficiencyof the ZIFB is very low because triiodide can directly react with the zinc electrode via the spontaneousself-discharge reactionZn(s)+ I3− ⇀↽ Zn2++3I− · (R20)Experiments where a piece of Zn metal was left in a solution containing 0.1 M I−3 caused the zinc to becompletely consumed (via Reaction 20) overnight, highlighting the need to prevent as little triiodide fromcrossing over into the negative half-cell as possible. The CE of the F-1050 ZIFB is higher than that ofDaramic because of the negatively-charged moieties on F-1050 that enhance ion selectivity by excluding thenegatively-charged I−3 . Although the VE of Daramic was comparable to that of F-1050, significant crossoverwas visually evident from the red colour in the zinc-side electrolyte even after an hour of operation whichresulted in a much lower coulombic efficiency compared to F-1050 (Figure 2.3c), and the ZIFB constructedwith a F-1050 membrane was able to operate for several days at a time without any noticeable I−3 crossover.2.3 SpectroscopyUV-Vis experiments were performed with a Beckman spectrophotometer and 4.5 cm3 volume pipettes.These experiments were performed to investigate the solubility of iodine in different electrolytes, which in-cluded samples with or without iodide in order study the iodine and triiodide species separately. For samplesthat contained iodide (either from KI or ZnI2), the ratio of iodine:iodide was consistently kept at 2:1. In theneat aqueous solvent, iodine gives rise to a peak at approximately 450 nm, while iodide/triiodide appears astwo peaks at 250 and 380 nm Figure 2.4. When iodide is added to the electrolyte, the peak corresponding to312.4. Catalyst Preparationiodine vanishes due to the iodine being converted to triiodide. Scanning electron microscopy (SEM) imageswere performed with a Hitachi S2600 environmental/low vacuum variable pressure SEM.00.511.522.530 200 400 600 800Absorbance / AUWavelength / nmI2KI I2ZnI2I2I2 + KII2 + ZnI2 Figure 2.4: UV-Vis spectra for aqueous samples containing 0.1 M I2, 0.1 M I2 + 0.05 M KI, and 0.1 M I2 +0.025 M ZnI2.2.4 Catalyst PreparationCatalyst inks were prepared from different carbon active materials (Vulcan XC-72R, Ketjenblack EC-300J, Ketjenblack EC-600J, and multi-walled carbon nanotubes) (Table 2.1). 75 mg of carbon catalyst wasadded to 14.26 mL of isopropyl alcohol (IPA) as solvent and 0.74 mL of Nafion® solution (5 wt%) was usedas a binder, and the catalyst layer (CL) had a dry wt% composition of 70:30 carbon:Nafion®. For half-cellexperiments, a freshly polished glassy carbon electrode GCE (geometric area = 0.2 cm2) was used as thesubstrate for drop-casting of the catalyst ink. 10 µL of ink was carefully drop-casted onto the GCE andallowed to dry at room temperature, resulting in a catalyst loading of 0.25 mg cm−2. The electrochemicallyactive surface area (ECSA) for the various catalyst layers on the glassy carbon was calculated by measuringthe double layer capacitance CDL and specific capacitance Cs of the electrodes Section 2.5. For flow batteryexperiments, AvCarb G280A graphite felts (2.8mm thick and bulk density 0.08 g cm−3) were treated with322.5. Determination of the Electrochemically Active Surface Areathe catalyst inks. The felts were first cleaned by thoroughly sonicating in water and then dried at 60◦ C for24 hours. The catalyst inks were then deposited onto the felt by one of two methods: 1) soaking the felt inthe ink under ultrasonication for 15 minutes to uniformly coat the entire three-dimensional substrate, or 2)directly spraying the ink onto the surface of the felt in a region that is facing (and closest to) the membrane.The felts were dried overnight at 60◦C to remove any residual IPA solvent from the electrode, and the massof the felts was measured before and after the ink treatments to determine the catalyst loading. All treatedfelts had a catalyst loading of 2 mg cm−2 unless stated otherwise. Scanning electron microscopy (SEM)images were performed with a Hitachi S2600 environmental/low vacuum variable pressure SEM.Table 2.1: Overview of carbon catalysts used in this study, with physical data provided from the suppliers.Catalyst Abbreviation Supplier Bulk Density /kg m−3BET SurfaceArea / m2 g−1Vulcan XC-72R VXC-72R Cabot 96 250Ketjenblack EC-300J Ketjen-300 Nouryon 125-145 800Ketjenblack EC-600J Ketjen-600 Nouryon 100-120 1400Multiwalled CarbonNanotubesMWCNT Cheaptubes 280 602.5 Determination of the Electrochemically Active Surface AreaThe electrochemically active surface area (ECSA) of an electrode can be calculated from the doublelayer capacitance CDL [F or C V−1] and specific capacitance Cs [F cm−2]:ESCA=CDLCs(2.1)The value of CDL can be determined by performing a series of cyclic voltammograms at various scan ratesin a potential range where no faradaic reactions occur [99]. The double-layer charging current icap is equalto the product of the scan rate ν and the double layer capacitance CDL:icap = vCDL (2.2)By plotting the charging current in the non-faradaic region as a function of the scan rate, a linear plot shouldbe observed with a y-intercept close to zero and a slope corresponding to the value of CDL (Figure 2.5). The332.6. Calculation of Kinetic Parameterspotential range used for measuring the capacitance of the electrodes in the iodide electrolytes was 0.05 to0.2 V (vs. Ag/AgCl [sat’d KCl]). The capacitance of a material is defined as the charge stored in responseto a potential differenceC =∆Q∆V(2.3)which can be normalized by the surface area of the electrode material Aelectrode:Cs =∆Q∆VAelectrode(2.4)Since ∆Q= I∆t and scan rate ν = E/t, the specific capacitance can be expressed asCs = ACV/(2Aelectrodeν∆V ) (2.5)where ACV is the integral area enclosed by the cyclic voltammogram taken in the non-faradaic region ofpotential window ∆V and scan rate ν [100]. The experimentally calculated values ofCDL,Cs, and ECSA aresummarized in Table 2.2.Table 2.2: Values of ECSA and roughness factor RF (= AECSA/Ageom) determined by calculating the double-layer capacitance CDL and Cs via equations Equation 2.2 and Equation 2.5 respectively.Double-Layer Capacitance/ mFSpecific Capacitance / mFcm−2ECSA / cm2 RFGlassy Carbon 1.1 5.0 0.21 1.1VXC-72R 0.91 2.3 0.38 1.75Ketjen-300 1.9 4.7 0.40 1.95Ketjen-600 3.1 7.6 0.41 2MWCNT 0.13 0.39 0.33 1.552.6 Calculation of Kinetic ParametersIntroduced in Subsection 1.2.3, the current-overpotential relationship in a well-mixed electrolyte isexpressed as Equation 1.19i= i0[eαa fη − e−αc fη ]342.6. Calculation of Kinetic Parameters-0.15-0.1-0.0500.050.10.150 50 100 150 200 250Current / mAPotential / mV (vs. Ag/AgCl Sat'd KCl)     Charge Current / mAScan Rate V s-1R2 = 0.98-0.2-0.100.10.20.30 50 100 150 200 250Current / mAPotential / mV (vs. Ag/AgCl Sat'd KCl) ǇсϮ͘ϬϴϯǆZϸсϬ͘ϵϲϭϭ  Charge Current / mAScan Rate V s-1R2 = 0.97-0.3-0.2-0.100.10.20.30.40.50 50 100 150 200 250Current / mAPotential / mV (vs. Ag/AgCl Sat'd KCl) Ǉсϯ͘ϰϭϱϳǆZϸсϬ͘ϵϲϬϯ  Charge Current / mAScan Rate V s-1R2 = 0.97-0.1-0.0500.050.10.150 50 100 150 200 250Current / mAPotential / mV (vs. Ag/AgCl Sat'd KCl) Ǉсϭ͘ϬϰϭϯǆZϸсϬ͘ϵϱϱϮ  Charge Current / mAScan Rate V s-1R2 = 0.97-0.04-0.0200.020.040 50 100 150 200 250Current / mAPotential / mV (vs. Ag/AgCl Sat'd KCl)  y = 0.1613xR² = 0.991700.0050.010.0150.020 0.05 0.1Charge Current / mAScan Rate V s-1R2 = 0.99Figure 2.5: Cyclic voltammograms of the capacitance current on various different electrode materials (glassycarbon, Ket-300, Ket-600, VXC-72R and MWCNT) for calculating the double layer capacitanceCDL. Inletsshow linear fit of icap as a function of v. (electrolyte: 0.05 M ZnI2 + 0.5 M NaClO4). Performed at scanrates of 5, 10, 25, 50, 75, 100, and 200 mV s−1.At small overpotentials, Equation 1.19 can be estimated by the low-field approximation:i= i0 fη (2.6)Since the iodide oxidation reaction is complicated by mass-transfer effects from the iodine film at moder-ately low anodic overpotentials, the low-field approximation is more appropriate than the Tafel equation for352.7. Materials and Reagentscalculating the exchange current density i0 for this system. The ratio of η/i is referred to as the charge-transfer resistance Rct , and can be calculated by rearranging Equation 2.6:Rct =RTFi0(2.7)2.7 Materials and ReagentsTable 2.3: List of reagants and solvents used in this study, including purity/grade and supplier data.Supplier Reagent/SolventSigma Aldrich Zinc Iodide (ZnI2, 98%) Acetonitrile (ACN)Potassium Bromide (KBr, ≥ 99.0%) Dimethyl sulfoxide (DMSO)Sodium Perchlorate (NaClO4, 99%) N-Methyl-2-pyrrolidone (NMP)Zinc nitrate (Zn(NO3)2·6H2O, 99%) Dimethoxyethane (DME)Zinc sulphate (ZnSO4·7H2O, 98%) Methanol (MeOH)Ethanol (EtOH)2-propanol (IPA)Methanesulfonic acid (MSA)ThermoFischer Zinc Bromide (ZnBr2, 98%)Zinc Chloride (ZnCl2, ≥ 99.0%)Potassium Iodide (KI, 99%)Potassium Chloride (KCl, 98%)Potassium Nitrate (KNO3, 99%)Supelco Dimethylformamide (DMF)Acros Organics Propylene carbonate (PC)Pyridine (PYR)Ethylenediamine (ETDI)36Chapter 3Interplay of I2-Ion, I2-Solvent, andTemperature Effects on I2/I−3 /I− RedoxCouples3.1 The Iodine ElectrodeAs described in Subsection 1.3.2, the relevant electrochemical reactions at the positive electrode of theZIFB are Reaction 6I3−+2e− ⇀↽ 3I− E0 = 0.536VSHEand Reaction 5I2(s)+2e− ⇀↽ 2I− E0 = 0.621VSHEThe former is the predominant reaction at the electrode when there is sufficient iodide at the electrode surfaceto form triiodide through the electrolyte-phase complexation reaction (Reaction 1):I2(s)+ I− ⇀↽ I3− ·During prolonged charge the local concentration of iodide at the electrode surface decreases causing Re-action 5 (and/or Reaction 7) to be the primary oxidation reaction(s) resulting in iodine being the primaryoxidation product instead of triiodide. The solubility of iodine in water is low without I−, and when the localconcentration of iodine generated at the electrode surface reaches its saturation point, the onset of iodine373.1. The Iodine Electrodefilm formation occurs with precipitation of solid iodine from the surface (Figure 3.1). 2I- I2(s) Film Precipitation when C S + C + C     2e-I2(solv) I3-(solv) I2X-DiffusionRDSI2(solv) e-I- X-   Dissolution until C = SI2 I2High SoCI2 I2 I3 I2XElectrode    Iodine Film C = SI2 I2-≥ -(solv)I- I2Figure 3.1: Schematic of iodine film. Formation of the iodine film occurs when the local concentrationof iodine CI2 exceeds its solubility SI2 and the amount of iodine present as triiodide I−3 (CI−3 ) or any otheriodine-charge transfer complex I2X− (CI2X−)From Equation 1.25 and Equation 1.26, the current density arising from the oxidation of iodide (as-suming fast electrode kinetics) can be calculated asinF=DI−δ[CbI−−CsI− ] orDI2δ[CsI2 −CbI2 ] (3.1)where D is the diffusivity of iodide/iodine, Cb and Cs are the concentration of iodide/iodine in the bulkelectrolyte and at the electrode surface, respectively, and δ is the thickness of the Nernst diffusion layer(the length of the concentration gradient). For a typical electrode reaction the current will continue to risewith the growing concentration gradient until the surface concentration of reactant is depleted, i.e. when thereaction is proceeding at a rate high enough that iodide is oxidized immediately upon reaching the electrode383.1. The Iodine Electrodesurface (CsI− = 0):il = 2FDI−δCbI− = 2FDI2δ[CsI2 −CbI2 ] (3.2)Due to the iodine film, the oxidation of iodide is limited by the diffusional flux of product iodine away fromthe electrode and not the arrival of reactant iodide, and the actual limiting current of iodide oxidation iss iscalculated asiss = 2FDI−δ[CbI−−CsI− ] = 2FDI2δ[CsI2 −CbI2 ] = 2FDI2δ[SI2 −CbI2 ] (3.3)where SI2 is the maximum solubility of I2 at the conditions of the electrode-electrolyte interface, and is themaximum possible local concentration of iodineCsI2 before supersaturation occurs. The I2 film starts to format the electrode surface once the following inequality is reached:CsI2 ≥ SI2 +CsI−3 +CsI2X− (3.4)where SI2 is the local maximum solubility of iodine and CsI2 ,CsI−3, andCsI2X− are the amount of iodine presentas I2, I−3 (Reaction 1), and I2X− (Reaction 4), respectively, at the electrode surface [101].The driving force for iodide oxidation is the extent of unsaturation of iodine (SI2 - CbI2), and the criticalcurrent where this point is reached also depends on the concentration of iodide [102]. The actual steady-statecurrent iss for iodide oxidation is lower than the theoretical value of il calculated from Equation 3.2, withiss being reported as 40-60 % of il depending on the iodide concentration [101]. The film at the electrodesurface can physically block further oxidation of iodide, decrease the electrical conductivity at the electrodesurface, and increases the charge-transfer resistance of the iodide redox reactions [103]. This precipitatedI2 can still complex with the excess I− available in the bulk electrolyte, but the presence of the iodine filmis undesireable because it limits the rate of the entire oxidation process. Early studies on the iodine filmdescribed it according to a dissolution-precipitation model where the electrochemically generated iodine atthe electrode surface dissolves until reaching the local saturation point at the interface, which is followedby precipitation of solid iodine and establishment of equilibrium [104]. The iodine film is believed toconsist of a thick layer directly on the electrode surface along with a porous over-layer that extends into theelectrolyte [104, 105]. The diffusion of iodide through the porous over-layer occurs akin to the Grotthus(proton-hopping) mechanism seen in proton exchange membrane fuel cells, where analogously iodide ions393.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Filmcan ‘hop’ from one I2 molecule to another in the film by forming and breaking triiodide bonds akin to H-bonds [106]. After iodide travels through the film and reaches the electrode surface, it can be oxidized toiodine and high order polyiodides (I−5 and I−7 ) can form according toI−(f)+nI2( f ) ⇀↽ I−2n+1( f ) (R21)where subscript f indicates the species are within the iodine film. Fractional analysis at high concentrationsof ZnI2 on a platinum ultramicroelectrode (UME) showed that I−7 is the main species of I−2n+1( f ) present atthe UME-film interface where the concentration of iodine is high, but only I−3 is stable in the bulk electrolyte[32]. At the interface of the film and the bulk electrolyte (Figure 1.4), iodine can be reduced back to iodidevia a galvanic exchange reaction with iodide by [107]:I2( f )+3I-(aq) ⇀↽ 2I−(f)+ I−3(aq) · (R22)This can facilitate the oxidation reactions at the electrode by removing iodine and replenishing iodide, butwas stated to only significantly affect the electrode reactions at thin film thicknesses.Based on Equation 3.4, there are three logical approaches to alleviating the iodine film: (1) decrease theamount of free insoluble iodine by complexing it with something other than I−, 2) increasing the solubility ofiodine in the solvent by modifying its chemical properties through the addition of non-aqueous cosolvent, or3) increasing the operating temperature to increase the saturation point of I2. This chapter will explore theseeffects on the iodine film, the electrode and electrolyte-phase reactions, and the ZIFB as a whole. Section 3.2provides an initial look at the general characteristics of the iodine/iodide/triiodide redox reactions, includingthe iodine film, and introduces several of the common electrochemical methods that will be used further insubsequent parts of the dissertation. The effect of temperature is presented in Section 3.3 and the effect ofcharge-transfer species I2X− and non-aqueous cosolvent is described in Section 3.4.403.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Film-20-10010200 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)30 mV/s20 mV/s10 mV/s5 mV/s2 mV/sI2 + 2e- 2I-I3- + 2e- 3I-3I- I3- + 2e-2I- I2 + 2e- Figure 3.2: Cyclic voltammograms for the I2/I− system (electrolyte composition: 0.05 M ZnI2 + 0.5 MNaClO4) at a variety of scan rates on a freshly polished glassy carbon electrode. Boldface arrows indicatescan rate direction.3.2 Electrochemistry of Iodine/Iodide/Triiodide and the Iodine FilmThe cyclic voltammograms at various scan rates in Figure 3.2 provides an initial look at the iodinehalf-cell reactions. During the forward (anodic) scan, a shoulder peak is observed before reaching the peakanodic current density ipa. This anodic shoulder corresponds to the initial oxidation of iodide to triiodide(Reaction 6) when there is abundant I− available at the vicinity of the electrode. The main anodic peakcorresponds to the oxidation of iodide to iodine (Reaction 5) where the local concentration of I− is too lowfor significant I−3 formation (Reaction 1). With I2 now being the primary product, solid iodine eventuallystarts to precipitate from the electrode/film interface once the local concentration has reached saturation.Following the anodic peak, the current density decreases to a steady-state value iss,CV where the rate ofiodide oxidation and the rate at which iodine is solubilized (either as molecular I2 or through formation ofI−3 ) are equal and at equilibrium [108, 109]. At fast scan rates the value of iCV,ss is close to zero because of413.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Filmthe rate at which iodine is dissolved is slower than the rate at which it is being generated at the electrode andthe iodine film significantly impedes further oxidation past ipa. The value of iss,CV increases with decreasingscan rate because the iodine film has more time to dissolve and diffuse away from the electrode, allowingsome iodide oxidation to continue. On the reverse (cathodic) scan, a single reduction peak (ipc) is observedthat is ascribed to the reduction of the iodine formed during the forward scan back to iodide. Since mosttriiodide formed on charge is further oxidized to iodine (Reaction 6), only the reduction peak correspondingto iodine is seen during the reverse scan except at very low scan rates where sufficient time has passed forthe iodine film to partially dissolve and allow triiodide to reach the electrode surface, which leads to theformation of a second peak (see 5 and 2 mV s−1 in Figure 3.2). The ratio of ipa/ipc at 30 mV s−1 is closeto unity (ipa/ipc = 0.9) which indicates a nearly reversible system but becomes increasingly irreversible asthe scan rate decreases, evidenced by ipc becoming much lower than ipa. At slower scan rates, the iodinegenerated during charge dissolves and diffuses/settles away from the electrode before it can be reduced onthe reverse scan, decreasing ipc. These CVs verify that the oxidation of iodide occurs through a precipitation-dissolution mechanism, with the value of ipa dependent on the saturation point of iodine, and the value ofipc depending on the amount of iodine/triiodide retained at the electrode surface.Collection experiments were performed on a rotating ring disk electrode (RRDE) where the potentialon the central glassy carbon disk was swept linearly to increasing oxidative potentials while a constantreduction potential (0 V vs. Ag/AgCl [sat’d KCl]) was held at the concentric platinum ring in order toreduce iodine/triiodide produced from the disk (Figure 3.3). The linear sweep voltammograms (LSVs)taken on the RRDE can be divided into five regions based on the value of the disk (iD) and ring (iR) currentdensities:1. iD and iR both increasing; at these low applied anodic overpotentials, iodide is being oxidized totriiodide since there is ample supply of I− at the disk. The highly soluble I−3 is mass transferred to thering and reduced back to iodide.2. iD increasing, iR decreasing; the decrease in iR is due to the onset of iodine film formation where muchof the iodine being produced remains as a solid at the disk electrode and does not reach the ring. Thevalue of iodide iD is still increasing with potential because I− is being oxidized to I2 over I−3 .3. iD decreasing iR increasing; the iodine film at the disk inhibits further oxidation by physically blockingiodide from reaching the electrode surface. Over time, some of the I2 formed at the disk is able to423.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Film-40-200204060200 600 1000 1400Current Density / mAcm-2Potential / mV vs. Ag/AgCl (Sat'd KCl)         Figure 3.3: Rotating ring disk experiments investigating the formation of the iodine film during iodide oxi-dation (electrolyte composition: 0.05 M ZnI2 + 0.5 M NaClO4). The GCE disk was polarized to increasinglyanodic potentials at 50 mV s−1 while the platinum ring was held at a constant reduction potential of 0 V vs.Ag/AgCl (sat’d KCl) at a rotation rate of 1000 rpm.dissolve and complex as I−3 and reach the ring.4. iD increases; after more time, a sufficient amount of I2 in the non-porous region of the film at theelectrode surface has dissolved allowing iodide to reach the surface. Iodide is reported to have ahigh diffusivity through the porous overlayer region of the iodine film because it can ‘hop’ from oneiodine molecule to the next by forming/breaking triiodide bonds [104], analogous to the H+ hoppingGrotthus mechanism for PEM fuel cells.5. steady state iD and iR; the rate of the iodide oxidation reaction(s), complexation of iodine as triiodide,and dissolution of the iodine electrochemically generated at the electrode surface are equal and inequilibrium.The anodic and cathodic polarization behaviour for the iodine electrochemical reactions on a glassycarbon RDE at various rotation rates is displayed in Figure 3.3. On a rotating disk electrode, the theoretical433.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Film0204060300 600 900 1200Current Density / mA cm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)           (a)00.010.020.030.040.050.060.070 2 4 6i pa/ mA cm-2(ω)0.5 / (mV s-1)0.5R2 = 0.9949(b)-25-20-15-10-50-650 -350 -50 250Current Density / mA cm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)            (c)-0.025-0.02-0.015-0.01-0.00500 2 4 6i pc/ mA cm-2(scan rate)0.5 / (mV s-1)0.5R2 = 0.9921(d)Figure 3.4: Rotating disk experiments for iodide oxidation and (a) and iodine/triiodide reduction (c) atvarious rotation rates, with plots of steady state current as a function of ω0.5 with linear regression and R2value for the anodic (b) and cathodic (d) polarization data. Electrolyte composition: 0.05 M ZnI2 + 0.01 MI2 + 0.5 M NaClO4. The disk was swept from the experimental determined open-cell potential to 1000 mVvs. Ag/AgCl [sat’d KCl] for the oxidation branch, and to -450 mV for the cathodic branch.limiting current density can be calculated by the Koutecky-Levich equation:il = 0.62nFD2/3ω1/2ν−1/6k Cb (3.5)where D is the diffusivity (cm2 s−1), ω is the rotation rate of the electrode (rad s−1), νk is the kinematicviscosity (cm2 s−1) and Cb is the concentration of reacting species in the bulk electrolyte (mol cm−3).Equation 3.5 assumes that the rate of the electrode reaction is fast enough that no reactant accumulates atthe surface, hence Cs = 0. Since the oxidation of iodide is limited by the flux of product iodine away from443.2. Electrochemistry of Iodine/Iodide/Triiodide and the Iodine Filmthe electrode, the steady state current density iss at the RDE would be more accurately calculated asiss = 0.62nFD2/3ω1/2ν−1/6k (CsI2 −CbI2). (3.6)This can be visualized in Figure 3.3a where the current density reaches a peak value closer to the value ofil predicted by Equation 3.5 before decreasing to the value of iss governed by Equation 3.6. The impactsof the iodine film on iodide oxidation are observed at high rotation rates as well, demonstrating that forcedconvection alone (such as from the circulating electrolyte of the ZIFB) cannot resolve the issue of theiodine film. The Koutecky Levich equation (both Equation 3.5 applicable to iodine/triiodide reductionand Equation 3.6 that describes iodide oxidation) predicts that a plot of the limiting current density as afunction of ω0.5 should yield a straight line, and this is observed for both the anodic Figure 3.4b and cathodicFigure 3.4d reactions. The Levich equation holds only when mass transfer is limiting the electrode reaction,and so the linear fit observed for iodide oxidation verifies that the process is still mass transfer limited, evenif it is limited by the diffusion of product iodine away from the electrode instead of arrival of iodide to theelectrode.0501001502002500 20 40 60 80Current Density / mA cm-2Time / s0.5 M ZnI20.5 M ZnI2 + 0.05 M I20.5 M ZnI2 + 0.1 M I20.5 M ZnI2 + 0.25 M I20.5 M ZnI20.5 M ZnI2 . I20.5 M ZnI2 . I20.5 M ZnI2 . I2       (a)-80-60-40-2000 50 100 150Current Density / mA cm-2Time / s0.5 M ZnI2 + 0.05 M I20.5 M ZnI2 + 0.1 M I20.5 M ZnI2 + 0.25 M I20.5 M ZnI2 . I20.5 M ZnI2 . I20.5 M ZnI2 . I2       (b)Figure 3.5: Chronoamperometry experiments on a glassy carbon electrode in 0.5 M ZnI2 and increasing I2,with a constant applied electrode potential of 1 V (Ag/AgCl [sat’d KCl]) (a) and -0.5 V (b)Chronoamperometry (CA) experiments were performed with electrolytes of 0.5 M ZnI2 containingincreasing amounts of iodine, using an applied voltage of 1 V (Ag/AgCl [sat’d KCl]) for iodide oxidation(Figure 3.5a) and -0.5 V for iodine/triiodide reduction (Figure 3.5b). These applied voltages lay in the mass-transfer limiting regime for the oxidation and reduction processes (respectively), and so the steady-state453.3. Effects of Temperature on the Zinc Iodide Flow Batterycurrents that result can be completely ascribed to that arising from mass-transfer only. For iodine reduction,this steady state current density increases with increasing I2 concentration which is to be expected becausethe cathodic current at this potential is proportional to the diffusional flux of iodine arriving to the surface(i.e. CbI2 - CsI2). For iodide oxidation, the steady state current decreases with increasing I2 content because itis proportional to the flux of iodine away from the electrode surface (CsI2)-CbI2). This once again illustrates theneed to enhance the removal of I2 from the electrode surface in order to support prolonged iodide oxidation,and various methods for doing so are described in the following Section 3.3 and Section 3.4.3.3 Effects of Temperature on the Zinc Iodide Flow Battery-5051015200 200 400 600 800 1000Current Density / mA cm-2Potential / mV (vs. Ag/AgCl Sat'd KCl) 18 C 10 mV/sec48 C 10 mv/secT = 22 oCT = 5 oCFigure 3.6: Cyclic voltammograms showing the effect of temperature on the iodine electrochemical reac-tions, performed at 22◦C and 50◦C (electrolyte composition: 0.05 M ZnI2 + 0.5 M NaClO4) at a scan rateof 10 mV s−1. Boldface arrows indicate scan rate direction.One way to enhance the low solubility and/or increase the rate of dissolution of product iodine isby increasing the operating temperature, and the iodine electrochemical reactions were studied at 22◦Cand 50◦C. In addition to affecting the solubility the iodine, increasing the operating temperature is alsoexpected to increase the rate of the electrode reactions and electrolyte-phase complexation reactions (i.e.463.3. Effects of Temperature on the Zinc Iodide Flow BatteryReaction 1 and Reaction 4 and the diffusivity of the reacting species. The triiodide complexation reactionis exothermic (∆Ho of -17.0 kJ mol−1 at 22◦C with a Kstab of 723 [29]) and the value of Kstab is expectedto decrease with temperature, and from the Van’t Hoff equation the predicted equilibrium constant at 50◦Cwould be approximately 425, still implying that only small amounts of free iodine will be in the electrolyteat equilibrium. CVs comparing the iodide half-cell at 22◦C and 50◦C are shown in Figure 3.6. At 50◦C, theanodic peak current density ipa had increased significantly which indicates an improvement in the oxidationof iodide due to the higher saturation point of iodine at these conditions. The plateau steady state currentiss,CV following ipa is higher at the elevated temperature (and close to zero at 22◦C), since the higher rate ofdissolution allows more iodide to reach the surface. The value of ipc is lower at the higher temperature dueto the higher rate of dissolution which decreases the amount of iodine available to be reduced during thereverse scan.-20020406080-500 -250 0 250 500 750 1000Current Density / mA cm-2Potential / mVT = 20 CT = 50 CT = 22 oCT = 50 oC o(a)-8-4048-30 -10 10 30Current Density / mA cm-2Overpotential / mV22 C50 CT = 22 oCT = 50 oCT = 22 oCT = 50 o(b)Figure 3.7: RDE showing the effect of temperature on the iodine electrochemical reactions, with LSVs takenat 22◦C and 50◦C at a scan rate of 50 mV s−1 and rotation rate of 1000 rpm (a). A zoom in of the low-fieldregion is shown in (b). (electrolyte composition: 0.05 M ZnI2 + 0.01 M I2 + 0.5 M NaClO4)Polarization data on a rotating disk electrode (1000 rpm) at 22◦C and 50◦C is presented in Figure 3.7.The mass-transfer limiting current in the cathodic region at 50◦C is higher than at 22◦C, which is to be ex-pected since the diffusivity of the reacting species increases with the accompanying decrease in electrolyteviscosity. The kinetics of the electrode reactions are also clearly enhanced at the elevated temperature Fig-ure 3.7b, and with Equation 2.6 the exchange current density i0 increased from 4.0 mA cm−2 at 22◦C to8.8 mA cm−2 at 50◦C, and the charge-transfer resistance Rct decreased from 6.3 to 3.1 mΩ cm2. Moreremarkable is the effect on the iodine film observed in the oxidation sweep, where the mass-transfer limiting473.3. Effects of Temperature on the Zinc Iodide Flow Batterycurrent iss after the transient region is enhanced relative to the lower temperature trial, which from Equa-tion 3.6 implies that the solubility of iodine was increased. The transient region in the anodic LSV prior toiss indicates that the iodine film still ultimately limits the oxidation process, but can nonetheless be enhancedat higher operating temperatures.00.020.040.060 40 80 120Current / ATime / sT = 20 CT = 50 CT = 22 oCT = 50 oC(a) oC oC00.10.20.30.40.50 50 100 150Potential / VTime / sT = 20 CT = 50 CT 22 oCT = 50 oC   }} = 2 oC 0 oC   (b)Figure 3.8: Chronoamperometry at 1 V (Ag/AgCl [sat’d KCl]) with no rotation (a) and chronopotentiom-etry at ±50 mA cm−2 and 1000 rpm (b) showing the effect of temperature on the iodine electrochemicalreactions. Experiments were performed on glassy carbon electrode in 0.5 M ZnI2 at 22◦C and 50◦C.The effect of temperature was also investigated through chronoamperometry (Figure 3.8a) and chronopo-tentiometry (Figure 3.8b). As described in Figure 3.5, the anodic steady-state current at this constant elec-trode potential (1 V vs. Ag/AgCl [sat’d KCl]) is proportional to the rate of diffusion of iodine away from thesurface and is higher at 50◦C because of the enhanced solubility of iodine that leads to a higher CbI2 . In thechronopotentiometry (CP) experiments (Figure 3.8b), the constant applied current density (±50 mA cm−2)causes the rate of reaction at the electrode (anodic or cathodic) to be constant, and the I2/I− concentrationratio at the electrode surface changes accordingly. The difference between the charge and discharge plateauswas 451 mV at 22◦C and reduced to 212 mV at 50◦C, once again due to the increased rates of reaction andmass transfer at higher temperatures.The effect of temperature on the ZIFB flow battery (1 M ZnI2, set-up as described in Section 2.2) isshown in Figure 3.9a/b. Similar to Figure 3.8b, these experiments were performed by applying a constantcurrent density (± 20 and± 40 mA cm−2) at the positive and negative electrodes. The voltage plateau duringcharge is much lower at the higher temperature primarily due to the higher solubility of I2 which decreasesηd associated with the slow diffusion of iodine. The higher operating temperature also increases the rate of483.4. Iodine-Ion and Iodine-Solvent Interactions in the Electrolyte0.91.11.31.50 2000 4000 6000 8000Potential / VTime / s20 C45 C45 ഷC20 ഷC20 mA cm-2(a)  T = 22 oCT = 50 oC0.811.21.41.60 2000 4000 6000 8000Potential / VTime / s20 C45 C45 ഷC20 ഷC40 mA cm-2(b)  T = 22 oCT = 50 oCFigure 3.9: Charge–discharge curves comparing zinc-iodine flow batteries at ±20 mA cm−2geom (a) and ±40mA cm−2geom (b) at 22 ◦C and 50 ◦C where 50 mL of 1 M ZnI2 electrolyte is circulated at 20 mL min−1.the electrode reactions and improves the mass transfer of the electroactive species, and these improvementscan be seen in both the charge and discharge curves. The improvements to the charge behaviour were muchmore significant than improvements to discharge (due to the effect of temperature on the iodine film), andimprovements in performance became more significant at higher operating current densities. At 40 mAcm−2, the overpotentials were reduced by approximately 60 mV and the voltage efficiency increased fromto 60% to 65%.3.4 Iodine-Ion and Iodine-Solvent Interactions in the ElectrolyteAs described in Subsection 1.3.1, molecular iodine tends to form charge-transfer complexes with anyspecies sufficiently nucleophilic to do so according to Reaction 4I2+X− ⇀↽ I2X− ·The electrochemical behaviour of the iodide electrode reactions in the presence of bromide and chloride aregiven in Figure 3.10. At a high scan rate of 50 mV s−1 Figure 3.10a, the peak cathodic current density ipcis much higher in the presence of Br− and Cl− due to the formation of the I2Br− and I2Br− complexes. Atslower scan rates of 10 mV s−1 (Figure 3.10b and c), the the value of ipc is significantly higher for the I2Br−493.4. Iodine-Ion and Iodine-Solvent Interactions in the Electrolyte-40-200200 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)ZnI2ZnI2 + ZnBr2 a) 50 mV/secZnI2 + ZnCl2 -15-55150 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)b) 10 mV/secZnI2ZnI2 + ZnBr2 ZnI2 + ZnCl2 -10-505100 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)5 mV/secc)ZnI2ZnI2 + ZnBr2 ZnI2 + ZnCl2 Figure 3.10: Cyclic voltammograms comparing I−3 , I2Br−, and I2Cl− electrochemical reactions at 50 mVs−1(a), 10 mV s−1(b), and 5 mV s−1(c). Electrolytes contained 0.05 M ZnI2 + 0.5 M NaClO4, red CVs alsocontained 0.05 M ZnBr2 and blue CVs also contained 0.05 M ZnCl2. Boldface arrows indicate scan ratedirection.system, but the value of ipc for the I2Cl− is only slightly higher than for the electrolyte only containingiodide. These differences are due to the more favourable interaction between iodine and bromide than withchloride. Cyclic voltammograms (at 10 mV s−1) for iodide electrolytes containing varying amounts ofBr−, Cl−, NO−3 , ClO−4 , and SO2−4 are given in Figure 3.11. All five of these species were not significantlyelectrochemically active in this potential range, evidenced by the CVs containing just the X− species withoutI− (brown CVs in Figure 3.11a-e) but their addition to the iodide electrolyte had significant effects onthe I2/I− and I2/I− reactions. In Figure 3.11a a second oxidation peak is observed after that of iodideoxidation, and is due to the oxidation of bromide to bromine, since the standard electrode potentials ofthe I2/I− and Br2/Br− are close (0.54 and 1.09 V vs. SHE, respectively). For the electrolytes containingBr−, only a small amount of bromide (2:1 [I]− : [Br]−) was needed to observe a significant increase in the503.4. Iodine-Ion and Iodine-Solvent Interactions in the Electrolyte-15-10-50510150 200 400 600 800 1000 1200Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)a) 0.05 M ZnI20.1 M ZnBr20.05 M ZnI2 0.025 M ZnBr20.05 M ZnI2 0.05 M ZnBr20.05 M ZnI2 0.1 M ZnBr2-20-10010200 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)b)0.05 M ZnI20.1 M ZnCl20.05 M ZnI2 0.05 M ZnCl20.05 M ZnI2 0.25 M ZnCl2-10-50510150 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)c)0.1 M Zn(NO3)20.05 M ZnI2 0.05 M Zn(NO3)20.05 M ZnI2 0.5 M Zn(NO3)20.05 M ZnI2-15-10-50510150 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)e) 0.05 M ZnI20.1 M ZnSO40.05 M ZnI2 0.05 M ZnSO40.05 M ZnI2 0.5 M ZnSO4-10-50510150 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)d) 0.05 M ZnI20.1 M NaClO40.05 M ZnI2 0.05 M NaClO40.05 M ZnI2 0.5 M NaClO4Figure 3.11: Cyclic voltammograms comparing I−3 (a), I2Br−(b), I2Cl−(c), I2NO−3 (d), and I2(SO−4 )2(e)electrochemical reactions at 10 mV s−1. Electrolytes contained 0.05 M ZnI2 and varying amounts of ZnBr2,ZnCl2, Zn(NO3)2, ZnSO4, and NaClO4. Boldface arrows indicate scan rate direction.cathodic current density ipc, which is ascribed to formation of the I2Br−. For the chloride CVs, significantchanges in the value of ipc were not observed until a large amount of Cl− was added (1:5 [I]− : [Cl]− ratio).513.4. Iodine-Ion and Iodine-Solvent Interactions in the ElectrolyteFor the NO−3 , ClO−4 , and SO2−4 CVs, the increase in ipc was not observed until large amounts of X− wasadded ([I]− : [X ]− of 1:10), and the increases in ipc were not as significant as was observed for the I2Br−system. These differences ascribed to the stronger interaction between iodine and Br− compared to the otheranionic species, and accordingly the equilibrium constant Kstab of Reaction 4 being higher for I2Br− thanthe analogous I2X− species. The SO2−4 species appeared to interact more strongly with iodine than NO−3 andClO−4 when at equal concentrations, but the high concentrations of X− needed could impair the solubility ofZnI2 which is one of the most desirable traits of the ZIFB.One recent study described a ZIFB containing potassium thioscyanate (KSCN) in order to improvebattery performance through formation of the I2SCN− [110], using an electrolyte containing 1 M KI + 3M KSCN. Using such a large concentration of KSCN to favour formation of I2SCN ultimately limits theconcentrations of I− in the flow battery electrolyte. Furthermore, the standard electrode potential of thereduction of thiocyanogen(SCN)2+2e− ⇀↽ 2SCN− 0.77VSHE (R23)lays inbetween that of I2 and Br2, which could lead to decreases in the coulombic efficiency. Using a non-aqueous cosolvent to alleviate the build-up of iodine on charge could potentially enhance the solubility anddissolution of iodine without having to use large amounts of other salts that decrease the concentrations ofiodide for the ZIFB. Instead of using complexing species to modify the solubility of iodine, we used the non-aqueous organic cosolvent acetonitrile (ACN). Adding 5 vol% to the primarily aqueous electrolyte led tosignificant changes in the electrochemical behaviour, as seen in Figure 3.12. Similar to the X− additives, thevalue of ipc increases with the addition of ACN which indicates an increase in the rate of the iodine and/ortriiodide reduction reactions. Increasing the amount of acetonitrile from 5 to 10 vol% further increased thevalue of ipc. Since the electrochemically generated iodine is more soluble in these dipolar solvents, theiodine film is lessened to some extent. The oxidation peak appears to be partially resolved into two peaks,due to the increased stability (Kstab of Reaction 1) of the I−3 species with addition of cosolvent.The charge-discharge performance of the ZIFB with increasing vol% of acetonitrile is given in Fig-ure 3.13. The voltage efficiency is significantly improved with the addition of 5 vol% ACN, increasing from59.4% to 70.0%. The ZIFB containing 10 vol% ACN performed worse than that containing 5 vol%, butwas still an improvement relevant to the ZIFB with no cosolvent at all. The viscosity of acetonitrile is much523.5. Summary and Conclusions-20-15-10-50510150 200 400 600 800 1000Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)ZnI2ZnI2 5 vol% ACNZnI2 10 vol% ACNFigure 3.12: Cyclic voltammograms showing electrochemical behaviour of the iodide electrolyte with in-creasing amount of acetonitrile (b). Electrolyte contained 0.05 M ZnI2 + 0.5 M NaClO4 with either 5 to 10vol% of acetonitrile (ACN).lower than water (0.341 and 0.890 cP, respectively) but the conductivity of acetonitrile is approximatelytwo orders of magnitude lower than water (6 × 10−10 and 6 × 10−8, respectively) [111] which increasesthe internal resistance of the battery. Therefore, there is a compromise between the improved solubility ofiodine/dissolution of the iodine film and the internal resistance of the cell.3.5 Summary and ConclusionsIn this chapter, the influence of operating temperature, formation of I2X−, and addition of the organiccosolvent acetonitrile on the solubility of iodine and the iodide half-cell redox reactions was investigated.Increasing the operating temperature is able to increase the maximum solubility of iodine as well as enhancethe rate of the oxidation and reduction reactions and diffusivity of the reacting species. Despite the highertemperature also increasing the rate of iodine generation and decreasing the stability of I−3 , the impacton the solubility of iodine was more significant and overall the impacts of the iodine film at 50◦C were533.5. Summary and Conclusions0.811.21.41.60 2000 4000 6000 8000Potential / VTime / s40 mAcm2 nocosolvent40 mAcm2 5volACN40 mAcm2 10volACNNo cosolve t5 vol% ACN10 vol% ACN40 mA cm-2    VE = 59.4 % cosolventVE = 70.0 %5 vol% ACNVE = 63.5 % vol% ACNFigure 3.13: Zinc-iodide flow battery charge-discharge tests with increasing vol% of acetonitrile where 50mL of 1 M ZnI2 electrolyte containing varying amounts of acetonitrile (ACN) is circulated at 20 mL min−1,performed galvanostatically at ± 40 mA cm−2.significantly diminished relative to operating at 22◦C. The addition of small amounts 5 vol% ACN was ableto significantly improve the energy efficiency of the ZIFB. These experiments with acetonitrile in the ZIFBelectrolyte serve as a basis for the studies described in Chapter 4 where additional cosolvents are screened forthis purpose and a more detailed investigation of the interactions between iodine and these various organiccosolvents is provided.54Chapter 4Organic Cosolvents for the Zinc IodineFlow Battery4.1 Iodine Electrochemistry in Non-Aqueous SolventsWater is the most widely used solvent for flow batteries because of its availability, non-toxicity, andhigh relative permittivity (εr = 78) that allows it to dissolve a variety of salts when preparing the batteryelectrolyte. In Chapter 3, it was shown that the iodine film is a result of the low solubility of I2 in water(Equation 3.3):iss = nFDI−δ[CBI−−CSI− ] = nFDI2δ[CsI2 −CBI2 ] = nFDI2δ[SI2 −C∗I2 ]and the addition of organic cosolvents to the primarily aqueous electrolyte could enhance the solubility ofiodine SI2 , thereby increasing the rate of the oxidation reactions. Research on RFBs with a non-aqueoussolvent has been a growing field, and a variety of organic compounds and room temperature ionic liquids(RTILs) has been studied as prospective solvents for the electrolytes [112]. Besides the versatility offered bythe different combinations of solvent and redox species, advantages non-aqueous solvents offer comparedto water for flow batteries include wider windows of electrochemical stability and potentially higher energydensities than what is achievable with an aqueous solvent [113, 114]. Although the solubility of I2 is muchhigher in non-aqueous organic solvents than in water, it is important to emphasize that the main advantageof the ZIFB is the theoretically high energy densities that are a direct result of the high solubility of iodidesalts (8-12 mol L−1). Therefore, it is not feasible to completely move away from aqueous solvents, but aprimarily aqueous solvent with an organic cosolvent could theoretically improve the solubility of I2 in theflow battery electrolyte without compromising the solubility of I− or negatively impacting the performance554.1. Iodine Electrochemistry in Non-Aqueous Solventsof the ZIFB. The electrochemical behaviour of the iodide/iodine/triiodide redox species has been extensivelystudied in a variety of non-aqueous solvents for almost a century. In 1991 a ground-breaking paper (thatwould lead to the inventor Michael Gratzel receiving the Millenium Technology Prize) was published onthe dye sensitized solar cell (DSSC) [115], which included a cathode consisting of the I−3 /I− redox coupledissolved in acetonitrile with a platinum electrode. Even decades later the I−3 /I− redox species in acetonitrileremains the gold standard for DSSCs today [116, 117] and is the primary reason why there has been a surgein research on the iodine half-cell in non-aqueous media.The electrochemistry of iodine has been studied in various organic solvents such as acetonitrile [118,119], DMSO [120], and propylene carbonate [121] and also in several room temperature ionic liquids(RTILs) [122, 123]. Voltammetric studies for the oxidation of iodide to triiodide (Reaction 6) and to io-dine (Reaction 5) in non-aqueous solvents give rise to two well-resolved separate steps, compared to theapparent single step observed in aqueous media. The Nernst equation for the I2/I− and I−3 /I− redox reac-tions are given byE = E0(I−3 /I−)+RT2Fln([I−3 ][I−]3) (4.1)E = E0(I2/I−)+RT2Fln([I2][I−]2) (4.2)and subtracting Equation 4.1 from Equation 4.2, a relationship between the separation in E0 for the I2/I−and I−3 /I− redox pairs and Kstab for triiodide formation (Reaction 1) is obtained:E0(I2/I−)−E0(I−3 /I−) =RT2Fln(Kstab) =RT2Fln([I−3 ][I2][I−]). (4.3)Equation 4.3 predicts that the higher the stability of the I−3 species, the greater the separation betweenthe E0 of Reaction 5 and Reaction 6. The I−3 ion is more stable in aprotic solvents because of H-bondnetworks that interact favourably with and stabilize I− in protic solvents and changes the position of thetriiodide equilibrium. Differences in the stability of the I−3 species between various aprotic solvents islargely dependent on the basicity of the solvent, as strong interactions between the solvent and Lewis acidiodine will decrease the Kstab for triiodide formation [124, 125]. Although the oxidation of iodide takesplace through these two sequential oxidation steps in aqueous media as well, voltammetric studies are onlyable to resolve a single peak (as seen previously in Figure 3.2). The oxidation of iodide is also reported to564.1. Iodine Electrochemistry in Non-Aqueous Solventsproceed through this one-step oxidation in solvents like methanol, 2-propanol, and acetic acid [126] likelybecause they are weaker bases compared to the solvents where the two-step oxidation is observed.Figure 4.1: Schematic of hybrid aqueous/non-aqueous zinc iodine flow battery.In this chapter a variety of organic solvents are screened as potential cosolvent additives to the pri-marily aqueous ZIFB electrolyte in order to increase the solubility of I2, thereby alleviating the iodine film(Figure 4.1). The solubility of iodine with varying amounts of different cosolvents was investigated withUV-Vis spectrometry, and the impact of the cosolvent on the iodide electrochemical reactions was charac-terized through electrochemical half-cell and ZIFB charge-discharge tests. Prior to discussing the resultsof the study, a more in-depth review of relevant solvent properties for this purpose and preferential solva-tion in mixed solvents is provided in Subsection 4.1.1, and the specific solvent interactions between theiodine/iodide/triiodide species and their effects on solubility are described in Subsection 4.1.2.4.1.1 Solvent and Mixed Solvent PropertiesWhen selecting a (co)solvent for the ZIFB, there are several properties that must be considered includ-ing (but not limited to) the electrochemical window of stability, freezing/boiling temperatures, viscosity,574.1. Iodine Electrochemistry in Non-Aqueous Solventstoxicity, and cost [127, 128]. For the ZIFB specifically, one of the most desirable properties of a solventwould be to readily solubilize I2 and inhibit formation of the iodine film while still being able to achievehigh concentrations of I−. Iodine is a non-polar electrophile whereas iodide is a charged polar nucleophile,and so the former is expected to be more soluble in a non-polar basic solvent while the latter would be moresoluble in a polar and acidic solvent. The relative permittivity εr of a solvent is a measure of its ability toweaken the electrostatic interactions between charges, and solvents with εr values above 15 or 20 are gen-erally classified as polar. It is primarily the very high relative permittivity of water (εr = 78) that makes theI− and I−3 ions so soluble in this media. The polarity of a solvent can also be quantified through its dipolemoment µ , which is the vector sum of charge separation in a molecule. However, the dipole moment canlargely underestimate the polarity of smaller molecules (for example, the dipole moment of water being verylow despite its high εr Table 4.1) because its value depends on the distance between the positive and negativecharges in the molecule [111]. The viscosity of the solvent η is another important property impacting flowbattery performance as the viscosity is inversely proportional to the ionic conductivity of the electrolyte.Higher viscosities increase pumping costs of the flow battery and will also decrease the diffusivities of theelectroactive species. It is for this reason that RTILs were avoided in this study, as these species typicallyhave viscosities that are at least a magnitude of order higher compared to what is typical of molecular sol-vents [129]. A list of solvents used in this study along with their relevant physical properties is given inTable 4.1.Table 4.1: Properties of solvents investigated throughout this study [111].Solvent Abbreviation εrDipole/ DDN ANρ / gcm−3 η / cPWater H2O 78.39 1.85 18 54.8 0.997 0.892-Propanol IPA 19.9 1.66 21.1 33.5 0.781 2.04Acetonitrile ACN 35.9 3.53 14.1 18.9 0.776 0.341N-Methyl-2-pyrrolidinone NMP 32.2 4.09 27.3 13.3 1.026 1.67Dimethylformamide DMF 36.7 3.24 26.6 16 0.944 0.802Dimethyl sulfoxide DMSO 46.5 4.06 29.8 19.3 1.095 1.99Propylene Carbonate PC 64.9 4.94 15.1 18.3 1.195 2.531,2-Dimethoxyethane DME 19.2 1.71 23.9 10.2 0.864 0.455Pyridine PYR 12.9 2.37 33.1 14.2 0.978 0.884Ethylenediamine EN 12.9 1.90 50 9.4 0.893 1.54Mixtures of solvents (containing 2 or more components) have been employed in several different ap-584.1. Iodine Electrochemistry in Non-Aqueous Solventsplications to achieve solvent properties not offered by any neat (i.e. single component) solvent. Addingone solvent to another will cause the new solution to have physical and chemical properties different fromits pure constituent solvents, such as the viscosity and boiling/freezing temperatures. The mixed solventwill also have a different relative permittivity, dipole moment, and modified acid/base properties relative tothe individual components, which in turn affects its ability to solvate different solutes. In theory, adding anon-aqueous cosolvent to the primarily aqueous solvent could result in the enhancement of the solubilityof iodine without having to seriously compromise the solubility of iodide and would also largely avoid thesafety and cost concerns associated with a completely non-aqueous electrolyte. In a mixed solvent, a solutewill interact more strongly with one component than the other resulting in the composition of the solventimmediately surrounding the solute molecule to be different than that of the bulk solvent, a phenomenonreferred to as preferential solvation [130]. That is, the addition of cosolvent to water does not just affect thesolubility of a given solute by modifying bulk properties such as the relative permittivity, but also throughthe strength of the solute-cosolvent interactions (relative to that with water). The strength of these interac-tions between cosolvent and iodine have a significant impact on its solubility in the ZIFB electrolyte, andare detailed in the following Subsection 4.1.2.4.1.2 Interactions Influencing Iodine Species SolubilityThe interactions between a solute and solvent can be divided into electrostatic (i.e. coulombic) andnon-electrostatic interactions. The electrostatic portion of the solvation energy ∆Gel can be estimated fromthe Born equation:∆Gel =−69.4z2r(1− 1εr) (4.4)where z is the ionic charge, r is the ionic radius (nm), and εr is the relative permittivity of the solvent. Thecontribution of ∆Gel to the overall Gibbs energy of solvation is high for the charged I− and I−3 ions, butabsent for the uncharged non-polar iodine species. However, these electrostatic interactions still influencethe amount of free iodine in solution by influencing the solvation of iodide and its stability relative to I−3which in turn affects the position the triiodide equilibriumI2(s)+ I− ⇀↽ I3− ·594.1. Iodine Electrochemistry in Non-Aqueous SolventsThe value of Kstab for triiodide formation (at 22 ◦C) is several orders of magnitude higher in non-aqueoussolvents such as dimethylformamide (107.35) and dimethyl sulfoxide (104.9) than that in water (102.86). Thevalue Kstab is lower in water because of the high relative permittivity of the solvent which strongly stabilizesI− in solution, but the less favourable interaction with I− in solvents of lower relative permittivity shifts theequilibrium further towards I−3 which also removes free iodine I2 from solution.Non-electrostatic interactions include the interactions between the occupied and unoccupied orbitalsof the iodine and solvent species. Being a fairly strong Lewis acid electrophile, molecular iodine interactsfavourably with nucleophilic Lewis bases that donate electron density to the empty LUMO of I2. Iodine canbe further classified as a soft Lewis acid, and would therefore be solvated more favourably by Lewis basesthat are also soft in accordance to HSAB (hard soft acid base) theory. The hardness H of a chemical speciescan be described as how easily its number of electrons can be changed, approximated byH =12(IE−EA) = 12(ELUMO−EHOMO) (4.5)where IE is the ionization energy (ease of removing an electron) of a species and EA is its electron affinity(tendency to gain an electron). The decrease in energy resulting from the transfer of electron density from adonor species to an acceptor species can be estimated by [127]∆E =(χA−χB)24(HA+HB)(4.6)where χ is the electronegativity of the solute/solvent species and H is their hardness. From Equation 4.6, theinteraction between solvent and solute is driven by their difference in electronegativity and inhibited by thesum of their hardness. If the solute-solvent pair are both soft (i.e. HA + HB in denominator of Equation 4.6is small), the interaction will be most favourable.To quantify the Lewis acid or base strength of a solvent, the acceptor number (AN) (obtained bymeasuring the P-NMR chemical shift of triethylphosphine oxide in a given solvent) and donor number(DN) (calculated calorimetrically as the standard enthalpy change for the 1:1 adduct formation betweena given solvent and antimony pentachloride) are commonly used, and values of AN and DN for varioussolvents is given in Table 4.1). A high value of DN indicates that the species is a strong Lewis base, and a604.1. Iodine Electrochemistry in Non-Aqueous Solventshigh AN signifies a strong Lewis acid. An issue with using these values of AN and DN as a sole indicatorof acidity/basicity is that they are measured with respect to a common reference (triethylphosphine oxideand antimony pentachloride) when it is well known that the strength of interaction between an acid-basepair depends on the individual identity of both partners, and so the acidity/basicity of a species cannot bedefinitively represented by a single parameter (i.e. DN or AN). One of the most widely accepted two-parameter models for describing the strength of interaction between Lewis acid and base is the EC (orECW) model [131, 132], which expresses the enthalpy change associated with interaction between soluteand solvent as:−∆H = EAEB+CACB+W. (4.7)The E parameter represents the tendency of an acid/base to form electrostatic interactions (highest for hardchemical species) while the C parameter represents the tendency to form covalent interactions (highest forsoft species), and the A/B subscripts denote the value for the acid or base. The W parameter represents anyother energy changes associated with the formation of charge-transfer adduct between acid and base (suchas that arising from the cleavage of a dimeric acid or base) but is taken as 0 in most cases, including allsolvents screened in this study.To visualize the strength of interaction between iodine and various basic solvents, Equation 4.7 (withW = 0) can be rearranged to:−∆HCA+EA=CB+EB2+CB−EB2CA−EAEA+CA(4.8)and the value of −∆HCA+EA can be plotted as a function ofCA−EAEA+CA(commonly designated by RA) (Figure 4.2).The ordinate represents the strength of interaction between an acid and base while the abscissa expresses thenature of the interaction, with values ranging from -1 to +1 (-1 signifying a purely electrostatic interactionand +1 signifying a purely covalent interaction). The value of RA for iodine is close to 0.6 (labeled witha dashed line Figure 4.2) which indicates that iodine is a relatively soft species, and the values of −∆HCA+EAat this value of RA shows the strength of interaction between iodine and these solvents. The solubility ofiodine is highest in the amine solvents (pyridine, diethylamine) because of the basic and soft lone pair ofelectrons on their nitrogen atoms. It is interesting to note that the magnitude of −∆HCA+EA appears to be thelowest in acetonitrile compared to the other solvents, despite this being a common organic solvent for other614.2. Iodine Solubility & Spectroscopy in Non-Aqueous Cosolvents13579-1.4 -1 -0.6 -0.2 0.2 0.6 1 1.4െοࡴ࡯࡭൅ࡱ࡭ࡾ࡭                           RA for I2 = 0.6       Figure 4.2: EC-model graphical representation of the strength of interaction between Lewis bases and acids.The value of RA for iodine is 0.6, and the value of −∆HCA+EA for each solvent at this value quantifies theirstrength of interaction between iodine.applications involving iodine such as in DSSCs. There is not EC data available for every solvent that wasused in this study, but the trends in the strength of interaction between solvent and iodine predicted fromFigure 4.2 for the selected solvents is largely similar to that from the values of solvent donor number.Another non-electrostatic factor influencing the solvation of a solute are those related to the breakingand reforming of intermolecular solvent bonds that occur when a solute is introduced. The solvent net-work must create a cavity to accommodate the solute, and strong interactions between neighboring solventmolecules and/or weak interactions with the solute in the final solution inhibit this process. For water inparticular, the large network of hydrogen bonding is the primary reason why the non-polar iodine species ishighly insoluble in water despite water being a strong Lewis base.624.2. Iodine Solubility & Spectroscopy in Non-Aqueous Cosolvents00.20.40.60 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% DMF7.5 vol% DMFI2 10vol%DMF x5010 vol% D F00.20.40.60 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% NMP7.5 vol% NMP10 vol% NMP0123450 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% DMSO7.5 vol% DMSO10 vol% DMSO0123450 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% DME7.5 vol% DME10 vol% DMEa) b)c) d)00.511.520 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% Pyridine0123450 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% Ethylenediaminee) f)Figure 4.3: UV-Vis spectra for aqueous samples containing 0.1 M I2 and varying amounts of cosolvent:dimethylformamide (a), N-methyl-2-pyrrolidone (b), dimethyl sulfoxide (c), dimethoxyethane (d), pyridine(e), and diethylamine (f). NMP and DMF samples are diluted by factor 20x and DME samples are dilutedby factor 50x.4.2 Iodine Solubility & Spectroscopy in Non-Aqueous CosolventsThe solubility of iodine (in the absence of any added iodide salt) was studied in a variety of organiccosolvents in the amount of 5-10 vol% (Figure 4.3 and Figure 4.4). In the neat aqueous solvent, iodine givesrise to a peak at approximately 450 nm, while iodide/triiodide appears as two peaks at 250 and 380 nm.634.2. Iodine Solubility & Spectroscopy in Non-Aqueous Cosolvents00.511.522.530 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% ACN7.5 vol% ACN10 vol% ACNa)0123450 200 400 600 800Absorbance / AUWavelength / nmNo Cosolvent5 vol% PC7.5 vol% PC10 vol% PCb)Figure 4.4: UV-Vis spectra for aqueous samples containing 0.1 M I2 and varying amounts of cosolvent:acetonitrile (a), propylene carbonate (b).The solubility of iodine in the presence of different cosolvents evidenced by the intensity of the peaks in theUV-Vis spectra followed DME > DMF > NMP > IPA > DMSO > PC > ACN > H2O> EN > PYR (withH2O indicating purely aqueous with no cosolvent). The cosolvents DME, DMF, NMP, and DMSO had highdonor numbers (24-30, Table 4.1) and are also predicted to form relatively strong adducts with I2 based onthe EC analysis (Figure 4.2) and so the improvement in iodine solubility with their inclusion can be largelyattributed to stronger acid-base interactions relative to water. In the presence of DME cosolvent the peakcorresponding to I2 completely disappears, and this is also observed to a lesser extent in the spectra of iodinesolutions with DMF and NMP and attributed to the iodine/triiodide equilibrium shifting further towardstriiodide from the destabilizing of I− in solution even with only 5 vol% cosolvent. Pyridine and diethylaminewere the strongest Lewis bases investigated in this study (Table 4.1 and Figure 4.2), but the spectra takenwith these cosolvents are absent of any I2 peak (Figure 4.3e and Figure 4.3f) indicating little to no free I2 isdissolved in solution. These observations arise from the strongly protophilic cosolvents interacting with I2that stabilizes it as a solid out of solution, indicating that too strong of an acid-base interaction can actuallybe detrimental in the ZIFB electrolyte. These solutions containing 5 vol% pyridine and diethylamine werealso visibly absent of the deep red-black solution characteristic of dissolved iodine/triiodide in polar solventsand further verifies that cosolvents that are too basic are not suited for this application.Acetonitrile (Figure 4.4a) and propylene carbonate (Figure 4.4b) were also able to increase the solu-bility of iodine, but these solvents are protophobic and less basic than water and thus this improvement insolubility cannot be attributed to acid-base interactions. The cosolvents can also enhance the solubility of644.2. Iodine Solubility & Spectroscopy in Non-Aqueous Cosolventsiodine by disrupting the strong hydrogen bond network of the aqueous solvent structure, which reduces theamount of energy required to form the cavity in the solvent needed for accommodating I2. While cosolventscontaining lone pairs on nitrogen atoms can also form H-bonds, nitrogen can only form two H-bonds com-pared to oxygen which contains two lone pairs of electrons and can form up to four H-bonds. Furthermore,there is little steric hindrance impeding the H-bonding network between water molecules compared to thelarger cosolvents, and this further weakens H-bonding capabilities relative to pure water.The solubility of I2 and I−3 with cosolvent was also studied in the presence of I− (through addition ofKI or ZnI2, each solution containing an [I2]:[I−] ratio of 2:1) the results of which are shown in Figure 4.5.With iodide in the electrolyte, the peaks corresponding to iodine vanish due to all dissolved iodine beingconverted to triiodide. For the electrolytes containing KI, all solutions prepared with cosolvent increased theintensity of the I−3 peaks compared to the neat aqueous solution due to the increase in solubility of I2 (whichis subsequently converted to I−3 ), with the observed trend in solubility (at 5 vol%) being NMP > DMF> DME > PC > DMSO = ACN > IPA > H2O. The solubility of iodine with the addition of ZnI2 wasmuch different than the solutions that contained KI, with the solubility of I2 in the ZnI2 solutions at 5 vol%cosolvent following PC > ACN > IPA > DME > H2O > DMSO > DMF > NMP where the solubilityof I2 actually decreased with the addition of 5 vol% DMSO, NMP, and DMF (Figure 4.6) despite thesecosolvents previously improving the solubility of I2 without any iodide salt and when in the presence of KI.Since these solutions all contain the same concentration of I−, any differences in solubility are due to theidentity of the counter cation (i.e. K+ and Zn2+). In aqueous solution, the Zn2+ cation is typically solvatedas the hexahydrate aqua-cation [Zn(H2O)6]2+, but the water ligands in the primary solvation sphere canbe displaced by competing cosolvent molecules (and potentially I−) depending on their concentration andstrength of interaction with the Zn2+ cation. NMP, DMF, and DMSO have higher donor numbers and dipolemoments than water (Table 4.1), and thus able to interact more strongly with the Zn2+ cation than watermolecules. These strong protophilic abilities of these three cosolvents can also allow them to form strongsolvent interactions with H2O molecules in the secondary and ternary solvation shells of the Zn2+ cationby accepting H-bonds. Similar to I2, the Zn2+ species is a Lewis acid, and this can lead to competitionfor the nucleophilic cosolvents and explain in part why these cosolvents decrease the solubility of I2 inthe presence of Zn2+. The spectra containing no cosolvent (Figure 4.5a), acetonitrile (Figure 4.5b) and 2-propanol (Figure 4.5h) remain largely unchanged when comparing the KI and ZnI2 electrolytes. Acetonitrile654.2. Iodine Solubility & Spectroscopy in Non-Aqueous Cosolvents0123200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI01234200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2KIZnI201234200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI012345200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI012345200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI01234200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI01234200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KI0123200 300 400 500 600Absorbance / AUWavelength / nmKIZnI2ZnI2KIa) b)c) d)e) f)g) h)Figure 4.5: UV-Vis spectra for electrolytes containing 0.1 M I2 + 0.05 M KI or 0.1 M I2 + 0.025 M ZnI2 withno cosolvent (a), ACN (b), DMF (c), NMP (d), PC (e), DME (f), DMSO (g), and IPA (h) . The concentrationof KI is twice as high as ZnI2 to ensure that the amount of I− added to all cells is constant.and 2-propanol are poor Lewis bases, and therefore are not expected to interact too significantly with the zinccation in solution and it is likely this reason why they are able to solubilize I2 to the same extent regardlessof if KI or ZnI2 is added to the electrolyte. Compared to Zn2+, K+ has a lower charge and larger ionicradii (rion,K+ = 138 pm and rion,Zn2+ = 88 pm [133]), resulting in the potassium cation having a lower charge664.3. Non-Aqueous Iodine Electrochemistry0.1 M I20.05 M KI0.1 M I20.025 M ZnI20.1 M I20.05 M KI5 vol% DMF0.1 M I20.025 M ZnI25 vol% DMFFigure 4.6: Photos showing decrease in I2 solubility in the presence of Zn2+. Electrolyte composition fromleft to right was 0.05 M KI + 0.1 M I2, 0.025 M ZnI2 + 0.1 M I2, 0.05 M KI + 0.1 M I2 + 5 vol% DMF, and0.025 M ZnI2 + 0.1 M I2 + 5 vol% DMF.density, being a weaker nucleophile, and not interacting as strongly with protophilic solvents compared tothe zinc cation in solution.4.3 Non-Aqueous Iodine ElectrochemistryThe electrochemical behaviour of the I−3 /I− and I2/I− reactions containing KI or ZnI2 in the presence ofdifferent cosolvents was studied with cyclic voltammetry Figure 4.7. Similar to what was observed in (Fig-ure 4.5), there are significant differences when comparing the electrolytes containing DME/DMF/DMSOwith K+ or Zn2+. These protophilic solvents were observed previously to decrease the solubility of I2 inthe presence of Zn2+ and this decrease in solubility can explain the value of ipc being the highest for theelectrolytes containing these cosolvents with Zn2+, since iodine formed on charge precipitates out as a solidonto the electrode surface which increases the amount reduced on the reverse scan. This deposited iodineappears to increase the resistance at the electrolyte interface, since the value of these CVs have their ipcshifted to more negative potentials increasing the peak separation ∆Ep. This deposited iodine can also ex-674.3. Non-Aqueous Iodine Electrochemistry-50510150 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI waterZnI2 waterKIZnI2-10-50510150 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI CANZnI2 CANKI ZnI2-20-10010200 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI PCZnI2 PCZnI2KI-35-25-15-55150 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI DMFZnI2 DMFZnI2KI-30-20-10010200 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI DMEZnI2 DMEIZnI2-30-20-10010200 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI DMSOZnI2 DMSOKIZnI2-50510150 200 400 600 800 1000Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)KI IPAZnI2 IPAKIZnI2a)b) c)d) e)f) g)Figure 4.7: Cyclic voltammograms for 0.1 M KI + 0.5 M KNO3 and 0.05 M ZnI2 + 0.5 M NaClO4 withoutany cosolvent (a) and with 5 vol % of ACN (b), PC (c), DMF (d), DMSO (e), DME (f), and IPA (g) takenat 10 mV/sec on GCE. The black CV in all figures represents the purely aqueous electrolyte with no addedcosolvent.plain the appearance of a third anodic peak in the Zn2+ CVs that contained DMSO, DMF or DME. Forthe electrolyte containing PC, the adduct that formed on charge (for both K+ and Zn2+ electrolytes) wasa sparingly soluble oil-like complex that settles to the bottom of the electrolyte under its own weight andslowly dissolves with time. This oil-like complex has also been observed by others using a NaI electrolyte684.3. Non-Aqueous Iodine Electrochemistrywhich contained 20 vol % PC [134] and was ascribed to the large dipole moment of PC (µ = 4.94) forminga complex with I2 and Na+ with the hydrophobic portion of the PC molecule oriented inwards and the car-bonyl group and Na+/I−/I−3 inwards. For most of the studied cosolvents, the value of the steady state currentiCV,ss in the anodic portion is much higher than seen for the neat aqueous solvent solvent (where iCV,ss ≈ 0)implying that the iodine film had diminished enough to permit continued oxidation.01002003004000 20 40 60 80Current Density / mAcm-2Time / sAqueous5 vol% ACN5 vol% DMF5 vol% NMP5 vol% PC5 vol% DMSO5 vol% IPA5 vol% DME5 vol% EA5 vol% PYRKI(s) K+ + I-2I- I2 + 2e-0501001502002503000 20 40 60 80Current Density / mAcm-2Time / sAqueous5 vol% ACN5 vol% DMF5 vol% NMP5 vol% PC5 vol% DMSO5 vol% IPA5 vol% DME5 vol% EAZnI2(s) Zn2+ + 2I-2I- I2 + 2e-Figure 4.8: Chronoamperometry at 1 V for 1 M KI (a) and 0.5 M Zn2+ (b) with 5 vol% cosolvent.To further see the effect of the cosolvents on iodide charge behavior, chronoamperometry was per-formed at 1 V (vs. Ag/AgCl sat’d KCl) for the mixed solvents containing either 1 M KI or 0.5 M ZnI2,displayed in Figure 4.8a and Figure 4.8b, respectively. For the purely aqueous solvents, the current rapidlydecays to a steady state current; however, for all co-solvents tested except 2-propanol, this behavior is notobserved, and the steady state-current is quickly achieved without this transient. The magnitude of thesteady-state currents observed are directly proportional to the size of the I2 concentration gradient, indicat-ing that all the cosolvents tested except for 2-propanol and pyridine are in fact able to increase the diffusionof I2 away from the electrode surface by increasing the surface concentration of I2, with trends correlatingto the strength of the I2-cosolvent interaction. The CAs taken with 2-propanol are identical to that of thepurely aqueous solvent for both the K+ and Zn2+ electrolyte, which can be explained by the minimal inter-action between IPA and the I2 formed at the electrode surface. On the other hand, the addition of pyridineto the solvent causes the steady state current to immediately decay to 0 when the voltage is applied, whichonce again exemplifies that too strong of an interaction between I2 and cosolvent is detrimental to chargeperformance.694.3. Non-Aqueous Iodine Electrochemistry-40-200204060800 300 600 900 1200Current Density / mAcm-2Potential / mV vs. Ag/AgCl (Sat'd KCl)K+ aqK+ CANK+ DMFK+ PCK+ IPA       No cosolvent5 vol% ACN5 vol% DMF5 vol% PC5 vol% IPA(a) KI-40-200204060800 300 600 900 1200Current Density / mAcm-2Potential / mV vs. Ag/AgCl (Sat'd KCl)Zn2+ aqZn2+ CANZn2+ DMFZn2+ PCZn2+ IPA       No cosolvent5 vol% CN5 vol% DMF5 vol% PC5 vol% IPA(b) ZnI2Figure 4.9: Rotating ring disk experiments investigating the effect of cosolvent on iodide oxidation for KI (a)and ZnI2 (b) electrolytes. Electrolytes contained 0.05 M ZnI2 + 0.5 M NaClO4 or 0.1 M KI + 0.5 M KNO3with 5 vol% of acetonitrile (ACN), dimethylformamide (DMF), propylene carbonate (PC) or 2-propanol(IPA). The GCE disk was anodically polarized at 50 mV s−1 while the platinum ring was held at a constantreduction potential of 0 V vs. Ag/AgCl [sat’d KCl] at a rotation rate of 1000 rpm.RRDE experiments were performed to investigate the iodine-cosolvent interactions (Figure 4.9), wheresimilar to Figure 3.3 the disk is linearly swept towards increasingly oxidative potentials while the ring isheld at a constant reductive potential of 0 V. For both the K+ and Zn2+ electrolytes, there is a significantdecrease in the ring current when PC or DMF was added, due to the oxidation products remaining at thedisk electrode. In the Zn2+ electrolyte, the complex that forms between DMF and I2 is largely insoluble andpassivates the disk electrode, and repeat experiments had lower oxidative currents due to the solid depositimpeding oxidation at the disk. For the K+ electrolyte no such solid deposits were observed, and repeatexperiments were reproducible. The acetonitrile cosolvent was able to increase the iss compared to water(but less than DMF and PC), while the 2-propanol cosolvent had the least significant improvements whichis one again attributed to the weaker iodine-solvent adduct formation between the less basic cosolvents.Several of the cosolvents studied have higher viscosities and/or lower conductivities than that of water,and so electrochemical impedance spectroscopy (Figure 4.10) was performed to ellucidate any impacts of thecosolvent on the internal resistance. For both the KI (Figure 4.10a) and ZnI2 (Figure 4.10b) electrolytes, theelectrolyte with no cosolvent was observed to actually have a higher internal resistance than the electrolytescontaining 5 vol% of acetonitrile, dimethylformamide, 2-propanol or propylene carbonate. These decreasesin internal resistance could be in part due to the higher stability of I−3 in the presence of these cosolvents,704.3. Non-Aqueous Iodine Electrochemistry00.20.40.60.80 2 4 6Z'' / ohm ∙ cm2Z' / ohm ∙ cm2KI aqKI CANKI DMFKI IPAKI PCno cosolvent5 vol% ACN5 vol% DMF5 vol% IPA5 vol% PCKI(a)00.20.40.60.80 2 4 6Z'' / ohm ∙ cm2Z' / ohm ∙ cm2ZnI2 aqZnI2 CANZnI2 DMFZnI2 IPAZnI2 PCno cosolvent5 vol% ACN5 vol% DMF5 vol% IPA5 vol% PCZnI2(b)Figure 4.10: Nyquist plots for KI (a) and ZnI2 (b) electrolytes containg different cosolvents. Electrolytescontain 1 M KI or 0.5 M ZnI2 with 0.1 M I2 and various cosolvents.reducing free iodine in solution.The impact of the various cosolvents on the electrochemical behaviour of the zinc redox reactions(Reaction 11)Zn2++2e− ⇀↽ Zn E0 =−0.76VSHEwas also studied Figure 4.11. For the zinc CVs (Figure 4.11a), the value of ipc was nearly identical regard-less of the cosolvent, indicating the rate of zinc deposition was largely unaffected. All of the cosolventsincreased the value of ipa relative to the electrolyte without cosolvent, and this improvement in the rate ofzinc oxidation was highest in the presence of propylene carbonate. Research on zinc deposition and dis-solution in non-aqueous electrolytes has shown that Zn2+ interacts more strongly with oxygen containingcarbonyl functional groups compared to functional groups containing nitrogen [135], and can therefore beexpected to interact very favourably with propylene carbonate that contains a carbonyl group that is alsoflanked by two alkoxy groups. The electrochemical deposition of zinc in the presence of cosolvents wasstudied on the RDE (rotated at 1000 rpm) (Figure 4.11b), where there was not significant change in thecathodic polarization behaviour with or without cosolvent. Therefore it is expected that incorporating theorganic cosolvents into the electrolyte of the ZIFB will not negatively affect the reactions at the negativezinc electrode.714.4. Hybrid Aqueous-Nonaqueous ZIFB-20-10010203040-1.8 -1.3 -0.8 -0.3Current Density / mAcm-2Potential / V (vs. Ag/AgCl sat'd KCl)No Cosolvent5 vol% ACN5 vol% DMF5 vol% NMP5 vol% DMSO5 vol% PC5 vol% IPAZn Zn2+ + 2e-Zn2+ + 2e- Zna)-80-60-40-20020-2 -1.5 -1 -0.5 0Current Density / mAcm-2Potential / V (vs. Ag/AgCl sat'd KCl)No Cosolvent5 vol% ACN5 vol% DMF5 vol% NMP5 vol% DMSO5 vol% PC5 vol% IPAZn2+ + 2e- Znb)Figure 4.11: Cyclic voltammetry (a) and RDE (b) analyses of zinc electrochemistry in the presence ofvarious organic cosolvents. Electrolyte contained 0.05 M ZnI2 and 0.5 M NaClO4 with 5 vol% of cosolvent.RDE polarization data was obtained at 1000 rpm.4.4 Hybrid Aqueous-Nonaqueous ZIFB0.60.811.21.41.61.80 2000 4000 6000 8000Potential / VTime / s40 mAcm2 nocosolvent40 mAcm2 5 volACN40 mAcm2 5 volIPA40 mAcm2 PC40 mAcm2 5volDMFNo cosolvent5 vol% ACN5 vol% IPA5 vol% PC5 vol% DMF40 mA cm-2VE = 70.0 %l ACNVE = 59.4 %No cosolve tVE = 49.9% l DMFVE = 60.1 % l PCVE = 60.8 %l IPA     051015200 5 10 15 20 25 30 35Z'' / ohm ∙ cm2Z' / ohm ∙ cm2No cosolvent5 vol%ACN5 vol%DMF5 vol%IPA5 vol%PC cosolvent5 l N ol% MF ol% IPA vol% PCRct = 5.5 Ω cm2 Rct = 5.8 Ω cm2 Rct = 12 Ω cm2 Rct = 6.7 Ω cm2 Rct = 7.3 Ω cm2 Figure 4.12: ZIFB performance with various different organic solvents, where 50 mL of 1 M ZnI2 electrolytecontaining 5 vol% cosolvent is circulated at 20 mL min−1, performed galvanostatically at ± 40 mA cm−2.Flow battery experiments were performed with a cell containing 1 M ZnI2 in both half-cells with 5 vol%of different cosolvents Figure 4.12. The cell that contained acetonitrile performed the best by far comparedto the purely aqueous flow battery and other cosolvents, increasing the voltage and energy efficiency from59.4 to 70.0 %. Acetonitrile was observed previously to increase the solubility of I2 without forming strongadducts with iodine. The ZIFB that contained IPA or PC were only slightly improved compared to that of theZIFB with no cosolvent, whereas the flow battery containing DMF actually performed significantly worse724.4. Hybrid Aqueous-Nonaqueous ZIFBthan without any cosolvent. The worsening with inclusion of DMF can again be ascribed to the decrease iniodine solubility that results in solid iodine depositing onto and passivating the felt electrodes. Nyquist plotsof the ZIFB containing DMF (Figure 4.12b) also indicate a significant increase in the internal resistanceof the ZIFB with DMF. Previously in Figure 4.7 and Figure 4.7 the addition of IPA was seen to have littleimpact on the iodine redox reactions, and this is likewise reflected in the ZIFB containing IPA only modestlyincreasing the efficiency. While PC was previously shown to significantly affect the iodine redox reactions,it did not improve the performance of the ZIFB too noticeably likely due to the higher resistance of theoil-like complex with iodine that increases the charge-transfer resistance at the electrode and increase theviscosity of the electrolyte.0.60.811.21.41.61.80 2000 4000 6000 8000Potential / VTime / s40 mAcm2 nocosolvent40 mAcm2 5volDMFVE = 59.4 %No cosolventVE = 49.9%5 vol% DMF 40 mA cm-2  1 M ZnI20.60.811.21.41.60 2000 4000 6000 8000Potential / VTime / s2 M KI 40mAcm22 M KI 5volDMFVE = 58.9 %No cosolventVE = 65.4%5 vol% DMF 40 mA cm-2  2 M KIFigure 4.13: ZIFB performance with 5 vol% DMF for an electrolyte of 1 M ZnI2 (a) and 2 M KI (b), where50 mL of electrolyte containing 5 vol% cosolvent is circulated at 20 mL min−1, performed galvanostaticallyat ± 40 mA cm−2. For the ZIFB with 2 M KI in the iodide half-cell, 1 M ZnBr2 was used in the zinchalf-cell.The differences in the electrochemical behaviour of iodine in non-aqueous solvents with Zn2+ are veryinteresting to note, since a majority of the literature regarding the iodine electrochemical reactions in non-aqueous media have used KI, NaI, and similar salts instead of ZnI2 or other acidic metals cations. A ZIFBwas constructed containing 2 M KI of electrolyte on the iodine side posolyte (i.e. without any Zn2+) and 1M ZnBr2 was used for the zinc half cell negolyte (which requires Zn2+ in some form for electrodepositionof zinc on charge). This composition was chosen in order to balance out the amount of charge from eachspecies in both half-cells (2 M K+ and 2 M I− in the posolyte, 1 M Zn2+ and 2 M Br− in the negolyte) andminimize migration of Zn2+ into the iodine half-cell. The performance of this K+ ZIFB with and without 5vol% DMF is given in Figure 4.13 and it can be seen that there is a significant improvement in the voltage734.5. Summary and Conclusions0.40.60.811.21.41.60 200 400 600 800Potential / VTime / min40 mA cm-2(a)2 M KI + 5 vol% DMF0.60.811.21.41.60 1000 2000 3000 4000Potential / VTime / s40 mA cm-2(b)Cycle 1Cycle 152 M KI + 5 vol% DMFFigure 4.14: ZIFB performance with 5 vol% DMF for an electrolyte of 2 M KI over 15 cycles (a) and directlycomparing the 1st and 15th (b), where 50 mL of electrolyte containing 5 vol% cosolvent is circulated at 20mL min−1, performed galvanostatically at ± 40 mA cm−2. 1 M ZnBr2 was used in the zinc half-cell.efficiency with the addition of DMF, in contrast to what was observed in the K+ ZIFB (Figure 4.13a). Theimprovement in charge performance is much more significant than for discharge due to the effect of DMFon the solubility of iodine and thickness of the iodine film. This is a very significant result, since it allowsa wider range of cosolvents to be used for enhancing the solubility of iodine in the ZIFB electrolyte thatwere not possible in the presence of the zinc cation. Remarkably, there was no deterioration in the K+ ZIFBcontaing 5 vol% DMF after 15 charge-discharge cycles (Figure 4.14) and indicates that there was negligibleif any Zn2+ crossover into the iodine half-cell which would have otherwise led to increases in resistancethat were observed for Figure 4.13a. This further highlights the viability of a ZIFB based on K+ or othernon-transition metal salts with cosolvents that can enhance the solubility of iodine.4.5 Summary and ConclusionsThe effect of different organic cosolvents on the solubility of iodine and impact on iodine electro-chemistry and ZIFB performance was extensively studied. Cosolvents that were strongly basic and formedadducts with iodine (DMF, DMSO, NMP) were able to significantly enhance the solubility of iodine relativeto water even at only 5 vol%. Cosolvents that were weakly basic such as acetonitrile and 2-propanol werealso able to enhance the solubility of I2, primarily due to weakning the H-bond network and destabilizingI− relative to I−3 , and were able to do so without negatively impacting the zinc redox reactions in the neg-744.5. Summary and Conclusionsative half-cell. In the presence of Zn2+, the solubility of iodine is significantly reduced upon addition ofprotophilic cosolvents as they strongly interact with the acidic Zn2+ cation and decrease the solubility of I2which can then deposit onto and passivate the electrode surface. Protophobic solvents like acetonitrile donot interact significantly with the Zn2+ and are able to increase the solubility of iodine, partially alleviatethe iodine film and enhance the electrochemical reactions, and a ZIFB containing 5 vol% ACN was able toimprove the voltage efficiency of the ZIFB by 10.6% (at 40 mA cm−2) relative to the purely aqueous cell.A ZIFB using a catholyte containing KI instead of Zn2+ with DMF was also tested and was able to improvethe voltage efficiency from 58.9 % to 65.4 % without any significant crossover of Zn2+ that would lead toprecipitation of solid iodine deposits and no deterioration of performance even after prolonged cycling wasobserved. In conclusion, the addition of small amounts of cosolvent to the aqueous ZIFB is able to achievesignificant improvements in the cycling efficiency and being able to expand the amount of options availablethrough a ZIFB catholyte absent of Zn2+ provides opportunity for future improvements as well.75Chapter 5High Surface Area & Activity CarbonMaterials for the I2/I− Redox Couple in theZIFB5.1 Electrode Materials for the Iodine Redox ReactionsConstructing a ZIFB capable of providing large power densities primarily depends on the rate of theiodine electrode redox reactions Reaction 5 and Reaction 6:I2(s)+2e− ⇀↽ 2I− 0.621VSHEI3−+2e− ⇀↽ 3I− E0 = 0.536VSHEAccording to early studies by Vetter and Newson [136, 137], the iodine redox reactions can be describedaccording to:I2 ⇀↽ Iads+ Iads (Tafel reaction) (R24)Iads+ e− ⇀↽ I− (Volmer reaction) (R25)where the charge-transfer step Reaction 25 is rate determining. When iodide is available, the formation oftriiodide (Reaction 1)I2(s)+ I− ⇀↽ I3−765.1. Electrode Materials for the Iodine Redox Reactionsoccurs rapidly in comparison to the electrode reactions and the degree of reversibility increases as the con-centration ratio of triiodide to iodide is decreased. However, later studies performed by Dane [138, 139]supported a different mechanism, where Reaction 24 is replaced by:I2+ e− ⇀↽ Iads+ I− (Heyrovsky reaction) (R26)These authors stated that Reaction 26 was the rate determining step and that the degree of coverage of I2and I− on the electrode surface can significantly affect the mechanism, and had not been taken into accountin prior kinetic studies.A plethora of research has gone into studying the kinetics of the Br2/Br− redox reactions on numerouselectrode materials for the ZBFB [79], providing a basis for electrode research on the more recently devel-oped ZIFBs. Different studies of the bromine redox reactions on carbon electrodes (in aqueous electrolytes)have agreed that the bromine electrode reactions follow the mechanism:Br2+ e− ⇀↽ Brads+Br− (Heyrovsky reaction) (R27)Brads+ e− ⇀↽ Br− (Volmer reaction) (R28)where Reaction 27 is the rate determining step for bromine evolution and Reaction 28 is rate determiningfor the reduction reaction [77, 140–142]. The sluggish kinetics of the bromide redox reactions (particularlybromine reduction) lead to high activation overpotentials at the bromine electrode, and dozens of electrodematerials have been reported and characterized for the ZBFB [79]. Carbon based electrodes are by farthe most researched for the ZBFB because of their low cost, abundance, and electronic conductivity. Theactivity of the bromide redox reactions can vary greatly from one carbonaceous electrode to another becauseof differences in pore size/volume, surface area, surface-functionalities, and other properties [78, 143]. Somecarbon-based catalysts proposed for the positive electrode of the ZBFB include carbon blacks [80], carbonnanotubes [81, 82], hetero-atom (most commonly nitrogen and boron) doped carbons [83–86], and severalother novel materials [87, 88].Compared to the bromine electrode, there has been very little research on the kinetics of the iodineelectrode reactions in flow batteries and research on improving the performance of the ZIFB has focused775.1. Electrode Materials for the Iodine Redox Reactionsalmost exclusively on the electrolyte composition [45, 94, 96, 97]. The kinetics of the I−3 /I− redox reactionshave been studied extensively for applications in dye-sensitized solar cells (DSSCs) where a I−3 /I− half-cellis commonly implemented as the (photo)cathode. Many of these DSSCs employ transition metal-basedelectrocatalysts for the I−3 /I− [115, 117, 144], which if integrated into the ZIFB would significantly increasethe costs relative to the carbon-based electrodes common for most common RFB systems [21, 22]. DSSCsalso operate in completely non-aqueous electrolytes where the iodine film does not occur, and so there iscomparatively little literature available regarding the iodine electrode in aqueous media. Although funda-mentally different from the zinc-iodine flow battery, there have been various electrode studies for the iodinehalf-cell in non-flow static zinc-iodine batteries (ZISB) [145]. Initial ZISBs used a microporous carbon hostelectrode (such as a carbon felt/cloth) which can encapsulate iodine within its pores, inhibiting formationof I−3 and promoting Reaction 5 which increases the utilization efficiency of I− and the battery capacity[146, 147]. In these static zinc iodine batteries, iodine adsorbs strongly on the pore walls of the carbon hostelectrode and electrodes with sufficiently small pores can prevent the iodine species from diffusing backinto the electrolyte [148, 149]. In aqueous electrolytes where the solubility of molecular iodine is low, theadsorption of iodine on the internal pore walls is generally much more thermodynamically favourable thanits solvation in water which has been exploited to confine the reactions within the pores of the electrode andprevent the formation of the I−3 as an intermediate in ZISBs [146]. Further improvements to the iodine-sideelectrode has included modifying the porous carbon host structure or introducing catalysts onto its surfacethat enhance the adsorption of I2 and/or better catalyze the iodine redox reactions [150]. The most success-ful electrodes implemented in the static zinc-iodine battery have been those with large pore volumes (butsmall pore sizes) and electronic structures that promote the adsorption of iodine over its solvation in the elec-trolyte, such as hetero-atom doped carbons [151–153] and transition metals [154]. Similar to the research onI−3 /I− electrocatalysts in DSSCs, it is difficult to apply a majority the electrode materials proposed for thesesmaller-scale ZISB coin cells to the ZIFB that is intended for grid-scale applications, because a majority ofthem employ expensive metals like platinum or synthesis procedures ill-suited for scale up [144, 155] butnonetheless provide a foundation for the kinetics of the iodine reactions in flow batteries which until nowhas been a largely unexplored topic.The ZIFB is intended for applications in medium/large scale stationary energy storage, and electrodematerials must be inexpensive and abundantly available to be viable for any realistic implementation. In this785.2. Iodine Kinetics on Carbonaceous Materialsstudy, graphite felt electrodes are loaded with different carbon materials including high surface area carbonblacks (Vulcan XC-72R, Ketjenblack EC-300J, Ketjenblack EC-600J) and multiwalled carbon nanotubes(MWCNT) to improve the activity of the iodide/iodine/triiodide redox reactions and efficiency of the ZIFB.Identical electrochemical tests were also performed with the Br2/Br− system in order to directly comparethe rate of the bromine and iodine redox reactions on the various electrode materials studied. To improveutilization of the carbonaceous catalyst materials deposited onto the graphite felt electrodes, the preparedinks containing the dispersed carbon material were introduced onto the felt substrate by either uniformlycoating the entire three-dimensional porous structure with the catalyst ink, or the electrode was preparedby spraying the carbon ink into a concentrated layer on the surface of the electrode adjacent to the batteryseparator. In doing so, the electrochemical reaction sites are concentrated closer to the opposite electrode,reducing the ionic transport distance and decreasing the internal resistance of the cell [143, 156, 157].Thischapter bridges some of the gaps between the abundant research on the Br2/Br− electrode used in ZBFBsand the much less studied I−3 /I− and I2/I− reactions in zinc iodine flow batteries, and demonstrates thatsignificant improvements to the efficiency and power densities of the ZIFB can be achieved through low-cost carbonaceous electrodes.5.2 Iodine Kinetics on Carbonaceous MaterialsTo study the electrochemical behaviour and kinetics of the iodine redox reactions on different carboncatalysts, inks containing the catalyst were prepared as described in Section 2.4 and drop casted onto a glassycarbon electrode and dried at room temperature. Cyclic voltammograms comparing the electrochemicalkinetics of the iodine redox reactions on Vulcan XC-72R (VXC-72R), Ketjenblack EC-300J (Ketjen-300),Ketjenblack EC-600J (Ketjen-600), and multiwalled carbon nanotubes (MWCNT) to the glassy carbon elec-trode (GCE) as a standard are displayed in Figure 5.1. The cathodic peak current density (ipc) is significantlyhigher on all carbon materials studied compared to the bare GCE, indicating a significant improvement inthe reduction of iodine/triiodide. The value of ipc increases from 5 mA cm−2geom (based on the geometric elec-trode area) on the GCE to 20 mA cm−2geom on VXC-72R and 25 mA cm−2geom for the Ketjen-300 and MWCNTelectrodes, which can be explained in part by the increased surface area of the porous electrodes relative tothe planar GCE. When the cathodic current densities are normalized based on the electrochemically active795.2. Iodine Kinetics on Carbonaceous Materials-30-25-20-15-10-50510150 300 600 900 1200Current Density / mA cmPotential / mV vs. Ag/AgCl (sat'd KCl)-2 geom (a)-25-20-15-10-50510150 300 600 900 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)geom (b)-25-20-15-10-50510150 300 600 900 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)geom  (c)-20-15-10-50510150 300 600 900 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl)geom (d)Figure 5.1: Cyclic voltammograms measured in 0.05 M ZnI2 + 0.5 M NaClO4 at 10 mV s−1 on KetjenblackEC-300J (a), Vulcan XC-72R (b), multi-walled carbon nanotubes (c) and Ketjenblack EC-600J (d). Theblue CVs correspond to the dried drop-casted carbon catalysts while the black CVs are on a GCE at identicalconditions.surface area ECSA (Section 2.5), the same trend in cathodic current density is observed between the threeelectrodes indicating that surface area alone cannot account for the improvements observed in the kineticsof the iodine/triiodide reduction reactions. When comparing Ketjen-300 to an even higher surface area car-bon black Ketjen-600 (BET surface areas of 800 m2 g−1 and 1400 m2 g−1, respectively) the value of ipcdecreased, which further suggests that the improvements to the iodine reduction reactions are determinedby other factors in addition to surface area. It has been documented in zinc-iodine static batteries that con-fining molecular iodine inside the pores of the cathode electrode promotes the I2/I− reactions by inhibitingI−3 formation [146, 148], and so it is likely that the increase in cathodic peak observed in the CVs is due inpart to the enhancement of the I2/I− reaction within the electrode pores that occurs along with the triiodideand iodine redox reactions on the outer surface of the electrode.805.2. Iodine Kinetics on Carbonaceous Materials-40-30-20-10010200 400 800 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl) Cycle 1Cycle 10Cycle 50(a) -40-30-20-10010200 400 800 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl) Cycle 1Cycle 10Cycle 50(b) -40-30-20-10010200 400 800 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl) Cycle 1Cycle 10Cycle 50  (c)-40-30-20-10010200 400 800 1200Current Density / mA cm-2Potential / mV vs. Ag/AgCl (sat'd KCl) Cycle 1Cycle 10Cycle 50 (d)Figure 5.2: Cyclic voltammograms measured in 0.05 M ZnI2 + 0.5 M NaClO4 at 20 mV s−1 over 50successive cycles on Ketjenblack EC-300J (a), Vulcan XC-72R (b), multi-walled carbon nanotubes (c) andKetjenblack EC-600J (d) to demonstrate their durability over repeated cycling in the presence of the iodinefilm.The anodic peak current densities ipa (Figure 5.1) remain similar for all the carbon materials comparedto the GCE. Due to the previously discussed iodine film, at sufficiently high anodic potentials the overall rateof iodide oxidation is limited by the slow diffusional flux of I2 (which in turn limits the flux of I−) and notthe rate of the charge transfer reaction, explaining why ipa is relatively insensitive to changes in the electrodematerial. While the value of ipa remained relatively constant on the various electrodes, the onset potentialfor iodide oxidation occurs at more negative potentials on the Ketjen-300 (195 mV vs. Ag/AgCl [sat’d KCl])and VXC-72R (240 mV) electrodes compared to the GCE (320 mV). This indicates that these materials arestill effective at improving the rate of the iodide oxidation to triiodide at lower potentials where I−3 is theprimary oxidation product instead of I2. Figure 5.2 shows the CVs of the various carbon catalysts over 50successive cycles. The cathodic and anodic peak current densities remain essentially identical after repeatedcycling, indicating the stability and high durability of the dried catalyst layers even with the formation of815.2. Iodine Kinetics on Carbonaceous Materialsiodine film.-50-20104070-600 -300 0 300 600 900 1200Current Density / mA cm-2Potential / mV (vs. Ag/AgCl Sat'd KCl) geomI-/I3-       { {I-/I2   Figure 5.3: Linear sweep voltammograms at 1000 rpm comparing the cathodic and anodic polarizationbehaviour on glassy carbon and Ketjen-300. Electrolyte: 0.05 M ZnI2 + 0.01 M I2 + 0.5 M NaClO4.Catalyst loading was 0.2 mg cm−2geom for Ketjen-300 (22 ◦C).The kinetics of iodide redox reactions were also studied on a rotating disk electrode (at 1000 rpm),and a comparison of the polarization behaviour on the glassy carbon and Ketjen-300 electrodes is given inFigure 5.3. During the reduction scan, the cathodic current density on the Ketjen-300 catalyst increases untila peak value but subsequently decreases to its steady-state value, but this transient region prior to reachingthe steady state current density iss is not observed on the glassy carbon. This larger initial peak cathodiccurrent density is ascribed in part to adsorbed I2 within the pores of the electrode being reduced along withthe triiodide/iodine being reduced at the external surface, which decreases to a lower steady state value oncethe iodine confined within the pores of the electrode is reduced and the cathodic reactions occurring onthe external electrode surface predominate. During the anodic scan, the current density on the Ketjen-300electrode is higher than the GCE at all potentials until reaching the anodic iss where they become identical.As described in Section 3.2, the anodic current density reaches a peak value prior to decreasing to iss wherethe dissolution and removal of I2 limits the oxidation reaction. The identical value of iss between GCEand Ketjen-300 is a result of the rate of dissolution of iodine being identical regardless of the electrode,indicating that these porous carbon catalysts can improve the rate of the iodide oxidation reaction but only825.2. Iodine Kinetics on Carbonaceous Materialsup until the diffusion of product iodine away from the electrode becomes rate determining.    Current Density / mA cm-2Overpotential / mVGCE (previous)Ketjenblack EC-300JVulcan XC 72RMulti-Walled CNTsKet600 NEWCEVXC-72Rtj -300MWCNTsKetjen-600geom       Figure 5.4: A comparison of the low-field polarization data on the RDE for GCE and other carbon catalystmaterials. Electrolyte: 0.05 M ZnI2 + 0.01 M I2 + 0.5 M NaClO4. Catalyst loading was 0.2 mg cm−2geom forall drop-casted materials (22 ◦C).Table 5.1: Kinetics parameters of I2/I− on various electrodesGCE Vulcan XC-72 Ketjen-300 Ketjen-600 MWCNTi0 mA cm−2geom 3.1 6.4 9.5 8.1 6.4i0 mA cm−2ECSA 2.9 3.4 5.0 3.9 4.1Rct Ω cm2geom 8.2 4.0 2.7 3.1 4.0Rct Ω cm2ECSA 8.7 6.9 5.2 6.5 6.2The low-field polarization data for the iodide redox reactions on the GCE and other electrode materialsis given in Figure 5.4 along with the exchange current density i0 and charge transfer resistance Rct (pergeometric area and ECSA) calculated from this data in Table 5.1. The low-field approximation of theErdey-Gruz & Volmer (EGV) equation was used (Section 2.6, Equation 2.6 and Equation 2.7) for thesecalculations to avoid errors associated with data at higher anodic overpotentials where the oxidation reactionis limited by diffusion through the iodine film, and thus is not under kinetic control where the high-field Tafelequation would be applicable. The exchange current density increased significantly on all the materialsstudied compared to the GCE (Table 5.1), but to varying extents. The exchange current density was the835.3. Comparing the Kinetics of the Bromine and Iodine Electrodeshighest on Ketjen-300 when calculated from the geometric electrode area (increasing from 3.1 mA cm−2geomon the GCE to 9.5 mA cm−2geom) or the electrochemically active surface area (2.8 mA cm−2ECSA to 5.0 mAcm−2ECSA). Likewise, the charge-transfer resistance (both on a geometric and ECSA basis) was smallest forthe Ketjen-300 electrode. The improvements in the rate of the electrode reactions are likely due to a muchhigher adsorption capacity towards iodine on the different porous catalyst films compared to that of glassycarbon. The low solubility of iodine in water makes its adsorption out of solution more favourable, anddifferences in the values of i0 between the carbonaceous catalysts are likely due to differences in their porenetwork and structure.5.3 Comparing the Kinetics of the Bromine and Iodine Electrodes-15-10-5051015200 400 600 800 1000 1200 1400Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)GCE (10 mV/sec)VXC (10 mV/sec, no bubble)Ket300 (10 mV/sec, no bubble)MWCNT (10 mV/sec nobubbles)Ket600 (10 mV/sec, no bubble)VXC-72RGCEKetjen-300Ketjen-600MWCNTgeomBr2 + 2e- 2Br-2Br- Br2 + 2e- Figure 5.5: Cyclic voltammograms of the Br2/Br− redox couple on various electrode materials: glassycarbon electrode (GCE), Vulcan XC-72R (VXC-72R), Ketjen-300, Ketjen-600, and multi-walled carbonnanotubes (MWCNT) at 10 mV/sec. The electrolyte composition was 0.05 M ZnBr2 + 0.5 M NaClO4. Thecatalyst loading was 0.2 mg cm−2geom. Performed at a scan rate of 10 mV s−1 (22 ◦C).The electrochemical behaviour of the bromine redox reactions (Reaction 14 and Reaction 15)Br2+2e− ⇀↽ 2Br− E0 = 1.09VSHE (R29)845.3. Comparing the Kinetics of the Bromine and Iodine ElectrodesBr−3 +2e− ⇀↽ 3Br− E0 = 1.05VSHE (R30)were also studied on the different carbon catalysts (Figure 5.5). The value of ipc (corresponding to thereduction of bromine and/or tribromide) is significantly higher on all carbon catalysts tested compared tothe GCE while the value of ipa is not affected significantly, which is similar to what was observed for theiodine system at identical conditions (Figure 5.1). However, there is a clear decrease in the peak separation∆Ep for the bromide system on all the catalysts compared to that of the GCE, with ∆Ep decreasing from525 mV on GCE to 239 mV on Ketjen-300. This significant decrease in ∆Ep indicates an improvement inthe reversibility and decrease in the activation polarization of the Br2/Br− redox reactions by facilitating therate determining adsorption and desorption of bromine [80]. This is in contrast to what was observed forthe I2/I− couple, where the peak separation remains relatively constant (255 mV) for all catalysts studieddespite the significant increase in ipc.-30-150153045-200 200 600 1000 1400Current Density / mA cm-2Potential / mV (vs. Ag/AgCl Sat'd KCl) geom           (a)Figure 5.6: Linear sweep voltammograms at 1000 rpm and scan rate 10 mV s−1 comparing the cathodic andanodic polarization behaviour of the bromine/tribromide/bromide redox reactions on glassy carbon (GCE)and Ketjen-300. Electrolyte: 0.05 M ZnBr2 + 0.01 M Br2 + 0.5 M NaClO4. The catalyst loading was 0.2mg cm−2geom. (T = 22 ◦C)The polarization behaviour for the bromine redox reactions on both GCE and Ketjen-300 are presentedin Figure 5.6a. The kinetics of the bromide redox reactions on the glassy carbon electrode appear to be quitelow compared to on Ketjen-300, evidenced by the sigmoidal shape of the LSV curve at low overpotentials855.4. Electrode Impact on Flow Battery Performanceon the GCE and was not observed previously for the iodine cell at identical conditions (Figure 5.3) whichwould imply that the rate of reaction of the iodine redox reactions is higher than that of bromine on the planarglassy carbon. This is further verified by the experimentally determined exchange current densities whereon the GCE the rate of the iodine reactions are two orders of magnitude higher than bromine (Table 5.1 andTable 5.2). The tabulated values of i0 also show that the kinetics of the Br2/Br− and I2/I− redox reactions areboth improved on the Ketjen-300 electrode compared to the glassy carbon, but the improvement observedfor the Br2/Br− redox reactions is much more significant than that was observed for I2/I− indicating that thisredox couple is much more sensitive to the changes in the electrode than the I2/I− system. Although theseporous carbon materials can significantly improve the kinetics of the iodide and bromide redox reactions,the improvement is not as drastic because the kinetics of the I2/I− system on the planar glassy carbon isalready much higher than that for the Br2/Br− system where adsorption and desorption of bromide/brominehas a much higher energy barrier compared to iodide/iodine.Table 5.2: Kinetics parameters of Br2/Br− on various electrodesGCE Ketjen-300i0 mA cm−2geom 0.058 3.4i0 mA cm−2ECSA 0.051 1.9Rct Ω cm2geom 440 7.5Rct Ω cm2ECSA 484 13Cyclic voltammograms directly comparing the iodine and bromine redox reactions on identical elec-trode materials is presented in Figure 5.7. On the glassy carbon and the dried catalyst layers, the value of ipcfor the iodine CVs are much higher than that for bromine, showing that the rate of iodine/triiodide reductionis higher than that of bromine/tribromide. The values of ipa are identical for the I2/I− and Br2/Br− on thesame electrode. Since the anodic behaviour of iodide oxidation is limited by the formation of the iodinefilm, it is likely that bromide oxidation suffers from a similar issue where further oxidation is limited by thediffusion and dissolution of bromine formed at the electrode surface.865.4. Electrode Impact on Flow Battery Performance-10-50510150 200 400 600 800 1000 1200 1400Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)geom I2 + 2e- 2I-Br2 + 2e- 2Br-  (a) -25-15-55150 200 400 600 800 1000 1200 1400Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)geomBr2 + 2e- 2Br-I2 + 2e- 2I-(b) -20-15-10-50510150 200 400 600 800 1000 1200 1400Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)geom (c) I2 + 2e- 2I-Br2 + 2e- 2Br--20-15-10-50510150 200 400 600 800 1000 1200 1400Current Density / mAcm-2Potential / mV (vs. Ag/AgCl Sat'd KCl)geom  (d) I2 + 2e- 2I-Br2 + 2e- 2Br-Figure 5.7: Comparison of the Br2/Br− (red CVs) and I2/I− (violet CVs) redox reactions on a glassy carbonelectrode (a), Ketjen-300 (b), VXC-72R (c) and MWCNTs (d) at 10 mV s−1. The bromine electrolyteconsisted of 0.05 M ZnBr2 + 0.5 M NaClO4 and the iodine electrolyte contained 0.05 M ZnI2 + 0.5 MNaClO4. The catalyst loading was 0.2 mg cm−2geom. (T = 22 ◦C)5.4 Electrode Impact on Flow Battery PerformanceGraphite felts were treated with Ketjen-300 and prepared by either uniformly coating the catalyst layerover the entire 3-D surface of the porous felt, or spraying the same amount of the catalyst material (2 mgcm−2) onto a smaller region of the electrode surface that is adjacent to the battery separator. Scanningelectron microscopy (SEM) images of the untreated and treated electrodes are shown in Figure 5.8. Thefibres of the felt that was entirely coated in the CL (Figure 5.8b) appear to be uniformly covered without anycatalyst material in the spacing between fibers. For the felt that had the carbon material sprayed into a layeron the electrode surface (Figure 5.8c), the fibers are coated in a much thicker layer of catalyst comparedto the soaking treatment, but still without significantly obstructing the intra-electrode space between fibers.The catalyst layer adheres very strongly to the GF fibers, remaining on the surface after 20 charge-dischargecycles operated in flow-through mode (Figure 5.8d) and demonstrating the durability of the treated electrode.875.4. Electrode Impact on Flow Battery Performancea) b)c) d)     Figure 5.8: SEM images of the pristine untreated GF (a), the GF that was uniformly covered in Ketjen-300by immersing in the catalyst ink (b), the GF prior to cycling that had Ketjen-300 concentrated on a layer onits surface (c) and the same electrode as (c) but after 20 charge/discharge cycles (d). All images were takenwith an accelerating voltage of 15 kV. Inlets show working diameter (WD) and magnification of the images.The performance of a zinc-iodine flow battery fabricated from an untreated graphite felt was com-pared to ZIFBs containing felts treated with VXC-72R or Ketjen-300 (uniformly over the entire substrateor concentrated onto the surface) at 20 mA cm−2 (Figure 5.9a) and 40 mA cm−2 (Figure 5.9b). All treatedelectrodes were able to improve the efficiency of the flow battery compared to the untreated felt, and a com-parison of the voltage efficiency (VE), energy efficiency (EE), and power density on discharge (PDdischarge)are summarized in Table 5.3. The felt with the concentrated Ketjen-300 layer on its surface performed thebest, decreasing the separation between charge-discharge plateaus by 103 mV relative to the untreated feltat 20 mA cm−2, and by 225 mV at 40 mA cm−2. At 40 mA cm−2, the energy efficiency of the untreatedZIFB was 63% and increased to 76% with the Ketjen-300 sprayed electrode.The electrochemical impedance spectroscopy (EIS) tests performed after cycling the battery once, andthe resulting Nyquist plots are given in Figure 5.10. The charge transfer resistance on the untreated felt ismuch higher than on the Ketjen-300 treated felts, evidenced by the large-radius semi circle in the Nyquist885.4. Electrode Impact on Flow Battery Performance0.80.911.11.21.31.41.50 2000 4000 6000 8000Potential / VTime / minBlank Graphite FeltKetjenblack (entiresubstrate)Ketjenblack (concentratedon surface)VXC Spray 20 mAcm2Graphite FeltVulcan XC-72R(deposited on surface)Ketjenblack EC-300J(deposited on surface)Ketjenblack EC-300J(enti e felt coated)20 mA cm-2(a)    0.60.811.21.41.61.80 2000 4000 6000 8000Potential / VTime / minBlank Graphite FeltKetjenblack (entiresubstrate)Ketjenblack (concentratedon surface)VXC Spray 40 mAcm2Graphite FeltVulcan XC-72R(deposited on surface)Ketjenblack EC-300J(deposited on surface)Ketjenblack EC-300J(entire felt coated)40 mA cm-2(b)  Figure 5.9: Charge–discharge curves comparing zinc-iodine flow batteries at 20 mA cm−2geom (a) and 40 mAcm−2geom (b) where 50 mL of 1 M ZnI2 electrolyte is circulated at 20 mL/min (T = 22 ◦C). The electrode for thepositive half-cell was either an untreated GF, a GF uniformly coated with Ketjen-300, or a GF containing thesame loading (2 mg cm−2) of Ketjen-300 but concentrated in a region on its surface adjacent to the batterymembrane.plot for the blank cell which is in line with the charge transfer resistances calculated previously in Table 5.1.The charge-transfer resistance of the two Ketjen-300 electrodes is similar, but the internal resistance is muchlower for the ZIFB with Ketjen-300 concentrated onto the felt surface. This lower internal resistance is dueto the iodine electrode reactions primarily occurring within this carbon layer that is immediately adjacent tothe battery separator, shortening the ionic transport distance compared to other sites on the porous electrode.895.4. Electrode Impact on Flow Battery Performance012343 6 9 12 15Z'' / ohm ∙ cm2Z' / ohm ∙ cm2Blank Graphite FeltKetjen-300 (entirefelt coated)Ketjen-300(concentrated onsurface)Graphite FeltKetjenblack EC-300J(entire felt coated)Ketjenblack EC-300J(deposi ed on surface)Figure 5.10: Impedance spectra at open circuit voltage of the ZIFB with and without Ketjen-300 on thegraphite felt electrode (22 ◦C).The Ketjen-300 layer also increases the electrical conductivity of the felt electrode, which contributes tofurther lowering the internal resistance of the battery.0.20.611.41.80 12000 24000 36000Potential / VTime / min200 400 60050607080901000 2 4 6 8 10 12 14 16 18 20Efficiency / %Cycle NumberCE VE EEFigure 5.11: Cycling performance at 40 mA cm−2geom for the zinc-iodine flow battery with Ketjen-300 sprayedon the surface of the iodine-side electrode (a) and the coulombic efficiency (CE), voltage efficiency (VE)and energy efficiency (EE) as a function of cycle number. 50 mL of 1 M ZnI2 was pumped at 20 mL/minthrough both half-cells for twenty 1 hour long charge-discharge cycles.The cycling performance of the ZIFB at 40 mA cm−2 with the concentrated Ketjen-300 layer on theiodine-side felt electrode are presented in Figure 5.11. The ZIFB maintained its energy efficiency of 76%(compared to 63% for the untreated felt) over 20 charge-discharge cycles, a significant improvement over the905.4. Electrode Impact on Flow Battery Performanceblank ZIFB, and with no indications of deterioration at this point. Furthermore, the discharge power density(at 40 mA cm−2) increased from 391 to 441 W m−2. The improvement in performance and efficiency is onceagain attributed to the high surface area, increase in electrical conductivity of modified electrode comparedto the blank GF, and ability to encapsulate some I2 within the pores of the CL material which promotes theI2/I− reaction over I3/I−.0.70.91.11.31.51.70 2000 4000 6000 8000Potential / VTime / sblankket300can40 mA cm-2Graphite Felt(5 vol% ACN)(a)Graphite Felt(no cosolvent)Ketjen EC-300J(no cosolvent)0.70.91.11.31.51.70 2000 4000 6000 8000Potential / VTime / sblankket300ketacnketacn TGraphite Felt(no cosolvent)Ketjen EC-300J(no cosolvent)Ketjen EC-300J5 vol% ACNKetjenblack EC-300J5 vol% ACNT = 50 C40 mA cm-2o(b)Figure 5.12: Charge–discharge curves comparing the individual effect of electrode and electrode modifica-tions for the zinc-iodine flow batteries at 40 mA cm−2geom (a) their combined effect (b). The electrode for thepositive half-cell was either an untreated GF or a GF containing 2 mg cm−2 of Ketjen-300 concentrated ina region on its surface adjacent to the battery membrane. The operating temperature was 22 ◦C for all testsunless specified at 50 ◦C.The combined effects of acetonitrile cosolvent and the Ketjen-300 treated felt are presented in Fig-ure 5.12. Comparing the individual effects of the Ketjen modified electrode and the electrolyte with 5 vol%acetonitrile, the changes to the electrode material improved the operating efficiency more than the additionof cosolvent but both were significant improvements compared to the ZIFB without electrode and electrolytemodifications Table 5.3. However, incorporating the organic cosolvent into the flow battery along with thecarbon-treated felt was able to improve the performance relative to futher lower the overpotential lossesduring both charge and discharge (Figure 5.12b). The voltage during discharge decreases at a much lowerrate with the inclusion of acetonitrile cosolvent, likely because the mixed-solvent electrolyte is better ableto soak the carbon catalyst layer relative to the purely aqueous cosolvent. Increasing the operating tem-perature from 22 to 50 ◦C was able to further increase the efficiency, with the energy efficiency increasingfrom 61.8% to 80.3% after combining the electrode and electrolyte modifications high the higher operatingtemperature.915.5. Summary and ConclusionsTable 5.3: ZIFB performance at an applied current density of ± 40 mA cm−2geom as a function of positiveelectrode, with and without 5 vol% acetonitrile cosolvent. The graphite felts were loaded with 0.2 g cm−2 ofcarbon catalyst by either being uniformly coating the felt in the catalyst ink (I) or having the catalyst layerconcentrated onto the surface of the felt (II). The operating temperature was 22 ◦C for all tests except thelast entry which was taken at 50 ◦C. Electrolyte: 50 mL of 1 M of ZnI2 circulated at a flow rate of 20 mLmin−1. The coulombic efficiency for all tests was 98%.Positive Electrode Average chargevoltage / VAverage dischargevoltage / %EE / % PDdis / W m−2Pristine Graphite Felt 1.55 ± 0.0009 0.98 ± 0.01 61.8 391Ketjenblack EC-300J (I) 1.51 ± 0.001 1.00 ± 0.004 64.8 401Ketjenblack EC-300J (II) 1.42 ± 0.004 1.09 ± 0.01 74.6 441Vulcan XC-72R (II) 1.47 ± 0.003 1.04 ± 0.009 69.2 416Ketjenblack EC-300J (II)+ 5 vol% ACN1.40 ± 0.005 1.11 ± 0.005 77.7 444Ketjenblack EC-300J (II)+ 5 vol% ACN at 50 ◦C1.38 ± 0.008 1.13 ± 0.004 80.3 4525.5 Summary and ConclusionsSimilar to the analogous ZBFB, the efficiency and cycling performance of the ZIFB can be improvedby utilizing high surface area and high activity carbon materials as the positive electrode for the flow battery.Half-cell and flow battery studies revealed that inexpensive high surface area carbons with adequate porestructure and conductivity can significantly reduce the activation overpotentials of the iodide/iodine/triiodideredox reactions and the internal resistance of the cell to various extents. The improvements in the ZIFBperformance can be ascribed to the more favourable adsorption of iodine on these carbon materials whichpromotes the I2/I− redox reactions over that of I−3 /I−, in addition to improving the rate of these reactions andoverall lowering the activation overpotentials at the iodine electrode. Out of the materials studied, Ketjen-300 was consistently the most active towards the iodide (and bromide) redox reactions and incorporating thisinexpensive carbon material into the ZIFB electrode was able to improve the energy efficiency from 61.8% to74.6% even after days of continuous cycling with no signs of performance fade. Incoporating the acetonitrilecosolvent into the ZIFB with the catalyst-coated electrode further increased the energy efficiency of theZIFB to 80.3% This study provides much needed insight into the kinetics and electrochemical behaviour ofthe iodide redox reactions in aqueous flow batteries, and demonstrates that carbonaceous catalysts can beused to improve ZIFB performance without expensive or complicated catalyst synthesis procedures. Furtherresearch should be directed towards better understanding and fine-tuning the pore structure of the electrode925.5. Summary and Conclusionsand how this can promote I2/I− reactions within its pores while still taking advantage of the high solubilityof the I−3 species which allows zinc iodine flow batteries to achieve their impressive theoretical capacities.These improvements to the positive electrode material along modifications made to the electrolyte (Chapter 3and Chapter 4) promote the viability of this technology for large-scale stationary energy storage.93Chapter 6Conclusions and Recommendations6.1 Summary of Key FindingsThe ZIFB is presently one of the most promising flow batteries and overall candidate for stationaryenergy storage. The findings of this study have provided progress and insight with regards to both the elec-trolyte and electrode composition and corresponding impact on the ZIFB performance. The iodine film thatforms on charge is a major issue for the technology, but can be alleviated through enhancing the solubilityof molecular I2. Increasing the solubility of I2 was achieved by increasing the operating temperature ofthe ZIFB, which was also observed to increase the rate of the iodine electrode reactions and increase theoperating efficiency of the ZIFB The addition of various cosolvents was observed to significantly increasethe solubility of I2 and enhance both the charge and discharge processes. A majority of the non-aqueouscosolvents screened were able to increase the solubility of iodine except for those that were strongly basicsuch as pyridine and ethylenediamine. The addition of 5 vol% acetonitrile to the electrolyte of the ZIFBwas able to significantly increase the operating efficiency from 59.4% to 70%. Cosolvents like DMF andDMSO that were sufficiently basic and with large dipoles were found to interact with Zn2+ in solution andactually decrease the solubility of I2, despite iodine being highly soluble with these cosolvents when absentof Zn2+. A ZIFB constructed with a KI posolyte absent of Zn2+ but with 5 vol% DMF was able to increasethe operating efficiency from 58.9% to 65.4% without the formation of any precipitated iodine and therewas no significant crossover of Zn2+ into the iodine half-cell after repeated cycling. Thus far the iodineelectrode composition for the ZIFB has not been explored heavily, and the studies in Chapter 5 reveal thatlow cost carbonaceous catalysts can improve the cycling performance and power densities of the ZIFB. Thecarbonaceous catalysts VXC-72R, Ketjen-300, Ketjen-600, and MWCNT were able to increase the rate of946.2. Future Recommendationsthe iodine/triiodide reduction reactions significantly compared to the glassy carbon electrode. The ZIFBconstructed with a felt electrode sprayed with Ketjen-300 was able to operate at an energy efficiency of74.6% over repeated cycling, and increased the power density from 391 A m−2 on the pristine graphite feltto 441 A m−2 on the Ketjen-300 electrode. Combining the electrode and electrolyte modifications enhancedthe energy efficiency of the ZIFB by over 20% relative to the initial ZIFB.6.2 Future Recommendations(i) Further characterization of the iodine film through electrochemical quartz crystal nanobalancetechniquesThe electrochemical quartz crystal microbalance (EQCM) is used to measure the mass deposited ontoan electrode during the course of an electrochemical reaction. Recording the changes in mass observed overthe course of oxidation in the presence of cosolvents and other species in the electrolyte could help betterelucidate their interactions between iodine.(ii) In situ electrochemical spectroscopy of the iodine film and iodine charge transfer complexesIn this study, the solubility of iodine in various solutions was analyzed through ex-situ UV-Vis spec-troscopy. Analysis of the iodine film and its interactions with cosolvent species is not possible with thismethod because the film is transient and dissolves with time. Spectroelectrochemical methods can monitorchanges in the electrolyte during experiments which is particularly useful for this application where theiodine film and its properties/composition greatly impact the ZIFB performance. Besides UV-Vis, otherspectrometry techniques (in- and ex-situ) such as IR and Raman can also provide further valuable informa-tion regarding the I2X− and I2-cosolvent species in the electrolyte.(iii) Physical characterization of the carbon electrode materialsThe various carbonaceous materials investigated in this study showed varying activity towards theiodide/iodine redox reactions, and further characterization of the pore morphology will better elucidate theexact mechanism. Research on static zinc-iodine batteries shows that the size and structure of the pores956.2. Future Recommendationsaffect their ability to encapsulate reactant iodine as well as the adsorption capacity of the various iodinespecies. N2 sorption isotherms can be performed to obtain the pore structures, and the Brunauer-Emmett-Teller and Barrerr-Joyner-Halenda methods can be used to determine the specific surface areas and pore sizedistributions, respectively.(iv) Techno-economic analysis to evaluate scale-up potentialThe two main commercialized flow batteries at present are the VRB and the ZBFB. To further assessthe viability of ZIFB in the current market, techno-economic analyses can be performed to make furthercomparisons between the three flow batteries. The cost of iodine is lower than that of vanadium but ishigher than that of bromine but does not require the costly auxiliary components necessary when workingwith corrosive bromine electrolytes. 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Thesis/Dissertation
eng
Chemical and Biological Engineering
Vancouver : University of British Columbia Library
Attribution-NonCommercial-NoDerivatives 4.0 International
Graduate
Gyenge, Előd
Development of a polyiodide aqueous/non-aqueous flow battery
Text