Fundamental Flow Battery Studies

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FUNDAMENTAL FLOW BATTERY STUDIES:

ELECTRODES AND ELECTROLYTES

MALLORY A. MILLER

Submitted in partial fulfillment of the requirements for the degree of Doctor of Philosophy

Dissertation Advisor: Dr. Robert F. Savinell

Department of Chemical and Biomolecular Engineering CASE WESTERN RESERVE UNIVERSITY

May, 2017

CASE WESTERN RESERVE UNIVERSITY

DISSERTATION APPROVAL

We hereby approve the thesis/dissertation of

______Mallory A. Miller______

Candidate for the Doctor of Philosophy degree*

______Professor Robert Savinell______

______Professor Jesse Wainright______

______Professor Burcu Gurkan______

______Professor Mark DeGuire______

Date _____December 13, 2016_____

*We also certify that written approval has been obtained for any proprietary material contained therein.

Dedication

To my mother for always believing in me and for her encouragement and support.

III

Table of Contents

Acknowledgements ...... VIII

List of Tables ...... IX

List of Figures ...... XII

Abbreviations ...... XXIV

Abstract ...... XXVI

Chapter 1. Introduction and Background ...... 1

1.1 Redox Flow Batteries ...... 1

1.2 Specific Motivations and Scope ...... 5

1.2.1 Ionic Liquids - A Different Approach to Non-Aqueous ...... 5

1.2.2 Electrode Kinetics ...... 7

1.3 Concluding Remarks ...... 9

1.4 Publications ...... 10

Chapter 2. Prospect of Ionic Liquids as Redox Flow Battery Electrolytes ...... 12

2.1 Introduction and Literature...... 12

2.1.1 Ionic Liquids ...... 12

2.1.1.i Ionic Liquids for Electrochemical Applications ...... 16

2.1.1.ii Ionic Liquids for Flow Batteries ...... 18

2.1.1.iii Iron Electrochemistry - an Introduction ...... 21

2.1.1.iv Physical Properties of Deep Eutectic Solvents (DES) ...... 22

2.2 Iron-Containing Ionic Liquid Electrolytes ...... 25

2.2.1 Iron Electrokinetics ...... 27

2.2.2 Electrodeposition of Iron...... 30

2.2.2.i Nucleation Mechanism ...... 31

2.2.3 Physical Properties of the Electrolytes ...... 40

2.2.3.i Effect of Chloride to Iron Ratio on Conductivity ...... 40

2.2.3.ii Walden Plot ...... 42

2.2.4 Iron Speciation ...... 44

2.2.4.i X-ray Absorption Near Edge Structure Measurements ...... 44

2.2.4.ii X-ray Photoelectron Spectroscopy ...... 47

2.2.4.iii Raman Spectroscopy ...... 51

2.3 Iron IL Flow Battery - Proof of Concept ...... 54

2.4 Conclusions ...... 56

Chapter 3. Technique Development, Carbon Fiber Microelectrodes ...... 58

3.1 Introduction and Literature...... 58

3.2 Technique Development - Carbon Fiber Microelectrode ...... 61

3.2.1 Preconditioning ...... 62

3.2.2 Electrochemically Active Area ...... 65

3.3 Potential Applications ...... 67

3.3.1 Non-Aqueous Systems - Understanding Kinetics and Degradation ...... 67

3.3.2 Metal Electrodeposition for Hybrid RFBs ...... 72

3.4 Conclusions ...... 75

Chapter 4. Carbon Electrode Pretreatment Effects ...... 77

4.1 Introduction and Literature...... 77

4.2 Carbon Electrode Treatments ...... 78

4.3 Surface Analysis Results ...... 79

4.3.1 Scanning Electron Microscopy ...... 80

4.3.2 Electrochemical Capacitance...... 81

4.3.3 XPS Analysis...... 82

4.3.5 Contact Angle Characterization ...... 87

4.4 Applications to the All-Vanadium System - Carbon Fiber Electrodes for Understanding of Electrokinetics ...... 89

4.4.1 Significant Results ...... 90

4.4.1.i Kinetic Parameters from EIS and LSV ...... 92

4.4.1.ii Effects of Treatment Potential ...... 94

4.4.1.iii Development of a Mechanistic Model ...... 98

4.5 Concluding Remarks ...... 106

Chapter 5. Conclusions and Future Work ...... 108

5.1 Conclusions ...... 108

5.1.1 Ionic Liquids as RFB Electrolytes ...... 108

5.1.2 The CFME, Technique Development ...... 109

5.1.3 Carbon Electrode Pretreatment Effects ...... 110

5.2 Future Work ...... 111

5.2.1 Other Halides and Hydrogen Bond Donors ...... 111

5.2.2 Investigation of Alternative Active Metal Species ...... 111

5.2.3 Aqueous All-iron Flow Battery with Choline Chloride Supporting Electrolyte (and

Choline as an Additive for Plating) ...... 112

5.2.4 Ionic Liquids for High Power Energy Storage Applications (Supercapacitors)...... 112

Appendix A1. Halides and Hydrogen Bond Donors ...... 114

A1.1 Introduction ...... 114

A1.2 Effects of the Halide ...... 114

A1.3 Effects of the Hydrogen Bond Donor (HBD) ...... 118

Appendix A2. Investigation of Alternative Active Metal Species ...... 123

A2.1 Introduction and Background ...... 123

A2.2 Electrodeposition Reactions ...... 124

A2.2.1 Nickel ...... 125

A2.2.2 Tin ...... 126

A2.2.3 Zinc ...... 128

A2.2.4 Iron-zinc Alloy ...... 129

A2.3 Redox Reactions ...... 135

A2.3.1 Vanadium Electrolytes ...... 140

A2.3.2 Iron-vanadium Mixed Electrolyte ...... 142

VII

Appendix A3. Aqueous All-iron Flow Battery with Choline Chloride Supporting

Electrolyte ...... 145

A3.1 Introduction and Background ...... 145

A3.2 Three electrode cell - Preliminary Data ...... 145

A3.3 Small Scale Flow Cell ...... 149

References ...... 153

VIII

List of Tables

2+ + Table 1-1. Kinetic rate constants for the VO /VO2 reaction at carbon electrodes...... 9

Table 2-1. Common DES formed by mixing choline chloride with various hydrogen bond donors in a 1:2 molar ratio [87]...... 15

Table 2-2. A comparison of aqueous vs. non-aqueous electrolyte-solvent systems...... 17

Table 2-3. A summary of flow battery systems that utilize ionic liquids...... 20

Table 2-4. Examples of the synthesized electrolytes evaluated with corresponding iron concentration. Solutions containing 1:2:0 and 2:0:4 could not be created using only FeCl2...... 27

Table 2-5. Summary of exchange current densities for the Fe2+/Fe3+ reaction estimated from the slope of the LSV potential-current density data in the linear region...... 29

Table 2-6. Coulombic plating efficiency for the IL electrolytes of this study at 80 oC. Iron was plated at −1.0 V vs Ag/AgCl for 5 minutes then stripped for 5 minutes at +0.2 V...... 39

Table 2-7. X-ray Absorption Near Edge Structure (XANES) data obtained from deconvolution of the Fe pre-edge peaks for samples prepared with (a) FeCl2 and (b) FeCl3...... 46

Table 2-8. Summary of speciation data obtained from deconvolution of high resolution XPS. . 51

Table 3-1. Equations governing the diffusion current; three different electrode geometries are compared...... 60

Table 4-1. Electrochemical capacitance values for various CFMEs tested in 3 M H2SO4; each data point represents an average of 5 different electrodes...... 81

Table 4-2. Chemical spectra obtained from curve fitting the high resolution C1s peaks...... 84

Table 4-3. Chemical spectra obtained from curve fitting the high resolution C1s peaks...... 86

Table 4-4. Contact angles for various felt samples; a decrease in contact angle indicates an increase in felt wettability...... 88

Table 4-5. Comparison of kinetic rate constants, ko, after treatment at +1.5 V (MSE) to that after

2+ 3+ 2+ + o treatment at -2.0 V (MSE) for both V /V and VO /VO2 reactions on CFMEs at 25 C; rate constants were estimated from both LSV and EIS and each value is the average of four different

CFMEs...... 93

Table 4-6. Parameters used to fit experimental EIS data for CFME in 1.5 M total vanadium (1:1

+ + molar ratio of VO /VO2 ) and 3 M H2SO4. Note that Rsoln is the high frequency solution resistance,

Rct is the charge transfer resistance, and Ra is the resistance due to adsorption; all have units of

2 P-1 2 W ×cm . CPE T (units of Fs /cm ) is related to capacitance of a constant phase element and CPE

P is the constant phase exponent which relates to deviation of the straight capacitive line from 90o by angle α=90o(1-P). Note that when CPE P=1, the constant phase element acts as an ideal capacitor. Subscripts “dl” and “a” indicate processes associated with double layer charging or adsorption...... 103

Table 4-7. Parameters used to fit experimental EIS data for CFME in 1.5 M total vanadium (1:1

2+ 3+ molar ratio of V /V ) and 3 M H2SO4. Note that Rsoln is the high frequency solution resistance,

Rct is the charge transfer resistance, and Ra is the resistance due to adsorption; all have units of

2 P-1 2 ×cm . CPE T (units of Fs /cm ) is related to capacitance of a constant phase element and CPE

Table A1-1. Conductivity and kinetics (at platinum) for electrolytes containing 1:1:4

o FeX2:ChX:EG, where X=Cl, Br, or I at 80 C, N2-atmosphere. Note that only the exchange current

X estimated from LSV is provided; this is due to the fact that only Fe2+ ions were present and thus, ko could not be easily estimated (cannot assume equal molar reduced and oxidized species). .. 115

Table A1-2. Results from elemental analysis of the electrodeposited iron from selected ILs. Note that the carbon balance is not included...... 122

Table A2-1. Standard reduction reactions with the corresponding equilibrium potentials [175,

214]...... 124

Table A2-2 Physical and electrochemical properties of the synthesized electrolytes at 80 oC; all solutions contained MCl2:choline chloride:ethylene glycol in a 1:1:4 molar ratio...... 128

Table A2-3 Physical and electrochemical properties of the synthesized electrolytes at 80 oC; all solutions contained MCl2:choline chloride:ethylene glycol in a 1:1:4 molar ratio...... 130

Table A2-4. Selected standard reduction reactions with the corresponding equilibrium potentials

[175, 214]...... 136

Table A2-5. Summary of initial data for metals tested in a choline chloride and ethylene glycol based electrolyte...... 137

Table A2-6. Exchange current densities estimated from EIS (OCV ± 15 mV AC) and LSV at

o platinum and glassy carbon in 0.1M VCl3 solutions; 50 C, N2-atmosphere...... 142

Table A2-7. Summary of conductivity, kinetic, and cell potential data for 1:1:4 MClx:ChCl:EG

o electrolytes; 80 C, N2-atmosphere. Exchange current densities were estimated from LSVs at glassy carbon or platinum electrodes...... 144

Table A3-1. Hydrogen evolution current at an iron rod in each of the different electrolytes (50 oC)

2 showing potential (iR corrected) at which the H2 evolution current reached 1 mA/cm . All experiments had been performed in a sealed three electrode cell which was purged with N2 gas.

...... 146

List of Figures

Figure 1-1. Schematic of an aqueous all-vanadium redox flow battery. Reactive species are dissolved in an electrolyte and stored in external reservoirs, then pumped though the electrolytic stack, where the reactions occur; at the negative electrode, the reactive species are reduced and at the positive electrode they are oxidized on charge. The negative and positive half cells are often separated by an ion-selective membrane or microporous separator...... 2

Figure 1-2. The vanadium (III) acetylacetonate metal-ligand complex ...... 4

Figure 1-3. Maximum solubility of V(acac)3 in acetonitrile as a function of supporting salt concentration at 25oC. Figure is reprinted from reference [46]...... 6

Figure 1-4. Potential windows for various electrolytes shown on a platinum disk electrode at

100mV/s and 25oC. Cyclic voltammetry of non-aqueous electrolytes were performed under a dry

N2-atmosphere; 0.1 M TEABF4 in acetonitrile (ACN) represents a typical organic electrolyte and reline and ethaline are two deep eutectic type ionic liquids...... 7

Figure 2-1. Possible complex formation in a 1:2 choline chloride:urea deep eutectic solvent.

Figure was obtained from reference [86]...... 13

Figure 2-2. A mixture of choline chloride and urea in a 1:2 molar ratio is shown (a) prior to heating and (b) after being placed in a vacuum oven at 150oC for four hours and cooled back down to room temperature. The image on the right hand side is the resulting deep eutectic, reline, at 25 oC. ... 14

Figure 2-3. Various hydrogen bond donors which when mixed with choline chloride for DES. 15

Figure 2-4. Schematic of iron flow battery reactions at carbon felt electrodes upon charge...... 21

Figure 2-5. (a) Conductivity and (b) viscosity vs temperature for two DES...... 23

Figure 2-6. Conductivity vs temperature for DES synthesized from different precursors...... 24

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Figure 2-7. Cyclic voltammetry of 0.51 M FeCl2 in ethaline (a 0.25:2:4 molar ratio

o FeCl2:ChCl:EG) at a platinum electrode; 10 mV/s, 50 C, N2-atmosphere...... 25

Figure 2-8. LSV (at 5 mV/s) in 1:1:4 (equal molar Fe2+/Fe3+) at a platinum microelectrode; 80 o C, N2-atmosphere...... 29

Figure 2-9. Cyclic voltammetry at a glassy carbon electrode in (a) 0.1:2:4 (b) 0.5:1:4 or (c) 1:1:4

o FeCl2:ChCl:EG at 50 mV/s, 80 C, in a N2-atmosphere. Current was normalized to allow for a more direct comparison. Inset: comparison of the 1:1:4 and 1:0:4 solutions...... 31

Figure 2-10. Current vs time transients for iron electrodeposition at −1.15 V vs Ag/AgCl in

(a) 0.25:2:4 and (b) 1:1:4 FeCl2:ChCl:EG electrolytes...... 32

2 Figure 2-11. Nondimensional plots (I/Imax) vs t/tmax for experiments of iron electrodeposited at

−1.15 V vs Ag/AgCl in (a) 0.25:2:4 and (b) 1:1:4 FeCl2:ChCl:EG electrolytes. Dashed and dotted lines represent progressive and instantaneous nucleation, respectively...... 32

2 Figure 2-12. Nondimensional plots (I/Imax) vs t/tmax for iron electrodeposited at (a) selected potentials (V vs Ag/AgCl) and (b) at −1.15 V vs Ag/AgCl onto different substrates. All experiments had been performed in a 1:1:4 FeCl2:ChCl:EG electrolyte. Dashed and dotted lines represent progressive and instantaneous nucleation, respectively...... 34

2 Figure 2-13. Nondimensional plots (I/Imax) vs t/tmax for zinc electrodeposited onto glassy carbon at −1.0 V vs Ag/AgCl in a 1:1:4 ZnCl2:ChCl:EG electrolyte. Dashed and dotted lines represent progressive and instantaneous nucleation, respectively...... 35

2 Figure 2-14. Nondimensional plots I/Imax vs t/tmax for iron electrodeposited at −1.15 V from 1:1:4

FeX2:ChX:EG, where X=Cl, Br, or I. Dashed and dotted lines represent progressive and instantaneous nucleation, respectively...... 36

XIII

Figure 2-15. Plating 10 mAh @ 5 mA/cm2 onto a glassy carbon substrate (80 oC) from (a) 0.25:2:4 and (b) 1:1:4 FeCl2:ChCl:EG electrolytes. Initial potential was ~−1.15 V vs Ag/AgCl for deposition and levelled out at −0.9 and −1.1 resp. Elemental analysis (energy dispersive spectroscopy) using a Helios NanoLab™ 650 confirmed the samples to be iron, however, the iron plated from a 0.25:2:4 electrolyte exhibited more oxides (91.9 ±0.4 iron, 6.9 ±0.3 oxygen, and carbon balance compared to 100 ±0 for a deposit from the 1:1:4 electrolyte)...... 37

Figure 2-16. Example plot for estimating coulombic plating efficiency. The plating potential was held at −1.0 V for 500 s then the metal deposit was stripped at + 0.2 V for 500 s at a glassy carbon

o substrate in 1:1:4 FeCl2:ChCl:EG electrolyte; N2-atmosphere, 80 C...... 38

Figure 2-17. H-type cell setup for studying electrolyte degradation products. This photograph was taken after bulk electrolysis of a 1:2 M ChCl:EG electrolyte for 5 hours at 22 mA/cm2, 80 oC. . 63

Figure 2-18. Conductivity of ILs as a function of concentration of (a) iron (b) choline or (c) ethylene glycol at 80oC...... 41

Figure 2-19. Walden plot for selected iron (Fe2+) containing ILs compared to 2:4 M ChCl:EG electrolyte as well as the ideal solution line; the red line indicates a deviation of ΔW = 1...... 43

Figure 2-20. Normalized pre-edge spectra (Fe K-edge) for selected electrolytes containing (a) Fe2+

3+ and (b) Fe compared to standard, FeCl2•4 H2O (FeCl2 salt) and FeCl3•6 H2O (FeCl3 salt) with the corresponding models of best fit...... 46

Figure 2-21. High resolution Fe 2p spectra for the iron metal before and after sputtering for

4 minutes to remove the oxide layer...... 48

Figure 2-22. High resolution XPS for the (a) Fe 2p and (b) C 1s peaks. High power mode is compared to the 100 µm spot, highlighting the effects of sample charging...... 49

Figure 2-23. High resolution XPS for the (a) Fe 2p and (b) N 1s peaks...... 50

XIV

Figure 2-24. Raman spectra for various iron(III)chloride containing IL electrolytes compared to the spectra for ethylene glycol and a mixture of ethylene glycol and choline chloride. The presence

- -1 [FeCl4] tetrahedral complex was confirmed by large peaks located at 333 cm [148, 149, 154].52

Figure 2-25. Raman spectra for ethylene glycol compared to 1:4 M ChCl:EG as well as 0.5:1:4 and 1:0:4 FeCl2:ChCl:EG electrolytes; vibrational modes for EG are labeled [158, 159]...... 53

Figure 2-26. Proposed Fe-EG complex...... 53

Figure 2-27. Schematic of the small-scale flow battery hardware...... 55

Figure 2-28. One hour charge-discharge cycles for a RFB utilizing 0.5:1:4 FeCl2:ChCl:EG after

2 o an initial charge to 10% SOC; current density of 5 mA/cm , flow rate 25 mL/min, 80 C, N2- atmosphere...... 56

Figure 3-1. (a) Photograph and (b) corresponding SEM micrograph to show the structure of typical carbon electrodes used in flow batteries...... 59

Figure 3-2. (a) Exchange current densities from TFFRDEs for two different active species concentrations (equimolar FeCl2 and FeCl3); supporting electrolyte of 1 M NaCl was used.

(b) Pulse voltammetry (250ms) polarization of TFFRDEs at four different loadings in 0.02 M total iron and 1 M NaCl (equimolar concentrations of FeCl2 and FeCl3). Reprinted with permission from reference [173]...... 59

Figure 3-3. SEM micrographs of a typical carbon fiber used ...... 61

Figure 3-4. Schematic of carbon fiber microelectrode (CFME)...... 61

Figure 3-5. Schematic of a three electrode cell employing a platinum counter electrode, reference electrode (situated inside a Luggin capillary), and CFME working electrode...... 62

Figure 3-6. Selected CV cycles at a CFME in H2SO4 at 100 mV/s to show the electrochemical pretreatment procedure used to obtain a stable surface...... 63

Figure 3-7. SEM micrographs of (a) an untreated fiber compared to (b) a preconditioned fiber.

Images (c) and (d) represent untreated felts coated with palladium so that the insulating material on the surface could be detected using SEM-EDS techniques (electron interaction depth was

0.8 µm and radius of 0.4 µm); (d) shows the general size of the particles (<100 nm diameter). . 64

Figure 3-8. Comparison of typical CVs (100 mV/s) at a preconditioned CFME and a glassy carbon electrode in 3 M H2SO4. The currents were normalized to the peak current at ~0 V (MSE)...... 65

Figure 3-9. (a) CVs of a CFME in 3 M H2SO4 at varying scan rates (20 – 300 mV/s) which can be used to determine capacitance when (b) the capacitive current, ic, is vs scan rate...... 66

Figure 3-10. Cyclic voltammetry of non-aqueous vanadium reactions at 500 mV/s comparing two different electrodes (0.01 M V(acac)3, 0.1M TEABF4 in acetonitrile)...... 68

Figure 3-11 Cyclic voltammetry of non-aqueous vanadium reactions at 100 mV/s in 0.01 M

V(acac)3, 0.1M TEABF4 in acetonitrile at (a) a CFME and (b) a platinum fiber electrode.

o Experiments were performed in a N2-atomosphere (>3 ppm water) at 25 C with a N2 blanket . 69

Figure 3-12 Cyclic voltammetry at 100 mV/s in 0.1M TEABF4 in acetonitrile with either 0.01 M

V(acac)3 (which had been contaminated with oxygen/moisture) or 0.0075 M VO(acac)2 at a

o CFME. Experiments were performed in N2-atomosphere (3 ppm water), 25 C with a nitrogen blanket above the electrolyte...... 70

Figure 3-13 SEM micrographs showing a CFME which had been cycled 200 times at 50 mV/s between −0.9 V and 1.5 V (vs Hg/Hg2SO4) and subsequent potential holds alternating between

−2 V and +1.5 V until the electrode broke. Images are in order, moving along the length of the fiber where (e) is closest to where the fiber broke during the test...... 72

Figure 3-14. Cyclic voltammetry showing the iron reactions at a CFME (10 mV/s) in a 1:1:4

o FeCl2:ChCl:EG electrolyte. Experiments had been performed in a N2-atmosphere at 80 C...... 73

XVI

Figure 3-15 Coulombic plating efficiency as a function of potential for iron electrodeposited from

o 1:1:4 FeCl2:ChCl:EG at 80 C...... 74

Figure 3-16. 10 mAh of iron plated at 5 mA/cm2 onto a CFME (80 oC) from a 1:1:4

FeCl2:ChCl:EG electrolyte. A section along the (a) length as well as the (b) end are shown; the end had been broken off the CFME assembly which caused some iron to flake off...... 75

Figure 4-1.SEM micrographs show surface morphology of PAN-based fibers after treatment. . 80

Figure 4-2. Survey scans for the treated felt samples, compared to the untreated sample...... 83

Figure 4-3. Typical XPS survey scans of carbon felt samples after electrochemical treatments in

3 M H2SO4 after reduction at ‒2.0 V (blue) compared to oxidation at +1.5 V (red). The two spectra were normalized with respect to the intensity of the C1s peak...... 85

Figure 4-4. Curve fitted high-resolution C1s and O1s spectra for the (a) untreated sample compared to felt samples that had been oxidized at +1.5V and (c) reduced at −2.0V...... 87

Figure 4-5. (a) Schematic and (b) photograph showing contact angle measurements of 1.0 µL water droplets on felt; (c) water droplets on the surface of an untreated (as-received) felt...... 88

Figure 4-6. LSVs (5 mV/s) at a CFME in (a) V2+/V3+ electrolyte performed using the same electrode in the same electrolyte but after two different treatments; reduction at -2.0 V (MSE) or

2+ + oxidation at +1.5 V (MSE). And (b) LSVs (5 mV/s) at a CFME in VO /VO2 electrolyte performed using the same electrode in the same electrolyte but after two different treatments; reduction at -2.0 V (MSE) or oxidation at +1.5 V (MSE)...... 91

Figure 4-7. Nyquist plots for a CFME in V2+/V3+ electrolyte after (a) oxidation and (b) reduction treatments at the potentials indicated. The dashed blue (a) and dashed red (b) lines correspond to the initial (baseline) treatments at − 2.0 V and +1.5 V (MSE), respectively. Each spectrum is represented by a continuous line. Selected frequencies: 100 Hz ( ), 10 Hz ( ), 1 Hz ( )...... 94

XVII

2+ 3+ Figure 4-8. Rate constants, ko, for V /V plotted against oxidation (O) and reduction (R) treatment potentials. The data in O and R are from Nyquist plots such as those in Fig. 7(a) and

7(b), respectively. Arrows show the oxidation or reduction treatment potential relative to the initial baseline potential...... 96

2+ + Figure 4-9. Rate constants, ko, for VO /VO2 plotted against oxidation (O) and reduction (R) treatment potentials. The arrows show the oxidation or reduction treatment potential relative to the initial baseline potential...... 97

2+ + Figure 4-10. Normalized activity for VO /VO2 as a function of treatment potentials for various carbon electrodes (GC: glassy carbon, F: CFME, RVC: reticulated vitreous carbon, CP: carbon paper, Gr: graphite). This figure is from ref [164]...... 98

2+ + Figure 4-11. Electronic structures of (a) vanadyl (VO ) and (b) vanadate (VO2 ) [202-204] . 122

Figure 4-12. Nyquist plot showing impedance of a typical electrocehmical reaction but with

(a) semi-infinite linear diffusion to a flat plate compared to diffusion to (b) a micro cylindrical electrode and (c) spherical diffusion to a micro disk electrode ...... 100

Figure 4-13. Nyquist and Bode plots for CFME in 1.5 M total vanadium (1:1 molar ratio of

2+ + 2+ 3+ (a) VO /VO2 or (b) V /V ) and 3 M H2SO4 (please note: the inset in the Nyquist plot for

(b) represents a zoomed in version). EIS was performed using the under the same conditions but after oxidation (+1.5V vs Hg/Hg2SO4) or reduction (−2.0V vs Hg/Hg2SO4) pretreatments. This set of experiments had been repeated/cycled 8 times to ensure surface area was not changing, only the data from cycle 6 are shown for clarity...... 101

Figure 4-14. Equivalent circuit model for a faradaic reaction involving one adsorbed species with subsequent desorption; this model was used to fit all experimental data. Note that Rsolution is the high frequency solution resistance, Rct is the charge transfer resistance, and RA is the resistance

XVIII due to adsorption. Subscripts “dl” and “A” indicate processes associated with double layer charging or adsorption ...... 102

Figure 4-15. Nyquist and phase angle plots for CFME in 1.5 M total vanadium (1:1 molar ratio

2+ + of VO /VO2 ) and 3M H2SO4. Simulated values (solid line) in the range of 0.2 to 20,000 Hz are compared to experimental data (open circles) for both (a) reduced and (b) oxidized electrodes; refer to Table 5 for values used in simulations...... 104

Figure 4-16. Nyquist and phase angle plots for CFME in 1.5 M total vanadium (1:1 molar ratio

2+ 3+ of V /V ) and 3 M H2SO4. Simulated values (solid line) in the range of 0.2 to 20,000 Hz are compared to experimental data (open circles) for both (a) reduced and (b) oxidized electrodes; refer to Table 5 for values used in simulations...... 105

o Figure A1-1. Cyclic voltammetry 1:1:4 FeX2:ChX:EG, X=Cl, Br, or I at 50 mV/s, 80 C...... 116

Figure A1-2. Coulombic plating efficiency as a function of voltage at a glassy carbon substrate for solutions containing 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I. Iron was plated at a selected potential for 5 minutes then stripped for 5 minutes at a potential more positive of the electrodissolution peak (0.2 V for Br and Cl containing electrolytes and 0.4 V for I); efficiency was estimated from the number of coulombs passed during stripping divided by coulombs during plating. Temperature was maintained at 80 oC...... 116

Figure A1-3. SEM micrographs of 4 mAh iron electrodeposited (at 0.2 mA/cm2) onto a glassy

o carbon substrate at 80 C from solutions containing 1:1:4 FeX2:ChX:EG, X=Cl, Br, or I...... 117

Figure A1-4. Conductivity of 1:n:4 FeClz:ChCl:X where n=1 or 2, X = ethylene glycol or urea,

o and Z = 2 or 3; 80 C, dry N2-atmosphere...... 119

Figure A1-5. Cyclic voltammetry at a glassy carbon electrode in 1:1:4 FeCl2:ChCl:X where

X = ethylene glycol, urea, or water at 50 mV/s, 80 oC...... 119

XIX

Figure A1-6. Coulombic plating efficiency for the synthesized electrolytes at 80 oC. Iron was plated at a selected potential for 5 minutes then stripped for 5 minutes at a potential more positive of the electrodissolution peak (0.2 V); efficiency was estimated from the number of coulombs passed during stripping divided by coulombs during plating...... 120

Figure A1-7. Plating 10 mAh @ 5 mA/cm2 onto a glassy carbon substrate (80 oC) from 1:1:4

FeCl2:ChCl:X electrolytes where X = (a) ethylene glycol and (b) urea...... 121

Figure A1-8. Plating 10 mAh @ 5 mA/cm2 onto a glassy carbon substrate (80 oC) from 1:1:4

FeCl2:ChCl:urea. The red arrow indicates an area that has been oxidized (rust)...... 121

o Figure A2-1. The same nickel electrolyte (1:1:4 NiCl2:ChCl:EG) at (a) 27 C compared to

(b) 80 oC stored in a dry, nitrogen-filled glovebox...... 126

Figure A2-2. Comparison of CVs (100 mV/s) at a 100 µm platinum disk in electrolytes containing different metal species in a 1:1:4 ratio; all solutions had been thermostatted to 80 oC and kept in a dry nitrogen-filled glovebox...... 127

o Figure A2-3 Zinc deposited from (a) 1:1:4 ZnCl2:ChCl:EG (10 mAh, 80 C) compared to zinc deposited from (b) a typical aqueous electrolyte (0.1 M ZnO + 4.0 M KOH, room temperature) at

12.5 mA/cm2. Fig 3 (b) has been taken from [221]...... 129

Figure A2-4 Comparison of CVs (100 mV/s) at a platinum disk in electrolytes containing either

(a) ZnCl2, (b) equal molar FeCl2/ZnCl2 or (c) FeCl2 only in a 1:1:4 ratio; all solutions had been thermostatted to 80 oC and kept in a dry nitrogen-filled glovebox...... 131

Figure A2-5 Example of X-ray photoelectron spectroscopy survey scan from metal

o electrodeposited onto a glassy carbon substrate from 1:1:4 FeCl2/ZnCl2:ChCl:EG at 80 C. .... 132

Figure A2-6. Metal composition determined from XPS survey scans for Iron-Zinc

o electrodeposited onto a glassy carbon substrate from 1:1:4 FeCl2/ZnCl2:ChCl:EG at 80 C. All

XX samples contained 10 mAh electrodeposited metal. Two different areas had been analyzed for the sample where metal was deposited at −1.05 V...... 132

Figure A2-7 Iron-Zinc (10 mAh) electrodeposited onto a glassy carbon substrate from 1:1:4

o 2 FeCl2/ZnCl2:ChCl:EG at 80 C at (a) 0.5, (b) 2.5, (c) 4, and (d) 5 mA/cm ...... 134

Figure A2-8 Iron-Zinc (10 mAh) electrodeposited onto a glassy carbon substrate from a 1:1:4

o 2 FeCl2/ZnCl2:ChCl:EG at 80 C, 5 mA/cm ...... 135

o Figure A2-9 Cyclic voltammetry at platinum (25 mV/s) in 0.1 M CrCl2 in ethaline, 50 C ..... 137

Figure A2-10 Cyclic voltammetry at a platinum electrode (100 mV/s) in 0.1 M MnCl2 in

(a) ethaline and (b) reline based electrolytes at 50 oC...... 138

Figure A2-11 Cyclic voltammetry (100 mV/s) at a platinum disk in an electrolyte containing 1:1:4

o InCl3:ChCl:EG thermostatted to 80 C and kept in a dry nitrogen-filled glovebox. The arrows indicate which oxidation/reduction processes are coupled with one another...... 139

Figure A2-12 Cyclic voltammetry for a 1:1:4 (YbCl3:ChCl:EG) electrolyte at a platinum microelectrode, 150 mV/s, 80 oC. Inset shows the electrolyte paste...... 140

o Figure A2-13 Cyclic voltammetry at GC in 0.01 M VCl3 (in ethaline); 100 mV/s, 50 C ...... 141

Figure A2-14. Cyclic voltammetry 0.1 M VCl3 in (a) ethaline and(b) reline on GC (10 mV/s) 141

Figure A2-15. Cyclic voltammetry of 1:1:4 VCl3:ChCl:EG and 1:1:4 (FeCl2/FeCl3:ChCl:EG) at a platinum electrode, 80 oC, 50 mV/s ...... 143

o Figure A2-16. Cyclic voltammetry of 1:1:4 FeCl2/VCl3:ChCl:EG at platinum, 80 C,50 mV/s143

o Figure A3-1. Cyclic voltammetry showing H2 evolution current at an iron rod (50 C)...... 146

Figure A3-2. Coulombic plating efficiency as a function of voltage at a glassy carbon substrate for a typical aqueous electrolyte (0.5 M FeCl2•4H2O and 1M KCl in water) compared to solutions containing 1:1:4 FeCl2:ChCl:EG (prepared and stored in a dry N2-atmosphere), a 1:1:4

XXI

FeCl2•4H2O:ChCl:EG electrolyte, and a 1:1:4 FeCl2•4H2O:ChCl:water electrolyte. Temperature was maintained at 80 oC with a nitrogen gas purge...... 147

Figure A3-3. Cyclic voltammetry at a glassy carbon electrode in 1:1:4 FeCl2:ChCl:EG (“EG”)

o compared to 1:1:4 FeCl2•4H2O:ChCl:water at 50 mV/s, 80 C...... 148

Figure A3-4. Schematic of the small-scale flow battery hardware...... 149

Figure A3-5.1 hour charge/discharge cycles in a small-scale flow cell employing a 1:1:4

2 2 o FeCl2•4H2O:ChCl:water (cycles 1- 4, 25 mA/cm and 5-40, 10 mA/cm ); N2 purge, 50 C. .... 150

Figure A3-6. Coulombic, voltaic, and energy efficiency as a function of cycle number for 1 hour charge/discharge cycles in a small-scale flow cell employing a 1:1:4 FeCl2•4H2O:ChCl:water

o electrolyte; N2 purge, 50 C ...... 151

Figure A3-7. Photograph of the negative electrode after assembling the cell after a charge at

10 mA/cm2. The arrow indicates where the electrolyte had thickened, blocking flow out of the cell (from the negative electrode compartment)...... 151

Figure A3-8. Battery electrodes (KFD carbon felt bonded to the graphite plates using Acheson carbon ink) and Daramic separator are shown after the battery was disassembled after a 1 hour

2 o charge at 10 mA/cm ; 50 C, N2 gas purge. Arrows mark the direction of flow. Electrodeposited iron can be seen on the surface of the negative electrode...... 152

XXII

Acknowledgements

First off, I would like to thank my advisors, Dr. Jesse Wainright and Dr. Robert Savinell. I have learned SO MUCH over the past few years under your guidance. I am thankful to have had the opportunity to work with two individuals who truly value their students and put their best interests first. Jesse Wainright, I have learned a great deal from our conversations and from working in the lab with you, as well as being a teaching assistant for your class. Robert Savinell, I thank you for providing me with this amazing opportunity, but more than anything for believing in me and for mentoring me. Words cannot express my gratitude.

I would like to thank the members of the electrochemical engineering and energy lab (EEEL) team.

I am so fortunate to have had the chance to work in such a collaborative environment, learning from each and every one of you. Many of you were mentors to me and I truly looked up to you.

You all have made my time a CWRU an enjoyable experience.

I am thankful for collaborations with Dr. Noel Buckley (University of Limerick) and his team of researchers. I was able to gain new insights and a different perspective while working alongside him and his students. I am also grateful for Dr. Levi Thompson (University of Michigan) and his team of researchers for their insight.

This research used resources of the Advanced Photon Source, a U.S. Department of Energy Office of Science User Facility operated for the DOE Office of Science by Argonne National Laboratory.

Furthermore, I would like to thank the National Science Foundation, Sustainable Energy Pathways

Program (NSF-1230236) for funding this work as well as our collaborators at the University of

Michigan.

XXIII

Abbreviations

A Area (cm2)

α Transfer coefficient

C Concentration of active species (mol/L)

2 Cd Differential capacity (uF/cm )

2 Cdl Capacitance of the double layer (uF/cm )

CE Coulombic efficiency

CPE P Constant phase exponent, relates to deviation of the straight capacitive line

from 90o by angle: α=90o(1-P). Note that when P=1, the constant phase

element acts as an ideal capacitor.

P-1 2 CPE T Parameter related to capacitance of a constant phase element (Fs /cm )

D Diffusion coefficient (cm2/s)

E Applied potential (V)

Ê Energy density (Wh/L)

EE Energy efficiency

Eo Standard equilibrium potential (V)

F Faraday’s constant i Current density (mA/cm2)

2 ic Differential capacitance current density (mA/cm )

2 io Exchange current density (mA/cm ) k Boltzmann Constant ko Heterogeneous rate constant (cm/s) n Number of electrons transferred

XXIV

OCV Open Circuit Voltage (V)

P Power density (W/cm2)

Q Charge (Coulombs)

R Universal gas constant

2 Ra Resistance due to adsorption ( W/cm )

2 Rct Charge transfer resistance ( /cm )

2 Rsoln Electrolyte solution resistance ( /cm )

T Temperature (oC)

Vcell Cell potential (V)

VE Voltaic efficiency

ΔW Vertical deviation from the ideal solution line on the Walden plot

σ Conductivity of the electrolyte (mS/cm)

η Viscosity (cP)

ηa Activation overpotential (V)

ηohm Ohmic overpotential (V)

XXV

Abstract Fundamental Flow Battery Studies:

Electrodes and Electrolytes Abstract

MALLORY A. MILLER

Redox flow batteries (RFBs) are flexible in design due to de-coupling of power and energy, thus making them attractive for large scale energy storage applications. This technology is capable of interfacing with renewable energy sources and provides an alternative solution to balancing power consumption and generation. Despite their advantages, RFBs have not yet been widely commercialized. This work addresses two major issues with RFB systems. First, new non- aqueous based electrolytes were explored and developed. This approach is based on low cost ionic liquids that are safe, non-volatile, and allow for high concentrations of active species. The electrochemistry of multi valent iron species has been examined in a deep eutectic solvent (DES) system. The iron redox reactions reactions appear to be fast and reversible, however, RFB performance is limited by the sluggish kinetics of the iron electroplating reaction as well as poor deposit quality. Thus, iron deposition was studied in detail and an understanding of the roles of ionic species as well as the deposition mechanism is discussed. Second, novel carbon fiber microelectrodes (CFME) were developed as a way to isolate electrokinetics effects occurring within RFB electrodes. This enabled exploration of effects of surface treatments on reaction kinetics in a controlled fashion. The all-vanadium chemistry is on the way to commercialization and was thus chosen for this study. This work shows that the addition of surface oxides improves

2+ 3+ 2+ + the kinetics of the V /V reaction by an order of magnitude and hinders VO /VO2 by a factor of four, which contrasts most contemporary views within the literature.

XXVI

Chapter 1. Introduction and Background

1.1 Redox Flow Batteries

Renewable energy sources such as wind and solar have been gaining interest due to the potential for clean energy and a decreased dependence on fossil fuels. Unfortunately, in 2015 the USA relied on renewable sources for only ~13.4% of its total primary energy supply [1]. Due to intermittency, there is a need for a viable energy storage system with load-levelling capabilities.

The development of such storage technologies is a key challenge facing large-scale implementation of renewables. Batteries offer robust and reversible energy storage along with the load-leveling capabilities necessary to facilitate grid penetration of stochastic renewables. Unlike pumped hydro, batteries can be used in virtually any location.

Redox flow batteries (RFB) are similar to traditional batteries in that there is a direct conversion of chemical to electrical energy. RFBs potentially offer a lower cost energy storage solution that is flexible in design due to separation of power and energy. Unlike conventional batteries, dissolved reactive species are stored in tanks and flowed through electrodes in an electrolytic stack, where the electrochemical reactions occur, allowing scalability (Figure 1-1). The energy storage capacity is thus dependent on the size of the reservoir. Load-leveling capabilities allow this technology to interface with renewable energy sources and provide an alternative solution to balancing power consumption and generation. RFBs have relatively high efficiencies and are safe due to their simplified thermal management. Several reviews can be found [2-6].

Figure 1-1. Schematic of an aqueous all-vanadium redox flow battery (upon charge). Reactive species are dissolved

in an electrolyte and stored in external reservoirs, then pumped though the electrolytic stack, where the reactions

occur; at the negative electrode, the reactive species are reduced and at the positive electrode they are oxidized on

charge. The negative and positive half cells are often separated by an ion-selective membrane or microporous

separator.

Various RFB chemistries have been explored including the iron-chromium [7] systems and the all-vanadium chemistry [8-10], the latter being developed on a large-scale. In order to prevent complications such as permanent capacity loss, a single active species with three or more oxidative states is desired. The aqueous all-vanadium redox flow battery represents the most widely studied system over the past 30 years [8, 11, 12]; however, more cost efficient chemistries such as iron

[13], copper [14], and metal-free systems utilizing quinones [15] are also under investigation.

Hybrid redox flow batteries such as the all-iron [13], all-copper [14], and zinc halide [16] systems involve metal deposition/dissolution at one electrode which limits the amount of energy that can be stored (storage capacity is related to the amount of metal electrodeposited within the stack).

The use of flowable carbon slurry electrodes [17] is under investigation as a means to resolve this

2 limitation as metal can be electrodeposited onto the particles and stored in the tanks external to the stack.

Aqueous based systems are not only limited to operation within 0 – 100 oC, but also by the thermodynamic potential window of water (0 V to 1.229 V vs NHE). In recent years, non-aqueous electrolytes have gained interest due to the possibility of enhancing energy and power densities of current RFB technology. With the use of polar aprotic solvents, operating voltage ranges can be increased. For example, the potential range for acetonitrile with the addition of tetrabutylammonium salts (supporting electrolyte) is −2.9 to 3.4 V vs SCE [18]. Anhydrous acetonitrile (ACN) and propylene carbonate (PC) are widely used within non-aqueous electrochemical systems due to their wide potential windows and high dielectric constants

(allowing them to dissolve organic salts) [19-21]. Organic non-aqueous electrolytes have the ability to support cell reactions with potentials > 5 V, more than four times larger than the thermodynamic potential window of water [18, 22, 23]. This is significant, power density (P) and energy density (Ê) scale linearly with potential (Vcell) as shown below in Equations 1 and 2 where n is the number of electrons transferred, F is Faraday’s constant, C is concentration of active species, and i is the current density. In addition to higher energy density, the increased voltage allows for larger tolerance of overpotentials (i.e. more energy needed to drive the reaction than thermodynamically expected) for the same voltaic efficiency. Furthermore, an absence of water results in elimination of the parasitic hydrogen evolution reaction, thus increasing coulombic efficiency and pH stability.

1 (1) Ê = ⁄2 푛 퐹 퐶 푉푐푒푙푙

푃 = 푖 푉푐푒푙푙 (2)

General non-aqueous redox reactions often includes a transition metal (M) complexed with a ligand (L), such as shown below in Equations 3. There are several reports of non-aqueous redox flow battery systems that utilize chemistries based on ruthenium [19, 24, 25], nickel and iron [26,

27], cobalt [28], chromium [29], manganese [30], uranium [31], and vanadium [32-36].

(z-y)+ - z+ [MLn] + ye  [MLn] (3a)

(z+x)+ z+ - [MLn]  [MLn] + xe (3b)

Presently, organic systems utilizing vanadium (III) acetylacetonate (V(acac)3) in acetonitrile

(ACN) (Figure 1-2) as the active metal species represents the most widely studied non-aqueous system [32-34, 36-39]. With use of cyclic voltammetry, two dominant redox couples are reported which correspond to the oxidation and reduction of V(acac)3. Other characteristics of the non- aqueous organic vanadium system include a large thermodynamic potential of 2.2 V, much larger than the 1.4 V typically observed in the aqueous vanadium system [32].

Figure 1-2. The vanadium (III) acetylacetonate metal-ligand complex

More recent work has focused on functionalization of the acetylacetonate ligands in order to enhance solubility of the active metal complex. Other, more exotic chemistries are currently being explored and have been demonstrated in small-scale RFBs such as polyoxometalates [40] and a lithium organic (hybrid) system using TEMPO (2,2,6,6-tetramethylpiperidine-1-oxyl), a radical compound which has a solubility of up to 5.2 M in the EC/PC/EMC solvent [41, 42]. Other all- organic RFB chemistries such as DBBB (2,5-di-tert-butyl-1,4- bis(2-methoxy ethoxy)benzene) and molecules derived from quinoxaline have been investigated as potential active materials [42].

1.2 Specific Motivations and Scope

1.2.1 Ionic Liquids - A Different Approach to Non-Aqueous

Organic solvents have safety hazards due to their volatile and flammable nature. Contamination by trace amounts of moisture or oxygen can have detrimental effects on battery performance [36].

For example, a fade in capacity which is associated with side reactions of active species with trace amounts of water and oxygen has been reported for small scale flow cells utilizing the V(acac)3 chemistry [39, 43, 44]. Furthermore, the solubility of active species in these types of electrolytes are often very low, limiting the energy storage capacity [45]. V(acac)3 in acetonitrile is limited to

0.59 ± 0.02 M in the absence of the supporting electrolyte [33]; this value is lowered when supporting salt, such as TEABF4 is added to the electrolyte as show in Figure 1-3. Current work is ongoing to improve solubility via functionalization of the ligands as well as using mixed solvent systems [35]. Despite the wide potential windows and fast electrode kinetics of many non-aqueous systems, limitations such as low electrolyte conductivity, low active species solubility, poor stability of the active species as well as high volatility of the organic solvents must be overcome prior to commercialization.

o Figure 1-3. Maximum solubility of V(acac)3 in acetonitrile as a function of supporting salt concentration at 25 C.

Figure is reprinted from reference [46].

Ionic liquids (IL) offer a different approach to increase energy and power densities of current non- aqueous RFB technologies. They possess wide electrochemical windows similar to organic solvents (as shown by the CVs in Figure 1-4) but with the advantages of high thermal stability and low volatility [47-49]. High solubility of transition metals (> 6 M) can be achieved when the electroactive material becomes part of the ionic liquid; one such example is the

+ [Cu(MeCN)4][Tf2N] system where copper is part of the cation as such, Cu(MeCN)4 [50-53].

Thus, this research has a focus on applying a potentially very low cost ILs as an electrolyte for energy storage applications [54-60]. A type of moisture stable ionic liquids called “deep eutectic solvents” (DES) are formed by combining organic halide salts such as choline chloride with a hydrogen bond donor, such as ethylene glycol or urea [61, 62]. DES can be prepared and used under ambient conditions and are an order of magnitude less expensive than traditional ILs; this is attributed to their ease of synthesis and wide availability of the non-toxic pre-cursor materials.

DES represent interesting electrolytes for flow battery applications and are explored in further detail in Chapter 2. The all-iron chemistry is used, and its feasibility is demonstrated in a small scale flow cell. An understanding of the effect of speciation on the physical and electrochemical 6 properties is discussed in detail. One peer-reviewed article has been published on this work with another manuscript recently submitted to the Journal of the Electrochemical Society [52, 63].

1 M H2SO4 (water)

0.1 mA/cm2 Ethaline

Density Reline

0.1 M TEABF4 in ACN Norm Current Current Norm

-3 -2 -1 0 1 2 Applied Potential (V vs Ref)

Figure 1-4. Potential windows for various electrolytes shown on a platinum disk electrode at 100mV/s and 25oC.

Cyclic voltammetry of non-aqueous electrolytes were performed under a dry N2-atmosphere; 0.1 M TEABF4 in

acetonitrile (ACN) represents a typical organic electrolyte and reline and ethaline are two deep eutectic type ionic

liquids. These electrolytes will be discussed in detail in chapter 2.

1.2.2 Electrode Kinetics

Non-aqueous electrolytes offer the potential for enhancing energy and power densities of RFBs.

A major drawback of current systems is slow reaction kinetics which lead to large overpotentials.

Energy efficiency of a flow battery depends on both coulombic and voltaic efficiencies. Voltaic inefficiencies are the result of sluggish electrode kinetics. Furthermore, the larger the overpotential, the greater the risk of undesired side reactions (i.e. electrolyte breakdown). Thus, in order to improve RFB performance, it is important to assess factors that influence electrokinetics and to understand the factors that might effect reaction rates.

A full understanding of electrokinetics at carbon/graphite electrodes has yet to be demonstrated.

Redox flow batteries employ high surface area carbon felt electrodes. Due to the 3-dimensional structure of the carbon felt electrodes, complex current distributions make kinetic analysis difficult. Thus, a new technique to study reaction kinetics is discussed in detail in Chapters 3 and

4. Chemical modifications via thermal treatments, acid oxidation, or electrochemical oxidation are believed to enhance electrochemical activity [64-73]. Concrete knowledge of the effects of surface modification on reaction kinetics will allow for optimization of electrodes, a reduction of the overall cost, and extended lifecycles for RFB applications.

Since the all-vanadium chemistry represents the most widely studied RFB system, this will be used as an example application for electrokinetic studies. Within the literature, the presence of oxygen containing functional groups on the surface of carbon electrodes has been reported by some groups to enhance kinetics [64-68, 70, 72, 74] while others have reported that kinetics are hindered [75,

2+ + 2+ 3+ 76]. One group reported that kinetics for the VO /VO2 are faster than the V /V reaction at a plastic formed carbon electrode, but found the opposite when using pyrolytic graphite [77]. Values

2+ + reported for the VO /VO2 reaction differ in orders of magnitude between different research groups (Table 1-1). It is clear from the literature that the type of carbon as well as the preparation of the carbon electrode surface affect kinetics, however, a full understanding is lacking.

The development of a new tool, a carbon fiber microelectrode, to understand electrokinetics at flow battery electrodes is described in Chapter 3. This is demonstrated using the aqueous vanadium reactions in Chapter 4 and subsequently a reaction mechanism is proposed, which addresses the controversy within the literature. Four peer-reviewed articles have been published on this work; these papers are in collaboration with Professor Buckley’s group at the University

8 of Limerick in Ireland [78-81]. I am the first author on one of these papers, and I made major contributions related to the technique and interpretation of data of the other articles.

2+ + Table 1-1. Kinetic rate constants for the VO /VO2 reaction at carbon electrodes. Rate Constant, ko Electrode Material (cm/s) Ref. Glassy Carbon 8 ×10-4 [11] Glassy Carbon 6 ×10-5 [82] Graphite 8 ×10-7 [83] Pyrolytic Graphite 1 ×10-4 [82] Carbon Composite 9 ×10-4 [82]

1.3 Concluding Remarks

The main conclusions from this work will be summarized and discussed in Chapter 5.

Furthermore, future work will also be briefly discussed. This will touch on other potential redox couples in an ionic liquid type electrolyte as well as exploration of different halides and hydrogen bond donors. Details as well as preliminary data for these experiments are located in the Appendix.

1.4 Publications

Manuscript submitted to the Journal of the Electrochemical Society, Iron Electrodeposition in

Ionic-Liquid Electrolytes for Flow Batteries; Mallory A. Miller, Jesse S. Wainright, and Robert

F. Savinell

Manuscript submitted to the Materials Research Society Advances, Influence of Pretreatment on

Kinetics at Carbon Electrodes and Consequences for Flow Battery Performance; D. N. Buckley,

A. Bourke, R. P. Lynch, N. Quill, M. A. Miller, J. S. Wainright, and R. F. Savinell

Kinetic Study of Electrochemical Treatment of Carbon Fiber Microelectrodes Leading to In-situ

Enhancement of Vanadium Flow Battery Efficiency; Mallory A. Miller, Andrea Bourke, Nathan

Quill, Robert Lynch, D. N. Buckey, Jesse Wainright, and Robert Savinell. Journal of the

Electrochemical Society, 163 (9) A2095-A2102 (2016).

Novel Iron Ionic Liquid Electrolytes for Redox Flow Battery Applications; Mallory A. Miller,

Jesse S. Wainright, and Robert F. Savinell. Journal of the Electrochemical Society, 163 (3) A578-

A579 (2016).

An Investigation into Factors Affecting the Iron Plating Reaction for an All-Iron Flow Battery;

Krista L. Hawthorne, Tyler J. Petek, Mallory A. Miller, Jesse S. Wainright, and Robert F.

Savinell; The Journal of The Electrochemical Society, 162 (1) A108-A113 (2015)

Effect of Cathodic and Anodic Treatments of Carbon on the Electrode Kinetics of VIV/VV

Oxidation-Reduction; A. Bourke, M. A. Miller, R. P. Lynch, J. S. Wainright, R. F. Savinell, and

D. N. Buckley; The Journal of The Electrochemical Society, 162(8):1547-1555, (2015)

Electrode Kinetics of Vanadium Redox Flow Batteries: Contrasting Response of VII/VIII and

VIV/VV to Electrochemical Pretreatment of Carbon; Andrea Bourke, Mallory A. Miller, Robert

Lynch, Xin Gao, James Landon, Jesse Wainright, Robert Savinell, and D. N. Buckey; The Journal of the Electrochemical Society, 163(1): 5097-5105 (2015).

Electrode Kinetics in All-Vanadium Flow Batteries: Effects of Electrochemical Treatments; A.

Bourke, M. A. Miller, R. P. Lynch, J. S. Wainright, R. F. Savinell, and D. N. Buckley, ECS

Transactions, 66 (8) 181-211 (2015).

Chapter 2. Prospect of Ionic Liquids as Redox Flow Battery Electrolytes

2.1 Introduction and Literature

2.1.1 Ionic Liquids

Classical electrolyte solutions are created by dissolution of salts in a solvent. However, solutions may exist that contain only salts and no solvent. Ionic liquids (ILs) are ion pairs (contain an anion and cation) with frustrated lattices, enabling them to be liquids below 100 oC. ILs can exist at room temperature and below. Room temperature ionic liquids date back to the late 1800’s [84], however, recently there has been a renewed interest due to their unique and tunable properties.

Physical and chemical properties can be tuned by using different anions or cations or by functionalizing the ions [85]. ILs often possess high intrinsic conductivity, thus eliminating the need for a supporting electrolyte, making them attractive for electrochemical applications [86].

Large thermodynamic potential windows bracketing electrolyte/solvent decomposition makes them attractive for electroplating or depositing metals [87].

In the 1980s most of the IL electrochemistry focused on chloroaluminates complexed with pyridinium and imidazolium cations [88], however, these mixtures had limited applications due to their reactivity with trace amounts of moisture (resulting in HCl gas formation). In 1992 room- temperature ILs (RTILs) that employed hexafluorophosphate and tetrafluoroborate anions transformed IL research, which allowed for the formation of air and water stable mixtures, rendering them useful as replacements for environmentally-hazardous, volatile organic solvents

[89]. Subsequent research on ILs focused on these types of mixtures. Drawbacks for ILs as a solvent for large-scale energy storage devices are limited availability, high cost, as well as

12 relatively high viscosities. Thus another class of non-aqueous solvents called “deep eutectic solvents” (DES) are being explored [53, 90, 91].

Deep eutectic solvents are defined as a type of ionic liquid which forms a eutectic with a melting point that is much lower than either of its individual components due to hydrogen bonding [90].

DES are systems with the ability to form neutral, ionic, and H+ bonded complexes; an example of the dominant species is provided for the choline chloride and urea system (Figure 2-1). In this system, the quaternary ammonium salt, choline chloride, is complexed (or has a strong interaction) with the hydrogen bond donor, urea, as such:

+ - [Choline]Cl + 2 urea ↔ [Choline] + [Cl•Urea2] (1)

In general, the hydrogen bond donors (HBDs) complex with the halide ion resulting in a delocalization of charge of the anion which decreases the freezing point of the resulting mixture

[92]. DES are derived from readily available chemicals such as choline chloride and urea, which are simply combined, heated to approx. 150 oC and stirred until a transparent liquid is formed

(Figure 2-2).

Figure 2-1. Possible complex formation in a 1:2 choline chloride:urea deep eutectic solvent. Figure was obtained

from reference [93].

(a) (b)

Figure 2-2. A mixture of choline chloride and urea in a 1:2 molar ratio is shown (a) prior to heating and (b) after being placed in a vacuum oven at 150 oC for four hours and cooled back down to room temperature. The image on

the right hand side is the resulting deep eutectic, reline, and is shown at room temperature.

It has been reported that complexing hydrogen bond donors with salts of symmetrical cations such

o as NH4Cl and NMe4Cl do not form liquids below 200 C, however, by decreasing the symmetry,

o freezing points were significantly decreased (ex. Urea and NMe3EtCl freezes at 53-55 C) [53].

Furthermore, additional functional groups could be added to further decrease freezing points, however, the role that these functional groups play has not yet been determined. From the

o literature, choline chloride ([NMe3C2H4OH]Cl, or ChCl) has the lowest freezing point (23 C) and was thus recommended for further study [53].

A list of melts composed of choline chloride and different alcohols, their complexing agent (or

HBD) and their resulting freezing points are shown in Table 2-1 (data from [94]). Other early melts include DES with quaternary ammonium salts with hydrogen bond donors such as amines and carboxylic acids or metal chlorides [61, 90, 95]. Some commonly used HBDs are shown in

Figure 2-3. Possible applications for DES include purification of biodiesel, metal cleaning prior to electroplating, electropolishing and electroplating of zinc [48, 87, 94, 96-101]. These ionic melts not only have a wide potential window, but also readily dissolve chloro metallic compounds, making them extremely useful for electrodeposition of metals and alloys [96].

Table 2-1. Common DES formed by mixing choline chloride with various hydrogen bond donors in a 1:2 molar ratio [94].

Complexing Freezing Product Agent/HBD Point (oC) Reline Urea (Amide) 12 Oxaline Oxalic Acid (Acid) 13 Maline Malonic Acid (Acid) 3 Ethaline Ethylene Glycol (Alcohol) < -13 Glyceline Glycerol (Alcohol) -20

Figure 2-3. Various hydrogen bond donors which when mixed with choline chloride for DES.

DESs have lower viscosities than most typical ILs, are often biodegradable, recyclable, non-toxic, and non-reactive with water [62, 102, 103]. Furthermore, DES tend to be more cost efficient than other non-aqueous electrolytes, with their ease of synthesis also an advantage. Organic solvents

15 such as acetonitrile range in price of about $2-3/kg without the addition of supporting electrolyte; whereas typical ionic liquids, for example 1-ethyl-3-methylimidazolium trifluoromethanesulfonate, are much more expensive at $180-200/kg [104]. The cost to prepare deep eutectic solvents is significantly lower: choline chloride is approximately $2/kg and urea and ethylene glycol are around $0.30/kg [105, 106]. These materials are already manufactured on a large scale; choline chloride is used as an additive in chicken feed, urea is used in fertilizers, and ethylene glycol is used in the fabric industry as well as in automobiles (“antifreeze”) [107]. The potential for low cost thus makes DES attractive for use in large scale electrochemical applications, such as redox flow batteries (RFB). Recent reviews on ionic liquids and DES can be found [62,

94, 108-114].

2.1.1.i Ionic Liquids for Electrochemical Applications

The properties of ionic liquid electrolytes, their electrochemical windows as well as other properties, are compared to aqueous and organic electrolytes in Table 2-2. From the table it is clear that the organic and the ionic liquid electrolytes have wide potential windows compared to the typical aqueous electrolyte. It should be noted that upon the addition of supporting salts, such as TEABF4 or TEAPF6, electrolyte breakdown occurs at less extreme potentials. Ionic liquids and the organic electrolytes ACN, DMF, PC, γ-Butyrolactone, and DMSO provide opportunities for wider operating temperature ranges compared to a typical aqueous electrolyte. In addition, the non-volatile and non-flammable nature of these ILs (and DES) adds attractive advantage over the organic solvents from a safety standpoint.

Fluid properties such as conductivity and viscosity are important for electrochemical applications.

Ionic liquids are quite viscous, however, the DES type ILs have lower viscosities which is

16 beneficial for redox flow battery systems where the electrolyte must be pumped. When compared to the acidic aqueous electrolyte, conductivities are quite low for the ionic liquid and organic electrolytes. Acidic (or basic) aqueous electrolytes contain H+ and OH- ions that may participate in a sort of “hopping” mechanism (Grotthuss mechanism) for transport. This allows for unusually fast ion conduction rates [115]. A similar conduction mechanism (relating to “halide hopping”) is believed to occur in certain ILs (for example SnCl2 and Urea [116], other examples can be referenced [86, 117, 118]), however, since this lies outside the scope of this study, they will not be discussed in detail here.

Table 2-2. A comparison of aqueous vs. non-aqueous electrolyte-solvent systems. Density, viscosity, conductivity are reported for electrolytes at 25 oC with the exception of reline and ethaline (reported values are at 80 oC).

Solvent Decomposition Potential Temperature Density Viscosity Conductivity Solvent Material Ref. Voltage at Pt (V vs SCE) Window (V) Range (oC) (g/mL) (cP) (mS/cm)

[119] *This Water 0.24 to 1.47 1.2 0 100 1.00 1.00 48.5b work ACN −2.9 to 3.4 6.3 (4.3 with −46 82 0.78 0.34 49.6c [46, 119-121] TEABF4) *This work DMF −3 to 1.6a 4.6 −61 153 0.94 0.80 22.8c [120, 121] PC −3 to 3.6a 6.6 −49 241 1.19 2.53 10.6c [119-122] γ-Butyrolactone −3 to 5.2a 8.2 −43 202 1.13 1.75 14.3c [120] DMC −2.8 to 3.3a 6.1 4 90 1.07 0.59 2.0d [122] DMSO −2.9 to 1.5a 4.4 18.6 189 1.10 2.24 13.9c [46, 120, 121]

Imidazolium ILs [114, 123- + − e [BuMeIm] [PF6] −2.3 to 3.4 5.7 −8 391 1.36 312 1.4 126] + − f [EtMeIm] [N(CF3SO2)2] −2 to 2.5 4.5 22 436 1.52 34 8.8 [114, 123] Reline (80oC) −1.5 to 1.1g 2.6 12 1.17 28.1 9.7 *This work Ethaline (80oC) −1 to 1.1g 2.1 −66 1.09 9.1 28.1 *This work a Glassy carbon working electrode. b With addition of 0.1 M H2SO4; note that at operating conditions for aqueous system (2 M H2SO4), system conductivity is around 700 mS/cm [127]. c With addition of 0.65 M TEABF4 d With addition of 0.65 M TBABF4 e Potential is referenced vs a platinum wire pseudo reference. f - - Potential is referenced vs I /I3 g Potential is referenced vs (non-aqueous) Ag/AgCl

Ionic liquids and deep eutectic solvents offer a different approach to increase energy and power densities of current non-aqueous redox flow battery (RFB) technologies. As shown above in Table

2-2 they possess wide electrochemical windows similar to organic solvents but with the advantages of high thermal stability and low volatility. High solubility of transition metals (> 6 M) can be achieved [52]. Thus, this research has a focus on applying DES as an electrolyte for energy storage applications.

2.1.1.ii Ionic Liquids for Flow Batteries

Recently, there has been a growing interest for utilizing ILs (and DES) as flow battery electrolytes

(examples from the literature can be found in Table 2-3). An all-vanadium system utilizing

V(acac)3 was among the first chemistries explored in DES [57, 60]. Despite poor solubility (<0.25

M), they reported 12 charge/discharge cycles for V(acac)3 with 0.5 M TEABF4 in ethaline using an H-cell with an anion-exchange membrane [60]. Coulombic and energy efficiencies were 49-

52% and 25-31%, respectively. Ferrocene as well as other metallic acetylacetonate salts had been investigated as alternatives (chromium, manganese, cobalt, copper, nickel, iron, zinc) in quaternary ammonium and phosphonium based DES using cyclic voltammetry (CV) along with solubility studies [58-60]. However, the metallic acetylacetonate species have poor solubility (<2 mM) with the exception of manganese and zinc and irreversible or quasi-reversible kinetics at 50 oC. Future research is ongoing to study effects of complexing agents such as cyanide or ammonia for enhancement of solubility of the active species.

The all-copper (chloride) chemistry been demonstrated in an ionic liquid RFB [128]. In the ionic

+ liquid RFB, the copper was part of the cation (Cu(MeCN)4 ) enabling high concentrations of active species (>3 M). However, it was speculated that the dendritic and non-adherent nature of the

18 electrodeposited metal in addition to crossover at the membrane are the cause of low energy efficiencies as determined from charge/discharge experiments in an H-type cell (10-30%).

The all-copper chemistry was also studied in a DES type electrolyte (ethaline) and demonstrated in a small-scale flow cell [54]. However, due to mass transport limitations and poor conductivity of the electrolyte, energy efficiencies of only 52% and 62% at current densities of 10 and

7.5 mA/cm2 respectively had been achieved. This group also explored the possibility of iron chloride and zinc chloride and obtained an energy efficiency of 78% (at 0.5 mA/cm2) [55]. The

Fe2+/Fe3+ redox reaction was used as the positive couple with zinc deposition and dissolution at the negative electrode; electrodeposition of iron as the negative electrode reaction had been abandoned due to its sluggish kinetics, whereas zinc is known to exhibit fast electrodeposition and dissolution kinetics.

From the literature it is clear that ILs have demonstrated promise for future energy storage applications through enhanced potential windows for higher power density and the ability to operate under ambient conditions as well as a wide range of temperatures. However, electrodeposition in these hybrid type RFB systems has not been widely studied. There has been little investigation of the Zn, Fe, or Cu electrodeposition reactions or the factors that affect electrodeposition. Furthermore, as with most non-aqueous systems, conductivity must be improved in order to decrease ohmic overpotentials and enhanced fluidity is desired to eliminate mass transfer limitations. In addition, for higher energy density systems, higher solubility must be achieved.

Table 2-3. A summary of flow battery systems that utilize ionic liquids.

Solubility of Cell C/D current Operating Redox IEM Efficiency Electrolyte Active Potential Cell type Electrode density Temperature (oC) Ref. System (IEM Type) species (M) (V) (mA/cm2) %CE %VE %EE Morgane 10- [Cu(ACN) ][Tf N] >3 0.9 H-Type Cell 2 mA (Area?) 90 [128] 4 2 Membrane 30 1:2 Choline Flow type Jellified 10 52 [CuCl ]2- >0.5 0.72 Carbon cloth 50 [54] 4 Chloride:Ethylene Glycol cell Separator 7.5 97 62 1:2 Choline Side-by-side Tungsten Jellified [FeCl ]- >0.5 1.2 0.5 50 37 [55] 4 Chloride:Ethylene Glycol Cell Wire Separator - [FeCl4] / 1:2 Choline Side-by-side Tungsten Jellified - >0.5 1.1 0.5 50 95 78 [55] [ZnCl4] Chloride:Ethylene Glycol Cell Wire Separator 1:2 Choline AMI-7001S Graphite 0.1/0.01 49- 25- V(acac) Chloride:Ethylene Glycol + <0.25 2 H-Type Cell Anion- [60] 3 Rods (Chg/Dchg) 52 31 0.5 M TEABF4 Exchange Mem

2.1.1.iii Iron Electrochemistry - an Introduction

For the purpose of understanding plating and stripping as well as redox reactions in DES, the all- iron system is chosen. The all-iron chemistry is an attractive alternative to the all-vanadium system initially proposed in 1981 by Hruska and Savinell [13]. Similar to the all-vanadium system, the availability of multiple oxidation states allows the all-iron system to use a single active species.

Thus, crossover can be dealt with in a more simplified manner; electrolytes can merely be remixed and brought back to their original states. The all-iron flow battery chemistry consists of the

Fe2+/Fe3+ redox reaction as well as the electrodeposition/dissolution reaction (Fe2+/Fe0) at the negative electrode (Figure 2-4). Iron is the fourth most abundant material with ores containing approximately 70% Fe, allowing for simple processing [129]. Due to its domestic availability, iron is more than 250 times less expensive than vanadium (assuming an ore grade of 50% for vanadium) [130-132]. Furthermore, vanadium is fairly toxic whereas iron is considered to be environmentally benign.

Figure 2-4. Schematic of iron flow battery reactions at carbon felt electrodes upon charge.

2.1.1.iv Physical Properties of Deep Eutectic Solvents (DES)

Conductivity of ionic liquids can be enhanced dramatically with an increase in temperature (and viscosity can be decreased). All conductivities were determined using a conductivity probe (Fisher

Scientific, Accumet) with a known cell constant (“푙/퐴”). Measuring the AC impedance of the circuit, the solution conductivity can be determined using the following relationship:

푙 휎 = (2) 푅퐴

Where σ is the conductivity (S/cm) and R is the measured resistance (Ω). Viscosity of ethaline and reline was determined using a rheometer (Brookfield). The viscosity for all other electrolytes was determined using a Cannon SimpleVIS glass viscometer (size 2, 10-100 centistokes with a cell constant of 0.1045 cSt/s) which was submersed inside a thermostatted mineral oil bath inside a nitrogen-filled glovebox; all values represent an average of three trials. Conversion from centistokes to centipoise was done using the densities. Densities were estimated by simply weighing a set volume of electrolyte (extracted after it had been held at the desired temperature for a minimum of 20 minutes); all densities are an average of 5 measurements. Temperatures had been allowed to equilibrate for a minimum of 20 minutes prior to making any measurements.

Figure 2-5 shows plots of conductivity and viscosity as a function of temperature for two DES,

“ethaline” (1:2 M ChCl:EG) and “reline” (1:2 M ChCl:Urea). Unlike the typical aqueous electrolytes, the DES electrolytes have an Arrhenius type dependence on temperature, which is described in the literature [133]1. This originates from the idea that the conductivity and viscosity

1It should be noted that typical ionic liquids exhibit a Vogel-Tammann-Fulcher type relationship in the form σ = A exp(B/[T−To]), where A is constant and proportional to the number of charge carriers, To is also constant and related to the glass transition temperature, and B is a fitting parameter related to activation energy, Ea = R×B, where R is the ideal gas constant [134] A. Angell, "Oxide Glasses in Light of the “Ideal Glass” Concept: I, Ideal and 22 are related to the free volume and the probability of finding holes large enough for the ions to move into [135]. This probability increases with temperature due to an increase in molecular vibrations. Thus, by increasing the temperature to 80 oC, the conductivity of reline increases by a factor of ~12 (0.81 compared to 9.7 mS/cm) and viscosity decreases by a factor of ~60 (1648 and

28.1 cP respectively). Furthermore, the hydrogen bond donor (HBD) is shown to influence fluid properties as well. The two DES compared differ in that the HBDs are an amide (urea) versus an alcohol (ethylene glycol). This affects the conductivity by a factor of ~2 and the viscosity by two orders of magnitude.

30 1800 60 Reline Ethaline Ethaline

24 Reline 1500 50 /cm) 1200 40 mS 18 900 30 12 600 20

Conductivity ( Conductivity 6

Viscosity of Viscosity Reline (cP) 300 10 Viscosity Viscosity of Ethaline(cP) 0 0 0 25 35 45 55 65 75 85 95 15 25 35 45 55 65 75 85 Temperature (oC) o Temperature ( C)

Figure 2-5. (a) Conductivity and (b) viscosity as a function of temperature for two deep eutectic solvents, “ethaline”

and “reline”.

Figure 2-6 shows an example of how changing the ratios of the precursor materials or the hydrogen bond donor or even the halide can alter solution conductivity. This is not widely studied within the literature and will be explored in further detail to gain an understanding of what factors affect

Nonideal Transitions, and Departures from Ideality," Journal of the American Ceramic Society, vol. 51, no. 3, p. 117, 1968. 23 fluid properties as well as the electrochemistry of metals in these types of electrolytes. This work will focus on electrolytes containing choline chloride and ethylene glycol due to their superior conductivity and viscosity compared to other electrolytes tested. The effects of the halide as well as other hydrogen bond donors is explored in further detail in the Appendix A1.

35 1:4 ChCl, EG 30 2:4 ChCl, EG 2:4 ChBr, EG 25 2:4 ChCl, Urea

Conductivity(mS/cm) 10

0 0 25 50 75 100 125 o Temperature ( C)

Figure 2-6. Conductivity as a function of temperature for DES type ILs synthesized from different precursor

materials.

This work explores all-iron non-aqueous based systems employing DES type ILs with the anticipation that the insights gained will provide a fundamental framework and design principles for similar systems. Ethaline (1:2 molar ratio of choline chloride and ethylene glycol) is a starting point due to its wide electrochemical window, relatively low viscosity, and high conductivity compared to other DES and ILs. Figure 2-7 shows initial CVs for the iron reactions in ethaline.

The open circuit voltage for the system is maintained (1.21 V). Furthermore, the electrolyte breakdown does not occur until around −1.2 V vs Ag/AgCl which is an added benefit since it is negative of the iron plating/stripping reaction. Therefore, high coulombic plating efficiencies are expected for these systems; this will be discussed in further detail in this chapter.

Ionic liquids are often referred to as “tailorable electrolytes” where both the physical and electrochemical properties can be tuned when the molar ratios of the precursor materials are altered. This work focuses on understanding factors that influence iron deposition as well as the nature of iron species present and how speciation is responsible for observed changes in the physical properties of the electrolytes. One peer-reviewed article has been published on this work

[52] with another manuscript recently submitted to the Journal of the Electrochemical Society

[63].

6 ) 2 3

-3

-6 1.21 V

Current Density (A/m Density Current -9

-12 -1.5 -1 -0.5 0 0.5 1 Potential (V vs Ag/AgCl)

Figure 2-7. Cyclic voltammetry of 0.51 M FeCl2 in ethaline (a 0.25:2:4 molar ratio FeCl2:ChCl:EG) at a platinum

2+ 0 2+ 3+ o electrode showing the negative (Fe /Fe ) and positive (Fe /Fe ) reactions; 10 mV/s, 50 C, N2-atmosphere. The open circuit voltage (OCV) in a DES is similar to that of the expected OCV in typical aqueous electrolytes, 1.21 V.

2.2 Iron-Containing Ionic Liquid Electrolytes

In order to investigate the effects of electrolyte composition on the electrochemical properties of the iron containing ILs, the molar ratios of the salts were varied. The ILs were prepared by combining iron chloride (FeClx), choline chloride (ChCl), and ethylene glycol (EG) and stirring at

80 oC until all solids dissolved. For the more concentrated iron electrolytes, it took several days

25 for all solids to dissolve. It is important to note that electrolytes synthesized with FeCl2 contain two moles of chloride per mole of iron; additional chlorides come from any added choline chloride.

Electrolytes synthesized with FeCl3 contain three moles of chloride per every one mole iron (again, additional chloride comes from added choline chloride).

All chemicals were analytical grade, used without further purification; iron (II) chloride

(anhydrous), iron (III) chloride (anhydrous), choline chloride, and ethylene glycol (anhydrous) were obtained from Alfa Aesar. Choline chloride had been dried at 150 oC inside a vacuum oven for a minimum of 12 hours. In order to eliminate possible effects due to water or oxygen, experiments were performed in a nitrogen-filled glove box (<15 ppm water). To investigate parameters affecting solution-related properties such as conductivity, kinetics, and nature of the electroplated iron, the molar ratios of the precursor materials were varied. As previously discussed, these electrolytes are quite viscous at room temperature, thus, elevated temperatures are necessary for flow battery applications. All data is shown for the electrolytes at 80 oC, higher temperatures are not desirable for large scale energy storage and would require development of new battery hardware and membranes and thus are avoided for this study.

Examples of different electrolytes are shown in Table 2-4. The concentrations of active species range from 0.2 up to 8.4 M iron which is advantageous in enhancing energy densities. It is important to note that a solution containing iron chloride, choline chloride and ethylene glycol in a 2:0:4 molar ratio can be synthesized using FeCl3, or equal molar FeCl2 and FeCl3. However, a

o 2:0:4 solution synthesized using only FeCl2 was not liquid at temperatures below 100 C and thus would not be desirable for use in a battery. Similarly, solutions containing FeCl3 and choline

o chloride in a 1:2 molar ratio were liquid at 80 C but not the FeCl2 analogue. All other electrolytes were liquid at or below 80 oC, independent of the iron oxidation state, and were studied in detail.

Table 2-4. Examples of the synthesized electrolytes evaluated with corresponding iron concentration. Solutions

containing 1:2:0 and 2:0:4 could not be created using only FeCl2.

Molar Ratio Concentration of c (FeClx:ChCl:EG) Iron (M)

1:2:0 b 8.4 2:0:4 b 6.3 1:0:4 a 3.8 1:1:4 a 2.6 0.5:1:4 a 1.3 1:2:4 a 1.9 0.1:2:4 a 0.2 a Electrolytes synthesized using equal molar FeCl2 and FeCl3

b Electrolytes synthesized with only FeCl3 c Molarity had been calculated based on the number of moles of iron per solution volume assuming that the volume

did not change upon mixing and thus this represents a conservative estimate (since total solution volumes visually

decreased upon mixing).

2.2.1 Iron Electrokinetics

Electrochemical experiments were performed in a thermostatted three electrode glass cell with a potentiostat/galvanostat (Solartron SI 1287) and a frequency response analyzer (Solartron SI 1260) employing a platinum microelectrode or glassy carbon coupon as the working electrode and a graphite rod as the counter electrode. A Ag/AgCl reference was used for all experiments; the reference electrode was constructed by inserting a chloridized silver wire into a fritted glass capillary filled with 1:2 molar ratio choline chloride to ethylene glycol. The protocol for chloridizing the silver wire was followed as described in the literature [23]. Furthermore, this reference electrode configuration has been shown to be stable in these types of electrolytes [96].

It should be noted that ferrocene, a typical internal reference, was not soluble in these electrolytes.

The Nernst equation (equation 3) provides a relationship between the electrode potential and the concentrations of oxidized and reduced species in solution, Co and CR, respectively. However, the

Nernst equation describes the thermodynamics of the system and does not provide insight on the rate at which these reactions occur (electrokinetics). Thus, the Butler-Volmer model (equation 4) is used to correlate how reaction kinetics depend on the applied potential.

푅푇 퐶 (3) 퐸 = 퐸표 + 푙푛 표 푛퐹 퐶푅

푛퐹 푛퐹 표 −훼 (퐸−퐸표) (1−훼) (퐸−퐸표) (4) 푖 = 퐹퐴푘 [퐶표(0, 푡)푒 푅푇 − 퐶푅(0, 푡)푒 푅푇 ]

Combining the Butler-Volmer and the Nernst equation allows for the derivation of a relationship between io, the exchange current and the potential (equation 5). At equilibrium (or open circuit potential) the net current is zero, thus, by evaluating either the forward/oxidation or the reverse/reduction reaction separately, one can obtain io, which allows for quantification of reaction kinetics.

푛퐹 푛퐹 (5) 퐶표(0, 푡) −훼 (휂) 퐶푅(0, 푡) (1−훼) (휂) 푅푇 푅푇 푖 = 푖표 [ 퐵푢푙푘 푒 − 퐵푢푙푘 푒 ] 퐶표 퐶푅

Slow scan (5 mV/s) linear sweep voltammetry (LSV) was used to investigate reaction kinetics at the surface of the working electrode in this case, a platinum microelectrode (Fig 2-8 shows an example of this). Since the kinetics in the linear region are of interest for flow battery applications, the overpotential is small (±20 mV) and thus, the linearized Butler-Volmer kinetic model (ex ≈

1+x) which relates current to overpotential can be used. For a simple electron transfer reaction,

푘 [Oxidized species] + 1 electron ↔ [Reduced species], this relationship is shown to be:

nF (6) 푖 = −푖 ( ) 휂 표 RT

푅푇 (7) 푖 = ∗ 푆푙표푝푒 표 푛퐹

500

) 400 2 300 200 100 0 -100

Current Density (A/m Density Current -200 -300 -1.00 -0.50 0.00 0.50 1.00 η (V)

2+ 3+ o Figure 2-8. LSV (at 5 mV/s) in 1:1:4 (equal molar Fe /Fe ) at a platinum microelectrode; 80 C, N2-atmosphere.

When current is plotted as a function of the overpotential, the slope (- 휂/i) has units of resistance and can thus be thought of as the charge transfer resistance. The exchange current density, io was estimated from the slope as shown by equation 7. The exchange current densities are reported in

Table 2-4. IR-corrections were not made and expected to be very small at these microelectrodes.

As shown by Table 2-5, the kinetics are influenced by the electrolyte composition. Exchange currents do not follow the normal trend (increasing with increasing iron concentration), and this is attributed to the formation of different complexes in each of the different electrolytes which will be discussed in detail in the section entitled Iron Speciation. Overall, the kinetics for the Fe2+/Fe3+ reaction are fast and likely not limiting battery performance. The Fe2+/Fe0 reaction is known to exhibit sluggish kinetics and thus, the effects of electrolyte composition on the plating reaction were investigated in further detail.

Table 2-5. Summary of exchange current densities for the Fe2+/Fe3+ reaction estimated from the slope of the LSV

potential-current density data in the linear region.

o Molar Ratio i 2 (FeClx:ChCl:EG) (A/m at Cref)

2:0:4 5.2 1:0:4 17 1:1:4 94 0.5:1:4 96 1:2:4 203 0.1:2:4 3.9

2.2.2 Electrodeposition of Iron

Cyclic voltammetry was performed in iron containing ILs with varying molar ratios of

FeCl2:ChCl:EG (Figure 2-9). The solutions contained only FeCl2 due to the corrosive nature of the ferric ions on iron deposits. It is important to note that an iron rod was used as a counter electrode for all electrodeposition experiments so as to maintain constant ferrous concentrations within the electrolyte (and not introduce ferric ions at the counter electrode). The CVs for 1:1:4,

0.5:1:4 and 0.1:2:4 are compared. It should be noted that the CVs shown are the 5th scan and represent the steady state response. The current response is similar to what would be expected for

Fe2+/Fe0 in a typical aqueous based electrolyte [136]; iron deposition/dissolution displays sluggish kinetics as indicated by the wide separation between the potential for deposition and the dissolution peaks. The 1:1:4 electrolyte contains 3 chloride ions for every 1 iron ion, the 0.5:1:4 electrolyte contains 4:1 Cl:Fe and the 0.1:2:4 represents a solution containing chloride in excess of iron (Cl:Fe is 22:1). As the chloride to iron ratio is increased (from 3:1 up to 22:1), the onset for dissolution is shifted to a more negative potential and a reduced separation between reduction and oxidation is observed, implying faster kinetics for dissolution of iron. It should be noted that the CVs for

30 the 1:0:4 and 1:1:4 electrolytes overlay (Figure 2-9 inset), indicating that the addition of choline does not largely affect the kinetics for the deposition/dissolution reaction. Furthermore, the CVs for the 0.5:1:4 and 1:2:4 electrolytes agree (both contain 4:1 Cl:Fe) and the CVs for the two electrolytes where chloride is in excess (0.25:2:4 and 0.1:2:4) also agree with one another. Thus, there exists three regimes: that of iron deposited from solutions containing the (curve “c”) low chloride, (a) an excess of chloride, or (b) a region in between where chloride to iron ratio is 4:1.

This indicates that the iron deposition/dissolution reaction depends on the ratio of chloride to iron.

(a) (b) (c)

Normalized Current 1:1:4'

0.5:1:41:0:4' 1:1:4' 0.1:2:4 -1.05 -0.55 -0.05

-1.05 -0.85 -0.65 -0.45 -0.25 -0.05 0.15 Applied Potential (V vs Ag/AgCl)

Figure 2-9. Cyclic voltammetry at a glassy carbon electrode in (a) 0.1:2:4 (b) 0.5:1:4 or (c) 1:1:4 FeCl2:ChCl:EG at

o 50 mV/s, 80 C, in a N2-atmosphere. Current was normalized to allow for a more direct comparison. Inset:

comparison of the 1:1:4 and 1:0:4 solutions.

2.2.2.i Nucleation Mechanism

In order to understand how iron electrodeposits, it is important to first understand the initial stages of deposition and nucleation. As shown by the CVs in Figure 2-9, the electrodeposition of iron

31 changes with the Cl:Fe ratio, thus, it is of interest to compare nucleation cases in which there is an excess of chloride to that of iron deposited from an IL containing a low Cl:Fe ratio.

0.025 0.025

a) 0.005 b) 0.01

)

) 2 2 (풊풎풂풙, 풕풎풂풙) 0.02 0.02

0.0025 0.005

A/cm

- - 0.015 0.015 0 0 0 100 200 300 400 500 0 1 2 0.01 0.01

0.005 0.005

Current Density ( Current Density (

0 0 0 10 20 30 40 50 0 10 20 30 40 50 Time (s) Time (s)

Figure 2-10. Current vs time transients for iron electrodeposition at −1.15 V vs Ag/AgCl in (a) 0.25:2:4 and (b)

1:1:4 FeCl2:ChCl:EG electrolytes.

1.2 1.2 a) b) 1 1

0.8 0.8 2

0.6 2 0.6

max)

max) (I/I 0.4 (I/I 0.4 Progressive Progressive 0.2 Instantaneous 0.2 Instantaneous Experimental Experimental 0 0 0 1 2 3 0 1 2 3 t/tmax t/tmax

2 Figure 2-11. Nondimensional plots (I/Imax) vs t/tmax for experiments of iron electrodeposited at −1.15 V vs Ag/AgCl

in (a) 0.25:2:4 and (b) 1:1:4 FeCl2:ChCl:EG electrolytes. Dashed and dotted lines represent progressive and

instantaneous nucleation, respectively.

The current transient for the case of high chloride (0.25:2:4) in Figure 2-10a shows a typical response where the current decays quickly upon relaxation of the double layer charging, then

32 displays an increase (~0.2 s as shown in the inset, Fig 2-10a) due to growth of independent nuclei

(or simultaneously an increase in the number of independent nuclei) without the effect of overlapping diffusion boundary layers. In Figure 2-10a, this part of the transient is shown to exhibit an increase in current with t2, which is characteristic of a progressive nucleation mechanism [137].

After a certain point, a maximum is reached after the independent nuclei have grown close enough together for the diffusion boundary layers to overlap. At this point diffusion to the surface changes from hemispherical to linear mass transfer. Interestingly the current response for the case of low chloride (1:1:4 electrolyte) does not reach a maximum and does not decay even after 500 s (inset

Figure 2-10b). In fact, the current response appears to be invariant with time (Figure 2-10b). It is important to note that electrolyte breakdown was not observed at these potentials (especially for those containing low chloride concentrations) and thus would not contribute to the current transients.

Nucleation studies have been carried out in ionic liquid media using the theoretical model proposed by Scharifker and Hills [138, 139]. This model describes the two mechanisms for formation of a monolayer: (1) instantaneous where nucleation spreads out on the substrate from nuclei formed at t=0 (ie nucleation rate is fast compared to growth rate) or (2) progressive nucleation, where nuclei appear randomly on the substrate (slow nucleation rate). When there is an excess of chloride, iron electrodeposition was found to follow progressive nucleation as shown in Figure 2-11a. However, for electrodeposition in solutions containing low chloride, nucleation cannot be explained using the traditional S-H model (Figure 2-11b).

The nucleation mechanism may be expected to change as a function of applied potential as well as substrate material. However, iron electrodeposition from solutions containing low chloride exhibited similar nucleation behavior regardless of applied potential (ranging from −0.75 to

−1.35 V, as shown by Figure 2-12a) or substrate material as shown by Figure 2-12b which compares a carbon fiber microelectrode (whose fabrication is discussed in Chapter 3), platinum, and gold electrode. It is also interesting to note that the time invariant deposition behavior is not characteristic of iron electrodeposition or even electrodeposition of metals that exhibit sluggish

2 kinetics. Figure 2-13 shows the nondimensional plots (I/Imax) vs t/tmax for a kinetically facile electrodeposition reaction, zinc, from a 1:1:4 ZnCl2:ChCl:EG electrolyte. Similar time invariant deposition behavior has been reported in the literature for zinc deposition onto a platinum substrate from a choline chloride and ethylene glycol based electrolyte as well as deposition of Ru, Rh, and

Pd onto a stainless steel substrate from a 1-butyl-3-methylimidazolium chloride IL [140, 141]. It has been proposed that the cations and the halides present within the double layer effect the nucleation mechanism, and thus it is important to determine speciation (i.e. the composition of various ionic species present in the electrolyte).

1.2 1.2 a) b) 1 1

0.8 0.8

2 ) 0.6 ) 0.6

-0.75

max max

0.4 -0.95 (I/I (I/I 0.4 -1.05 Carbon Fiber 0.2 -1.25 0.2 Gold -1.35 Platinum 0 0 0 1 2 3 0 1 2 3 t/tmax t/tmax

2 Figure 2-12. Nondimensional plots (I/Imax) vs t/tmax for iron electrodeposited at (a) selected potentials (V vs

Ag/AgCl) and (b) at −1.15 V vs Ag/AgCl onto different substrates. All experiments had been performed in a 1:1:4

FeCl2:ChCl:EG electrolyte. Dashed and dotted lines represent progressive and instantaneous nucleation,

respectively.

1.2

0.8 2

max) 0.6 (I/I 0.4

0.2

0 0 1 2 3 t/tmax

2 Figure 2-13. Nondimensional plots (I/Imax) vs t/tmax for zinc electrodeposited onto glassy carbon at −1.0 V vs

Ag/AgCl in a 1:1:4 ZnCl2:ChCl:EG electrolyte. Dashed and dotted lines represent progressive and instantaneous

nucleation, respectively.

Halide adsorption has been believed to play a significant role in the electrodeposition of transition metals. For example, adsorption of the halide onto a copper substrate (in aqueous media) has been shown to increase in the order Cl- < Br- < I- [142]. Figure 2-14 shows the nondimensional plots for iron deposited from electrolytes containing 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I. It is interesting to note that the time invariant deposition behavior is more pronounced for the electrolyte containing iodide. A more detailed investigation is needed to understand the mechanism but this evidence suggests that the halide is playing a role.

1.2

0.8 2

max) 0.6 I- (I/I Cl- 0.4 Br-

0.2

0 0 0.5 1 1.5 2 2.5 3 t/tmax

2 Figure 2-14. Nondimensional plots I/Imax vs t/tmax for iron electrodeposited at −1.15 V from 1:1:4 FeX2:ChX:EG,

where X=Cl, Br, or I. Dashed and dotted lines represent progressive and instantaneous nucleation, respectively.

From the early stages of deposition (nucleation), it is clear that the mechanism for electrodeposition changes with chloride to iron ratio. This also influences the nature of the electrodeposited iron. As shown by the SEM micrographs in Figure 2-15, iron deposited from an electrolyte containing high chloride was rough and almost dendritic compared to iron deposited from an electrolyte containing low chloride. These rough deposits present an issue for a system where the iron must be reversibly plated and stripped upon charge/discharge of the battery and thus it is of interest to investigate the coulombic efficiency for the deposition/dissolution reaction with solution composition.

a) b)

Figure 2-15. Plating 10 mAh @ 5 mA/cm2 onto a glassy carbon substrate (80 oC) from (a) 0.25:2:4 and (b) 1:1:4

FeCl2:ChCl:EG electrolytes. Initial potential was ~ −1.15 V vs Ag/AgCl for deposition and levelled out at −0.9 and

−1.1 resp. Elemental analysis (energy dispersive spectroscopy) using a Helios NanoLab™ 650 confirmed the

samples to be iron, however, the iron plated from a 0.25:2:4 electrolyte exhibited more oxides (91.9 ±0.4 iron,

6.9 ±0.3 oxygen, and carbon balance compared to 100 ±0 for a deposit from the 1:1:4 electrolyte).

Plating experiments were performed under static conditions on a polished glassy carbon substrate.

For the coulombic plating efficiency experiments, a controlled plating potential was held for 5 minutes, then the deposit was stripped for 5 minutes at a voltage positive of the stripping reaction

(found from CV analysis). Figure 2-16 shows an example of the current vs time response.

Coulombic plating efficiencies had been determined by integrating under the curve to get the charge, QFe, for electrodeposited iron which was compared to the amount of charge from stripping the iron back off the substrate. Thus, if for example, electrolyte breakdown were occurring while metal was being deposited, then QFe, for metal stripped off would be smaller resulting in a lower efficiency. In between consecutive experiments, the working electrode had been held at a positive overpotential for 5 minutes to ensure that the iron was completely stripped.

0.8 0.6 0.4

0.2 푄퐹푒 0

-0.2 푄퐹푒 Current (mA) -0.4 푄퐹푒, 푆푡푟푖푝 -0.6 푃푙푎푡푖푛푔 퐸푓푓푖푐푖푒푛푐푦 = × 100 푄퐹푒, 푃푙푎푡푒 -0.8 0 200 400 600 800 1000 Time (sec)

Figure 2-16. Example plot for estimating coulombic plating efficiency. The plating potential was held at −1.0 V for

500 s then the metal deposit was stripped at + 0.2 V for 500 s at a glassy carbon substrate in 1:1:4 FeCl2:ChCl:EG

o electrolyte; N2-atmosphere, 80 C.

Coulombic plating efficiencies at for iron deposited at −1.0 V and stripped at +0.15 V vs Ag/AgCl are compared for the various compositions of iron DES electrolytes (Table 2-6). Efficiencies are surprisingly low for the 0.1:2:4 and 0.25:2:4 electrolytes, where chloride is in excess of iron.

Inefficiencies have been attributed largely to the nature of the electrodeposited iron (i.e. if the deposit does not stay attached to the substrate, losing electrical contact) [52] as well as formation of a passive film [143] or electrolyte decomposition. Iron deposited from a 1:2:4 FeCl2:ChCl:EG solution was found to be non-adherent and easily cracked and flaked off the surface whereas iron deposited from a 1:1:4 electrolyte was strongly attached to the surface in the form of a smooth deposit. The 0.1:2:4 and 0.25:2:4 electrolytes behave similarly to the 1:2:4 electrolyte where the deposits do not adhere well, thus, inefficiency is due to loss of metal from the surface of the electrode. It is apparent that electrolytes that contain an excess of chloride (compared to iron) exhibit this deposit morphology.

Table 2-6. Coulombic plating efficiency for the IL electrolytes of this study at 80 oC. Iron was plated at −1.0 V vs Ag/AgCl for 5 minutes then stripped for 5 minutes at +0.15 V.

Plating Molar Ratio Efficiency at (FeCl2:ChCl:EG) −1.0 V (%) 1:0:4 92 1:1:4 98 0.5:1:4 77 1:2:4 69 0.25:2:4 20 0.1:2:4 10

In addition to poor deposit quality, electrolyte breakdown was found to occur in the solutions containing an excess of chloride which contributed to coulombic inefficiencies observed. To study electrolyte degradation products, electrolytes underwent bulk electrolysis for 5 hours at 22 mA/cm2

(normalized using the membrane area) in a sealed glass H-type cell which was thermostatted to

80 oC using a salt water bath. The headspace was swept with argon gas; differential pumping was used to introduce the gas to the mass spectrometer. The H-type cell utilized high surface area platinum mesh electrodes and each half-cell was separated by a Daramic 175 microporous separator (exposed area of 0.8 cm2) that had been pre-soaked in a 2:4 choline chloride:ethylene glycol solution (24 hours). The cell was assembled inside a nitrogen filled glovebox.

Hydrogen gas evolution was confirmed using mass spectrometry after bulk electrolysis of a 1:2

ChCl:EG electrolyte. It is interesting to note that no chlorine gas evolution had been observed at the positive electrode, however, instead the electrolyte turned from clear to orange (as shown by

Figure 2-17). This observation has been reported in the literature and may be an indication of the presence of acetaldehyde, methylene chloride, dichloromethane, and/or chloroform [144]; additional degradation products have been reported for a 1:2 choline chloride and ethylene glycol solution after bulk electrolysis [144]. 39

Figure 2-17. H-type cell setup for studying electrolyte degradation products. This photograph was taken after bulk

electrolysis of a 1:2 M ChCl:EG electrolyte for 5 hours at 22 mA/cm2 (membrane area) and 80 oC.

The association of anions or cations plays an important role in electrochemical stability. The fact that electrolyte breakdown was not observed within the potential range of interest for electrolytes containing low concentrations of chloride suggests that the reduction potential of the hydrogen bond donor (ethylene glycol) was lowered to a more negative potential. This implies that the hydrogen bond donor is interacting more strongly within the ILs containing low chloride, which increases the lowest unoccupied molecular orbital (LUMO), therefore increasing the electrochemical stability window (i.e. it is more difficult to put an additional electron in this orbital).

2.2.3 Physical Properties of the Electrolytes

2.2.3.i Effect of Chloride to Iron Ratio on Conductivity

In Figure 2-18, conductivity is plotted as a function of (a) mole fraction of iron, (b) mole fraction choline and (c) ethylene glycol. When referencing Fig. 2-18a, the conductivity decreases dramatically when the mole fraction of iron is greater than 0.17 (as indicated by the dashed vertical line); this trend is consistent for electrolytes containing Fe2+ or Fe3+. It is useful to make note of the moles of chloride per moles of iron (Cl:Fe ratio), independent of oxidation state of the iron.

Interestingly, solutions containing 0.17 mole fraction iron correspond to those with a chloride to

40 iron ratio of 4:1. Thus, it is apparent that a chloride to iron ratio of at least 4:1 is required for enhanced conductivity. Increasing this ratio to 5:1 and even higher does not appear to significantly alter electrolyte conductivity. It is known that the addition of choline, a bulky inorganic cation, will improve conductivity [53]. However, when referencing Fig 2-18b, the conductivity of these iron containing electrolytes appears to depend more strongly on the chloride to iron ratio; the data can be grouped into two clusters. Those with ≥4:1 ratio Cl:Fe exhibit a high conductivity whereas those with <4:1 ratio Cl:Fe have low conductivities. Similarly, these two clusters are present when referencing Fig 2-18c where electrolytes with ≥4:1 ratio Cl:Fe exhibit a high conductivity and those with <4:1 ratio Cl:Fe have low conductivities. Furthermore, the amount of ethylene glycol does not largely affect the electrolyte properties.

35 35

Fe3+ Fe3+ 30 30 Fe2+ Fe2+

25 25

C (mS/cm)

o o 20 20

15 15

10 10

Conductivity at 80

Conductivity at 80 5 5

0 0 0.0 0.1 0.2 0.3 0.4 0.0 0.1 0.2 0.3 0.4 Mole Fraction Iron Mole Fraction Choline (a) (b) 35 Fe3+ 2+ 30 Fe Figure 2-18. Conductivity of iron 25

C (mS/cm) containing ILs as a function of o 20

15 concentration of (a) iron, (b) choline, or

o 10 (c) ethylene glycol at 80 C.

Conductivity at 80 5

0 0.5 0.6 0.7 0.8 Mole Fraction EG

Viscosity is also influenced by electrolyte composition, for example, the 1:1:4 IL was 22.0 cP which is almost twice that of ILs which had Cl:Fe in a 4:1 ratio (viscosity of 1:2:4 and 0.5:1:4 ILs was 10.2 and 14.4 cP at 80 oC, respectively). A decrease in viscosity, η, is expected to enhance

푘푇 diffusion, D, of active species when considering the Stokes-Einstein relation (퐷 = ) where k 6휋푟휂 is the Boltzmann constant, r is the radius of the diffusing species, and T is temperature. However diffusion coefficients are not reported here. Due to the fact that the active species also acts as the supporting electrolyte in the higher concentration ILs, the effective diffusion coefficient can be further enhanced or hindered by electric field migration as known to occur in conventional binary electrolytes [145]. Furthermore, strong ion pairing and complexation may occur which makes it difficult to determine active species concentrations. Ion-ion interactions such as ion pairing and agglomeration can be explained using the Walden plot.

2.2.3.ii Walden Plot

Within the IL literature, the use of the Walden Rule has been applied to explain physical properties of ILs and to provide insight on the degree of ionicity [86]. The Walden rule suggests that the molar conductivity (conductivity per mole of charge) is proportional to the fluidity, or inverse viscosity (Λη=Constant). For example, in a dilute aqueous solution (0.01 M KCl), the ions are completely dissociated and ion mobility is solely effected by solution drag thus, the relationship is linear and passes through the origin. However, many ILs fall below this line on the Walden plot.

This is not surprising since the ionic interactions are expected to occur in very concentrated ionic media that would yield a lower molar conductivity. These ionic interactions in the form of ion pairs or aggregates can strongly influence the conduction mechanism. This is shown by the

Walden plot (Figure 2-19) for selected iron containing ILs where the log of the molar conductivity

42 is plotted as a function of the log of the inverse viscosity. The 0.5:1:4 FeCl2:ChCl:EG solution contains Cl:Fe in a 4:1 ratio, whereas the 1:1:4 FeCl2:ChCl:EG IL does not (Cl:Fe of 3:1). The

0.5:1:4 and the 0:1:2 M ChCl:EG solutions show a similar trend, implying that there are some interactions between the ions in solution but also that the iron does not largely affect IL conductivity (compared to a traditional IL). However, the 1:1:4 solution line is much lower implying that the ions are interacting very strongly, hindering solution conductivity. In fact, the line for the solution containing 1:1:4 FeCl2:ChCl:EG deviates significantly from the “ideal solution line”; a vertical deviation, ΔW, of 1 indicates an ionic liquid which exhibits only 10% of the ionic conductivity expected (implying that it is only 10% ionized) [146]. This suggests that there are strong ion-ion interactions and possible aggregation in electrolytes containing low concentrations of chloride which hinder solution conductivity.

1.2

1:2:4  0.5:1:4 0.9 0:2:4 1:1:4

) KCl

-1 0.6 Increasing Temperature

mol

2 0.3

/Scm

 0.0 "Ideal Solution Line"

Log ( Log

-0.3

-0.6

-0.6 -0.3 0.0 0.3 0.6 0.9 1.2 -1 -1 Log ( /Poise )

Figure 2-19. Walden plot for selected iron (Fe2+) containing ILs compared to a neat 2:4 M ChCl:EG electrolyte as

well as the ideal solution line (0.01 M KCl in water). The red line indicates a deviation of ΔW = 1.

2.2.4 Iron Speciation

Thus far we have shown that the electrodeposition kinetics and also the fluid properties of the electrolyte change when the chloride to iron is above or below a 4:1 molar ratio. Electrochemical reactivity of the solute ions as well as the physical properties of the electrolyte are controlled by speciation of the metal moieties in solution. To understand what ions are responsible for the observed properties in these iron IL electrolytes, chemical speciation was investigated using X-ray

Absorption Near Edge Structure (XANES) measurements, X-ray photoelectron spectroscopy

(XPS), and Raman spectroscopy.

2.2.4.i X-ray Absorption Near Edge Structure Measurements

Liquid samples for X-ray absorption studies were contained using polyimide tubing (diameter was dependent on the desired path length) that had been prepared, sealed with epoxy, and stored inside a nitrogen-filled glovebox. X-ray absorption experiments were conducted at beamline 10 BM of the Advanced Photon Source at Argonne National Laboratory. All experiments had been performed using transmission mode with beam widths of 0.5 to 0.9 mm to constrain the beam to be narrow enough to fit inside the narrow width of the tube while still allowing for enhanced signal to noise. Iron foil was placed in line with the second ionization chamber to allow for an internal energy calibration between each test (Fe K-edge 7112 eV). In addition, the iron foil as well as

FeCl2 and FeCl3 (hydrates) salts ground to fine powders and sandwiched between teflon tape had been used as standards. Analysis of the data was performed using Demeter software [147] and

Origin Lab. A spline function was used to interpolate the background using data from several eV before and after the pre-edge feature. Deconvolution of the pre-edge features were modeled using a Voigt fit in Origin Lab.

X-ray absorption fine structure (XAFS) spectroscopy is sensitive to the iron oxidation states as well as the coordination geometry, especially in the pre-edge region located ~10 eV before the main K-edge absorption peak. Intensity of the pre-edge in this region is related to the centrosymmetry or coordination geometry and the centroid position is dependent on oxidation state. In Figure 2-20, the normalized pre-edge spectra for the Fe K-edges for an electrolyte containing a 4:1 ratio of Cl:Fe (0.5:1:4 FeCl2:ChCl:EG) is compared to an electrolyte which contains a 3:1 ratio Cl:Fe (1:1:4 FeCl2:ChCl:EG) as well as a salt standard. The peak intensity for the salt standard (aq) shows the lowest intensity, characteristic of a centrosymmetric complex

2+ (likely octahedral). This is since iron exists as an octahedral complex, Fe(H2O)6 in aqueous solution [148]. In contrast, the DES containing a 4:1 Cl:Fe ratio shows the highest intensity, characteristic of a lack of inversion symmetry (likely tetrahedral). Table 2-7 gives a summary of the data obtained from deconvolution of the pre-edge peaks. Peaks for electrolytes containing Fe2+ that is 4-coordinated can be satisfactorily fit using two components centered around 7112.4 and

7114.0 eV, which is consistent with literature [149, 150]. When the chloride to iron ratio is greater than 4:1, the iron is 4-coordinated with tetrahedral or square planar geometry. Within the literature

- 2- the tetrahedral complexes [FeCl4] and [FeCl4] have been shown to be the dominant species in iron containing ILs [151-154]. Considering the stiochiometric ratios, it is not surprising that for

Cl:Fe ratios ≥4:1 we see evidence of these complexes as well.

a) 0.5:1:4'0.5:1:4 b) 1:2:4'1:2:4 1:1:4'1:1:4 1:1:4'1:1:4 FeClFeCl22 Salt Salt 0.5:1:4'0.5:2:4 1:0:4'1:0:4

FeClFeCl33 Salt Salt

Absorption, Normalized xµ(E) Normalized Absorption, Absorption, Normalized xµ(E) Normalized Absorption,

7108 7110 7112 7114 7116 7118 7110 7112 7114 7116 7118 Energy (eV) Energy (eV)

Figure 2-20. Normalized pre-edge spectra (Fe K-edge) for selected electrolytes containing (a) Fe2+ and (b) Fe3+ compared to standard, FeCl2•4 H2O (FeCl2 salt) and FeCl3•6 H2O (FeCl3 salt) with the corresponding models of best

fit.

Table 2-7. X-ray Absorption Near Edge Structure (XANES) data obtained from deconvolution of the Fe pre-edge peaks for samples prepared with (a) FeCl2 fit using two components centered around 7112.4 and 7114.0 eV and (b)

FeCl3 fit using1 component centered around 7114.0 eV (with the exception of the FeCl3 salt which had an additional

peak located at 7115.0 eV).

(a) (b)

2.2.4.ii X-ray Photoelectron Spectroscopy

Due to low vapour pressure of ionic liquids, XPS analyses can be performed on IL electrolytes without the need for a specially modified XPS chamber. Small droplets of each IL (less than 1

µL) had been placed atop a scratched (and thoroughly cleaned) glassy carbon surface which was in contact with the XPS stub. Scratching the surface of the glassy carbon with sand paper allowed for the liquid to distribute more evenly on the surface and not bead up. Samples were left to dry in the transfer chamber for a minimum of 12 hours under vacuum, although due to their low volatilities, they remained liquid.

Survey scans in the range of 0 – 1000 eV were carried out to identify surface elements. High resolution XPS of the Fe 2p, Cl 2p, C 1s, N 1s and O 1s peaks was performed on the liquid samples as well as FeCl2 and FeCl3 salt standards. Choline chloride salt as well as iron metal had also been used as standards. The oxide layer on the iron metal had been removed by sputtering the sample using argon ions for a total of 4 minutes. The resulting high resolution spectra for the Fe 2p metal is shown by Figure 2-21. The two peaks are associated with Fe 2p3/2 located at 706.9 eV and the

Fe 2p1/2 located 720.1 eV (which gives a separation, Δ = 13.2 eV) [155].

After Sputtering, 4 min

Initial Spectra Norm Counts Norm

732 722 712 702 Binding Energy (eV)

Figure 2-21. High resolution Fe 2p spectra for the iron metal before and after sputtering for 4 minutes to remove the

oxide layer.

Care must be taken to prevent sample charging when examining ILs using XPS [156]. Even though the high power mode facilitates higher counts (enabling better signal/noise in less time), it was not used on the liquid or salt samples. The electron flood gun and neutralizer are unable to cover the entire area of interest due to rastering which causes charging of the sample (noted by the broad peak when referring to the high resolution C 1s spectra as shown by Figure 2-22b). Charging of the sample complicates analysis since the peaks are not only broad but also may appear in different locations (as shown by Figure 2-22 a and b). So although sampling with the 100 µm spot size takes longer, it is necessary to prevent charging of the ionic liquids. Figure 2-22 compares high resolution spectra for the Fe 2p and C 1s peaks in high power mode to spectra taken with the

100 µm spot size. Note that the number of scans for the 100 µm spot size is 1.75 times that of the spectra obtained with the high power mode in order to get comparable resolution.

a) b)

100 µm Spot

NornCounts NormCounts

HP (100 µm)

732 722 712 702 288 286 284 282 Binding Energy (eV) Binding Energy (eV)

Figure 2-22. High resolution XPS for the (a) Fe 2p and (b) C 1s peaks. High power mode is compared to the 100

µm spot, highlighting the effects of sample charging.

The C 1s peak located at 284.5 eV was used as charge reference in determining the binding energies. XPSPEAK software [157] had been used to fit the data to determine binding energies of each of the peaks. Sample spectra for the high resolution Fe 2p peaks are shown by Figure 2-23.

An electrolyte that contains 4:1 Cl:Fe is compared to an electrolyte containing less than 4:1 Cl:Fe.

2- 1- The peaks located at ~709.1 eV and ~711.5 eV confirm the presence of [FeCl4] and [FeCl4] , respectively [155, 158]. Other prominent peaks located at 710.8, 709.7, and 711.2 eV indicate the

+ presence of FeCl2, FeCl2 and FeCl3, respectively [158]. A summary of the data is provided in

Table 2-8. An additional peak is present for the chloride poor electrolytes at 706.2 eV. This specific compound could not be identified within the literature but indicates that an organometallic iron complex is forming (which causes the iron to be in a reduced state). Thus, it is of interest to probe the N 1s peak to monitor how the choline is interacting.

It is interesting to note that all N 1s peaks aligned at 402.4 eV (Figure 2-23b) which is the binding energy for choline [158]. This implies that the choline cation electronic structure is constant and thus not interacting differently when the molar ratios are varied. The concentration of choline showed little to no effect on the iron deposition/dissolution kinetics or on the solution conductivity.

a) b)

ChCl Salt

Iron Metal 1:2 ChCl:EG

1:1:4 1:1:4 FeCl :ChCl:EG

2 FeCl2:ChCl:EG NormCounts NormCounts

1:1:4 1:1:4 FeCl3:ChCl:EG FeCl3:ChCl:EG

1:2:4

FeCl2:ChCl:EG

1:2:4

FeCl3:ChCl:EG

732 722 712 702 410 405 400 395 Binding Energy (eV) Binding Energy (eV)

Figure 2-23. High resolution XPS for the (a) Fe 2p spectra for selected electrolytes as well as the (b) N 1s peaks.

Table 2-8. Summary of speciation data obtained from deconvolution of peaks from high resolution XPS.

Molar Ratio Identified Species (FeClx:ChCl:EG) Fe3+ Fe2+

Salt FeCl3 FeCl2

- + 2- 1:0:4 FeCl3, [FeCl4] , [FeCl2] FeCl2, [FeCl4] +peak at low B.E.a

- 2- 2- 1:1:4 FeCl3, [FeCl4] , [FeCl4] FeCl2, [FeCl4] +peak at low B.E.a

- + 2- 1:2:4 FeCl3, [FeCl4] , [FeCl2] FeCl2, [FeCl4] a B.E. stands for binding energy

2.2.4.iii Raman Spectroscopy

Liquid samples for Raman measurements had been contained using 1 dram glass vials that had been prepared, sealed, and stored inside a nitrogen-filled glovebox. Raman spectra of the liquids were recorded using a SENTINEL® Raman spectrometer from Bruker Optics. Results obtained from Raman spectroscopy support the observations from XANES and XPS that for solutions containing Cl:Fe ratios ≥4:1, the dominant species are the tetrahedral chloride complexes.

- Example spectra can be referenced in Figure 2-24 which confirms the presence of the [FeCl4] tetrahedral iron complex by the large peaks located at 333 cm-1 [153, 154, 159].

Ethylene glycol

'0:1:4 Norm Counts (a.u.) (a.u.) NormCounts '1:2:4 '0.5:1:4 '1:1:4 '1.5:1:4 '1:0:4 '2:0:4 290 490 690 890 1090 1290 1490 1690 Raman Shift (cm-1)

Figure 2-24. Raman spectra for various iron(III)chloride containing IL electrolytes compared to the spectra for

- ethylene glycol and a mixture of ethylene glycol and choline chloride. The presence of the [FeCl4] tetrahedral iron

complex was confirmed by the large peaks located at 333 cm-1 [153, 154, 159].

For electrolytes containing low amounts of chloride (ie Cl:Fe ratios less than 4:1) it is of interest to probe the interactions with ethylene glycol. Figure 2-25 compares selected Raman spectra in the region from 800 to 1500 cm-1. Ethylene glycol is compared to a solution of choline chloride and ethylene glycol (1:4 M ChCl:EG) as well as two electrolytes that contain iron. Upon addition of choline chloride, the Raman spectra for ethylene glycol do not change but additional peaks from the choline appear at 957, 1140, and 1342 cm-1. For the case of electrolytes containing ≥4:1 Cl:Fe

(ie 0.5:1:4 FeCl2:ChCl:EG), the spectra resemble that of the choline chloride and ethylene glycol.

However, when the chloride:iron ratio is low (i.e. 2:1) there is a dramatic difference in the spectra

52 suggesting that the ethylene glycol is interacting with the iron. The hydrogen bond donor has been shown to play a significant role in the coordination of metal ions. Furthermore, ethylene glycol is known within the literature to form complexes with transition metal chloride salts [160-162]. The proposed Fe-EG complex is shown in Figure 2-26. Thus, when the chloride to iron ratio is low

(regardless of oxidation state of iron) the tetrahedral iron complexes are still formed, however, the dominant species exists as an Fe-EG complex.

C-C stretching CH bending C-O stretching 2 CH2 rocking CH2 wagging

Ethylene Glycol

Choline chloride & ethylene glycol

Counts(mW/s) 0.5:1:4

1:0:4

800 900 1000 1100 1200 1300 1400 1500 Raman Shift (cm-1)

Figure 2-25. Raman spectra for ethylene glycol compared to 1:4 M ChCl:EG as well as 0.5:1:4 and 1:0:4

FeCl2:ChCl:EG electrolytes. Vibrational modes for ethylene glycol are labeled [163, 164].

Figure 2-26. Proposed Fe-EG complex.

2.3 Iron IL Flow Battery - Proof of Concept

Previously, it was shown that a specific chloride to iron ratio (4:1) gave the best electrochemical and physical properties, thus, an electrolyte containing 0.5:1:4 FeCl2:ChCl:EG was selected for testing in a small-scale flow cell. Reasonable coulombic plating efficiencies (77% at −1.0 V), fast kinetics (94 A/m2), high concentration of active species (1.3 M), as well as acceptable fluid properties had been discussed in this chapter. The conductivities for the 0.5:1:4 FeClx:ChCl:EG electrolyte containing only FeCl3 or only FeCl2 were similar (25.5 and 19.2 mS/cm, respectively) and thus, large changes in conductivity and viscosity are not expected to occur upon charge and discharge.

Flow battery testing was performed using a 4 cm2 cell (membrane area) such as the one shown by the schematic in Figure 2-27. SIGRACELL® carbon felts from SGL were used as electrodes,

KFD 2.5 for the negative and GFD 4.6 for the positive. A less conductive felt was used for the negative electrode to avoid metal deposition adjacent to the membrane [165, 166]. GFD felt is often employed in typical aqueous vanadium redox flow batteries; if any degradation of the felt electrode were to occur it is expected to be less for the all iron system. The felts were bonded to impervious graphite plates using carbon ink (Acheson) and separated by a Daramic 175 microporous separator that had been pre-soaked in a 1:4 choline chloride:ethylene glycol solution

(2 days). Maximum compression of the felts had been limited to 10% by using Teflon gaskets.

Electrolytes were circulated by a peristaltic pump at a flow rate of 25 mL/min. The positive reservoir contained 35 mL of electrolyte and the negative contained 17.5 mL of electrolyte (since two electrons are transferred at the negative electrode, half the volume is required). The cell was

2 o charged and discharged at 5 mA/cm , at 80 C in a N2-filled glovebox (<15 ppm water).

Figure 2-27. Schematic of the small-scale flow battery hardware.

The flow cell had been charged to 10% state of charge, then cycled for one hour charge-discharge cycles at 5 mA/cm2. Figure 2-28 compares the first three cycles to cycles 14-16. These results demonstrate the feasibility of an iron-containing IL. It should be noted, however, that the capacity does begin to fade even after 14 cycles. This is likely attributed to crossover of Fe3+ at the separator into the negative half cell (lowering the coulombic efficiency); other possibilities include passive film formation at the electrodes or electrolyte breakdown (lowering voltaic efficiency, as well as coulombic efficiency). Further work is needed to optimize the cell hardware as well as understand long-term degradation in battery performance after more extensive cycling. However, the results presented here show that ILs are a promising alternative to traditional non-aqueous electrolytes due to their ease of fabrication, high concentration of active species, reasonable plating efficiency and fast redox kinetics and should be explored further.

1.5

1.2

0.9

0.6

Cycle 1 Cycle 2 Cycle 3 Cycle 14 Cycle 15 Cycle 16 Cell Potential (V) CE 99.4 % CE 99.1 % CE 99.2 % CE 69% CE 68% CE 69% 0.3 VE 49.6 % VE 48.2 % VE 47.7 % VE 35% VE 34% VE 33% EE 49.3 % EE 47.8 % EE 47.3 % EE 24% EE 23% EE 23%

0.0

6 8 10 34 36 Time (hours)

Figure 2-28. One hour charge-discharge cycles for a RFB utilizing 0.5:1:4 FeCl2:ChCl:EG after an initial charge to

2 o 10% SOC; current density of 5 mA/cm , flow rate 25 mL/min, 80 C, N2-atmosphere.

2.4 Conclusions

Moisture-stable ionic liquids show promise for energy storage applications. These electrolytes are not only easy to synthesize but also made from safe and cost effective materials. Iron containing

ILs were synthesized by combining iron chloride, choline chloride, and a hydrogen bond donor, such as ethylene glycol. An understanding of the factors that affect the electrochemistry, specifically that of the iron deposition and dissolution reaction, as well as how iron speciation influences the physical properties of the electrolytes was demonstrated. For example, it was shown that the chloride to iron ratio dominates both electrochemical and physical properties of these

56 electrolytes. X-ray absorption spectroscopy (XAS), X-ray photoelectron spectroscopy (XPS), and

- Raman spectroscopy data indicate the presence of the tetrahedral iron chloride complexes, [FeCl4]

2- and [FeCl4] for solutions containing ≥ 4:1 molar ratio chloride to iron. However, for low chloride electrolytes (ie <4:1 Cl:Fe), the ethylene glycol forms a complex with the iron. This EG-Fe complex allows for enhanced plating efficiencies due a decrease in the H2 evolution side reaction and a smoother iron deposit, however, the physical properties of the solution are hindered (i.e. conductivity for a 3:1 electrolyte is an order of magnitude lower that that of the 4:1 electrolyte).

The Walden plot suggests that this is due to strong ion-ion interactions which hinder solution conductivity. Furthermore, iron plating in solutions containing an excess of chloride follow progressive nucleation at low overpotentials. However, for electrodeposition in solutions containing low chloride, there is a change in the mechanism and nucleation cannot be explained using the traditional S-H model.

An electrolyte containing 0.5:1:4 FeClx:ChCl:EG (where the Cl:Fe ratio is 4:1) was selected for validation in a small-scale flow battery due to reasonable coulombic plating efficiencies (77% at

2 −1.0 V), fast kinetics (io = 94 A/m ), high concentration of active species (1.3 M), as well as acceptable fluid properties. The battery had been cycled at 5 mA/cm2 and energy efficiencies of

~ 48 % are reported for the first three cycles. This demonstrates promise for the iron IL as an electrolyte for flow batteries. One peer-reviewed article has been published on this work [52] with another manuscript recently submitted to the Journal of the Electrochemical Society [63].

Chapter 3. Technique Development, Carbon Fiber Microelectrodes

3.1 Introduction and Literature

Some electrode reaction kinetics depend on the nature of the electrode surface, Fe2+/Fe3+ [167,

2+ + 2+ 3+ 2+/3+ 2+/3+ 168], VO /VO2 [78-81, 167, 169, 170], V /V [80, 81, 169, 171], Eu [168, 172], Cr

[172], and hydroquinone/quinone [167] represent a few examples. For instance, the presence or absence of oxygen-containing functional groups on the surface of the electrode as well as the carbon material can alter redox reaction kinetics by over an order of magnitude [80, 81, 169].

Furthermore, electrode material has been known to influence electrodeposition and nucleation

[173]. Thus, it is important to study the electrode-electrolyte interface at electrode materials used in practical redox flow batteries.

Investigations have been carried out using flow through electrodes, such as carbon felt [174] or carbon paper [175]. However, due to the 3-dimensional structure of these electrodes (shown in

Figure 3-1), complex current distributions make analysis difficult. Thin fiber film rotating disc electrodes (TFFRDEs) were considered as a means for evaluating electrochemical kinetics [176].

They have been shown to yield variable results and highly disordered surfaces; ohmic and mass transfer within the fiber mat complicates kinetic analysis. For example, increasing exchange current densities are expected to correlate with increased loading, due to an increase in active area.

However, as shown by Figure 3-2a, there is no correlation. Pulse voltammetry experiments also did not exhibit a dependence on fiber loading (Figure 3-2b). Consequently, due to their variable performance, TFFRDEs at this time cannot be recommended for studying electrochemical processes at carbon felt electrodes.

(a) (b)

Figure 3-1. (a) Photograph and (b) corresponding SEM micrograph to show the structure of typical carbon

electrodes used in flow batteries.

)

2 (a) (b)

-3 -5 4.0x10 0.05 M Fe 0.01 M Fe -6 3.0x10-3

ln(i) -7 2.0x10-3 7.6 mg 6 mg -3 -8 1.0x10 3 mg

Exchange Current Density (A/cm 2.4 mg -9 0 1 2 3 4 5 6 7 8 0.0 0.1 0.2 0.3 FIber Loading (mg)  (V) a

Figure 3-2. (a) Exchange current densities from TFFRDEs for two different active species concentrations

(equimolar FeCl2 and FeCl3); supporting electrolyte of 1 M NaCl was used. (b) Pulse voltammetry (250ms) polarization of TFFRDEs at four different loadings in 0.02 M total iron and 1 M NaCl (equimolar concentrations of

FeCl2 and FeCl3). Reprinted with permission from reference [176].

Single carbon fiber microelectrodes which have well-defined geometries have been investigated as a more promising alternative for characterizing interfacial redox electrochemical reactions. For this study, novel microelectrodes fabricated using single carbon fibers extracted from bulk carbon

59 felt have been developed to gain a better understanding of the electrode-electrolyte interface at application-relevant electrode surfaces. CFMEs are convenient for use in resistive media, such as non-aqueous electrolytes due to their small currents; thus, ohmic overpotentials remain small. Due to enhanced diffusion to microelectrodes [23, 177, 178], CFMEs are not limited by mass transfer

(or ohmic overpotentials), thus kinetic analysis can be performed at these stationary electrodes.

For example, linear diffusion occurs at large planar electrodes; the time dependent diffusion current follows the Cottrell equation (as shown in Table 3-1) where n, is the number of moles of electrons, F is Faradays’ constant, A is the electrode area, D is the diffusion coefficient, and CR is the bulk reactant concentration. In the case of a cylinder, an additional radial diffusive component is present which is parallel to the surface of the CFME. Furthermore, this effect is enhanced as the radius of the cylinder, r, decreases as shown by the equations below for the current-time relationship (Table 3-1). The current transient for a CFME does not reach a steady state like that of the spherical case since it is still time dependent. However, a quasi-steady state can be achieved since at long times the current decays slowly (since it is an inverse logarithmic function).

Table 3-1. Equations governing the diffusion current; three different electrode geometries are compared (derivations

can be found in reference [22]).

Electrode Planar Spherical Cylindricala Geometry 1 1 1 1 푛퐹퐴퐷퐶푅 2exp (−0.05휋2휏2) 푖 = 푛퐹퐴퐷퐶 [ + ] 푖 = [ Diffusion 푑 푅 ( )1/2 푟 1 1 푛퐹퐴퐷1/2퐶 휋퐷푡 푟 휋2휏2 Current 푖 = 푅 푑 휋1/2푡1/2 Equation 푛퐹퐴퐷퐶푅 1 푖 = 푖 + + ] 푑 푑(푃푙푎푛푎푟) 푟 1 ln (5.2945 + 0.749휏2) Diffusion 푛퐹퐴퐷퐶 2푛퐹퐴퐷퐶푅 Current, 푖 = 0 푖 = 푅 푖 = 푑 푑 푟 푞푠푠 푟 ln(τ) t→∞ aτ = 4Dt/r2

3.2 Technique Development - Carbon Fiber Microelectrode

Carbon fiber microelectrodes (CFME) were prepared from carbon felts (GrafTech). Single fibers were extracted from the bulk felt at random and were typically 5 mm in length and 15±3 μm in diameter with an approximately circular cross-section. The morphology of a typical fiber is shown by the SEM (Hitachi SEM) micrograph in Figure 3-3; the solid core is obvious from the inset.

Electrical contact was made to copper wires on both ends of the fiber using silver ink2 (EPO-TEK

H20S). The copper wires were placed inside a glass capillary and sealed along with the junctions, which include both ends of the fiber, using non-conductive epoxy (Loctite E-05MR Hysol) so that only the fiber was exposed to solution (Figure 3-4). All CFMEs used in this study were fabricated from the same batch of carbon felt.

Figure 3-3. SEM micrographs of a typical carbon fiber used Figure 3-4. Schematic of carbon fiber microelectrode (CFME).

2 It is important to note that silver ink had been used in the CFME construction since it is easy to detect exposure by CV, whereas carbon ink could pose a problem; if carbon ink was exposed, it would be undetectable and thus difficult to determine what surface the reactions are taking place at. 61

Initial cyclic voltammograms (CV) of the CFMEs were performed in H2SO4 to test for good contact. Cyclic voltammetry was performed in a jacketed (25 oC) three electrode cell (refer to the sketch in Figure 3-5) with a potentiostat/galvanostat (Solartron SI 1280) and an impedance phase analyzer (Solartron SI 1260) employing a working electrode (in this case the CFME), a platinum counter electrode, and a reference electrode situated inside a Luggin capillary. It is important to note that a saturated Hg/Hg2SO4 reference (mercury/mercurous sulfate electrode, MSE) was used in the three electrode cell to eliminate possible contamination due to chloride ions, which becomes important when the electrode geometry is miniaturized [23].

Figure 3-5. Schematic of a three electrode cell employing a platinum counter electrode, reference electrode (situated

inside a Luggin capillary), and CFME working electrode.

3.2.1 Preconditioning

Carbon electrodes often require preconditioning or cleaning processes in order to obtain a stable surface response, and many treatment procedures can be found in the literature [11, 179-184].

Prior to electrochemical treatments and kinetic measurements, the CFMEs in this study were

“preconditioned” by scanning from ‒0.5 V to +1.2 V for 50 cycles at 100 mV/s (Figure 3-6). In

Figure 3-6, CVs overlaid after ~30 cycles, implying a stable surface response. Some fiber electrodes took more than 30 cycles to achieve a stable surface response, however, all had stabilized after 50 cycles and thus, 50 cycles were used for all electrodes for consistency. In addition, to minimize any effects of electrode history, electrodes were also subjected to a normalization treatment (as described by Bourke et al. for glassy carbon electrodes [78, 170]) between experiments, consisting of three cycles, each of 60 s at ‒2.0 V and 60 s at +1.5 V.

8 Initial 6 20th 30th 4 45th 47th 2 50th

0 Current (µA) -2

-4 -0.55 -0.15 0.25 0.65 1.05 Potential (V vs Hg/Hg2SO4)

Figure 3-6. Selected CV cycles at a CFME in H2SO4 at 100 mV/s to show the electrochemical pretreatment

procedure used to obtain a stable surface.

The untreated felts had contaminants present on the surface, however, after electrochemical preconditioning treatments, fewer contaminants were observed (Figure 3-7). SEM-XEDS analysis revealed that the surface of the untreated felt was contaminated with an insulating aluminum- containing compound, which was not present after preconditioning. Prior to SEM-XEDS analysis, the felts had been sputter coated with palladium so as to allow for detection of the insulating material on the surface of the fibers. These particles were approximately 100 nm in diameter or smaller and could be a byproduct of carbon fiber processing (from the ovens during pyrolysis of the fibers, for example).

(a) (b)

Figure 3-7. SEM micrographs of (a) an untreated fiber compared to (b) a preconditioned fiber. Images (c) and (d) represent untreated felts coated with palladium so that the insulating material on the surface could be detected using

SEM-EDS techniques (electron interaction depth was 0.8 µm and radius of 0.4 µm); (d) shows the general size of the aluminum-containing particles (<100 nm diameter).

After preconditioning, the CFMEs were ready for use. In Figure 3-8 a typical CV is shown for a preconditioned CFME compared to that of a standard glassy carbon electrode in 3M H2SO4. Both the CFME and the glassy carbon electrode exhibit typical carbon response, with peaks present around 0 V (MSE) which are associated with hydroquinone species on the carbon surface [185].

The hydroquinone species are a surface oxide which is typically present on the surface of carbon electrodes. Furthermore, the CV in 3M H2SO4 does not show peaks for silver or copper oxidation

(from the silver ink or copper wire) indicating that only the carbon fiber is exposed to the solution.

Thus, from the CVs, fabrication methods have been verified.

Normalized Current Normalized CFME -1 GC

-0.6 -0.2 0.2 0.6 Potential (V) vs. MSE

Figure 3-8. Comparison of typical CVs (100 mV/s) at a preconditioned CFME and a glassy carbon electrode in 3 M

H2SO4. The currents were normalized to the peak current at about 0 V (MSE).

3.2.2 Electrochemically Active Area

The active surface area of CFMEs was monitored by electrochemical capacitance. Capacitance was estimated by CVs in H2SO4, where capacitive current is plotted as a function of scan rate, the slope of the line yields the capacitance (sample data is shown in Figure 3-9)3. It is important to note that the window for CV scans should be performed where no reactions are occurring so that only the current from double layer charging is measured. For example, the CVs shown in Figure

3-9a extend from +0.1 to +0.5 V vs MSE so as to avoid the regions for which the

3 Differential capacity, Cd is equal to the rate of change of surface charge density of the double layer with potential (Cd = dQ/dE). For cyclic voltammetry the scan rate, v, represents a change in potential with a change in time (v = dE/dt). The differential capacitance current, ic, is equal to dQ/dt and thus, Cd = ic/v.

65 quinone/hydroquinone oxidation/reduction is observed as well as oxygen evolution and hydrogen evolution. Capacitance gives a good estimate of the electrochemically active surface area as opposed to the Brunauer–Emmett–Teller method (BET) which relies solely on physisorption of nitrogen. Furthermore, BET requires a large amount of surface area and cannot be performed at a single CFME. Typical capacitance values for glassy carbon electrodes in acid media range from

10-25 µF/cm2 [179]. Values were compared to the untreated fibers in order to obtain a relative roughness factor which can be used to normalize kinetic data, for example.

It should be noted that after electrochemical preconditioning at various numbers of cycles, it was found that the capacitance increases with preconditioning from a value of 24 ± 3 µF/cm2 for an untreated electrode, eventually reaching a steady state value of 277 ± 19 µF/cm2 after ~30 cycles.

It is important to note that the capacitance did not change during subsequent experiments, indicating a stable surface and a stable surface area. Thus, the electrochemical preconditioning treatment yields a larger active surface area after removing impurities in addition to developing a reproducible surface.

(a) (b) 3.E-07 2.5E-07 Scan Rate 2.E-07 2.0E-07 1.E-07 1.5E-07 0.E+00 1.0E-07

Current (A) -1.E-07 5.0E-08

-2.E-07 Capacitive Capacitive Current (A) -3.E-07 0.0E+00 0.1 0.2 0.3 0.4 0.5 0 0.1 0.2 0.3 0.4 Applied Potential (V vs MSE) dV/dt (scan rate in V/s)

Figure 3-9. (a) CVs of a CFME in 3 M H2SO4 at varying scan rates (20 – 300 mV/s) which can be used to

determine capacitance when (b) the capacitive current, ic, is plotted as a function of the scan rate.

3.3 Potential Applications

3.3.1 Non-Aqueous Systems - Understanding Kinetics and Degradation

Carbon fiber microelectrodes offer a tool to study non-aqueous systems. Due to their low currents, ohmic overpotentials remain small. Non-aqueous electrolytes are often an order of magnitude less conductive than typical aqueous electrolytes (i.e. the conductivities of 2 M H2SO4 compared to

0.65 M TEABF4 in acetonitrile are 700 mS/cm [127] and 50 mS/cm [46]). An added advantage of using a microelectrode is simplification of the experimental setup. Enhanced diffusion to

CFMEs allows kinetic analysis to be performed at stationary electrodes (eliminating the need for a rotating disk setup inside a glovebox). Lastly, small solution volumes can be studied which becomes important when studying a typical ionic liquid or non-aqueous electrolyte for which large quantities are not yet commercially available.

Non-aqueous systems show promise for the future, however, little is known on a fundamental level. In general, there is a lack in understanding the electrode/electrolyte interface and the effects of different solvents and electrolytes. Vanadium (III) acetylacetonate (V(acac)3) in acetonitrile

(ACN) represent the most widely studied non-aqueous system due to its wide potential window of

2.2 V [32-34, 36-39] (equations 4a and 4b). However, issues with capacity fade are one factor limiting its use in RFBs. This fade in capacity is associated with side reactions of active species with trace amounts of water and oxygen [33]. Particularly, V(acac)3 is highly sensitive to environmental contaminants that can yield undesirable and less soluble species such as VO(acac)2

[43, 44] (equation 4c). Any water present in the supporting electrolyte, solvent, or electrodes will negatively affect the performance of the non-aqueous V(acac)3 RFB. Carbon felt electrodes used in flow batteries often have oxygen containing functional groups on the surface. It is of interest to

67 investigate whether this could be a cause for capacity fade observed within the flow cells and thus

CFMEs provide a useful tool for understanding capacity fade.

+ - + V(acac)3  V(acac)3 + e +0.44 V vs Ag/Ag (4a)

- - + V(acac)3 + e  V(acac)3 −1.76 V vs Ag/Ag (4b)

+ - + VO(acac)2  VO(acac)2 + e +0.75 V vs Ag/Ag (4c)

For these experiments, a platinum fiber electrode was used as a control. The platinum fiber electrode had been constructed in the same manner as the CFME and was of the same characteristic dimension (15 µm diameter). The working electrodes had been soaked in acetonitrile for a minimum of 10 minutes prior to experiments in non-aqueous electrolytes. Cyclic voltammetry

+ performed at these electrodes is shown in Figure 3-10; the V(acac)3/V(acac)3 and

+ V(acac)3/V(acac)3 reactions can be observed. As shown by Figure 3-10 the CFME is in agreement with the platinum fiber.

1.5 p 1

0.5

-0.5 Platinum Fiber -1 Carbon Fiber

-1.5 Normalized Current Density, i/i Density, Current Normalized -2.1 -1.6 -1.1 -0.6 -0.1 0.4 0.9 Potential (V vs Ag/Ag+)

Figure 3-10. Cyclic voltammetry of non-aqueous vanadium reactions at 500 mV/s comparing two different

electrodes (0.01 M V(acac)3, 0.1M TEABF4 in acetonitrile).

From Figure 3-10 it is clear that the initial CV cycles at platinum and CFMEs overlay. However, over time the current response at the CFMEs in the non-aqueous electrolyte was not stable. Figure

3-11 shows CVs for 0.01 M V(acac)3, 0.1M TEABF4 in acetonitrile at a CFME and a platinum

+ fiber electrode. Even after 50 CV scans using a CFME, the peak currents for V(acac)3/V(acac)3

+ had decreased by 20 % and V(acac)3/V(acac)3 had decreased by 21 %. The CV response at the platinum fiber remained constant even after 150 CV (Figure 3-11b).

(a) (b)

75 40

)

2 )

2 30 45 20 15 10 0 -15 10th -10 50th 10th -20 -45 100th 150th

Current Density (A/m -30

Current Density (A/m 150th -75 -40 -2.3 -1.3 -0.3 0.7 -2.3 -1.3 -0.3 0.7 Applied Potential (V vs Ag/Ag+) Applied Potential (V vs Ag/Ag+)

Figure 3-11. Cyclic voltammetry of non-aqueous vanadium reactions at 100 mV/s in 0.01 M V(acac)3, 0.1M

TEABF4 in acetonitrile at (a) a CFME and (b) a platinum fiber electrode. Experiments were performed in a N2-

atomosphere (>3 ppm water) at 25 oC with a nitrogen blanket above the electrolyte.

It should be noted that if this observed change in the current response at the CFMEs was due to degradation of the electrolyte from reacting with oxygen-containing functional groups then a larger current due to oxidation and reduction of the VO(acac)2 should be observed over time in Fig 3-11a

(due to a higher concentration of this contaminant). However, the 10th cycle shows larger currents than the 150th implying that the effect is not simply due to oxygen/moisture contamination. Figure

3-12 compares the CV response at new CFMEs for an electrolyte containing 0.01 M V(acac)3

69 which had been exposed to trace amounts of moisture and oxygen compared to an electrolyte which contains 0.0075 M VO(acac)2. The peaks located at ~ 0.75 V are for the reduction and oxidation of VO(acac)2 and the measurements in the two electrolytes agree.

1.5E-07

-5.0E-08

Current (A) Current VO(acac)2 V(acac)3

-2.5E-07 -2.3 -1.8 -1.3 -0.8 -0.3 0.2 0.7 + Applied Potential (V vs Ag/Ag )

Figure 3-12. Cyclic voltammetry at 100 mV/s in 0.1M TEABF4 in acetonitrile with either 0.01 M V(acac)3 (which had been contaminated with oxygen/moisture) or 0.0075 M VO(acac)2 at a CFME. Experiments were performed in

o a N2-atomosphere (3 ppm water) at 25 C with a nitrogen blanket above the electrolyte.

After extended cycling in the non-aqueous vanadium electrolyte (>150 cycles), CFMEs can be cleaned by cycling in neat electrolyte. The response improves but does not return to the original.

Thus, this affect is in addition to the formation of a passive film as suggested previously [39] and could be an indication of degradation of the fiber. In fact, capacitance measurements in 3M H2SO4 before use in non-aqueous compared to after cycling, indicate that there is a loss in active surface area. The initial capacitance values for two fiber electrodes compared to those found after 500 cycles in the non-aqueous electrolyte were 49 vs 37 and 77 vs 61 µF/cm2 (not that these values

70 cannot be compared to capacitance values for carbon in aqueous H2SO4 since the ions present in the double layer are very different). This represents a 25% and 21% loss for the two fiber electrodes, respectively. Unfortunately, handling of the delicate CFMEs proved difficult for post testing analysis as many did not stay intact after testing inside the glovebox and thus only two values are reported. Furthermore, many CFMEs broke while testing in the non-aqueous electrolyte.

It is known that extreme potentials can cause degradation of carbon felt electrodes [186-188].

Figure 3-13 shows an example of this where a CFME had been cycled 200 times at extreme potentials (between −0.9 to 1.5 V vs Hg/Hg2SO4) then held at extreme potentials, alternating between −2 and 1.5 V for 60s each in 3 M H2SO4. The potential holds had been performed 3 times, after which the CFME broke. The SEM micrographs had been taken along the length of the fiber and show that near where the fiber broke, the surface was heavily pitted and destroyed. This experiment had been performed in an aqueous electrolyte, however, it is believed that similar phenomena are occurring in non-aqueous electrolytes where the carbon electrodes are cycled between extreme potentials which cause the carbon to degrade. It is thus recommended that different electrode materials be explored for non-aqueous flow battery systems.

(a) (b)

Figure 3-13. SEM micrographs showing a CFME which had been cycled 200 times at 50 mV/s between −0.9 V and

1.5 V (vs Hg/Hg2SO4) and subsequent potential holds alternating between −2 V and +1.5 V until the electrode broke. Images are in order, moving along the length of the fiber where (e) is closest to where the fiber broke during

the test.

3.3.2 Metal Electrodeposition for Hybrid RFBs

The all-iron redox flow battery chemistry consists of the Fe2+/Fe3+ redox reaction at the positive half-cell and iron plating and stripping at the negative [13]. In Chapter 2, it was shown that the iron electrodeposition reaction was limiting flow battery performance for a system utilizing iron

(II) chloride, choline chloride, and ethylene glycol in a 1:1:4 molar ratio. Three-electrode cell testing showed promise for iron electrodeposition/dissolution onto a glassy carbon substrate.

However, after charging a small-scale flow cell, no electrodeposited iron could be detected at the negative (carbon felt) electrode. Thus, it is of interest to study the iron electrodeposition and dissolution reaction using a CFME to determine whether the iron is electrodepositing onto the fibers within the felt electrodes and if this process is as efficient as was observed using a typical glassy carbon electrode.

350

250

) 2 150

-50 Current Density (A/cm -150

-250 -1 -0.5 0 0.5 1 Applied Potential (V vs Ag/AgCl)

Figure 3-14. Cyclic voltammetry showing the iron reactions at a CFME (10 mV/s) in a 1:1:4 FeCl2:ChCl:EG

o electrolyte. Experiments had been performed in a N2-atmosphere at 80 C.

Initial CVs for the iron reactions were run using CFMEs utilizing the same thermostatted three- electrode cell described previously (with an iron rod counter electrode to maintain constant Fe2+ concentrations). The Fe2+/Fe3+ as well as the electrodeposition/dissolution (Fe2+/Fe0) are observed in Figure 3-14. Note that the solution contained only FeCl2 and thus the limiting current for oxidation of Fe2+ to Fe3+ is much higher than that of the reduction of Fe3+ to Fe2+, as expected.

The CV for the Fe2+/Fe0 reaction appears to be sluggish, with a wider separation between plating

73 and stripping which is similar to electrodeposition/dissolution onto a glassy carbon electrode. The corresponding exchange current densities estimated from linear sweep voltammetry for plating at a CFME and glassy carbon were similar (0.35 ± 0.1 and 0.36 A/m2, respectively) indicating that the substrate material did not affect kinetics.

100%

80%

60%

40% Glassy Carbon PlatingEfficiency Untreated CFME 20% Electrochem Pretreated CFME

0% -1.80 -1.60 -1.40 -1.20 -1.00 -0.80 Plating Potential (V vs Ag/AgCl)

Figure 3-15. Coulombic plating efficiency as a function of potential for iron electrodeposited from 1:1:4

o FeCl2:ChCl:EG at 80 C.

Coulombic plating efficiencies, calculated by integrating under the current-time curve as described in Chapter 2, are shown as a function of potential for iron electroplated from the 1:1:4

FeCl2:ChCl:EG electrolyte. Two CFMEs, one untreated (“as received”) as well as an electrochemically pretreated fiber are compared to a glassy carbon substrate (Figure 3-15). All three are in good agreement and show a relatively high plating efficiency. Furthermore, the presence of electrodeposited iron was confirmed using scanning electron microscopy. Figure 3-

16 shows 10 mAh of iron deposited onto a CFME at 5 mA/cm2 where the iron deposit is strongly attached to the surface of the substrate material. The foregoing results suggest that iron

74 electrodeposition/dissolution is not dependent on the nature of the carbon surface. By using carbon felt electrodes inside a small-scale flow cell iron deposition is expected to occur upon charge and the use of carbon felt electrodes is not the cause for battery failure.

(a) (b)

2 o Figure 3-16. 10 mAh of iron plated at 5 mA/cm onto a CFME (80 C) from a 1:1:4 FeCl2:ChCl:EG electrolyte. A section along the (a) length as well as the (b) end are shown; the end had been broken off the CFME assembly which

caused some iron to flake off.

3.4 Conclusions

It is important to characterize electrochemical processes at carbon felt electrodes, particularly since they are widely used in flow batteries. This is often difficult due to their porous structure and the complex geometry of their surface. Therefore, electrodes fabricated from single carbon fibers

(CFMEs) have been developed to gain a better understanding of the electrode-electrolyte interface.

Due to enhanced diffusion to microelectrodes, CFMEs provide a useful tool for studying electrokinetics. CFMEs are convenient for use in resistive media, such as non-aqueous electrolytes due to their low currents; thus, ohmic overpotentials remain small.

Two applications for CFMEs have been demonstrated here and a third application will be presented in detail in Chapter 4. CFMEs were used to investigate capacity fade in a typical non- aqueous RFB electrolyte consisting of 0.01M V(acac)3 and 0.1M TEABF4 in acetonitrile. It was shown that over time, the current response at the CFME decreases. Felts used in non-aqueous systems are exposed to extreme potentials which can degrade the felt contributing to capacity fade observed with battery cycling. Thus, it is recommended that new electrode materials are explored for these types of systems.

Secondly, CFMEs had been used to study iron deposition from an ionic liquid type electrolyte consisting of 1:1:4 FeCl2:ChCl:EG. The exchange current densities and coulombic plating efficiency for iron deposition at a CFME was similar to that of a typical glassy carbon electrode indicating that the carbon surface does not play a major role. Furthermore, SEM imaging revealed that the iron deposit was smooth and strongly attached to the surface of the fiber. Thus, iron deposition in a RFB system is not hindered by using carbon felt electrodes.

Chapter 4. Carbon Electrode Pretreatment Effects

4.1 Introduction and Literature

Recently, attention has been directed toward improvement of the electrochemical properties of carbon based electrode materials. Modifications via thermal treatments [65, 70], chemical oxidation [68, 69, 71, 73], acid treatments [64, 66, 67], or electrochemical oxidation [72, 78-80,

169, 170] are thought to enhance electrochemical activity. The presence of oxygen containing functional groups has been shown to directly affect the kinetics; surface oxides resulting from the aforementioned treatments are thought to act as active sites, catalyzing flow battery reactions.

The all-vanadium flow battery chemistry is the most widely studied and a good example of the effects of electrode pretreatments. Some researchers have reported an increase [64-68, 70, 72, 74] while others reported a decrease [75, 76] in activity upon functionalization of carbon electrodes.

Many of these studies have been conducted using glassy carbon [76, 189-191], graphite [192, 193], carbon paper [175], multi-walled carbon nanotubes [194], or carbon composites [192, 195]. One

2+ + 2+ 3+ group concluded that the kinetics for the VO /VO2 are faster than the V /V reaction at a plastic formed carbon electrode, but found the opposite when using pyrolytic graphite [77].

It is apparent that kinetics depend on the nature of the electrode surface; thus, it is important to investigate factors affecting kinetics at electrode materials used in practical redox flow batteries

[179, 196, 197]. Thus, single carbon fiber microelectrodes discussed in Chapter 3 which have well-defined geometries are used to study the effects of surface treatments on the vanadium reactions.

4.2 Carbon Electrode Treatments

Modifications to carbon electrodes via thermal treatments, chemical oxidation, or acid oxidation were quantitatively assessed using single carbon fiber electrodes, along with surface characterization of bulk felt samples. These experiments were performed in order to evaluate effects of treatments found in the literature. Furthermore, electrochemical treatments such as those described previously employing glassy carbon electrodes [78, 170] were explored in depth since these types of treatments can be performed reversibly in-situ (ie electrodes remained in the electrolyte).

Prior to constructing CFMEs, bulk samples (4 mm thick GrafTech PAN-based) of carbon felts Untreated were given different treatments in an attempt to alter surface chemistry. These treatments had Thermal been previously reported in the literature. Two samples were refluxed in either nitric (70 wt%) or Nitric sulfuric acid (98 wt%) for five hours; longer treatment times have been shown to increase Sulfuric electronic resistance due to severe degradation of the felts [64, 66]. The acid oxidized felts were Fenton’s then rinsed to a neutral pH using deionized water and left to dry at room temperature. Fenton’s reagent was investigated as a means to chemically oxidize the surface: the felt sample had been submerged in 12 wt% H2O2 solution and 2 mM ferrous sulfate was added to catalyze the decomposition reaction; this procedure is outlined in the literature [68]. The reaction was allowed to proceed until all bubbling had ceased (approx. 3 hours). The felt was then sonicated in 0.1 M sulfuric acid solution for 30 minutes to rid the surface of iron, then rinsed to a neutral pH and dried at room temperature. One sample had been treated thermally in air at 400 oC for 30 hours; as described in the literature [70].

Electrochemical treatments were performed in a three electrode cell after fabrication of the CFME;

CFMEs were submersed in the electrolyte (3 M H2SO4 with or without 1.5 M vanadium) and held

78 at selected potentials ranging from −2.0 V up to +1.5 V vs MSE for 60 s; similar to experiments performed at glassy carbon electrodes [78, 170]. Before use, electrodes were “preconditioned” as described in Chapter 3. Carbon electrodes often require preconditioning or cleaning processes in order to obtain a stable surface response, and many treatment procedures can be found in the literature [11, 179-184]. The acid, chemical, and thermal processes described above represent a few examples of “preconditioning” or cleaning processes. Thus, CFMEs constructed with fibers that had been treated chemically, thermally, or with acid did not undergo electrochemical preconditioning.

4.3 Surface Analysis Results

Enhancement of electrochemical performance is often attributed to the addition of oxygen containing functional groups to the carbon surface, however, factors such as surface roughness and wettability may also play a role. Thus modifications of the electrode surfaces have been studied using scanning electron microscopy (Hitachi SEM), electrochemical capacitance measurements,

X-ray photoelectron spectroscopy (XPS) (PHI Versaprobe 5000 Scanning X-Ray Photoelectron

Spectrometer), and contact angle measurements (using a goniometer). In order to investigate the effects of treatment on wettability and functionalization of the carbon felt surface, bulk samples were needed for contact angle measurements as well as XPS. Thus, to evaluate electrochemical treatment of the felts, ‘test strip” electrodes were constructed using 2 cm×2 cm felt samples bonded with carbon ink (Acheson) to long grafoil strips (UCAR Carbon Co.). The back and edges of the grafoil strips were electrically insulated using non-conductive epoxy so that only the carbon felt was exposed to solution. The felt was then submerged in 3 M H2SO4 for treatment while the grafoil had been used to make electrical contact. All experiments were performed in a standard three-electrode cell employing a platinum counter electrode and saturated Hg/Hg2SO4 reference.

A treatment time of 20 minutes was used to ensure treatment of the bulk sample, however, XPS had also been performed on samples treated for only 60 seconds; results did not show a significant difference in the O/C ratio.

4.3.1 Scanning Electron Microscopy

Untreated Sulfuric Nitric

Fenton’s Thermal Electrochem. Precond.

Figure 4-1. SEM micrographs to show surface morphology of PAN-based fibers after various surface treatments.

Different treatments were shown to affect the physical state of the surface (Figure 4-1). Both acid and chemical treatments appeared to increase surface roughness; pitts or small holes can be observed for fibers treated with nitric or sulfuric acid and fibers treated using Fenton’s reagent appeared as if layers on the surface were peeling. The untreated felts were contaminated with an insulating aluminum-containing compound, which was not present after preconditioning or the thermal treatment (refer to discussion in Chapter 3). Electrochemical, acid, and thermal treatments

80 to carbon, such as those described here, have been known to cause damage within the crystal lattice; such defects also expose more surface sites [198]. While SEM provides qualitative insight, the active surface area can be quantified using electrochemical capacitance measurements.

4.3.2 Electrochemical Capacitance

The electrochemically active surface area of CFMEs was monitored by electrochemical

4 capacitance. As discussed in Chapter 3, capacitance was estimated by CVs in H2SO4 . In Table

4-1, values for the CFMEs are compared to the untreated fibers in order to obtain a relative roughness factor which can be used to normalize kinetic data as shown. The capacitance value for the electrochemically preconditioned CFMEs is shown for comparison. All values listed represent an average of 5 different CFMEs.

Table 1. Electrochemical capacitance values for various CFMEs tested in 3 M H2SO4; each data point represents an

average of 5 different electrodes.

Treatment Average Capacitance Relative (µF/cm2) Capacitance Untreated 29 ± 7 1.0 Thermal 290 ± 65 9.9 Nitric Acid 490 ± 110 17 Sulfuric Acid 23 ± 4 0.8 Fenton’s Reagent 27 ± 5 0.9 Electrochem. Preconditioning 280 ± 19 9.6

4 Differential capacity, Cd is equal to the rate of change of surface charge density of the double layer with potential (Cd = dQ/dE). For cyclic voltammetry the scan rate, v, represents a change in potential with a change in time (v = dE/dt). The differential capacitance current, ic, is equal to dQ/dt and thus, Cd = (1/v)(dQ/dt) = ic/v.

From Table 4-1, both electrochemical preconditioning and thermal treatments increased the capacitance (or assumed active electrochemical surface area) by 10 times that of CFMEs fabricated using the untreated fibers. Furthermore, the CFMEs fabricated from fibers that underwent nitric acid treatment showed an increase in capacitance that was 17 times that of the untreated electrodes; implying that there is a much larger increase in active surface area. This is not surprising in that

SEM revealed roughened surfaces for fibers that underwent thermal, electrochemical, and nitric acid treatment. However, CFMEs fabricated using fibers from felt treated using sulfuric acid and

Fenton’s reagent showed no significant change in active surface area, despite the fact that the SEM micrographs appeared to show roughened fiber surfaces. One explanation for this is that sulfuric acid and Fenton’s reagent treatments were not effective in “preconditioning” or cleaning the surface to rid the carbon of non-conductive alumina particles discussed in Chapter 3.

4.3.3 XPS Analysis

Survey scans in the range of 0 – 1000 eV were carried out to identify surface elements; example spectra are shown in Figure 4-2. High resolution XPS of the C1s and O1s peaks was performed on felt samples held a different potentials to investigate the effects of treatments on the oxygen to carbon (O/C) ratio. The C1s peak located at 284.8 eV was used as charge reference in determining the binding energies.

Figure 4-2. Survey scans for the treated felt samples, compared to the untreated sample.

A summary for the chemical spectral data obtained from curve fitting the high resolution XPS spectra for the carbon and oxygen peaks is shown in Table 4-2. A large increase in the oxygen to carbon ratio (O/C) for the felt sample treated using Fenton’s reagent is observed (32 %, compared to 9 % for untreated); however, analysis of the full survey indicates that this due to iron oxides present on the surface and not oxidation of the carbon surface (Figure 4-2). Thermal treatment (at

400 oC) decreased the O/C ratio to less than 3 %; this is not surprising in that decomposition of

o functional groups occurs at temperatures above 250 C [179]; the CVs in 3 M H2SO4 utilizing

CFMEs fabricated from thermally treated fibers lacked oxidation and reduction peaks for the quinone/hydroquinone surface groups, which further supports this claim. The two acid treatments did not significantly alter the surface chemistry. Electrochemical preconditioning did appear to greatly increase the amount of oxygen containing functional groups present on the carbon surface, therefore investigated in further detail.

Table 4-2. Chemical spectra obtained from curve fitting the high resolution C1s peaks.

Treatment O/C ratio Untreated 9.0 Thermal 2.9 Sulfuric Acid 7.1 Nitric Acid 9.1 Fenton’s Reagent 32* Electrochem. Precond. 34

The electrochemical preconditioning treatment (i.e. 50 cycles scanning from ‒0.5 V to

+1.3 V (MSE) at 100 mV/s) was shown to significantly increase the amount of surface oxides present on the carbon. These preconditioning CV cycles ended at a positive potential (they were not scanned negative again after scanning up to +1.3 V). Bourke et al. [78, 170] reported that kinetics of the vanadium reactions at glassy carbon depend on the applied potential for which glassy carbon electrodes were treated. Consequently, it is of interest to investigate electrochemical treatment as a function of applied potential.

The oxygen to carbon ratios for electrodes treated at different potentials are summarized in Table

4-3. Note that these electrodes all underwent the preconditioning treatment to clean the surface and then were subsequently treated by holding the specified potential for 20 minutes. As expected, electrodes that were treated at high positive potentials showed a greater percentage of oxygen present on the surface. When held at a potential of + 1.5 V (MSE) the oxygen to carbon (O/C) ratio increased to 0.49 compared to 0.15 for an electrode held at −2.0 V (MSE). This large change is further illustrated in Figure 4-3 where the survey scans are compared (intensities had been normalized to the C1s peak height). The amount of oxygen present on the felt treated at +1.5 V was clearly much larger compared to the felt treated at −2.0 V. The untreated felts had a low O/C ratio (0.09) which was increased after preconditioning as the preconditioning scans ended at a

84 potential of +1.3 V (MSE). The slightly higher O/C ratio of the reduced samples compared to the untreated sample may be attributed to surface oxides that were not reduced after preconditioning.

Thus, oxidation treatment increases the surface oxygen content while reduction treatment decreases it.

High resolution XPS of the C1s and O1s peaks (Figure 4-4) provide some insight onto the change in surface chemistry and the functional groups present on the carbon surface after treatment at high positive overpotentials. From the high resolution C1s fits, the emergence of two peaks located at

286.5 eV and 289 eV indicate the presence of phenolic and carbonyl groups [199] on the surface which was oxidized at +1.5 V ( MSE). The high resolution O1s peaks also signify the presence of phenolic and carbonyl groups (peaks at 532.5 eV and 532 eV). These peaks were not observed for the untreated or the reduced electrodes.

1.5

+1.5 V -2.0 V

1.0

0.5

Normalized Intensity

0.0 280 285 290 295 525 530 535 540 Binding Energy (eV)

Figure 4-3. Typical XPS survey scans of carbon felt samples after electrochemical treatments in 3 M H2SO4 after reduction at ‒2.0 V (blue) compared to oxidation at +1.5 V (red). The two spectra were normalized with respect to

the intensity of the C1s peak.

Table 4-3. Chemical spectra obtained from curve fitting the high resolution C1s peaks.

Sample Elemental Composition (atomic %) C O O/C ratio Untreated 91.5 8.5 0.09 Activated 74.8 25.2 0.34 Oxidized at +1.5V 67.2 32.8 0.49 Oxidized at +1.3V 71.5 28.5 0.40 Oxidized at +1.0V 88.2 11.8 0.13 Reduced at -0.5V 86.3 13.7 0.16 Reduced at -0.9V 87.4 12.6 0.14 Reduced at -2.0V 86.9 13.1 0.15

Intensity (a.u.) Intensity

295 290 285 536 532 528 Binding Energy (eV) Binding Energy (eV) (a)

Intensity (a.u.) Intensity

295 290 285 536 532 528 Binding Energy (eV) Binding Energy (eV) (b)

Intensity (a.u.) Intensity

Intensity (a.u.)

295 290 285 536 532 528 Binding Energy (eV) Binding Energy (eV) (c) Figure 4-4. Curve fitted high-resolution C1s and O1s spectra for the (a) untreated sample compared to felt samples

that had been oxidized at +1.5V vs Hg/Hg2SO4 and (c) reduced at -2.0V.

4.3.5 Contact Angle Characterization

To evaluate the effect of applied potential on the wettability of the electrode surface, the contact angle of 1.0 µL water droplets on the felt surface were measured (as shown by Figure 4-5). From

Figure 4-5c it is obvious that the untreated felt samples were hydrophobic and in Table 4-4 they were shown to exhibit an average contact angle of 126o. However, upon electrochemical oxidation, wettability of the felt is increased and becomes hydrophilic when held at a potential of

+1.5 V vs Hg/Hg2SO4. Electrochemical preconditioning of the electrodes has also been shown to increase felt wettability and thus, increase effective surface area. An increase in hydrophilicity could indicate either a presence of polar groups on the surface or a decrease in surface tension due to structural damage to the carbon. This increase in wettability would allow for better utilization of felt electrode surface area which would improve battery performance. At −2.0 V, the angle was measured to be lower than the untreated sample; this could be attributed to reduction of sulfate which adheres to the surface of the felt.

(a) (b) (c)

Figure 4-5. (a) Schematic and (b) photograph showing contact angle measurements of 1.0 µL water droplets on the

felt surfaces; (c) shows three water droplets on the surface of an untreated (as-received) piece of felt.

Table 4-4. Contact angles for various felt samples; a decrease in contact angle indicates an increase in felt

wettability.

Treatment Angle (degrees) Untreated 126 ± 1.8 Activated Hydrophilic Oxidized at +1.5 V Hydrophilic Oxidized at +1.0 V 102 ± 1.5 Reduced at − 0.9 V 132 ± 3.1 Reduced at − 2.0 V 42 ± 5.4

It should be noted that thermally treated felts were hydrophilic whereas those treated using

Fenton’s reagent and nitric or sulfuric acid did not yield reproducible results. For example, the wettability of the felts was not homogeneous and varied with location on the felt surface (most areas were hydrophilic but others did not absorb the water droplets). This further emphasizes the need for using a simplified electrode such as the CFME to study reaction kinetics.

Factors such as wettability or the presence of contaminants on the bulk felt samples highlight the importance for using a CFME where the surface of the electrode is well-defined and these factors can be controlled. The foregoing results suggest that treatments to carbon felt electrodes may

88 enhance active surface area as shown by capacitance measurements after removing impurities from the surface or damaging the carbon (pitting the surface) to expose more area (although pore size was difficult to determine due to the very small size). X-ray photoelectron spectroscopy results suggest that some electrode treatments alter the surface chemistry. For example, the thermal treatment decreased the amount of surface oxides present. The presence of surface oxides could be controlled using electrochemical treatments and thus will be explored as a means to understand what role surface oxides play in catalyzing the vanadium reactions.

4.4 Applications to the All-Vanadium System - Carbon Fiber Electrodes for Understanding of Electrokinetics

Improvements to the all-vanadium flow battery, such as enhancement of reaction kinetics, can be implemented once a concrete knowledge of the reaction mechanism and electroactivity is established. The effect of electrochemical treatment on electrochemical kinetics at individual fibers of a carbon felt electrode were investigated. The novel CFMEs described in Chapter 3 were

2+ 3+ 2+ + used to study the kinetics of V /V and VO /VO2 using slow-scan linear sweep voltammetry

(LSV) and electrochemical impedance spectroscopy (EIS) (these electrochemical reactions are shown below by equations 1 and 2). This work was done in collaboration with Dr. Noel Buckley’s group at the University of Limerick, Ireland.

V2+ ¬¾®V3+ + e- (1)

+ - + 2+ VO2 + e + 2H VO + H2O (2)

Before use, electrodes were preconditioned as discussed previously. In addition, to minimize any effects of electrode history, electrodes were also subjected to a normalization treatment as described by Bourke et al. [78, 80, 170] between experiments, consisting of three cycles, each of 89

60 s at ‒2.0 V and 60 s at +1.5 V. All electrochemical experiments were performed in a standard three-electrode cell with a potentiostat/galvanostat (Solartron SI 1287) and an impedance phase analyzer (Solartron SI 1260). The cell had a platinum counter electrode and saturated Hg/Hg2SO4 reference (mercury/mercurous sulfate electrode, MSE) inside a Luggin capillary. The electrolyte was thermostatted at 25 °C and deaerated by purging with nitrogen for 20 minutes; a nitrogen blanket was maintained over the electrolyte during experiments.

All chemicals were analytical grade and were used without further purification. The VO2+ electrolyte solutions were prepared using sulphuric acid and vanadyl sulphate (VOSO4) supplied

2+ + by Sigma Aldrich. Solutions containing V and VO2 were prepared by reducing or oxidizing the

VO2+ solution in a flow cell, which were then used to make mixed electrolyte solutions. The details of this procedure have been previously reported [200]. Vanadium electrolytes were titrated using reagent grade KMnO4 (Sigma Aldrich) to verify the concentration of vanadium species.

Experiments had been carried out at various concentrations (from 0.025 M up to 1.5 M total vanadium). However, as expected, the rate constants were not largely affected by a change in concentration; 1.5 M vanadium was chosen for this study because it most closely represented flow battery electrolyte.

4.4.1 Significant Results

To study the effect of electrochemical treatment on the kinetics of V2+/V3+ the electrodes underwent reduction at −2.0 V (MSE) and oxidation at +1.5 V (MSE) as described above. The resulting LSV traces are shown in Figure 4-6a where the oxidized electrode surface (red, solid lines) exhibits higher currents and steeper slopes when compared to the reduced surface (blue, dashed lines).

2+ + The same CFME was tested in a VO /VO2 electrolyte after undergoing either oxidation at

+1.5 V (MSE) or reduction at −2.0 V (MSE) as described above. The resulting LSV traces are shown in Figure 4-6b where the reduced electrode surface (blue, dashed lines) exhibits higher currents and steeper slopes when compared to the oxidized surface (red, solid lines). These results suggest that the kinetics of V2+/V3+ reaction are enhanced by oxidation treatment while the kinetics

2+ + of VO /VO2 are enhanced by reduction treatment. Similar results were obtained using EIS to monitor the kinetics; EIS results will be discussed in detail in the proceeding section. Furthermore, in Figure 4-6 the electrode was alternately oxidized and reduced for a total of 6 iterations (one iteration consists of a reduction treatment and an LSV followed by an oxidation treatment and an

LSV; for clarity, only the LSVs from iterations 2, 4, and 6 are shown). This demonstrates excellent reproducibility and indicates that the effect is not the result of a change in surface area.

(a) (b) 8

) 30

) Reduced -2 -2 Reduced Oxidized Oxidized 4 20

10 0 0

-4 -10

Current Density (mA cm

Current Density (mA cm -20 -8 -0.2 -0.1 0.0 0.1 0.2 -0.2 -0.1 0.0 0.1 0.2 Overpotential (V) Overpotential (V)

Figure 4-6. LSVs (5 mV/s) at a CFME in (a) V2+/V3+ electrolyte performed using the same electrode in the same

electrolyte but after two different treatments; reduction at -2.0 V (MSE) or oxidation at +1.5 V (MSE). And (b)

2+ + LSVs (5 mV/s) at a CFME in VO /VO2 electrolyte performed using the same electrode in the same electrolyte but after two different treatments; reduction at -2.0 V (MSE) or oxidation at +1.5 V (MSE). The LSVs from cycles 2, 4,

and 6 are shown.

4.4.1.i Kinetic Parameters from EIS and LSV

When using EIS as a tool to investigate reaction kinetics at the surface of the working electrode, the open circuit voltage is used as the DC potential. Since the chosen amplitude for the AC perturbations is small (<20 mV), the linearized Butler-Volmer kinetic model (ex ≈ 1+x) which relates current to overpotential can be used (this concept had been introduced in Chapter 2 for determination of kinetics from LSV). For a simple electron transfer reaction,

푘 [Oxidized species] + 1 electron ↔ [Reduced species], this relationship is shown to be:

nF (6) 푖 = −푖 ( ) 휂 표 RT

RT (7) 푅푐푡 = ( ) nFi표

Where - 휂/i has units of resistance and can thus be thought of as the charge transfer resistance,

Rct. Furthermore, if the concentrations of the oxidized and reduced species are equal, the current- overpotential relationship can be used to determine ko, the heterogeneous rate constant.

푏푢푙푘 (1−훼) 푏푢푙푘 (훼) (8) 푖표 = FAk표퐶표 퐶푅

푏푢푙푘 푖표 = FAk표퐶 (9)

Similarly, rate constants using LSV (at 5 mV/s) were estimated from the linear region (±20 mV overpotential); current is plotted as a function of the overpotential, for which the slope of the line

o is yields the exchange current density, io as shown by equation 10. The rate constant, k , can be calculated once exchange current density is known using equation 9.

푅푇 (10) 푖 = ∗ 푆푙표푝푒 표 푛퐹

The kinetic rate constants estimated from LSV and EIS are summarized in Table 4-5 where each value reported is the average of four different CFMEs. The actual values calculated using the two different techniques vary slightly; however, the relative values are consistent. Rate constants obtained using both techniques indicate that the kinetics of V2+/V3+ are approximately an order of

2+ + magnitude faster on an oxidized surface. Conversely, kinetics of VO /VO2 are slower on an oxidized surface: treatment at +1.5 V (MSE) decreases the rate constant by a factor of four. These results are consistent with previous findings from our group for various carbon electrode materials

[78, 80] and they demonstrate that electrochemical treatments also apply to felt electrodes used in typical redox flow batteries.

Table 4-5. Comparison of kinetic rate constants, ko, after treatment at +1.5 V (MSE) to that after treatment at -2.0 V

2+ 3+ 2+ + o (MSE) for both V /V and VO /VO2 reactions on CFMEs at 25 C; rate constants were estimated from both

LSV and EIS and each value is the average of four different CFMEs.

2+ 3+ 2+ + V /V VO /VO2 -6 -6 Electrode Treatment ko (10 cm/s) ko (10 cm/s) LSV EIS LSV EIS +1.5 V 5.8 ± 1.9 21 ± 1.3 17 ± 6.0 36 ± 7.0 ‒2.0 V 0.6 ± 0.3 1.0 ± 0.4 73 ± 29 170 ± 40

As shown by XPS, treatment at +1.5 V results in the addition of oxygen containing functional

2+ + groups. The results presented are counterintuitive since the VO /VO2 reaction involves the transfer of oxygen. Due to a more complex ligand structure, functionalization of the carbon has been believed to enhance kinetics, however, the opposite effect is being reported here in that surface oxides seem to hinder the kinetics. Furthermore, this effect depends on the applied potential as first noted by Dr. Andrea Bourke for glassy carbon electrodes [78, 170]. The effects

2+ 3+ 2+ + of treatment potential on electrode kinetics for the V /V and VO /VO2 reactions had been studied in collaboration with Noel Buckley’s group at the University of Limerick in Ireland. This work is summarized below.

4.4.1.ii Effects of Treatment Potential

To examine the effect of oxidation treatment potential on V2+/V3+, an electrode which had initially been reduced at − 2.0 V (MSE) was then treated at selected positive potentials. Typical results are shown in Figure 4-7(a) where the dashed line represents a typical Nyquist plot for the initial reduced (“baseline”) electrode surface and the solid lines represent the spectra for that same electrode after treatment at selected positive potentials. It can be seen that as the treatment potential is made progressively more positive, the charge transfer resistance becomes smaller, indicating faster kinetics for V2+/V3+.

(a)

) 2 10

 -0.7 V -0.3 V -0.5 V -Z'' ( +1.5 V 0 0 10 20 30 40 50 60 2 Z' ( cm )

(b)

)

2 10 -2.0 V

 -1.7 V -1.5 V

-Z'' ( -1.2 V 0 0 10 20 30 40 50 60 2 Z' ( cm ) Figure 4-7. Nyquist plots for a CFME in V2+/V3+ electrolyte after (a) oxidation and (b) reduction treatments at the potentials indicated. The dashed blue (a) and dashed red (b) lines correspond to the initial (baseline) treatments at

− 2.0 V and +1.5 V (MSE), respectively. Each spectrum is represented by a continuous line. Selected frequencies

are indicated: 100 Hz ( ), 10 Hz ( ), and 1 Hz ( ).

The effect of reduction treatment potentials on V2+/V3+ kinetics were investigated in a similar manner; the corresponding EIS results are shown in Figure 4-7(b), where the dashed line represents a typical Nyquist plot for the initial oxidized (baseline) electrode, and the solid lines represent curves after treatment at selected negative potentials. As the treatment potential is made progressively more negative, the charge-transfer resistance becomes larger indicating slower kinetics.

The electrochemical rate constant ko was estimated from the charge transfer resistance (as shown by equations 7 and 9) for each of the Nyquist plots in Figure 4-7. These rate constants are plotted as a function of treatment potential as shown in Figure 4-8. For oxidation treatment (O), at potentials more positive than −0.7 V the activity begins to increase; this trend continues as the potential is made more positive with a rapid increase in activity between −0.5 V and +0.5 V until the effect appears to approach saturation at around +1.0 V. For reduction treatment (R), at potentials more negative than −0.8 V the activity begins to decrease; this trend continues, as the potential is made more negative with a rapid decrease in activity between −0.8 V and −1.5 V.

8 R O 6

)

-1

Oxidation

cm s 4

-5

Reduction

(

k 2

V2+/V3+ 0 -2 -1 0 1 Treatment Potential (V) vs. MSE

2+ 3+ Figure 4-8. Rate constants, ko, for V /V plotted against oxidation (O) and reduction (R) treatment potentials. The

data in O and R are from Nyquist plots such as those in Fig. 7(a) and 7(b), respectively. The arrows show the

oxidation or reduction treatment potential relative to the initial baseline potential.

2+ + The effect of treatment potential on VO /VO2 kinetics was similarly investigated. The

2+ + electrochemical rate constants ko for VO /VO2 , estimated from the charge transfer resistances from EIS are plotted against treatment potential in Figure 4-9. As the treatment potential is made progressively more positive, the charge-transfer resistances from the EIS become progressively

2+ + larger, indicating a decrease in ko for VO /VO2 . These results contrast sharply with those in

Figure 4-8 where the opposite trend is observed for V2+/V3+. Furthermore, as the treatment potential is made progressively more negative, the charge-transfer resistances become

2+ + progressively smaller, indicating increased kinetics for VO /VO2 (contrasting with Figure 4-8 where the opposite trend is observed for V2+/V3+).

Figure 4-9 shows that for oxidation treatment (O), as the potential is made progressively more positive than +0.7 V, the electrode activity decreases. However, for treatment at potentials more negative than +0.4 V (R), the activity begins to increase; this trend continues as the potential is

96 made more negative with a rapid increase in activity between +0.1 V and −0.4 V until the effect appears to approach saturation near −0.6 V. These results are in direct contrast to those observed for the V2+/V3+ reaction (Figure 4-8).

R O

)

-1

10 Reduction

cm s

-5 Oxidation

(10

k 5

VO2+/VO + 2

0 -2 -1 0 1 Treatment Potential (V) vs. MSE

2+ + Figure 4-9. Rate constants, ko, for VO /VO2 plotted against oxidation (O) and reduction (R) treatment potentials.

The arrows show the oxidation or reduction treatment potential relative to the initial baseline potential.

The results reported here are not characteristic to CFMEs. In fact, these trends have been observed using other carbon electrode materials as reported by Bourke et al [78]. These include glassy carbon (GC), graphite (Gr), carbon paper (CP), and reticulated vitreous carbon (RVC). The

2+ + normalized activity for VO /VO2 as a function of treatment potentials for different carbons is shown by Figure 4-10. The trends in activity are consistent for each of the carbon electrode materials studied. These findings help to clear up the controversy within the literature for kinetics of the vanadium reactions at various carbon electrodes.

2+ + Figure 4-10. Normalized activity for VO /VO2 as a function of treatment potentials for various carbon electrodes

(GC: glassy carbon, F: CFME, RVC: reticulated vitreous carbon, CP: carbon paper, Gr: graphite). This figure is

from ref [78].

4.4.1.iii Development of a Mechanistic Model

Thus far, it has been shown that oxidation of the carbon electrode at +1.5 V has a much larger

2+ 3+ 2+ + 2+ 3+ effect on catalysis of V /V than on hindering the VO /VO2 redox couple. If the V /V reaction is assumed to be a single electron transfer reaction (outer-sphere reaction), one would expect minimal effects of surface catalysis. For the V2+/V3+ couple, both hex-aqua cations possess near perfect octahedral coordination of the first hydration shell [201-203], and thus structural reorganization energy would be expected have minimal effects [204] (Figure 4-11). Furthermore, a previous report on the dependence of inner-sphere electrode reactions on electrode surface supports these findings in that electrode material was shown to greatly influence oxidation kinetics of the aqueous V2+ cation [168]. In order to provide and understanding of the electrode kinetics, an impedance model is developed.

(a) (b)

2+ + Figure 4-11. Electronic structures of the (a) vanadyl (VO ) and (b) vanadate (VO2 ) ions.[205-207]

Electrochemical impedance spectroscopy is a technique that allows for an understanding of interfacial phenomena within an electrochemical system. As shown in the previous section, EIS is useful for evaluation of charge-transfer parameters; this section will focus on reaction kinetics at the electrode/electrolyte interface. EIS allows for in-situ analysis of the surface of electrodes which makes it a valuable tool for this work.

For all electrochemical experiments, a standard three electrode cell was employed. The “DC potential” of the working electrode was set as the open circuit voltage or equilibrium potential, which is determined by the ratio of oxidized or reduced species in solution. An AC sinusoidal component with peak-to-peak amplitude of 10 mV is chosen for these experiments. Larger perturbations often yield a more favorable signal to noise ratio, however, one must consider the effects of mass transfer. An amplitude of 10 mV is necessary for this analysis because deviations from the equilibrium potential are small, thus mass transfer limitations are negligible and the kinetics are limiting. This enables analysis within the linear region where current and overpotential are directly proportional (typically, ±20 mV from open circuit) which is necessary for determination of linear kinetics. Such a system exhibits first order harmonics in that the excitation frequency ω yields a current response at the same frequency, which cannot be said for EIS in regions of high overpotential (tafel region, for example). Typical redox reactions occur within the frequency range of 102 to 105 Hz, thus the frequency range of interest was 0.2 to 20,000 Hz.

Figure 4-13 shows the Nyquist and Bode plots for CFMEs after oxidation treatment at

+1.5 V (MSE) or reduction treatment at −2.0 V (MSE). It is obvious that the charge transfer

2+ + resistance (Rct) for VO /VO2 at a reduced surface (represented by the blue dashed curve on the

Nyquist plot in Figure 4-13a) and V2+/V3+ at an oxidized surface (represented by the red curve on the Nyquist plot Figure 4-13b) are much smaller, indicating faster kinetics as previously discussed.

What is interesting is that the EIS data presented here do indicate that the mechanism for the

V2+/V3+ couple is more complex than a single electron transfer reaction (Figure 4-13b). These results are not characteristic of a typical microcylinder electrode, especially that of the reduced surface [208]. In the case of cylindrical diffusion to a microelectrode one would expect the Nyquist plot to show impedance to be parallel to the real axis at low frequency due to enhanced mass transport (Figure 4-12). The charge transfer loop is defined, however, a second loop is present in the low frequency range. These results indicate that there are two different processes.

Figure 4-12. Nyquist plot showing impedance of a typical electrocehmical reaction but with (a) semi-infinite linear

diffusion to a flat plate compared to diffusion to (b) a micro cylindrical electrode and (c) spherical diffusion to a

micro disk electrode.

100

(a) (b)

500

Reduced 400 Oxidized

)

Reduced cm 300 Oxidized  20

-Z" ( 200

) 4

2 10

cm 100

 2 0 0 10 20 30 40

-Z" ( 0 0 0 2 4 6 8 10 12 14 16 18 20 0 200 400 600 800 2 Z' (cm2) Z' (cm )

40 80

30 60

20 40

10 20

Phase Angle (-Theta)

Phase Angle (-Theta) 0 0.1 1 10 100 1000 10000 0.1 1 10 100 1000 10000 Frequency (Hz) Frequency (Hz)

2+ + Figure 4-13. Nyquist and Bode plots for CFME in 1.5 M total vanadium (1:1 molar ratio of (a) VO /VO2 or (b)

2+ 3+ V /V ) and 3 M H2SO4 (please note: the inset in the Nyquist plot for (b) represents a zoomed in version). EIS was performed using the under the same conditions but after oxidation (+1.5V vs Hg/Hg2SO4) or reduction (-2.0V vs Hg/Hg2SO4) pretreatments. This set of experiments had been repeated/cycled 8 times to ensure surface area was

not changing - only the data from cycle 6 are shown for clarity.

In order to gain insight on the reaction mechanism, an equivalent circuit was fit to the data. The model shown below (Figure 4-14) produced the best fits for both vanadium reactions as determined by χ2, the sum of squares, and plots of the residuals (Table 4-6 and Table 4-7). This model is consistent with a mechanism that involves a single adsorbed species, which reacts on the surface of the electrode and subsequently desorbs [208-210]. It was assumed that all available sites are equivalent and that each site can only hold one adsorbing species. It also assumes that there are

101 no interactions between adsorbate molecules on adjacent sites (the adsorbed species follows a

Langmuir isotherm).

Figure 4-14. Equivalent circuit model for a faradaic reaction involving one adsorbed species with subsequent

desorption; this model was used to fit all experimental data. Note that Rsolution is the high frequency solution resistance, Rct is the charge transfer resistance, and RA is the resistance due to adsorption. Subscripts “dl” and “A”

indicate processes associated with double layer charging or adsorption. [208-210]

The equivalent circuit model discussed above (Figure 4-14) was applied to EIS data for oxidized

2+ + and reduced CFMEs in a VO /VO2 electrolyte solution and found to produce good fits (Table 4-

6). It should be noted that constant phase elements5 were used in the model instead of capacitors.

Constant phase elements are used when there is a frequency dispersion [205]. This is attributed to the fact that the surface of the fibers are inhomogeneous and defects are present within the carbon structure. These characteristics cause the impedance to be distributed, which is why the Nyquist plot does not resemble the perfect semi-circle illustrated by the Randle’s circuit.

5 P-1 2 CPE T Parameter related to capacitance of a constant phase element (Fs /cm ). CPE P Constant phase exponent, relates to deviation of the straight capacitive line from 90o by angle α=90o(1-P). Note that when P=1, the constant phase element acts as an ideal capacitor. Impedance of an ideal capacitor vs constant phase element: 1 1 푍 = vs 푍 = 퐶 (푗휔퐶) 퐶푃퐸 푇(푗휔퐶)푃

102

Table 4-6. Parameters used to fit experimental EIS data for CFME in 1.5 M total vanadium (1:1 molar ratio of

+ + VO /VO2 ) and 3 M H2SO4. Note that Rsoln is the high frequency solution resistance, Rct is the charge transfer

2 P-1 2 resistance, and Ra is the resistance due to adsorption; all have units of W ×cm . CPE T (units of Fs /cm ) is related to capacitance of a constant phase element and CPE P is the constant phase exponent which relates to deviation of the straight capacitive line from 90o by angle α=90o(1-P). Note that when CPE P=1, the constant phase element acts as an ideal capacitor. Subscripts “dl” and “a” indicate processes associated with double layer charging or adsorption.

Reduced Oxidized Element Value Error (%) Value Error (%) Rsoln 0.81 1.3 1.1 1.5 -3 -3 CPEdl T 6.0 × 10 9.3 1.7 × 10 3.9 CPEdl P 0.57 2.1 0.70 0.80 Rct 1.1 2.5 12.2 1.1 -2 -2 CPEa T 3.8 × 10 9.0 8.9 × 10 9.1 CPEa P 0.81 2.7 1 6.6 Ra 0.92 5.1 3.3 9.7 χ2 1.7 × 10-5 9.9 × 10-4 Sum of Squares 7.0 × 10-4 7.8 × 10-2

Table 4-7. Parameters used to fit experimental EIS data for CFME in 1.5 M total vanadium (1:1 molar ratio of

2+ 3+ V /V ) and 3 M H2SO4. Note that Rsoln is the high frequency solution resistance, Rct is the charge transfer resistance,

2 P-1 2 and Ra is the resistance due to adsorption; all have units of ×cm . CPE T (units of Fs /cm ) is related to capacitance of a constant phase element and CPE P is the constant phase exponent which relates to deviation of the straight capacitive line from 90o by angle α=90o(1-P). Note that when CPE P=1, the constant phase element acts as an ideal capacitor. Subscripts “dl” and “a” indicate processes associated with double layer charging or adsorption.

Reduced Oxidized Element Value Error (%) Value Error (%) Rsoln 0.72 N/A 0.67 N/A -5 -3 CPEdl T 4.3 × 10 0.88 1.0 × 10 6.6 CPEdl P 0.90 0.12 0.75 1.1 Rct 267 0.67 15 4.9 -4 -3 CPEa T 5.8 × 10 0.92 1.1 × 10 20.5 CPEa P 0.76 0.69 1 5.8 3 Ra 1.3 × 10 1.5 8.4 11.1 χ2 1.1 × 10-4 3.4 × 10-4 Sum of Squares 9.5 × 10-3 1.7 × 10-1

103

2+ + For the VO /VO2 reaction at an oxidized surface, all data was shown to be Kramers-Kronig (KK) compliant. This implies that the real part of the complex function can be calculated from the imaginary part and vice versa and that the data is linear, finite-valued, stable, and possess causality.

KK transformed data overlaid the original data and displayed no systematic differences [211, 212].

Simulations were in agreement with KK compliant experimental data for reduced and oxidized electrode surfaces (Figure 4-15).

2+ + For the VO /VO2 reaction at reduced surface, experimental data collected at lower frequencies in the range of 5 Hz to 0.2 Hz was not Kramers-Kronig compliant and was thus omitted from the analysis. Failure to comply is attributed to experimental noise. Nyquist and Bode phase angle plots obtained using simulated values indicated that this was a reasonable fit within the range of 20,000

Hz to 5 Hz (Figure 4-15).

)

2 0.4 2 4

 0.2



2 -Z" ( 0.0 -Z" ( 0.0 0.5 1.0 1.5 2.0 2.5 3.0 2 0 Z' (cm ) 0 5 10 15 Z' (cm2)

15 40

30 10

20 5 10

Phase Angle (-Theta)

Phase Angle (-Theta) 0 0 0.1 1 10 100 1000 10000 0.1 1 10 100 1000 10000 Frequency (Hz) Frequency (Hz) 2+ + Figure 4-15. Nyquist and phase angle plots for CFME in 1.5 M total vanadium (1:1 molar ratio of VO /VO2 ) and

3M H2SO4. Simulated values (solid line) in the range of 0.2 to 20,000 Hz are compared to experimental data (open

circles) for both (a) reduced and (b) oxidized electrodes; refer to Table 4-6 for values used in simulations.

104

500

)

2 ) 8 2 400

cm 6 300

  4 200

-Z" (

-Z" ( 2 100 0 0 0 200 400 600 800 1000 1200 1400 0 5 10 15 20 25 2 2 Z' (cm ) Z' (cm )

50 80 70 40 60 30 50 40 20 30 10 20

Phase Angle (-Theta) Phase Angle (-Theta) 10 0 0 0.1 1 10 100 1000 10000 0.1 1 10 100 1000 10000 Frequency (Hz) Frequency (Hz) Figure 4-16. Nyquist and phase angle plots for CFME in 1.5 M total vanadium (1:1 molar ratio of V2+/V3+) and 3

M H2SO4. Simulated values (solid line) in the range of 0.2 to 20,000 Hz are compared to experimental data (open

circles) for both (a) reduced and (b) oxidized electrodes; refer to Table 4-7 for values used in simulations.

When applying the same equivalent circuit (Figure 4-14) to the experimental data for the V2+/V3+ reaction at an oxidized electrode surface, the fit was good for frequencies in the range of 20,000

Hz to 5 Hz. Data collected at lower frequencies, in the range of 5 Hz to 0.2 Hz, was not KK compliant and was thus omitted from the analysis. Nyquist and Bode phase angle plots obtained using simulated values indicated that this was a reasonable fit within the range of 20,000 Hz to

5 Hz (Figure 4-16). Both the Nyquist and Bode phase angle plots for the reduced electrode surface in a 1.5 M vanadium V2+/V3+ solution are in agreement with simulated values found using the equivalent circuit model for a faradaic reaction involving one adsorbed species with subsequent desorption (Figure 4-16). All data were found to be KK compliant.

2+ + 2+ 3+ It can be concluded that both vanadium reactions, VO /VO2 and V /V consist of an electron transfer involving one adsorbed species. Unfortunately, EIS cannot be used to determine the

105 number of electron transfer steps or their individual rates. However, possible mechanisms for the

2+ + VO /VO2 reaction dealing with electron transfer coupled with adsorption have been proposed

[184].

2+ + It has been suggested that the rate determining step for the VO /VO2 reaction would be the reaction between monomeric vanadium(IV) and one of the dimeric vanadium complexes discussed

2+ + previously [213, 214]. This could be a potential explanation since rate constants for the VO /VO2 reaction are only affected by a factor of four when electrode surface is altered. This suggests that the rate limiting step may not be adsorption. However, the adsorption step for the V2+/V3+ reaction is likely to be rate limiting.

It is interesting that even at an oxidized electrode the V2+/V3+ reaction is slower than the

2+ + VO /VO2 reaction. Previous reports have presented evidence of mixed valence cation-cation

2+ + complexes for solutions of VO /VO2 in concentrated sulfuric acid [200, 213, 215]. Dimers such

4+ 3+ as V2O3 and V2O3 which contain two vanadium ions connected by an ‘oxygen bridge’ would allow for enhanced rates of electron transfer for inner-sphere mechanisms.

4.5 Concluding Remarks

The forgoing results strongly suggest that the nature of the surface of carbon electrodes has a significant effect on catalysis of the reactions occurring in a vanadium flow battery. Treatments to the felts change the surface tension which affect wettability and removal of surface contaminants enhance active surface area; both of which may improve the apparent kinetics and cause confusion within the literature. However, CFMEs which have a simple geometry had been used as a tool in this work (described in Chapter 3). Techniques had been developed to gain a fundamental understanding of reaction mechanisms including possible insight on reaction intermediates formed

106 at carbon electrodes. Results in aqueous vanadium electrolytes yield promise for CFMEs as a means to evaluate flow battery kinetics at the graphite-electrolyte interface.

It was shown that oxygen containing functional groups formed on the electrode during oxidation

2+ + at +1.5 V (MSE) hinder the reaction kinetics of the VO /VO2 redox couple while enhancing

2+ 3+ 2+ + V /V . Conversely, kinetics of the VO /VO2 couple are improved at a reduced electrode surface (-2.0 V) and the V2+/V3+ reaction is hindered. An oxidized surface had a significant effect on enhancement of V2+/V3+ but reduction of the surface did not have as large of an effect on

2+ + VO /VO2 . This was shown by kinetic analysis using two different techniques for four different fibers chosen at random (LSV and EIS). EIS suggests that the mechanisms are more complex than simple electron transfer and include subsequent adsorption and desorption steps.

These findings lead to improvement of the vanadium flow battery system and resulted in a publication in collaboration with Professor Noel Buckley’s group at the University of Limerick in

Ireland [81]. Charge and discharge cycles in a prototype flow cell were in agreement with findings using the CFME and demonstrated that simple electrode pretreatments in-situ resulted in decreased operating overpotentials and amongst other benefits, increased energy efficiency. Additional peer- reviewed articles that have been published from our group of which I was a co-author relating to this work are referenced as well [78-80].

107

Chapter 5. Conclusions and Future Work

5.1 Conclusions

Flow batteries are capable of interfacing with renewable sources such as wind and solar and thus, present an attractive solution for large scale energy storage. However, this technology is not widely commercially available. Two main areas for improvement have been investigated in this work. This chapter provides a summary of the three main results chapters (Chapters 2-4).

Recommendations for future work is also presented.

5.1.1 Ionic Liquids as RFB Electrolytes

Ionic liquids (ILs) were demonstrated as an alternative to current non-aqueous organic electrolytes.

These iron containing ILs are not only easy to synthesize but also made from safe and cost effective materials such as choline chloride and ethylene glycol. An understanding of the factors that affect the electrochemistry as well as how iron speciation influences the physical properties of the electrolytes was demonstrated. Specifically, solutions containing ≥ 4:1 molar ratio chloride to iron showed more desirable properties (lower viscosity and higher conductivity). X-ray absorption spectroscopy (XAS), X-ray photoelectron spectroscopy (XPS), and Raman spectroscopy data

- 2- indicate the presence of the tetrahedral iron chloride complexes, [FeCl4] and [FeCl4] in these electrolytes. However, when the chloride to iron ratio is less than 4:1, ethylene glycol complexes with the iron. Electrolytes containing this EG-Fe complex show an expanded window and thus, decrease in the H2 evolution side reaction which is beneficial for increasing coulombic efficiency for the electrodeposition/dissolution reaction. Furthermore, a smoother iron deposit was characteristic of metal deposited from such electrolytes, however, the physical properties of the

108 solution are hindered due to strong ion-ion interactions (i.e. conductivity for a 3:1 electrolyte is an order of magnitude lower than that of the 4:1 electrolyte).

As a result, an electrolyte containing 0.5:1:4 FeClx:ChCl:EG (where the Cl:Fe ratio is 4:1) was selected for proof-of-concept testing in a small-scale flow battery. This electrolyte showed reasonable coulombic plating efficiencies (77% at −1.0 V), fast kinetics (94 A/m2), high concentration of active species (1.3 M), as well as acceptable fluid properties. The battery had been cycled at 5 mA/cm2 and energy efficiencies of around 48 % are reported for the first three cycles. This demonstrates promise for the iron IL as an electrolyte for flow batteries and has been published in the Journal of the Electrochemical Society [52] with a second manuscript recently submitted.

5.1.2 The CFME, Technique Development

Carbon felt electrodes used in flow batteries have a porous structure and complex geometry. Thus, it is difficult to characterize electrochemical processes at their surfaces. Therefore, novel carbon fiber microelectrodes (CFME), created using single fibers extracted from the bulk carbon felt, were developed to gain a better understanding of the electrode/electrolyte interface. The CFME construction and cleaning process were discussed as well as potential applications.

For example, CFMEs are convenient for use in resistive media such as non-aqueous electrolytes due to their low currents; thus, ohmic overpotentials remain small. Thus, CFMEs could be used to investigate capacity fade within a typical non-aqueous electrolyte. It was shown that the current response for the fiber electrode which had been cycled in 0.01M V(acac)3 and 0.1M TEABF4 in acetonitrile decayed over time. However, a platinum electrode showed a stable current response.

It was found that the decay in current observed for the carbon electrodes was due to degradation

109 of the carbon. Thus, degradation of the carbon electrodes contributes to capacity fade observed with cycling of the V(acac)3 flow batteries.

Another application for the CFME was to study deposition of iron from an ionic liquid electrolyte.

Iron had been deposited onto CFMEs from a 1:1:4 FeCl2:ChCl:EG electrolyte. It was shown using

SEM that the iron deposit was smooth and strongly attached to the surface of the fiber.

Furthermore, the deposition/dissolution kinetics and coulombic plating efficiency at the fiber were similar to that of a glassy carbon electrode.

5.1.3 Carbon Electrode Pretreatment Effects

It is known within the literature that the nature of the carbon electrode surface influences reaction kinetics, however, this is not fundamentally understood. Thus, it is important to characterize reaction kinetics at electrode materials used in flow batteries. Use of the CFME allowed for exploration of effects of surface treatments on reaction kinetics at flow battery electrodes in a controlled fashion.

This work showed that oxygen containing functional groups formed on the electrode during

2+ + 2+ 3+ oxidation (+1.5 V vs MSE) hinder the reaction kinetics of VO /VO2 while enhancing V /V .

2+ + Conversely, kinetics of VO /VO2 are improved at a reduced electrode surface (−2.0 V vs MSE) and V2+/V3+ is hindered. In fact, surface treatments had a significant effect on V2+/V3+ where kinetics were an order of magnitude faster at an oxidized surface than that of the reduced surface.

Furthermore, EIS suggests that the mechanism for the V2+/V3+ reaction is more complex than simple electron transfer and includes subsequent adsorption and desorption steps.

These findings lead to improvement of the vanadium flow battery system and resulted in a publication in collaboration with Professor Noel Buckley’s group at the University of Limerick in

110

Ireland [81]. Additional peer-reviewed articles that have been published from our group of which

I was a co-author relating to this work are referenced as well [78-80].

5.2 Future Work

5.2.1 Other Halides and Hydrogen Bond Donors

By changing the ratio of the precursor materials in an iron-containing IL, specifically the chloride to iron ratio, the conductivity, viscosity, and electrodeposition of iron could be improved. Since the halide is believed to adsorb on the surface of electrodes, it is recommended to explore this in depth to gain insight on the double layer as well as the effects on metal deposition (additional metals, not just iron). Preliminary data for electrodeposition of iron from 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I can be round in Appendix A1.

Furthermore, other hydrogen bond donors should be explored. Urea and ethylene glycol are compared in Appendix A1. Additional alcohols are of interest due to their low freezing points, presenting the opportunity for a system with reasonable fluid properties at room temperature. For example, glyceline (1:2 ChCl:glycerol) has a freezing point of −20oC and is similar to ethylene glycol but with three –OH groups that can readily form bonds with the iron. This presents an interesting application for electrodeposition of metals in ionic liquids where the development of a fundamental understanding is needed.

5.2.2 Investigation of Alternative Active Metal Species

The all iron battery is limited by sluggish kinetics for the deposition/dissolution reaction at the negative electrode. A more facile electrodeposition reaction or even a redox reaction is desired.

Thus, exploration of alternative active metal species is recommended. Appendix A2 provides

111 preliminary results for metal chlorides such as tin, nickel, zinc, chromium, manganese, ytterbium, indium, vanadium as well as an iron-vanadium mixed electrolyte. Polyoxometalates which typically exhibit multiple oxidation/reduction reactions had been examined but were found to be insoluble in these ionic liquids; thus they are not recommended for further study.

Electrodeposition of tin, nickel, and zinc as well as an iron-zinc alloy are studied. Interestingly, electrodeposited zinc from a 1:1:4 ZnCl2:ChCl:EG electrolyte was not dendritic however, it did not adhere to the surface of the glassy carbon electrodes. Different substrate materials may be more appropriate or a different ratio.

5.2.3 Aqueous All-iron Flow Battery with Choline Chloride Supporting Electrolyte (and Choline as an Additive for Plating)

While the iron containing ionic liquids present an approach to non-aqueous flow batteries, it is recommended that choline chloride be explored as an additive for an aqueous iron flow battery.

Preliminary data is presented in Appendix A3. It is interesting to note that a

1:1:4 FeCl2•4H2O:ChCl:EG electrolyte can be created which shows an expanded potential window

(~750 mV), coulombic plating efficiencies near 100% down to ~−1.2 V vs Ag/AgCl, a high concentration of iron, as well as a reasonable conductivity without added acid at 50oC. It is thus recommended that this system be further explored. For example, effects of changing the molar ratios of the precursor materials, addition of acid to improve conductivity, possibly different organic cations with longer chain lengths (as well as NH4Cl).

5.2.4 Ionic Liquids for High Power Energy Storage Applications (Supercapacitors)

Another potential application for ionic liquids is electrochemical capacitors. Ionic liquids provide the opportunity to obtain a safer and more cost efficient alternative to contemporary non-aqueous

112 electrolytes as well as stability when operation under ambient conditions. Wide liquid phase range as well as low volatility are added benefits that could enable operation in extreme temperatures.

Few researchers have studied this class of electrolyte for supercapacitor applications. Physical and chemical properties of such systems have been studied, however, understanding of the double layer capacitance is lacking. Interaction with various carbon inks, carbon electrodes, and membranes has not been studied. This would require a deeper understanding of the double layer structure as well as further investigations to improve conductivity. CFMEs may provide a tool to understanding the electrode/electrolyte interface (i.e. surface tension and wettability effects).

Starting points for this work could include investigation of structural changes to the electrolyte as a result of anion and cation selection using a range of chemical characterization techniques such as XPS, XAS, IR and Raman spectroscopy. Electrochemical techniques such as voltammetry and

AC impedance could be used to gain an understanding of the effects of structural changes on conductivity, potential window and electrochemical capacitance. Furthermore, testing of the electrolyte in prototype cells would enable evaluation of its interaction with cell components such as carbon inks, electrodes, and membranes.

113

Appendix A1. Halides and Hydrogen Bond Donors

A1.1 Introduction

In Chapter 2, an ionic liquid containing iron chloride, choline chloride, and ethylene glycol is explored for use within a flow battery. It was shown that by changing the ratio of the precursor materials, specifically the chloride to iron ratio, the conductivity, viscosity, and electrodeposition of iron could be improved (example, a 4:1 Cl:Fe ratio was found to be best). Here the effect of the halide as well as the hydrogen bond donor are explored.

A1.2 Effects of the Halide

As expected, the electrolyte containing iodide exhibited enhanced fluid properties (iodide is a larger atom, and thus, there is more charge delocalization); the conductivity at 80oC was 4.6 mS/cm compared to bromide and chloride which are 2.8 and 2.4 mS/cm respectively. When comparing

2+ 3+ the exchange current densities for the Fe /Fe reaction in a 1:1:4 FeX2:ChX:EG, where X=Cl or

Br, the kinetics were faster in the Br electrolyte (estimated from LSV at a platinum microelectrode). However, this increase is not significant, implying that the mechanism is likely an outer sphere single electron transfer and not largely affected by the ligands bound to the iron

[204, 216]. It should be noted that the Fe2+/Fe3+ reaction could not be observed in an electrolyte containing 1:1:4 FeI2:ChI:EG due to a large oxidation current at ~0.71 V vs Ag/AgCl. This is

− − − likely attributed to the reaction I3 + 2e ↔ 3I which is known to occur at +0.74 V vs Ag/AgCl

[217].

114

Table A1-1. Conductivity and kinetics (at platinum) for electrolytes containing 1:1:4 FeX2:ChX:EG, where X=Cl,

o Br, or I at 80 C, N2-atmosphere. Note that only the exchange current estimated from LSV is provided; this is due to

2+ the fact that only Fe ions were present and thus, ko could not be easily estimated (cannot assume equal molar

reduced and oxidized species).

2+ o 2+/3+ Halide σ Fe only i Fe at Pt (mS/cm) (A/m2)

Chloride 2.4 93.9 Bromide 2.8 130.6 Iodide 4.6 n/a

Halide adsorption has been believed to play a significant role in the electrodeposition of transition metals. Adsorption increases in the order Cl- < Br- < I- [142] and thus plating experiments were carried out using electrolytes containing 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I. At glassy carbon electrodes it was found that the double layer capacitance (estimated from CV) was

17.8 µF/cm2 in the chloride containing electrolyte and 19.1 µF/cm2 in the bromide containing electrolyte. This is not surprising in that if the halide is present near the electrode surface, the Br- ion is larger than that of the Cl- ion and thus, expected to yield a larger double layer capacitance.

From the CVs in Figure A1-1 it is clear that iron deposition from the iodide electrolyte displayed sluggish kinetics compared to that of iron deposition/dissolution from chloride and bromide electrolytes. The onset for dissolution is shifted to a much more positive potential and an increase in separation between reduction and oxidation is observed. This is consistent with the fact that the iodide more strongly adsorbs which may hinder stripping of the iron deposit. Despite the fact that the kinetics for Fe2+/Fe0 are slowed, the plating efficiency is improved when switching from chloride to iodide as shown by Figure A1-2 where coulombic plating efficiency is plotted as a

115 function of plating voltage. In fact, even when iron was electrodeposited at −1.5 V (vs Ag/AgCl) the coulombic plating efficiency was near 100% in the iodide containing solution; the deposit was smooth and no electrolyte breakdown could be observed. However, the plating efficiency for iron deposited from 1:1:4 FeBr2:ChBr:EG is lower than that of the chloride or iodide containing electrolytes and decreases significantly at potentials lower than −1.3 V; this needs further investigation.

Normalized Current Normalized

-1.1 -0.8 -0.5 -0.2 0.1 0.4

Applied Potential (V vs Ag/AgCl)

o Figure A1-1. Cyclic voltammetry in solutions containing 1:1:4 FeX2:ChX:EG; X=Cl, Br, or I at 50 mV/s, 80 C.

100

Efficiency (%) Iodide 20 Chloride Bromide 0 -1.8 -1.6 -1.4 -1.2 -1.0 -0.8 -0.6 Plating Voltage (V vs Ag/AgCl)

Figure A1-2. Coulombic plating efficiency as a function of voltage at a glassy carbon substrate for solutions

containing 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I. Iron was plated at a selected potential for 5 minutes then

stripped for 5 minutes at a potential more positive of the electrodissolution peak (0.2 V for Br and Cl containing electrolytes and 0.4 V for I); efficiency was estimated from the number of coulombs passed during stripping divided

by coulombs during plating. Temperature was maintained at 80 oC.

116

From the SEM micrographs (Figure A1-3), it can be noted that the morphology of the electrodeposited iron does depend on the halides present in solution. This can be explained by the fact that the iodide solution was the most conductive with slower kinetics; ie if the activation overpotential is large compared to the ohmic overpotential, the current distribution will be uniform

[218] and thus the deposit will be more smooth and smaller features are expected. Nonetheless, it is interesting that different shapes are formed by simply changing the halide.

Figure A1-3. SEM micrographs of 4 mAh iron electrodeposited (at 0.2 mA/cm2) onto a glassy carbon substrate at

o 80 C from solutions containing 1:1:4 FeX2:ChX:EG, where X=Cl, Br, or I.

117

A1.3 Effects of the Hydrogen Bond Donor (HBD)

It is of interest to investigate the effects of the hydrogen bond donor on the electrochemical and physical properties of the electrolyte. Two hydrogen bond donors, urea (an amine) and ethylene glycol (an alcohol), are compared. Since in Chapter 2 it was shown that reline (1:2 ChCl:urea) is less conductive than ethaline (1:2 ChCl:EG), it is not surprising that the conductivity of electrolytes containing iron would exhibit a similar trend. For example, Figure A1-4 shows the conductivity as a function of temperature for electrolytes containing 1:2:4 FeClz:ChCl:X where X = ethylene glycol or urea and Z = 2 or 3 (for solutions created using FeCl2 vs FeCl3). The urea containing electrolytes have a much lower conductivity. However, it is interesting to note that the 1:2:4

FeCl2:ChCl:urea electrolyte is more conductive than the 1:2:4 FeCl3:ChCl:urea electrolyte. This is unexpected since it was shown that with the iron/ethylene glycol ILs, the electrolytes containing

FeCl3 were generally more conductive. The observed trend in Figure A1-4 implies that there might be different complexes forming in solutions containing urea and would be interesting to investigate in more detail since urea is known to be a stronger hydrogen bond donor [140]. As discussed in

Chapter 2, when the chloride to iron ratio is below 4:1 (example, for 1:1:4 FeCl2:ChCl:EG) the ethylene glycol formed a complex with the iron, this electrolyte had a low solution conductivity; the conductivity data for this electrolyte is very similar to that of the electrolytes prepared with urea as the hydrogen bond donor (Figure A1-4). The 1:1:4 FeCl3:ChCl:EG electrolyte which contains a Cl:Fe ratio of 4:1 has conductivity values similar to that of the 1:2:4 FeClx:ChCl:EG electrolytes (all have Cl:Fe ≥4:1).

118

25 Series3Fe2+, EG Series4Fe3+, EG 20 Series6Fe3+, Urea Series7Fe2+, Urea 15 Series2(1:1:4) Fe2+, EG Series5(1:1:4) Fe3+, EG 10

Conductivity(mS/cm) 5

0 20 40 60 80 100 Temperature (oC)

Figure A1-4. Conductivity of 1:n:4 FeClz:ChCl:X where n=1 or 2, X = ethylene glycol or urea, and Z = 2 or 3;

o 80 C, dry N2-atmosphere.

The CVs in Figure A1-5 are similar for iron electrodeposited from a urea containing electrolyte compared to that of iron deposition/dissolution from an electrolyte where ethylene glycol was the

HBD. The exchange current densities had been estimated at 80 oC using LSV at platinum microelectrodes and were 0.12 and 0.36 A/m2 for urea and ethylene glycol containing electrolytes, respectively. This implies that the hydrogen bond donor does not largely affect the activity (for the Fe2+/Fe0 reaction).

Urea

Normalized Current Normalized

-1.2 -0.7 -0.2 0.3 Applied Potential (V vs Ag)

Figure A1-5. Cyclic voltammetry at a glassy carbon electrode in 1:1:4 FeCl2:ChCl:X where X = ethylene glycol,

urea, or water at 50 mV/s, 80 oC.

119

The coulombic plating efficiencies as a function of plating potential are shown by Figure A1-6.

The efficiencies are lower for iron deposited from the urea electrolyte. No electrolyte breakdown was observed at negative potentials in the 1:1:4 FeCl2:ChCl:urea electrolyte and thus inefficiencies are likely attributed to deposit quality and not the occurrence of a side reaction. Figure A1-7 compares the morphology of 10 mAh iron deposited from 1:1:4 FeCl2:ChCl:X electrolytes where

X = (a) ethylene glycol and (b) urea. The metal deposited from the urea containing electrolyte was much rougher and easily flaked off the substrate material. Furthermore, the sample oxidized immediately upon exposure to ambient conditions. Figure A1-8 includes an arrow to show the flatter platelets that were confirmed to be an iron oxide (or rust) by elemental analysis (Table A1-

2). From elemental analysis, twice as much oxygen was present on the surface of the iron deposited from the urea containing electrolyte than that of the ethylene glycol analog. It should be noted that the two samples had been exposed to ambient conditions for roughly equal amounts of time.

100%

80%

60%

40%

Plating Efficiency Efficiency Plating Series11:1:4 EG 20% Series21:1:4 Urea 0% -1.8 -1.6 -1.4 -1.2 -1 -0.8 Plating Potential (V vs Ag/AgCl)

120

(a) (b)

2 o Figure A1-7. Plating 10 mAh @ 5 mA/cm onto a glassy carbon substrate (80 C) from 1:1:4 FeCl2:ChCl:X

electrolytes where X = (a) ethylene glycol and (b) urea.

2 o Figure A1-8. Plating 10 mAh @ 5 mA/cm onto a glassy carbon substrate (80 C) from 1:1:4 FeCl2:ChCl:urea. The

red arrow indicates an area that has been oxidized (rust).

121

Table A1-2. Results from elemental analysis of the electrodeposited iron from selected ILs. Note that the carbon

balance is not included.

Plated Metal Composition, Overall (wt%) IL Composition Fe O N Halide

1:1:4 (EG) 82.6a 11.9a -- 0.2a 100b ------1:1:4 (Urea) 74.5 21.8 2.9 0.8 2 2 aHigh current (5 mA/cm ) and blow current (0.2 mA/cm ) shown.

122

Appendix A2. Investigation of Alternative Active Metal Species

A2.1 Introduction and Background

Issues occurring at the negative electrode such as slow deposition kinetics and plating efficiencies below 100% are limiting in the all-iron ionic liquid RFBs, as reported in Chapter 2. It is thus of interest to investigate other active metal species. A more facile electrodeposition reaction that produces a stable and adherent deposit with high coulombic efficiency is required thus, electrodeposition of nickel, tin, and zinc as well as a zinc-iron alloy were investigated.

Furthermore, for RFB applications, the use of single-metal electrolytes is desired so that crossover at the membrane can be easily dealt with; electrolytes can be remixed and brought back to their initial states (as with the all-iron and all-vanadium systems [8, 13]). Thus, a metal with at least three accessible and stable oxidation states that lie within the electrolyte window is desired. First row transition metals typically exhibit these traits and were thus studied in detail in addition to polyoxometallates, which have been shown [219] to exhibit multiple electron transfer reactions in ionic liquids. Other metals such as ytterbium and indium, which are known to have very negative redox potentials in aqueous media, were investigated due to the possibility of increasing the cell potential when paired with the iron redox reaction.

Additional desired electrolyte properties include low viscosity, high conductivity, and high concentration of active metal species. Metal chlorides have shown good compatibility with ionic liquid electrolytes and the ability to achieve high active species concentrations [52]. More specifically, chloride salts are capable of complexing and interacting with the hydrogen bond donors via cholo-chloro or chloro-oxo coordination [162]. Ionic liquids containing chloride anionic compounds have been known to exhibit more favorable fluid properties as well and thus,

123 they have been chosen for this study. The following section focuses on assessing the physical and electrochemical properties of ionic liquid electrolytes formed with the different metal ions.

A2.2 Electrodeposition Reactions

Much of the recent literature focuses on aluminum [220, 221] and copper [51, 222] electrodeposition and will not be presented here. Furthermore, an all-copper ionic liquid battery has been demonstrated previously [54]. However, with a cell potential of only 0.75 V and large ohmic losses due to the low conductivity of the electrolyte, the copper ionic liquid was found to be unattractive for RFB applications. Instead nickel, tin, and zinc are compared to the iron electrodeposition reaction. The same ratio, 1:1:4 MCl2:choline chloride:ethylene glycol, was used in order to isolate the effect of the metal center. Table A2-1 compares the half reactions with the corresponding equilibrium potentials expected.

Table A2-1. Standard reduction reactions with the corresponding equilibrium potentials [178, 217].

Reduction half Potential, 퐸표 reaction V vs SHE V vs Ag/AgCl Fe3+ + e− ↔ Fe2+ + 0.77 + 0.57 Cu2+ + 2e− ↔ Cu0 + 0.34 + 0.14 Sn4+ + 2e− ↔ Sn2+ + 0.15 −0.05 Sn2+ + 2e− ↔ Sn0 −0.14 −0.34 Ni2+ + 2e− ↔ Ni0 −0.26 −0.46 Fe2+ + 2e− ↔ Fe0 −0.44 −0.64 Zn2+ + 2e− ↔ Zn0 −0.76 −0.96

124

A2.2.1 Nickel

Cyclic voltammetry was run for each of the electrolytes as shown in Figure A2-2. The first CV was run in the electrolyte containing nickel and displays an electrodeposition reaction with an open circuit voltage of −0.13 V vs Ag/AgCl. This potential is 300 mV negative of iron electrodeposition

(−0.43 V vs Ag/AgCl, Table A2-2). It is interesting to note that the electrodeposition reaction occurs at a more positive potential than what has been reported in the literature for aqueous nickel electrolytes at 25 oC (− 0.46 V vs Ag/AgCl) [217]. However, it is known that different ligands can shift the electrochemical potentials; thus, it is unsurprising that the electrodeposition potential for

2+ 2+ 2- the hexaquo nickel [Ni(H2O)6] differs from that of [Ni(EG)3] or [NiCl4] which are complexes known to form in choline chloride and ethylene glycol mixtures [162]. It would be interesting to investigate nickel speciation as a function of temperature; the solution color changed when heated to 80 oC as shown by Figure A2-1.

From the CV it can also be observed that the electrodeposition/dissolution reaction exhibits sluggish kinetics, even at elevated temperatures; the exchange current density, io, has been estimated from slow-scan linear sweep voltammetry at a platinum microelectrode to be 0.6 A/m2 which is not much of an improvement compared to iron (Table A2-2). An additional reduction peak is located at +0.06 V vs Ag/AgCl; the origin of this feature is unknown, however, it is only present when the potential is scanned positive of +1.25 V. No positive redox reaction was observed for nickel in this electrolyte. If the nickel electrodeposition reaction were coupled with the

Fe2+/Fe3+ reaction, the formal potentials indicate that an RFB system would yield a 0.59 V equilibrium cell potential. In combination with the high ohmic overpotentials expected (due to low solution conductivity of 6.4 mS/cm), this would not work as an RFB electrolyte. Furthermore, the electrodeposition kinetics are sluggish thus, nickel is not recommended for RFB applications.

125

(a) (b)

o o Figure A2-1. The same nickel electrolyte (1:1:4 NiCl2:ChCl:EG) at (a) 27 C compared to (b) 80 C stored in a dry,

nitrogen-filled glovebox.

A2.2.2 Tin

In the presence of high concentrations of chloride ions tin undergoes a reversible, two-electron transfer reaction, Sn2+/Sn4+ [223], as well as kinetically fast electrodeposition and was thus of interest for RFB applications. The CV for the 1:1:4 SnCl2:ChCl:EG electrolyte at a platinum mincroelectrode is shown by Figure A2-2. From the CV it can be shown that the electrodeposition/dissolution reaction is very fast compared to iron or nickel; the exchange current density is three orders of magnitude faster (as shown in Table A2-2). The electrodeposition reaction (Sn2+/Sn0) occurs at −0.35 V vs Ag/AgCl, which is unfortunately 80 mV positive of iron deposition. It is interesting to note the presence of two oxidation peaks coupled with tin reduction

(Sn2+/Sn0); it is expected that the current rapidly decay once the majority of the electrodeposited metal has been stripped and the surface concentration goes to zero, however, a second peak is

126 present at 0.01 V vs Ag/AgCl. By shifting the potential region examined it appears to be coupled with the electroplating reaction and is not as prominent if the potential is not swept more negative than −0.3 V. The Sn2+/Sn4+ reaction has been reported in (high chloride) aqueous electrolytes to have an OCV of −0.05 V vs Ag/AgCl [217], thus, the observed oxidation peak at +0.01 V is believed to be associated with this reaction. A third oxidation peak is present at 1.3 V vs Ag/AgCl, however, the reaction was not found to be coupled with a reduction peak and its origins are unknown at this time. The tin was light sensitive and unstable over time which poses additional challenges. For these reasons, tin was unfavorable for use in a RFB, despite its fast kinetics.

a) Nickel

b) Tin

Current Density (Norm) c) Zinc

d) Iron

-1.5 -0.5 0.5 1.5 Potential (V vs Ag/AgCl)

Figure A2-2. Comparison of CVs (100 mV/s) at a 100 µm platinum disk in electrolytes containing different metal

species in a 1:1:4 ratio; all solutions had been thermostatted to 80 oC and kept in a dry nitrogen-filled glovebox.

127

Table A2-2. Physical and electrochemical properties of the synthesized electrolytes at 80 oC; all solutions contained

MCl2:choline chloride:ethylene glycol in a 1:1:4 molar ratio.

Metal OCV (V vs Ag/AgCl) Concentration Plating Conductivity Notes Measureda vs (M) Kinetics (mS/cm) Expectedb (io A/m2)

Nickel − 0.13 − 0.46 2.6 0.6 6.4 *Undergoes color change with temperature

Tin − 0.35 − 0.34 2.5 400 *light sensitive, unstable

Zinc − 0.95 − 0.96 2.5 9.0 9.3

Iron − 0.43 − 0.64 2.7 0.4 2.4 aMeasurements taken after metal deposition at − 1.1 V for 60 s then 600 s rest at the open circuit

bData obtained from standard reduction potentials (ref [217])

A2.2.3 Zinc

Zinc has been known to exhibit fast electrodeposition kinetics in both aqueous and non-aqueous electrolytes [55, 96]. From the CV in Figure A2-2, it is apparent that this also holds true for the

1:1:4 ZnCl2:ChCl:EG electrolyte. As shown by the estimated values in Table A2-2, the electrokinetics are more than 20 times faster than that of iron electrodeposition. Furthermore, there is a 4 times enhancement in conductivity. The OCV for Zn2+/Zn0 reaction is comparable to what has been reported for aqueous electrolytes and is approximately 520 mV negative of the iron deposition reaction. No positive reactions were observed within the electrolyte stability window.

The Zn2+/Zn0 could be coupled with the Fe2+/Fe3+ reaction; the formal potentials indicate that an

RFB system would yield a 1.41 V equilibrium cell potential, which is attractive for increasing system power.

Battery systems utilizing zinc electrodes have been under investigation for the past 50 years due to low toxicity, high abundance of zinc, fast deposition kinetics as well as the OCV for the Zn2+/Zn0 reaction. However, issues with dendritic zinc deposits compromise battery life and safety [224].

128

Therefore it was important to characterize morphology of the electrodeposited zinc from these nonaqueous electrolytes. Figure A2-3 compares zinc deposited from the 1:1:4 ZnCl2:ChCl:EG electrolyte to zinc deposited from a typical aqueous electrolyte (0.1 M ZnO + 4.0 M KOH). From the SEM micrographs at the same magnification, it is clear that dendrites do not form on the surface of zinc deposited from the 1:1:4 ZnCl2:ChCl:EG electrolyte. However, the deposited metal did not adhere well to the glassy carbon substrate. Within a flowing battery system, this would pose a different challenge that must be addressed.

(a) (b)

o Figure A2-3. Zinc deposited from (a) 1:1:4 ZnCl2:ChCl:EG (10 mAh, 80 C) compared to zinc deposited from (b) a typical aqueous electrolyte (0.1 M ZnO + 4.0 M KOH, room temperature) at 12.5 mA/cm2. Fig 3 (b) has been taken

from [224].

A2.2.4 Iron-zinc Alloy

As discussed in the previous section, the Zn2+/Zn0 could be coupled with the Fe2+/Fe3+ reaction for an RFB system that would yield a 1.41 V equilibrium cell potential. This assumes that crossover of iron at the membrane would not occur, which is an unrealistic expectation given commercially available membranes. Therefore, an electrolyte containing 1:1:4 FeCl2/ZnCl2:ChCl:EG with equal

129 molar amounts of iron and zinc was investigated. From Table A2-3 it can be seen that the electrodeposition kinetics show a slight improvement for metal that was electrodeposited from the mixed electrolyte. The mixed electrolyte also exhibited an improvement in solution conductivity compared to the iron only solution. The OCV for metal deposition in the mixed electrolyte was

−0.58 V vs Ag/AgCl, which is in between the potentials for the pure iron or pure zinc electrolytes.

Table A2-3. Physical and electrochemical properties of the synthesized electrolytes at 80 oC; all solutions contained

MCl2:choline chloride:ethylene glycol in a 1:1:4 molar ratio.

Metal Measured OCV Plating Kinetics Conductivity (V vs Ag/AgCl)a LSV (mS/cm) (io A/m2) Zinc −0.95 9.0 9.3 Iron-zinc Alloy −0.58 0.8 5.0 Iron −0.43 0.4 2.4 aMeasurements taken after metal deposition at − 1.1 V for 60 s then 600 s rest at the open circuit

As shown by the CVs in Figure A2-4, the Fe2+/Fe3+ reaction potential remains unchanged with the addition of ZnCl2 at approximately +0.55 V. The formal potentials indicate that an RFB system would yield a 1.13 V equilibrium cell potential. This cell potential and the faster kinetics are comparable to findings reported by Lloyd et al. for iron-zinc in a similar electrolyte [55]; however, the OCV for the plating reaction was not as negative as what is being reported here. Furthermore,

Lloyd et al did not comment on the nature or composition of the metal that had been electrodeposited; the metal was assumed to be an alloy.

130

(a)

(b)

-1.5 -0.5 0.5 1.5 Potential (V vs Ag/AgCl)

Figure A2-4. Comparison of CVs (100 mV/s) at a platinum disk in electrolytes containing either (a) ZnCl2, (b)

o equal molar FeCl2/ZnCl2 or (c) FeCl2 only in a 1:1:4 ratio; all solutions had been thermostatted to 80 C and kept in

a dry nitrogen-filled glovebox.

Figure A2-5 shows a survey scan obtained from x-ray photoelectron spectroscopy of the glassy carbon electrode after electrodepositing 10 mAh of metal at 2.5 mA/cm2. It can be confirmed that the electrodeposited metal is an alloy as shown by the zinc 2p doublet located at 1021 eV (value for Zn 2p3/2) as well as the iron 2p doublet at 709 eV (2p3/2). Additional peaks for the zinc 3s, 3p, and 3d electrons can be found at 91, 98, and 10 eV and for the iron 3s and 2p at 92 and 53 eV, respectively. The zinc/iron ratio for this sample was determined from quantification of the survey scans to be 0.52, thus, at 2.5 mA/cm2 there is approximately twice the amount of iron as zinc present in the alloy. This was done for a range of current densities (or rather deposition potentials) and the results are summarized by Figure A2-6.

131

7 Zn 2p 6

4 5 10 × 4 Zn 3s, 3p, 3d O 1s (91, 89, 10 eV) 3 Fe 2p Fe 3s, 3p

2 Zinc (92, 53 eV) Norm. Counts Auger C 1s 1 Cl 2p

0 1000 800 600 400 200 0 Binding Energy (eV)

Figure A2-5. Example of X-ray photoelectron spectroscopy survey scan from metal electrodeposited onto a glassy

o carbon substrate from 1:1:4 FeCl2/ZnCl2:ChCl:EG at 80 C.

1.2

1.0 −1.05 V

0.8

0.6

Zinc/Iron ratio Zinc/Iron −0.73 V 0.4 −0.90 V

−0.96 V 0.2 −0.7 V vs Ag/AgCl −1.05 V

0.0 0 1 2 3 4 5 2 Current Density (mA/cm )

Figure A2-6. Metal composition determined from XPS survey scans for Iron-Zinc electrodeposited onto a glassy

o carbon substrate from 1:1:4 FeCl2/ZnCl2:ChCl:EG at 80 C. All samples contained 10 mAh electrodeposited metal.

Two different areas had been analyzed for the sample where metal was deposited at −1.05 V.

132

When referring to Table A2-3, the standard electrode potential of iron, 퐸0 = −0.43 V, is 520 mV more noble than zinc (퐸0 = −0.95 V). Thus, it is not expected that zinc codeposit readily with iron, however, significant amounts of zinc are present in the deposited metal. This observation has been encountered in the literature for aqueous systems as iron-zinc alloys are attractive for corrosion resistant coatings for steel [225]. Typically, the more noble metal deposits preferentially

[225]. In this case where the electrolyte contained equal molar amounts of iron and zinc, it is expected that the zinc/iron ratio be less than 1 since iron is the more noble of the two metals. This is observed at all current densities except for 5 mA/cm2. In the case of anomalous codeposition, at low current densities the concentration of the more noble metal is larger for the deposit, however, at higher current densities the system undergoes a transition and the concentration of the more noble metal (iron) becomes less than that of the less noble metal (zinc). An example of this can be referenced in the literature for an acidic aqueous sulfate electrolyte; they noted the transition to occur at 22.5 mA/cm2 for a bath at 90oC and 7.5 mA/cm2 at 75oC [226]. This could explain what is observed experimentally. It is also interesting to note that at two different locations on the sample, two different metal compositions were reported. This is explained with reference to the

SEM micrographs presented by Figure A2-7. As the current density is increased, the surface of the deposit became brittle, similar to that of the pure zinc deposit. Figure A2-8 shows low magnification SEM micrographs to note the peeling of metal as well as the presence of nodules. If surveys were run in different locations about the surface of the deposit and the metal was not homogeneous, then it is unsurprising that differences in composition were observed. Due to the fact that the surface became more brittle with increasing current density, it is difficult to determine whether the correlation in Figure A2-6 is truly representative for metal electroplated from the iron-

133 zinc electrolyte, however, the XPS results do confirm that the electrodeposited metal is indeed an alloy.

From the previous section it was shown that the morphology of the metal deposited from a 1:1:4

ZnCl2:ChCl:EG electrolyte did not contain dendrites but was non-adherent to the glassy carbon substrate. Iron deposited from a 1:1:4 FeCl2:ChCl:EG electrolyte is adherent and shows good plating efficiency (discussed in Chapter 2). Figure A2-7 compares the morphology of the electrodeposited metal at various current densities. Metal deposited at lower current densities, i.e.

0.5 and 2.5 mA/cm2 was smooth and adherent, however, the deposit became brittle and flaky at higher current density operations (>5 mA/cm2). This is apparent when referring to Figure A2-8 which shows SEM micrographs at lower magnification. Since higher current density operation is desirable for RFB applications this would not be feasible for a flowing system; it is expected that the coulombic efficiencies would be lower than those measured at a static electrode in the 3- electrode cell (which was 87% when metal was deposited at −1.15 V vs Ag/AgCl).

(a) (b) (c) (d)

Figure A2-7. Iron-Zinc (10 mAh) electrodeposited onto a glassy carbon substrate from 1:1:4 FeCl2/ZnCl2:ChCl:EG

at 80 oC at (a) 0.5, (b) 2.5, (c) 4, and (d) 5 mA/cm2.

134

(a) (b)

Figure A2-8. Iron-Zinc (10 mAh) electrodeposited onto a glassy carbon substrate from a 1:1:4

o 2 FeCl2/ZnCl2:ChCl:EG at 80 C, 5 mA/cm

A2.3 Redox Reactions

It is desired to eliminate the electrodeposition/dissolution reaction and instead have two redox reactions at the positive and negative half cells, respectively. The electrolyte must possess acceptable fluid properties as well as exhibit fast kinetics for the reactions and high solubility of the active metal species. Transition metals such as chromium, manganese, vanadium as well as polyoxometalates had been investigated since they are known to exhibit multiple oxidation states.

Ytterbium and indium are known to have very negative redox potentials in aqueous media, which are attractive for expanding the system voltage. The standard reduction reactions for selected reactions with the corresponding equilibrium potentials are provided by Table A2-4.

135

Table A2-4. Selected standard reduction reactions with the corresponding equilibrium potentials [178, 217].

Reduction half Potential, 퐸표 reaction V vs SHE V vs Ag/AgCl Cr2+ + 2e− ↔ Cr0 −0.90 −1.10 Cr3+ + e− ↔ Cr2+ −0.42 −0.62 Mn2+ + 2e− ↔ Mn0 −1.18 −1.38 Mn3+ + e− ↔ Mn2+ +1.50 +1.30 Yb3+ + e− ↔ Yb2+ −1.05 −1.25 In3+ + 3e− ↔ In0 −0.33 −0.53 In3+ + 2e− ↔ In+ −0.43 −0.63 V2+ + 2e− ↔ V0 −1.13 −1.33 V3+ + e− ↔ V2+ −0.26 −0.46

Table A2-5 provides a summary of the metals tested in a choline chloride and ethylene glycol based electrolyte. It was found that the two polyoxometalates, phosphomolybdic acid and phosphotungstic acid (H3PMo12O40 and H3PW12O40 ) were largely insoluble in both reline and ethaline (~10 mM) and thus, were ruled out for further study. Chromium appeared to have two redox processes occurring at around the same potential, however, no other reactions had been observed (Figure A2-9). Thus, Cr had also been abandoned. Manganese had been evaluated in both ethaline and reline due to the possibility of a 2.68 V widow for reduction and oxidation of

Mn2+ (Table A2-4). However, no reversible redox couples were observed in either electrolyte

(Figure A2-10) and in fact the oxidation and reduction processes observed were kinetically very sluggish. It should be noted that the experiments with manganese had been performed outside the glovebox (under ambient conditions). Perhaps if the experiments had been reproduced in a dry atmosphere, they would yield different results.

136

Table A2-5. Summary of initial data for metals tested in a choline chloride and ethylene glycol based electrolyte.

Metal Estimated OCV Notes (V vs Ag/AgCl) Chromium 0.08 Manganese −0.4 Sluggish Ytterbium −0.33 paste consistency Indium −0.40, −0.04, 0.81a Vanadium −0.25 POMs N/A Insoluble

2.E-05 Pt disk,

1.E-05

0.E+00

Current (A) -1.E-05

-2.E-05

-3.E-05 -0.5 0 0.5 1

Applied Potential (V vs Ag/AgCl)

o Figure A2-9. Cyclic voltammetry at a platinum electrode (25 mV/s) in 0.1 M CrCl2 in ethaline at 50 C

137

(a) (b)

4.E-05 5.E-05 2.E-05 2.E-05 0.E+00 -2.E-05

-2.E-05

Current (A) Current (A) -5.E-05 -4.E-05

-8.E-05 -6.E-05 -1.25 -0.5 0.25 1 -2 -1 0 1 2 Applied Potential (V vs Ag/AgCl) Applied Potential (V vs Ag/AgCl)

Figure A2-10. Cyclic voltammetry at a platinum electrode (100 mV/s) in 0.1 M MnCl2 in (a) ethaline and (b) reline

based electrolytes at 50 oC.

o Indium and ytterbium were both tested in ratios of 1:1:4 MCl3:ChCl:EG at 80 C in a N2- atmosphere. CVs in the indium chloride containing electrolyte exhibited multiple oxidation and reduction peaks (Figure A2-11). However, if the potential was swept between −0.5 V and 1.25 V, no reactions could be observed. The purple and green arrows are located on the CV to indicate which oxidation/reduction reactions are coupled. Furthermore, the reduction at ~ −0.53 V appears to be characteristic of an electrodeposition reaction. This is confirmed when referencing the standard reduction table; the reaction for indium reduction 퐼푛3+ + 3푒− ↔ 퐼푛(푠) occurs −0.53 V vs Ag/AgCl. The oxidation reaction at ~ 1.35 V was not reversible and is thought to be coupled with electrolyte breakdown. Consequently, (cost aside) indium does not appear to be suitable for use in a RFB system.

138

600

400 )

2 200

mA/m 0

-200

Current Density ( Density Current -400

-600

-800 -0.9 -0.4 0.1 0.6 1.1 1.6 2.1 Potential (V vs Ag/AgCl)

Figure A2-11. Cyclic voltammetry (100 mV/s) at a platinum disk in an electrolyte containing 1:1:4 InCl3:ChCl:EG

thermostatted to 80 oC and kept in a dry nitrogen-filled glovebox. The arrows indicate which oxidation/reduction

processes are coupled with one another.

Ytterbium was evaluated at a platinum microelectrode at 80oC and the corresponding CV is shown in Figure A2-12. One oxidation/reduction reaction was observed within the solvent window with an OCV ~ −0.5 V. Expanding the window to potentials more negative than – 1.0 V lead to passivation of the platinum electrode surface. Furthermore, this reaction was quite sluggish (0.7

A/m2, estimated from LSV) and even at 80 oC the electrolyte was a paste as opposed to a fluid with reasonable viscosity. The paste is shown as an inset in Figure A2-12.

139

) 2

-40

-80

-120

Current Density (A/m Density Current -160

-200 -1.5 -1 -0.5 0 0.5 Applied Potential (V vs Ag/AgCl)

Figure 1. Cyclic voltammetry for a 1:1:4 YbCl3:ChCl:EG electrolyte at a platinum microelectrode, 150 mV/s,

80 oC. Inset shows the electrolyte paste.

A2.3.1 Vanadium Electrolytes

Vanadium has already been successfully demonstrated in both the aqueous [8] and non-aqueous organic electrolytes [32], with cell potentials of 1.4 and 2.2 V, respectively. Thus, this chemistry was also investigated in an ionic liquid electrolyte. Vanadium (III) chloride was used due to good solubility of metal chlorides in reline and ethaline (recall V(acac)3 only has a solubility ~0.2 M in ethaline [60]). Initial CV scans had been run for 0.1 M VCl3 in ethaline (formed a dark purple

o solution) and reline (dark green solution) at a glassy carbon electrode (50 C, in a N2-atmosphere).

It should be noted that VCl4 was not used due to the safety hazards associated with it. The resulting scans are shown in Figure A2-14 where only one redox reaction was observed within the solvent window (example is shown by Figure A2-13); this is likely V3+ + e− ↔ V2+. Other oxidation peaks at ~0.73 and 1.15 V were not shown to be reversible. As an initial comparison between the reactions in ethaline vs reline, the currents are much smaller for the CV in reline, as expected due

140 to the higher viscosity of the electrolyte. Furthermore, the peak separation is much smaller in the

ethaline based electrolyte, implying faster kinetics. Current (A) Current

-2 -1.5 -1 -0.5 0 0.5 1 1.5 Applied Potential (V vs Ag wire)

o Figure A2-13. Cyclic voltammetry at glassy carbon in a 0.01 M VCl3 solution (in ethaline); 100 mV/s, 50 C

(a) (b)

4 0.5 0.4

) 3

) 2 2 0.3 2 0.2 1 0.1 0 0 -1 -0.1 Separation: 890 mV Separation: 125mV -0.2 -2 (Quasi-reversible) (Irreversible)

Current Density (A/m Density Current -0.3 -3 (A/m Density Current -0.4 -4 -0.5 -1 -0.5 0 0.5 1 -1.5 -1 -0.5 0 Applied Potential (V vs Ag) Applied Potential (V vs Ag)

Figure A2-14. Cyclic voltammetry of 0.1 M VCl3 in (a) ethaline and (b) reline on GC at 10 mV/s

Kinetics for the vanadium reactions in ethaline and reline are reported in Table A2-6. As expected, based off the CV, the exchange currents are much higher for V2+/V3+ in ethaline. In fact, these

141

2 exchange currents are comparable to the reduction of 0.01 M V(acac)3 in ACN on Pt (1.8 A/m at

o 25 C [37]) and much faster than the vanadium reactions in an aqueous electrolyte (0.1 M VOSO4

2 o in H2SO4 is 0.08 A/m at platinum, 25 C).

Table A2-6. Exchange current densities estimated from EIS (OCV ± 15 mV AC) and LSV at platinum and glassy

o carbon in 0.1M VCl3 solutions; 50 C, N2-atmosphere.

Exchange Current Density (A/m2) Reline Ethaline Pt (EIS/LSV) GC (LSV/EIS) Pt (EIS/LSV) GC (LSV/EIS) 0.02/0.09 0.002/0.004 1.3/1.0 3.5/0.8

A2.3.2 Iron-vanadium Mixed Electrolyte

The vanadium chemistry looks promising in ethaline, however, only one reversible redox reaction was observed. Nonetheless, this presents the possibility for coupling the V2+/V3+ with the

Fe2+/Fe3+ for a double redox system. Thus, the issues associated with the plating reaction in the iron system are avoided and instead, two kinetically facile reactions are employed. Example CVs for the individual reactions are shown by Figure A2-15 where the iron and vanadium are compared

o for electrolytes containing 1:1:4 MClx:ChCl:EG at 80 C at a platinum electrode. It should be pointed out that the current densities for the vanadium electrolyte are much lower than the CV shown in the iron IL; this is due to large differences in solution conductivity. One drawback of the vanadium IL electrolyte is that the solution conductivity is 1.6 mS/cm at 80oC which is about half that of the iron analog which is 2.4 mS/cm. However, in a practical system a mixed electrolyte may be more appropriate since crossover of active material is expected to occur over time at the

142 separator (especially for large concentration gradients – i.e. high concentration of vanadium or iron in one half cell and zero on the other). Thus, it is of interest to investigate the properties of a mixed iron-vanadium electrolyte; the CV for the mixed electrolyte is shown by Figure A2-16. The kinetics estimated from LSV at a platinum microelectrode for the vanadium and iron reactions are not largely affected. This data is summarized in Table A2-7.

800

) 2 400

-400 Current Density (A/m Density Current -800 -1 -0.5 0 0.5 1 1.5 Applied Potential (V vs Ag/AgCl)

Figure A2-15. Cyclic voltammetry of 1:1:4 VCl3:ChCl:EG and 1:1:4 (FeCl2/FeCl3:ChCl:EG) at a platinum

electrode, 80 oC, 50 mV/s

300

) 2

100 Current Density (A/m Density Current -100 -0.75 -0.25 0.25 0.75 Applied Potential (V vs Ag/AgCl)

o Figure A2-16. Cyclic voltammetry of a 1:1:4 FeCl2/VCl3:ChCl:EG mixed electrolyte at a platinum electrode, 80 C,

50 mV/s

143

Table A2-7 shows the conductivity for the 1:1:4 MClx:ChCl:EG electrolytes comparing the VCl3 with the FeCl2 only and the resulting mixture which contains equal molar iron and vanadium.

Interestingly, the conductivity for the mixed electrolyte is higher than that of the all-iron or all- vanadium. This increase in conductivity is likely attributed to the added chloride (three moles from VCl3 compared to two from the FeCl2). This value for conductivity unfortunately is still quite low and would cause high ohmic overpotentials in a flow battery system. Furthermore, the

OCV for the iron-vanadium system is only ~0.7 V. Further work is to be done to explore different ligands to increase the potential but not hinder the solubility. Additionally, different hydrogen bond donors could be evaluated as a means to enhance conductivity.

o Table A2-7. Summary of conductivity, kinetic, and cell potential data for 1:1:4 MClx:ChCl:EG electrolytes; 80 C,

N2-atmosphere. Exchange current densities were estimated from LSVs at glassy carbon or platinum electrodes.

Metal io (A/m2) Conductivity Potential (MClx) (mS/cm) Window (V) GC Pt

FeCl2 22.0 93.9 2.4 0.89

VCl3 2.1 7.6 1.6 N/A Fe-V 3.6(Fe), 31.5 (Fe), 2.8 0.7 3.7(V) 9.4(V)

144

Appendix A3. Aqueous All-iron Flow Battery with Choline Chloride

Supporting Electrolyte

A3.1 Introduction and Background

The standard reduction potential for the Fe2+/Fe0 reaction is −0.44 V vs SHE which is 440 mV negative of hydrogen evolution. The hydrogen evolution side reaction is a significant factor in lowering coulombic inefficiencies as well as pH imbalances (which cause iron precipitation) which are major challenges facing the aqueous iron battery [5]. Passive recombination reactors are one possible solution where the hydrogen gas is allowed to chemically react with excess ferric ions resulting in protons and ferrous ions in order to rebalance the electrolyte [227]. Other research explores additives such as ligands like glycine or glycerol as well as different supporting electrolytes [136]. In this chapter, choline chloride is explored as a supporting electrolyte for the aqueous iron system.

A3.2 Three electrode cell - Preliminary Data

Figure A3-1 shows the hydrogen evolution current at an iron rod in two ionic liquid electrolytes that contain choline chloride, an electrolyte containing choline chloride and water in a 1:4 molar ratio, as well as a typical aqueous electrolyte (1M KCl and 0.1M HCl). While it is expected that the hydrogen evolution current be small for the ionic liquid electrolytes (which possess wide potential windows), it is interesting to note that choline chloride and water showed a wider potential window (compared to the typical aqueous electrolyte). The potential for which an appreciable amount of hydrogen gas was formed (1 mA/cm2) at an iron rod in each of the

145 electrolytes is compared in Table A3-1. This shows that the choline chloride containing electrolyte expanded the water window by ~750 mV.

) 2 1:4 ChCl:Urea 1:4 ChCl:EG -0.02

-0.04 1:4 ChCl:Water

Current Density (A/cm 1M KCl, 0.1M HCl (pH 1) -0.06 -3 -2.5 -2 -1.5 -1 -0.5 Potential (V vs Ag/AgCl)

Figure A3-1. Cyclic voltammetry showing hydrogen evolution current at an iron rod in selected electrolytes (50 oC).

Table A3-1. Hydrogen evolution current at an iron rod in each of the different electrolytes (50 oC) showing potential

2 (iR corrected) at which the H2 evolution current reached 1 mA/cm . All experiments had been performed in a sealed

three electrode cell which was purged with N2 gas.

Electrolyte Potential at 1 mA/cm2 (V vs Ag/AgCl)

1 M KCl, 0.1 M HCl −0.5

5 M KCl −1.07

(1:4) ChCl:H2O −1.25

(1:4) ChCl:EG −1.39

(2:4) ChCl:EG −1.24

(1:4) ChCl:Urea −1.82

(2:4) ChCl:Urea −1.38

146

Iron was plated at a selected potential for 5 minutes then stripped for 5 minutes at a potential more positive of the electrodissolution peak; efficiency was estimated from the number of coulombs passed during stripping divided by coulombs during plating. Coulombic plating efficiency is plotted as a function of plating potential in Figure A3-2. As expected, the typical aqueous iron electrolyte exhibited the lowest plating efficiencies. However, the electrolyte utilizing water as the hydrogen bond donor, 1:1:4 FeCl2•4H2O:ChCl:water showed the highest plating efficiencies of those tested. Within the literature, choline chloride has been shown to hinder hydrogen evolution in aqueous electrolytes while simultaneously promoting other more desirable electrochemical reactions [228].

100%

80%

60%

40% Series50.5 M FeCl2•4H2O & 1M KCl in water 1:1:4 FeCl2•4H2O:ChCl:EG (wet) 20% Plating Efficiency Efficiency Plating Series31:1:4 FeCl2•4H2O:ChCl:water Series11:1:4 FeCl2:ChCl:EG 0% -1.8 -1.6 -1.4 -1.2 -1.0 -0.8 Plating Potential (V vs Ag/AgCl)

Figure A3-2. Coulombic plating efficiency as a function of voltage at a glassy carbon substrate for a typical aqueous

electrolyte (0.5 M FeCl2•4H2O and 1M KCl in water) compared to solutions containing 1:1:4 FeCl2:ChCl:EG

(prepared and stored in a dry N2-atmosphere), a 1:1:4 FeCl2•4H2O:ChCl:EG electrolyte, and a 1:1:4

o FeCl2•4H2O:ChCl:water electrolyte. Temperature was maintained at 80 C with a nitrogen gas purge.

147

In addition to enhanced plating efficiency, the electrolyte containing 1:1:4

2+ 0 FeCl2•4H2O:ChCl:water had faster kinetics for the Fe /Fe reaction when compared to the ethylene glycol analog (1.5 compared to 0.36 A/m2, respectively estimated from LSV at 80 oC).

The CVs for the water containing electrolyte and the ethylene glycol analog are compared Figure

A3-3. The decreased separation between the oxidation and reduction is observed for the 1:1:4

FeCl2•4H2O:ChCl:water electrolyte which is characteristic of faster kinetics. Furthermore, the water based electrolyte exhibited a reasonable conductivity (43 mS/cm vs compared to the 1:1:4

FeCl2:ChCl:EG electrolyte which has a conductivity of 0.3 mS/cm at 50 oC).

Water Normalized Current Normalized

-1.2 -0.9 -0.6 -0.3 0 0.3 Applied Potential (V vs Ag/AgCl)

Figure A3-3. Cyclic voltammetry at a glassy carbon electrode in 1:1:4 FeCl2:ChCl:EG (“EG”) compared to 1:1:4

o FeCl2•4H2O:ChCl:water at 50 mV/s, 80 C.

148

A3.3 Small Scale Flow Cell

Three electrode cell testing of the 1:1:4 FeCl2•4H2O:ChCl:water electrolyte showed high coulombic plating efficiencies, fast kinetics, as well as a reasonable conductivity and thus, this electrolyte had been tested in a small-scale flow cell. Flow battery testing was performed using a

4 cm2 cell (membrane area) such as the one shown by the schematic in Figure A3-4 employing a

1:1:4 FeCl2•4H2O:ChCl:water electrolyte. Carbon felt electrodes (KFD, supplied by SGL Carbon

Group; treated thermally at 400 oC for 1 hour) were bonded to impervious graphite plates using carbon ink (Acheson), compressed 30% using Teflon gaskets, and separated by a Daramic 175 microporous separator that had been pre-soaked in a 1:4 choline chloride:water solution

(overnight). Electrolytes were circulated by a peristaltic pump at a flow rate of 25 mL/min. The cell was charged and discharged at 25 mA/cm2 for the first four cycles then 10 mA/cm2 for the

o remaining cycles at 50 C. A N2-gas purge was used in both electrolyte reservoirs to prevent solution oxidation.

Figure A3-4. Schematic of the small-scale flow battery hardware.

149

1.9

1.4

0.9 Cell Cell Potential(V) 0.4

-0.1 0 5 10 15 20 25 30 35 40 Time (hours)

Figure A3-5. 1 hour charge/discharge cycles in a small-scale flow cell employing a 1:1:4 FeCl2•4H2O:ChCl:water

2 2 o electrolyte (cycles 1- 4 were at 25 mA/cm and 5-40 were at 10 mA/cm ); N2 purge, 50 C.

Figure A3-6 shows the resulting coulombic, voltaic, and energy efficiency as a function of cycle number for 1 hour charge/discharge cycles. The efficiencies were steady from cycles ~9-30 at

87 % CE, 49 % VE, and 42 % EE. Note that at cycle 4 the current density had been switched from

25 mA/cm2 to 10 mA/cm2, at cycle 7 the electrolyte reservoir had been replenished and a leak had been fixed, at cycle 31 the nitrogen had been disconnected due to pressure buildup within the cell, and at cycle 35 the electrolyte stopped pumping due to a clog in the cell. The electrolyte built up at the outlet of the negative electrolyte compartment of the cell due to an increase in viscosity with charging (shown in Figure A3-7). Pressure buildup was likely caused by the change in viscosity as well as hydrogen evolution side reaction that may still be occurring at the negative electrode on charge. No hydrogen bubbles had been observed, however, this possibility has not been ruled out.

Once the cell had been taken apart (after a charge), the iron was shown to be a strongly attached and stable deposit on the surface of the negative (felt) electrode (Figure A3-8). 150

KFD felt is a conductive carbon material (more conductive than the electrolyte) and thus, the iron electrodeposited near the separator (on the top of the felt) as shown by Figure A3-8. It is recommended that for future experiments, a low conductivity carbon felt is used for the negative electrode so that the iron deposits closer to the current collector (allowing for better utilization of the felt and enabling more iron to be deposited). It should be noted that no iron metal was found to be deposited in the membrane.

Figure A3-6. Coulombic, voltaic, and energy efficiency as a function of cycle number for 1 hour charge/discharge

o cycles in a small-scale flow cell employing a 1:1:4 FeCl2•4H2O:ChCl:water electrolyte; N2 purge, 50 C

Figure A3-7. Photograph of the negative electrode after assembling the cell after a charge at 10 mA/cm2. The arrow

indicates where the electrolyte had thickened, blocking flow out of the cell (from the negative electrode

compartment).

151

Figure A3-8. Battery electrodes (KFD carbon felt bonded to the graphite plates using Acheson carbon ink) and

2 o Daramic separator are shown after the battery was disassembled after a 1 hour charge at 10 mA/cm ; 50 C, N2 gas

purge. Arrows mark the direction of flow. Electrodeposited iron can be seen on the surface of the negative

electrode.

152

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