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Small organic molecules in all-organic redox flow batteries for grid-scale energy storage
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Small organic molecules in all-organic redox flow batteries for grid-scale energy storage
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Content
SMALL ORGANIC MOLECULES IN ALL-ORGANIC REDOX
FLOW BATTERIES FOR GRID-SCALE ENERGY STORAGE
by
Lena Hoober-Burkhardt
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirement for the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)
May 2017
Copyright 2017 Lena Hoober-Burkhardt
ii
DEDICATION
I dedicate this thesis to my family.
To the HB Clan – who have always, always, always supported me, even when I have made it difficult for
you. I wouldn’t be the person I am today without you and I am grateful every day that I am part of our
crazy, wonderful, wild, loving family. Thank you all for being such a wonderful part of my life. You four
will always be my foundation and my home.
And to my wonderful, smart, kind, patient fiancé, Nick. Thank you so much for, well, everything. Thank
you for the hugs, the support, the food, the trips, the adventures, the fun, the dancing, the laughter, and
most of all, the ever-present and unfading love. You have made every single moment of the past four
years of my life better.
iii
ACKNOWLEDGEMENTS
First, I would like to thank Professor Sri Narayan for his excellent advising and mentorship over
the past five years that I have been a member of his group. It was an honor to work with a man who is
so obviously passionate about his scientific pursuits and it was a privilege to be able to learn as much as I
could from his great, creative, and ever-optimistic mind. You are a scientist, in the truest sense of the
word, and I am incredibly thankful that I was able to work with you.
I will also always be especially grateful to all of my colleagues who worked on this project with
me – Dr. Bo Yang, Sankarganesh Krishnamoorthy, Advaith Murali, and Archith Nirmalchandar. Without
your help, knowledge, and support I would not have been able to accomplish what I have. Thank you for
working with me on such an exciting project.
Thank you to Dr. Robert Aniszfeld, Professor Surya Prakash, and the Loker Hydrocarbon Institute
for all of the support while working on this amazing ARPA-E project. I will always be grateful for the
funding that ARPA-E provided. It was an incredible and unique experience, and one which very few
graduate students can say that they’ve had. Because of your support on this project, I was able to see
another side of scientific research and I will always be grateful for that.
Thank you to all of the members of the Narayan group – past and present. You have some
awesome resources, a boss who gives you too many ideas to work with, and an amazingly collegial and
respectful atmosphere to work in. Thank you to all who supported me with my academic pursuits over
the years and I wish you all nothing but the best in your future careers.
I would also like to recognize Women in Chemistry – without the amazing support from these
women, graduate school would have been a lot harder. It was a privilege to meet, work with, and
iv
become friends with such amazing women over the course of the last five years. I can’t wait to see all of
the incredible things you do.
And, finally, thank you to all of my friends that I have made here in this crazy city. You took a
place that I was, at best, lukewarm about, and have made it feel like home. If it wasn’t for all of the trips,
the late nights, the margaritas, the happy hours, and the laughs, my graduate school experience would
have been a lot less fun and lot less rich. I am so happy and privileged to call you all my LA Family.
v
TABLE OF CONTENTS
DEDICATION.................................................................................................................................................. ii
ACKNOWLEDGEMENTS ............................................................................................................................... iii
LIST OF TABLES ........................................................................................................................................... vii
LIST OF FIGURES ........................................................................................................................................ viii
LIST OF ABBREVIATIONS AND SYMBOLS ................................................................................................... xiv
SUMMARY .................................................................................................................................................. xv
Chapter 1: Introduction ................................................................................................................................ 1
1.1 Grid-Scale Energy Storage .................................................................................................................. 1
1.2 Redox Flow Batteries: History and Overview .................................................................................... 6
1.3 Quinone Introduction ......................................................................................................................... 9
1.3.1 Charge-Transfer Processes ........................................................................................................ 11
1.3.2 Cell Voltage................................................................................................................................ 11
1.3.3 Mass Transport Processes ......................................................................................................... 12
1.3.4 Reactivity and Long-Term Cycling ............................................................................................ 13
1.3.5 Faradaic Efficiency ..................................................................................................................... 15
1.4 Focus on My Work ............................................................................................................................ 15
Chapter 2: Experimental Techniques and Methodology for Analysis of Results ..................................... 18
2.1 Electrochemical Characterization .................................................................................................... 18
2.1.1 Electrochemical Method of Analysis ......................................................................................... 19
2.2 Quantum Mechanics-Based Calculations ........................................................................................ 24
2.3. Synthetic and Chemical Characterization of Materials.................................................................. 24
2.4 Cell Set-Up ........................................................................................................................................ 25
Chapter 3: Kinetics of Electrochemical Reactions of Hydroquinones ...................................................... 29
3.1 Hydroquinone Electrochemical Characterization and Analysis ...................................................... 29
3.2 Sulfonic Acid Substituted Anthraquinones ...................................................................................... 43
3.3 Selection of a Benzoquinone and Anthraquinone Molecule for Initial Cycling .............................. 46
3.3.1 Initial Cycling Results and Discussion ....................................................................................... 46
3.3.2 The Issues Arising from the Michael Reaction of Water with BQDS ....................................... 50
3.4 The Design of a Michael-Reaction Resistant Benzoquinone ........................................................... 53
3.4.1 Surface Effects on CV Performance........................................................................................... 57
vi
3.4.2 Pyrogallol Compounds .............................................................................................................. 62
3.4.3 Methyl Substituted Benzoquinones .......................................................................................... 66
3.4.3.1 Proto-desulfonation ............................................................................................................... 68
3.5 Summary ........................................................................................................................................... 70
Chapter 4: Alkaline Studies ........................................................................................................................ 74
4.1 Introduction and Advantages of an Alkaline RFB System .............................................................. 74
4.2 Napthoquinones ............................................................................................................................... 74
4.3 Hydroxy Substituted Anthraquinones ............................................................................................. 77
4.4 Quinoxalines ..................................................................................................................................... 81
4.5 Riboflavin .......................................................................................................................................... 84
4.6 Hydroquinones ................................................................................................................................. 87
4.7 Summary ........................................................................................................................................... 92
Chapter 5: Full Cell Cycling of an Acid-Based Redox Flow Battery ........................................................... 95
5.1 AQDS/BQDS Full Cell Studies and Results ....................................................................................... 95
5.1.1 Solubility of the Acid Form and Sodium Salt of the Quinones ................................................. 95
5.1.2 Effect of Flow Field Modifications ............................................................................................ 97
5.1.3 Electrode Design ...................................................................................................................... 100
5.1.4 Power Density .......................................................................................................................... 101
5.1.5 Coulombic and Energy Efficiency ............................................................................................ 102
5.2 3,6-dihydroxy-2,4-dimethylbenzenesulfonic acid (DHDMBS) Studies in a Redox Flow Cell ........ 104
5.2.1. Improved Performance .......................................................................................................... 104
5.2.2 Confirmation of No Michael Reaction .................................................................................... 107
5.2.3 Crossover through the Nafion
Membrane ........................................................................... 110
5.3 Mixed Cell Studies and Potential for Bi-functional Redox Molecule ............................................ 112
5.3.1. New Membrane Testing ......................................................................................................... 112
5.3.2. Mixed Cell Studies .................................................................................................................. 114
5.3.2.1 Potential for Bifunctional Molecule ..................................................................................... 119
5.4 Summary ......................................................................................................................................... 120
Chapter 6: Conclusions ............................................................................................................................. 123
PUBLICATIONS, PATENTS, AND PRESENTATIONS ................................................................................... 127
References ................................................................................................................................................ 129
vii
LIST OF TABLES
Table 1: Electrochemical Energy Storage Technology for Stationary Applications
9,11-14
............................. 4
Table 2: Cyclic Voltammetry or Linear Sweep Voltammetry of Various Hydroquinones. All voltages are
against a MSE Reference Electrode (E° = +0.650 V), all quinone concentrations are 1 mM in 1 M sulfuric
acid. The working electrode was glassy carbon and the counter was a platinum wire. The scan rate was
50 mV/sec, and the rotation rate was 1500 rpm. ....................................................................................... 30
Table 3: Electrochemical Properties of Various Hydroquinones ................................................................ 37
Table 4: Electrochemical Properties of Di-sulfonic Acid Substituted Anthraquinones .............................. 44
Table 5: Kinetic Parameters and CVs of Methyl-Substituted Hydroquinones in Acid Media. All voltages
are against a MSE Reference Electrode (E° = +0.650 V), all quinone concentrations are 1 mM in 1 M
sulfuric acid. The working electrode was glassy carbon and the counter was a platinum wire. The scan
rate was 50 mV/sec. ................................................................................................................................... 66
Table 6: Effect of Temperature and Added Acid Concentration on Proto-desulfonation Rate ................. 70
Table 7: Cyclic Voltammograms of Naphthoquinoe Derivatives in Alkaline Media. All voltages are against
a MMO Reference Electrode (E° = +0.140 V), all quinone concentrations are 1 mM in 1 M potassium
hydroxide. The working electrode was glassy carbon and the counter was a platinum wire. The scan rate
was 50 mV/sec. ........................................................................................................................................... 76
Table 8: Structures and Bi-Functional Potentials of Hydroxy-Substituted Anthraquinone Derivatives in
Alkaline Media ............................................................................................................................................ 78
Table 9: Hydroquinone Molecules Tested in Alkaline Media. All voltages are against a MMO Reference
Electrode (E° = +0.140 V), all quinone concentrations are 1 mM in 1 M potassium hydroxide. The working
electrode was glassy carbon and the counter was a platinum wire. The scan rate was 50 mV/sec. ........ 88
Table 10: Double Layer Capacitance and Electrochemically Active Surface Area of Various Types of SGL
Carbon Felt (Geometric Area of 25 cm
2
) Determined from the Data in Figure 45. .................................. 100
Table 11: Redox Flow Battery Membrane Physical Characteristics.......................................................... 113
viii
LIST OF FIGURES
Figure 1: World net electricity generation by fuel type, 2012-2040. Adapted from U.S. Energy
Information Agency ...................................................................................................................................... 2
Figure 2: World net electricity generation from renewable power by fuel type, 2012-2040. Adapted
from U.S. Energy Information Agency .......................................................................................................... 3
Figure 3: Ragone plot of current battery technologies with stationary electrodes. Reprinted from
reference 15. Copyright 2011 Royal Society of Chemistry. .......................................................................... 5
Figure 4: Schematic of a redox flow cell ....................................................................................................... 6
Figure 5: A generalized “scheme of squares” for the two-proton, two-electron redox behavior of
benzoquinone, adapted from references 39 and 69 ................................................................................. 14
Figure 6: Pictures of the acidic and alkaline media rotating disk electrode and cyclic voltammogram set-
up ................................................................................................................................................................ 19
Figure 7: Rotating disk electrode set-up ..................................................................................................... 20
Figure 8: 1mM of BQDS in 1M sulfuric acid. The three-electrode set up consisted of a glassy carbon
rotating disk working electrode, a platinum counter electrode, and a MSE reference electrode. Rotation
rate is as indicated (in RPMs). Sweeps were performed from 0.1 V to 0.8 V. ............................................ 23
Figure 9: Calculations for BQDS to determine its rate constant, exchange current density, and diffusion
coefficient. All solutions were 1mM of BQDS in 1M sulfuric acid. Linear sweeps are shown in Figure 8. . 23
Figure 10: Picture of redox flow cell set-up ................................................................................................ 26
Figure 11: Theoretical calculations for the various hydroquinone molecules under study done by Dr.
Sankarganesh Krishnamoorthy ................................................................................................................... 41
Figure 12: Calculations of intramolecular hydrogen bonding on a benzoquinone ring with one or two
sulfonic acid group substituents ................................................................................................................. 42
Figure 13: Cyclic voltammetry (a) and rotating disk electrode linear sweep voltammetry (b) on a glassy
carbon disk of various anthraquinone derivatives. All scans were done at 50 mV/sec and at a rotation
rate of 1500 rpm. Redox couples were at a concentration of 1 mM dissolved in 1 M sulfuric acid. A MSE
reference electrode (E° = +0.650 V) and a platinum counter electrode were used. .................................. 45
Figure 14: 1H NMR at 400 or 500 MHz of BQDS samples taken after various numbers of cycles. Solutions
were diluted with deuterated water (D2O). ............................................................................................... 47
ix
Figure 15: a) Schematic of the transformation of BQDS during charging; b) Schematic of the Michael
reaction showing the nucleophilic addition of water followed by re-aromatization and proton exchange.
.................................................................................................................................................................... 48
Figure 16: Rotating disk electrode linear sweep voltammetry at glassy carbon disk (a) AQDS and (b)
BQDS at various states of charge during cycling. The %SOC values refer to the amount of charge with
reference to the theoretical discharge capacity based on BQDS. The theoretical capacity is based on the
total charge required to transform BQDS to its “cyclable” form. This value is tantamount to 6 Faradays to
charge every mole of BQDS. All scans were done at 50 mV/sec and at a rotation rate of 1500 rpm. Redox
couples were at a concentration of 1 mM dissolved in 1 M sulfuric acid. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. ........................................................................... 49
Figure 17: 25 cm
2
redox flow cell, 0.2 M BQDS and AQDS, charged and discharged at 2A (80 mA/cm
2
)
with graphite felt electrodes. a) Coulombic efficiency for the first 20 cycles when cycled between 1 V and
0 V (100 % depth-of-discharge); b) Charge curves over the first four cycles. ............................................ 52
Figure 18: CV and RDE of 1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid on a glassy carbon working
electrode. A MSE reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The
scan rate was 50 mV/sec. The quinone had a concentration of 1 mM in a 1 M sulfuric acid solution. ..... 54
Figure 19: a) Cyclic voltammograms of 1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid (THBDS, red),
and cycled BQDS (after 30 cycles – blue) and (after 100 cycles – yellow), and b) linear sweeps of cycled
BQDS material after 30 cycles (blue), and THBDS (red). A glassy carbon rotating disk working electrode, a
mercury/mercurous sulfate reference electrode and a platinum counter electrode were used. The scan
rate was 50 mV/sec and the rotation rate was 1500 rpm. Redox materials had a concentration of 1 mM
in 1 M sulfuric acid. ..................................................................................................................................... 55
Figure 20: For both compounds, a MSE reference electrode (E° = +0.650 V) and a platinum counter
electrode were used. The scan rate was 50 mV/sec. The quinone had a concentration of 1 mM in a 1 M
sulfuric acid solution. Blue: CV of 1,2,4,6-tetrahydroxy benzene; Green: CV of 1,2,4,6-tetrahydoxy, 3,5-
benzenedisulfonic acid................................................................................................................................ 56
Figure 21: 1 mM BQDS in 1 M sulfuric acid, with a graphite rod working electrode. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was as indicated, in
mV/sec. Inset shows the peak height variance with the square root of scan rate .................................... 58
Figure 22: 1 mM BQDS in 1 M sulfuric acid, with a graphite fiber working micro-electrode. A MSE
reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec. ....................................................................................................................................................... 59
x
Figure 23: Cyclic voltammograms of 1 mM benzoquinone di-tertbutyl in 1M sulfuric acid on a graphite
rod working electrode, at increasing scan rates (as indicated in the figure, in mV). A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. .................................................. 60
Figure 24: Cyclic voltammograms of 1 mM 2,6-dimethoxy, 1,4-benzoquinone in 1 M sulfuric acid on a
graphite rod (orange) and a glassy carbon (blue) working electrode. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec. .......................... 61
Figure 25: Cyclic voltammograms of various substituted pyrogallol molecules. The quinone
concentrations were 1 mM in 1 M sulfuric acid on a glassy carbon working electrode. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec. . 63
Figure 26: Cyclic voltammograms of molecules at different voltage cut-offs. The quinone concentration
was 1 mM in 1 M sulfuric acid on a glassy carbon working electrode. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec. .......................... 63
Figure 27: 3-hydroxy, 4,6-disulfo, o-benzoquinol cyclic voltammograms on various electrode surfaces
(labeled). The quinone concentration was 1 mM in 1 M sulfuric acid. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec. .......................... 64
Figure 28: Cyclic voltammograms of 3-hydroxy, 4,6-disulfo, o-benzoquinol (red) and 3-sulfo, 5-carboxyl
acid, o-benzoquinol (blue) on a glassy carbon working electrode. The quinone concentration was 1 mM
in 1 M sulfuric acid. A MSE reference electrode (E° = +0.650) and a platinum counter electrode were
used. The scan rate was 50 mV/sec. ........................................................................................................... 65
Figure 29: Intra-molecular hydrogen bonding effects after first proton-electron transfer for two different
pyrogallol molecules ................................................................................................................................... 65
Figure 30: Mechanism for Proto-desulfonation ......................................................................................... 69
Figure 31: Thermodynamic diagram of proto-desulfonation, based on pH. Adapted from reference 114
.................................................................................................................................................................... 69
Figure 32: Linear sweeps of various hydroxy-substituted anthraquinones on a glassy carbon working
electrode (see labels). The quinone concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec, and the rotation rate was 1500 rpm............................................................................................. 78
Figure 33: Linear sweeps of various hydroxy-substituted anthraquinones on a glassy carbon working
electrode (see labels). The quinone concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec, and the rotation rate was 1500 rpm............................................................................................. 80
xi
Figure 34: Scheme of quinoxaline reduction .............................................................................................. 81
Figure 35: Linear sweeps of quinoxaline (yellow) and 5-methyl quinoxaline (red) on a glassy carbon
working electrode. The quinoxaline concentrations were 1 mM in 1 M sulfuric acid. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec,
and the rotation rate was 1500 rpm. .......................................................................................................... 82
Figure 36: Scheme of quinoxaline reduction/oxidation at different pH. Adapted from reference 118 .... 83
Figure 37: Cyclic voltammograms of quinoxaline (red), 5-methyl quinoxaline (blue), and 6-carboxlyate
quinoxaline (green) on a glassy carbon working electrode. The quinoxaline concentrations were 1 mM in
1 M potassium hydroxide. A MMO reference electrode (E° = +0.140 V) and a platinum counter electrode
were used. The scan rate was 50 mV/sec. .................................................................................................. 84
Figure 38: Cyclic voltammograms of riboflavin (blue), riboflavin phosphate (red) on a glassy carbon
working electrode. The redox couple concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec. ....................................................................................................................................................... 85
Figure 39: Riboflavin Molecule ................................................................................................................... 85
Figure 40: 50 mM Riboflavin phosphate as the positive side material and 100 mM potassium
ferrocyanide as the negative side material in 1 M KOH. Cycling was done at a current density of 2
mA/cm
2
. ...................................................................................................................................................... 87
Figure 41: RDE of 2,6-dihydroxy, 9-10-anthraquinone and 2,3,6-trimethyl, 1,4-dihydroxy benzenesulfonic
acid and the potential difference between them. Rotation rate of 1500 rpm. 1mM concentration of
quinone in 100 mL of 1M sodium hydroxide. A MMO reference electrode (E° = +0.140 V) and a platinum
counter electrode were used. The scan rate was 50 mV/sec, and the rotation rate was 1500 rpm. ........ 92
Figure 42: Current-voltage curves for the flow cell: 1M of BQDS/AQDS in acid form (labeled A) and 0.2 M
of BQDS/AQDS in sodium salt form (labeled B). Both cells employed Toray® paper electrodes and a
“flow-by” flow field design. ........................................................................................................................ 96
Figure 43: Diagram of improved flow field arrangement with modified interdigitated flow field with
about 20% of “flow-by” properties. ............................................................................................................ 98
Figure 44: Polarization Curves (top figure) and power density curves (bottom figure) at 100% state-of-
charge for A) 1M AQDS/BQDS in 1M Sulfuric acid, with a “flow-through” flow field and carbon-coated
carbon felt electrodes and B) 1M AQDS/BQDS in 1M sulfuric acid, with a “flow-by” flow field and carbon-
coated Toray® paper electrodes. ................................................................................................................ 99
xii
Figure 45: The imaginary component of the impedance plotted against the reciprocal of the angular
frequency of excitation. The equations are for the line fits for the plotted data. ................................... 101
Figure 46: Coulombic efficiency and energy efficiency over 100 cycles for a 25 cm
2
flow cell with 1 M
AQDS/BQDS with a modified interdigitated flow field and carbon-coated graphite felt electrodes (flow
rate of 1 L/min). The cell was charged and discharged between 0 V and 1 V at 100 mA/cm
2
. Energy
efficiency was calculated based on voltage-time curves that were corrected for the internal resistance.
.................................................................................................................................................................. 103
Figure 47: Charge and discharge curves for 1 M DHDMBS as the positive electrolyte and 1 M 2,7-AQDS as
the negative electrolyte, each dissolved in 100 mL of 1 M sulfuric acid. The cell used interdigitated flow
fields and graphite felt electrodes. The membrane was Nafion® 117. Charge and discharge currents were
2.5 Amperes, and the cut-off voltage was 1.0 V during charge, and 0.005 during discharge. The dashed
lines represent the IR-corrected voltages; the solid lines the non-corrected voltage values. ................. 105
Figure 48: Cycling studies of 1 M DHDMBS and 1 M 2,7-AQDS when charged and discharged at 100
mA/cm
2
. (a) Coulombic efficiency of flow cell; (b) Current-voltage curve at the end of charge in the 25th
cycle. ......................................................................................................................................................... 106
Figure 49: Capacity of flow cell operating with 1 M DHDMBS and 1 M 2,7-AQDS when charged and
discharged at 100 mA/cm
2
. ....................................................................................................................... 108
Figure 50: 1H- NMR studies on samples of DHDMBS before and after 25 cycles. Imidazole was added
deliberately as an internal standard for estimating concentration. Electrolyte solutions were diluted with
deuterated-water (D2O). .......................................................................................................................... 109
Figure 51: Linear-sweep voltammogram of DHDMBS samples before and after 25 cycles. Both samples
were at a 1 mM concentration. The scan rate was 50 mV/sec and the rotation rate was 1500 rpm. The
working electrode was a glassy carbon rotating disk electrode, the counter electrode was a platinum
wire, and the reference electrode was mercury sulfate (MSE, E° = +0.65 V). .......................................... 110
Figure 52:
1
H- NMR studies on samples of 2,7-AQDS before and after 25 cycles. Imidazole was added
deliberately as an internal standard for estimating concentration. Electrolyte solutions were diluted with
deuterated-water (D2O). .......................................................................................................................... 112
Figure 53: Capacity and cycle life of various flow cells operating with 1 M DHDMBS and 1 M 2,7-AQDS
when charged and discharged at 100 mA/cm
2
. Various membranes were employed. Capacity fade rate
was normalized for thickness of membrane............................................................................................. 114
Figure 54: AQDS/DHDMBS mixed electrolyte cell in 1 M sulfuric acid solutions. Solutions were mixed and
split after capacity had faded 25%. A Nafion® 117 membrane was used. ............................................... 115
xiii
Figure 55: Capacity and cycle life of various flow cells operating a mixed (symmetric) cell with 0.5 M
DHDMBS and 0.5 M 2,7-AQDS on both sides. Charge and discharge current density was 100 mA/cm
2
.
Various membranes were employed. ....................................................................................................... 117
Figure 56: Linear-sweep voltammogram of DHDMBS/AQDS mixed cell samples before and after 216
cycles. Both samples were at a 1 mM concentration. The scan rate was 50 mV/sec and the rotation rate
was 1500 rpm. The working electrode was a glassy carbon rotating disk electrode, the counter electrode
was a platinum wire, and the reference electrode was mercury sulfate (MSE, E° = +0.65 V). ................ 118
Figure 57:
1
H- NMR studies on DHDMBS/2,7-AQDS mixed cell samples before and after 216 cycles.
Imidazole was added deliberately as an internal standard for estimating concentration. Electrolyte
solutions were diluted with deuterated-water (D2O). ............................................................................. 118
Figure 58: Sulfonated bifunctional molecule ............................................................................................ 119
Figure 59: Cyclic voltammogram (red) and RDE (blue) of the sulfonated bi-functional molecule on a
glassy carbon working electrode. The bi-functional molecule concentration was 1 mM in 1 M sulfuric
acid. A MSE reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan
rate was 50 mV/sec for both, and the RDE rotation rate was 1500 rpm. ................................................ 120
xiv
LIST OF ABBREVIATIONS AND SYMBOLS
EIA: U.S. Energy Information Agency
EES: Electrical Energy Storage
U. S. DoE: U.S. Department of Energy
RFB: Redox Flow Battery
ORBAT: Organic Redox Flow Battery
AQDS: 9,10-anthraquinone, 2,6-disulfonic acid
BQDS: 4,5-dihydroxybenzene-1,3-disulfonic acid
NHE: Normal Hydrogen Electrode
MMO: Mercury-Mercuric Oxide Reference Electrode
MSE: Mercury/Mercuric Sulfate Reference Electrode
RDE: Rotating Ring Disk Electrode
CV: Cyclic Voltammetry
ω: Rotation rate (radians/second)
RPM: rotations per minute
DHDMBS: 3,6-dihydroxy-2,4-dimethylbenzenesulfonic acid
xv
SUMMARY
This thesis is focused on all aspects of the design of an all-organic, aqueous, redox flow battery
(ORBAT). This kind of battery has the potential to be implemented on a grid-scale level, as it utilizes
environmentally-friendly and inexpensive materials. As more renewable energy power sources, such as
wind turbines and solar panels, are added to electricity grids across the world, it is of paramount
importance to design energy storage that is efficient, long-lasting, robust, cheap, and environmentally-
friendly. An organic redox flow battery like the one discussed in this thesis could be the solution to the
need for robust and long-lasting energy storage, as it has the potential to fulfill all of these
requirements. Another advantage of ORBAT is that the energy and power densities are de-coupled, thus
making it easily scalable for the required application.
This work covers many aspects of ORBAT, from the study of the properties of the redox-active
molecules to cell design. Knowing the physiochemical properties, understanding the mechanisms of the
electrochemical and chemical reactions, and measuring the interactions of the electrode surfaces with
the quinones is of paramount importance to effectively utilize these quinone molecules in large scale
energy storage systems. These properties are dependent on a number of factors, including solvent
system, geometry and molecular structure of the quinone, concentration, temperature, pH, electrode
material, and structure of the electrode surface. I have focused primarily on studying the effects of ring
substituents, pH, and electrode materials – I have tested 40 different quinone and quinone-related
structures for their reduction potentials, solubility, and electron and proton transfer kinetics. As a result,
I have been able to select the most suitable redox couples and have validated their performance in a
redox flow battery. I have also demonstrated major advances to the design of an all-organic aqueous
redox flow battery, which is the first of its kind.
xvi
In Chapter 1, I discuss the importance and need for grid-scale energy storage. As increased
amounts of renewable energy power sources are installed around the globe, it is essential that cheap,
efficient, long-lasting, and environmentally-friendly energy storage is installed. Redox flow batteries are
an attractive solution for this need, specifically aqueous, organic redox flow batteries. Quinones are a
promising group of molecules for use in redox flow batteries, as they can be used in aqueous media,
they have fast electrochemical reactions, and their reduction potentials can be tuned by changing
substituent position.
In Chapter 2, I describe all experimental techniques that were employed for this research. These
include electroanalytical methods (RDE, CV), quantum mechanical calculations, synthetic procedures of
the organic molecules under study, and flow cell set-up.
In Chapter 3, I discuss redox materials suitable for an acid-based redox flow battery. I studied 22
hydroquinone compounds for their use as positive side electrolyte materials. The effect of molecular
architecture on the electrochemical properties and cycleability is discussed. Some initial cycling results
(with anthraquinone complexes employed as negative side electrolytes) are presented, and the types of
chemical transformations that these molecules undergo are discussed. Understanding of the chemical
transformations and ways to prevent them from happening are also presented. It was also shown that
the electrode surface structure affects the electron transfer kinetics. Finally, I discuss the
electrochemical properties of sulfonic acid substituted anthraquinones for use as the negative side
electrolyte material.
In Chapter 4, alkaline-based redox flow batteries are discussed. There have been very few
examples of alkaline-based RFBs in the literature, as it is difficult to find compounds that are stable,
soluble, and give the required reduction potentials for this kind of RFB. Several different types of
xvii
molecules, including quinoxaline, riboflavin, and benzo-, naptho-, and anthra- quinones were researched
for their usefulness in this kind of system.
In Chapter 5, full cell cycling results were presented. Flow field design, electrode structure, and
membrane selection are important engineering parameters to optimize the performance of RFBs. It was
determined that a flow field with more flow-through characteristics used with carbon felt electrodes
was an optimal design for the interior of the flow cell. Mitigation of crossover is important to increase
the cycle life of the battery, and it was determined that Fumatech membranes could limit the transport
of the hydroquinone molecules because of their lower water content. The capacity fade rate could be
reduced by this approach.
Finally, I discuss the prospect of employing symmetric cells in a redox flow battery. One kind of
symmetric cell uses mixed electrolytes on both sides – different cycling protocols for this kind of system
are discussed. We have also shown that a single type of bifunctional molecule is viable for use as both
positive and negative electrolytes in ORBAT, which can also mitigate the effects of crossover. This
approach points to new avenues for future research.
Chapter 6 contains conclusions from my research and summarizes the understanding and
implications of the findings for future development of aqueous organic redox flow batteries.
xviii
1
Chapter 1
Introduction
1.1 Grid-Scale Energy Storage
As the global community grapples with climate change and its effects, increasing amounts of
renewable energy are being produced worldwide in an effort to move away from the use of fossil fuels.
Solar energy and wind energy are the most common renewable energy resources being tapped for
energy production. However, with the increased deployment of these renewable energy generators, the
variable and intermittent energy output of these systems is a grave concern for the stable operation of
the electricity grid. To match these inevitable surges in supply with the energy demand profile, it is
necessary to store electrical energy at a large scale to serve as a buffer. Such energy storage systems
must be capable of storing thousands of giga-watt hours of electricity per day to serve the global energy
demand, as per the U.S. Energy Information Agency (EIA). The EIA projects that even storing 20% of the
global electricity production of about 100 Terawatt hours/day
1
will require the deployment of 10-15
gigatons of batteries over a 15-year period, assuming a modest specific energy of 50 Wh/kg for storage
devices.
21.6 trillion kilowatthours of electricity were generated in 2012, and this number is expected rise by
69% by 2040, with 36.5 trillion kilowatthours projected to be produced in 2040.
2-3
Electricity is estimated
to be 80% of the total energy production by 2050.
4
About 22% of energy generated globally in 2012 was
from renewable sources (including biomass, wind, solar, geothermal, and hydroelectric) and that
number is expected to rise to 29% by 2040 (Figure 1).
5
In many states in the United States, the goal for
electricity production from renewable resources is 12% by 2020, and Europe has set an even higher
target for itself – 20% by 2020.
5
2
Figure 1: World net electricity generation by fuel type, 2012-2040. Adapted from U.S. Energy
Information Agency
However, the most environmentally friendly renewable energy sources (wind and solar) are also
the most intermittent. While these methods of generation constitute only a small fraction of produced
power compared to hydroelectric power, the global solar and wind capacity installed each year is
growing at the rate of 60% and 20%, respectively (Figure 2). This large growth is due to the push towards
reducing carbon emissions. However, the growing number of intermittent energy sources also demands
the mitigation of their fluctuating power output with energy storage systems. Therefore, electrical
energy storage (EES) is of crucial importance for the development of the smart grid and distributed
power generation.
3
Figure 2: World net electricity generation from renewable power by fuel type, 2012-2040. Adapted from
U.S. Energy Information Agency
The U.S. Department of Energy (U. S. DoE) has identified four major challenges to the
widespread implementation of EES: cost, safety and reliability, equitable regulatory environments, and
industry acceptance.
6
The U.S. DoE has also set a cost target for such a battery at $100/kWh.
7
The latter
two challenges are not something that scientists can quite solve – this is up to the market forces and
policy makers. However, scientists can focus on improving cost, safety, and reliability. It is important to
work with low-cost, safe, low-toxicity, environmentally-friendly, and abundantly-available materials to
make a long-lasting, efficient, and robust energy storage system for grid scale energy storage.
Rechargeable batteries are particularly attractive for EES because of their high energy efficiency
and scalability.
8-10
However, for such a large-scale application such as grid-scale energy storage, these
4
batteries must meet several very important criteria – low cost, robust, safe, and sustainable. None of
today’s commercially-available batteries can simultaneously meet all the performance and cost targets
for grid-scale energy storage. Today’s mature battery types include lead-acid, lithium ion, nickel-metal
hydride, and sodium-sulfur – all of which are either expensive, have safety concerns, are unsustainable,
or fail to provide the desired performance.
Table 1: Electrochemical Energy Storage Technology for Stationary Applications
9,11-14
Technology Typical Energy
Content
(MWh)
Storage
Cost
($/kWh)
Life Time
(cycles/years)
Efficiency
(%)
Drawbacks
Supercapacitors 0.001-10 500-3000 500,000/20 >90 Low energy
density, high
cost
Lead-Acid
Batteries
3-50 65-400 300-1000/1-5 70-90 Low energy
density, short
lifetime,
temperature
sensitive
Li-ion Batteries 0.1-50 400-600 500-2000/3-8 85-95 High cost, safety
risks, short
lifetime, self-
discharge,
temperature
sensitive
Sodium-sulfur
Batteries
0.25-50 300-500 2000-5000/6-
15
70-90 High cost, high-
temperature
operation, high
safety risks
Flow Batteries
(Vanadium)
0.5-40 150-2500 10000/5-15 60-85 Low energy
density, high
materials cost
Nickel-Metal
Hydride
0.001-0.1 1000 500-3000/3-7 85-90 Moderate
energy density,
high cost of
electrode
materials
5
Figure 3: Ragone plot of current battery technologies with stationary electrodes. Reprinted from
reference 15.
15
Copyright 2011 Royal Society of Chemistry.
The Ragone plot (Figure 3) depicts how power and energy in these current battery systems are
coupled. Furthermore, the cycle life of batteries is often dependent on the depth of discharge. Thus,
batteries may have to be sized with excess energy not only to meet the power requirements but also the
cycle life demands. These adjustments to the sizing of the batteries to simultaneously accommodate a
number of different requirements leads to higher than desired costs for many applications. This
situation has spurred a global search for a transformational solution. Addressing this challenge will open
up entirely new markets for energy storage – either for use by the utilities at the grid scale, at the
customer site (behind-the-meter) in developed countries, or for varied types of deployment with
renewable energy in under-developed countries as they continue to make gains towards energy
independence. Thus, cheap and reliable energy storage can truly transform how the world uses
renewable energy.
6
1.2 Redox Flow Batteries: History and Overview
A redox flow battery (RFB) is a particular type of battery configuration in which the energy is
stored in a solution or a slurry of two distinct redox couples. When the redox material flows through the
electrochemical cell anode and cathode chambers, these redox materials undergo electron-transfer
reactions at inert and conductive electrodes. The two sides of the cell are separated by an ion-
conducting polylmeric membrane. Unlike in conventional batteries, there is no physical transfer of
material across the electrode/electrolyte interface, as RFB’s rely on reversible solution-phase
electrochemical couples to store chemical energy. The solutions of these electrochemical couples are
stored in external tanks, as shown in Figure 4. Both the oxidized and reduced forms of each redox couple
are soluble in the electrolyte, and only the relative concentrations of oxidized and reduced forms change
during cycling.
16
Figure 4: Schematic of a redox flow cell
7
By virtue of the unique arrangement in RFBs, the electrodes in RFB’s do not undergo physical
changes during operation as all the changes occur in the dissolved reactants that react at the solid-
electrode surfaces. Consequently, RFBs have a simplified electrode design that ensures that their cycle
life is not directly influenced by depth-of-discharge or the number of cycles in the same way that
conventional rechargeable batteries are and the degradation of the electrode surface is not nearly as
much of an issue.
Typically, a number of individual electrochemical cells are connected in series to form a stack.
The electrolytes stored in separate tanks are pumped through the stack. RFBs were initially researched
by NASA in the 1970s with the goal of establishing energy storage for a base station on the Lunar
surface.
17-20
Thus, RFBs are relatively recent compared to the well-developed lead-acid battery.
21
RFBs
are not suitable for mobile applications because of their relatively low energy density. However, RFBs
have the advantages of independent scalability of energy and power, along with long cycle life
compared to conventional batteries.
22
The power density is determined by the cell and stack size and the energy content depends on
the difference in electrode potentials for the redox couples and the amount of material stored in the
tanks.
23-25
However, in a conventional battery, the amount of energy that can be stored is generally
limited by the amount of material attached to the electrodes and the effective path lengths for diffusion
and migration in the direction normal to the current collector. Thus, making an electrode thicker will add
to the amount of active material, but will also increase those effective path lengths and leading to
greater voltage losses from diffusional and ohmic processes. As shown in Figure 4, most RFB systems
require two separate electrolyte tanks: one for the anolyte and another for the catholyte. This
arrangement ensures that the potentials at each electrode are close to the reversible potential for each
of the half-cell reactions.
8
The main costs of RFBs include the active redox materials and the electrochemical cell
components. The costs of the cell scale with the total power requirement of the application, but these
costs are also directly related to the specific power of the device itself, i.e., how effectively the materials
are utilized. RFBs operate at relatively high current densities, as convection can be employed to deliver
reactants to the electrode surface. Thus, electrolyte management and cell design have a significant
impact on power density and cell material costs.
The most advanced and the most studied RFB is based on vanadium redox couples. It was first
invented in 1986 by Syllas-Kazacos et al.
26-28
The battery chemistry is based on the V
II
/V
III
redox reaction
in the anolyte, and the V
IV
/V
V
redox reaction in the catholyte (this reaction being considerably slower
and more complex) resulting in a cell voltage of approximately 1.3 V. The high cell voltage and high
solubility are the principal advantages of the vanadium RFB. The solubility for these compounds is up to
4M in varying concentrations of sulfuric acid, which yields a high theoretical specific energy of 60.5
Wh/kg.
29
However, this value of specific energy is far from being realized during operation, as the
precipitation of V 2O 5 at the required elevated operation temperatures severely limits solubility and
operation range. Even with this drawback, several companies are already extensively involved in the
development and commercialization of all-vanadium RFB’s. These companies include Sumitomo Electric,
UniEnergy Technologies, American Vanadium, Gildemeister AG, WattJoule, REDT, and Ashlawn Energy.
30
The discharge time and size of these systems vary from 3 to 10 hours and from tens of kWh to 7 MWh.
The major issue with this technology is the high cost of the vanadium-based redox materials; for
example – a cost analysis of a 300 kWh unit revealed that the vanadium-related electrolyte cost is
approximately 62% of the total cost.
31-32
Therefore, even though several companies are pursuing this vanadium-based RFB technology
and there has been much research done on this system, the high cost of the electrolyte is a difficult
hurdle to overcome. Additionally, the viscosity of these electrolyte systems is high, which results in
9
greater pumping losses. Furthermore, the electrolyte has a high rate of crossover through the polymer
electrolyte membrane, leading to decreased cycle life and capacity fade. Therefore, other redox couples
that could lower the cost of the electrolyte materials and the system cost are much needed.
1.3 Quinone Introduction
Quinones are among the most important and well-studied examples of organic redox molecules
over the last century.
33-36
Redox properties of quinones, including reaction mechanisms, effect of pH
conditions, kinetic parameters and physiochemical properties have been studied in great detail.
37-40
Certain quinones have been useful in medicinal and biochemical applications, including anthracyclines
that have been used to treat certain types of human cancer.
41
Quinones have found many industrial uses
such as in dye manufacturing
42
and as oxidizing and reducing agents.
43
A clear understanding of their
electrochemical and chemical properties was paramount to implementing quinones in all of these
important applications.
In the year 2011, we at the University of Southern California showed for the first time that
quinones can be employed in redox flow batteries suitable for large scale energy storage. Quinones –
with their fast electrochemical kinetics,
44-45
relatively low cost,
46
variety of molecular structures, and
non-toxic nature – are ideal candidate materials for designing a large-scale energy storage system. These
redox couples can undergo fast proton-coupled electron transfer in aqueous solutions and also provide
a large charge capacity. I have measured the heterogeneous rate constants for the charge-transfer
processes to be in the range of 10
-3
to 10
-4
cms
-1
.
41
The specific charge capacity for these molecules can
range from 200 to 500 Ah/kg. In addition, a variety of the molecular structures are possible with the
quinones, allowing the tuning of their properties; substituent groups can be used to modify solubility in
water, standard reduction potential, and the kinetics of the redox reactions.
47
The quinones are
expected to cost about $10-15/kg or $20-30/kWh, leaving ample scope for meeting the cost target for
the entire battery system of $100/kWh as set by US DoE, provided high power densities and voltage
10
efficiencies can be achieved.
48
Typical organic redox couples are 4,5-dihydroxybenzene-1,3-disulfonic
acid (BQDS) and anthraquinone-2-sulfonic acid (AQS) (Eqs. 1 and 2).
HO
3
S SO
3
H
OH
OH
+ 2e
-
+ 2H
+
O
O
HO
3
S SO
3
H
(1)
E
0
= + 0. 85 V
O
O
SO
3
H
OH
OH
SO
3
H
+ 2e
-
+ 2H
+
E
0
= + 0.09 V
(2)
In the last few years, there have been a number of studies investigating the uses of quinones in
redox flow batteries. In 2009, Xu et al studied the use of BQDS at the positive electrode and the
lead/lead sulfate couple for the negative electrode.
49
Aziz et al have reported on a flow cell using
aqueous solutions of anthraquinone disulfonic acid at the negative electrode and a conventional
bromine/tribromide couple at the positive electrode.
50-51
More recently, Aziz et al have also reported a
redox flow battery with an anthraquinone-based couple on the negative electrode with the
ferro/ferricyanide couple at the positive electrode in alkaline media.
52
Additionally, Aspuru-Guzik et al
published computational studies to identify quinone compounds that would be useful in an all-organic
redox flow battery.
53
While some of these molecules are predicted to show high positive electrode
potentials, their chemical stability and solubility in aqueous media and the propensity of the non-ionized
forms to crossover through the cation exchange membrane need to be considered.
Organic redox flow batteries (which we have termed ORBAT) have grown in popularity – even
within my time in graduate school – and present a rich opportunity for research. Although our focus at
11
USC has been on water-soluble quinone molecules, research in non-aqueous solutions of organic redox
substances has also been considered by the groups of Rasmussen,
54
Brushett et al,
55
and Abruna et al.
56
Researchers at the Pacific Northwest National Laboratory have investigated the redox behavior of the
stable free radical, TEMPO, in non-aqueous redox flow batteries.
57-59
1.3.1 Charge-Transfer Processes
The principal basis of an efficient rechargeable redox flow battery is the rapid kinetics of charge-
transfer at the positive and negative electrodes. Many organic redox couples, especially from the
quinone family, undergo rapid proton-coupled electron transfer without the need for dissociating high
energy bonds. Consequently, these redox couples have relatively high rate constants for the charge-
transfer process. In general, we expect those molecules with conjugated carbon-carbon bonds and keto-
and enol- groups that allow for the delocalization and rearrangement of the pi-electrons to undergo
these redox transformations with facility.
60
Such quinone-based redox couples have electron-transfer rate constants that are 2-3 orders of
magnitude greater than those of the reactions of vanadium redox couples in the commercial vanadium
redox battery system.
61-62
Therefore, overpotential losses from charge-transfer are expected to be low
with these organic redox couples. As dissociation and rearrangement of C-H bonds do not occur in these
electrochemical reactions, no precious metal catalytic electrodes are required; high-surface area
conductive electrode surfaces, such as those based on carbon black, are sufficient to support the
charge-transfer process. Selecting compounds with high rate constants (on the order of 10
-3
to 10
-4
cm/sec) is necessary when considering the technical viability of ORBAT.
1.3.2 Cell Voltage
The standard reduction potential of the organic redox couple is characteristic of the molecule
and its specific substituent groups. The difference in the standard reduction potentials determines the
12
cell voltage for the redox flow battery. This cell voltage is also determined by the difference in energy of
the highest occupied molecular orbital (HOMO) of redox couple used as the negative electrode material
and the lowest unoccupied molecular orbital (LUMO) of the redox couple used as the positive electrode
material. Previous experimental studies on quinones show that electron-withdrawing substituent groups
lower the energy levels of the HOMO and LUMO, while electron-donating substituents raise these
energy levels.
63
Thus, we can use substituents to selectively tune the standard reduction potential of the
quinone compounds to achieve the desired cell voltage. We can calculate the standard reduction
potential for a redox couple from the free-energy change of the reduction of the redox couple by
hydrogen. This free energy change may be calculated using quantum mechanics by a method called the
Density Functional Theory (DFT).
Being an aqueous battery, the voltage range for ORBAT is limited by the oxygen evolution
reaction at the positive electrode and the hydrogen evolution reaction at the negative electrode.
Consequently, based on the standard reduction potentials for hydrogen and oxygen evolution reactions,
we do not expect to attain a cell voltage more than 1.23 V at room temperature. However, by inhibiting
the kinetics of the hydrogen and oxygen evolution reactions, we may achieve higher cell voltages. In this
respect, the non-aqueous systems have an advantage of being able to access a wider range in cell
voltage. However, the inherent advantages of lower cost and higher levels of safety presented by
aqueous systems is particularly compelling for large-scale energy storage applications.
1.3.3 Mass Transport Processes
To keep the cost of the battery low, we must operate at high current densities without
compromising voltage efficiency. With the rapid charge transfer processes in quinones, the current
density will be limited by mass transport of the reactants and products. If a current density of 100 mA
cm
-2
is required at a reactant concentration of 1 M, we may use the steady-state Nernst diffusion layer
13
model to estimate the mass transport coefficients to be approximately 5x10
-4
cm/sec.
64
Consequently,
to maintain a current density of 100 mA cm
-2
with a redox couple that has a solubility of 1 M and a
diffusion coefficient of about 1x10
-6
cm
2
/sec, the solutions must be circulated past the electrodes so as
to maintain a thin diffusion layer of approximately 2x10
-3
cm. Such a diffusion layer thickness is in the
practical range of values observed with rapid circulation. Flow-through electrodes can lead to further
reduction in diffusion layer thickness.
65-66
Thus, the diffusion coefficient and solubility of the redox
couple are principal properties for which the values must be as high as possible for reaching the
performance and cost targets for large-scale energy storage applications.
In general, the un-substituted quinones have limited solubility in water. However, the solubility
of the quinones can be increased substantially by the addition of sulfonic acid and hydroxyl substituents.
For example, benzoquinone has a solubility of 0.1 M, while BQDS has a solubility of approximately 1.7 M
at 25
o
C. Further increases in solubility can be achieved by raising the temperature. Therefore, achieving
concentrations as high as 2 M is practical even within the quinone family of compounds. Other organic
redox couples with high aqueous solubility include carboxylic acids and aromatic heterocyclic
compounds.
43
1.3.4 Reactivity and Long-Term Cycling
Under acidic conditions, the electro-reduction of quinones occurs often by a concerted proton
transfer and electron transfer process,
32
and no radical species are produced as part of this redox
process. Sometimes, the mechanisms could involve sequential steps of protonation and electron
transfer.
67
All of these mechanisms are captured in the “scheme of squares” as shown in Figure 5. Thus,
a proton exchange membrane electrolyte with high ionic conductivity can be used in conjunction with
the quinones. Inexpensive hydrocarbon membranes such as polystyrenesulfonic acid, sulfonated
polyetheretherketone and sulfonated polyethersulfone can also be used instead of Nafion
.
68
14
Furthermore, the proton exchange membrane will inhibit the transport of any anionic chemical species
across the membrane. Therefore, any crossover of reactants from one side of the cell to another will be
avoided, thereby avoiding self-discharge. Consequently, the long-cycle life requirement of large-scale
energy storage systems is more likely to be realized with acidic systems.
Figure 5: A generalized “scheme of squares” for the two-proton, two-electron redox behavior of
benzoquinone, adapted from references 39 and 69.
39,69
As the scheme of squares shows (Figure 5), the mechanism for reduction or oxidation is highly
dependent on the pK a values associated with the intermediate species.
70
Under different pH and solvent
conditions, it is possible for the reduction or oxidation process to stop at either a lower protonated or
reduced state. It is known that in alkaline media, quinones are known to undergo stepwise one-electron
transfer reactions. However, there are several system-level advantages that an alkaline system can
offer, including the use of an inexpensive separator. For these reasons this type of alkaline flow battery
is also worth pursuing, if an appropriate set of redox couples can be found.
For continuous operation at temperatures as high as 60
o
C the selected organic compounds
must have sufficient hydrolytic stability. The quinones are generally quite stable. However, contact with
15
air will cause the organic redox couples to be rapidly oxidized, although no permanent loss of material
properties is anticipated at these temperatures. In contrast, the vanadium systems are very sensitive to
high temperature operation because of the reactions between water and the vanadium oxo ions. In
some cases, insoluble products such as vanadium pentoxide are produced that reduce the capacity and
cycle life.
63
1.3.5 Faradaic Efficiency
Faradaic efficiency losses are generally associated with the negative electrode in these types of
systems due to the proximity of the electrode potential to that of the hydrogen evolution reaction. In
the case of anthraquinone and its derivatives, the standard reduction potential is at least 100 millivolts
positive to that of the normal hydrogen electrode (NHE). Consequently, hydrogen evolution cannot
occur readily during charge. Additionally, the hydrogen evolution reaction is substantially inhibited on
carbon materials. This situation in ORBAT is to be contrasted with that in the vanadium redox battery;
the vanadium(III)/vanadium(II) couple operates at a potential approximately 0.350 volts negative to NHE
and hydrogen evolution occurs during normal charging of the battery.
71
Furthermore, the deposition of
metallic impurities on the electrode facilitates hydrogen evolution, thus lowering the overall efficiency
and changeing the pH of the solutions. By avoiding the use of any soluble metal ions and only using
carbon-based electrodes, ORBAT avoids the faradaic efficiency losses that arise from the hydrogen
evolution reaction.
1.4 Focus of My Work
By choosing the appropriate organic redox couples for the positive and negative electrodes, we
projected that a cell voltage as high as 1.0 V can be achieved.
47, 72
ORBAT is unique in that it does not use
any heavy metals such as vanadium, chromium or zinc, and also avoids volatile organic solvents such as
those used in lithium batteries. Additionally, the redox reactions will not require any expensive precious
16
metal catalysts. With the potential of being simultaneously inexpensive and environmentally-friendly,
such a battery presents an attractive solution for grid-scale energy storage. Therefore, designing a
commercialize-able ORBAT is the ultimate goal of this project.
Knowing the physiochemical properties, understanding the mechanisms of the electrochemical
and chemical reactions, and measuring the interactions of the electrode surfaces with the quinones is of
paramount importance in order to effectively utilize these quinone molecules in large scale energy
storage systems. These properties are dependent on a number of factors including solvent system,
geometry and molecular structure of the quinone, concentration, temperature, pH, electrode material,
and structure of the electrode surface. I have focused primarily on studying the effects of ring
substituents, pH, and electrode materials; I have tested 40 different quinone and quinone-related
structures for their reduction potentials, solubility, and electron and proton transfer kinetics. As a result,
I have been able to select the most suitable redox couples and have validated their performance in a
redox flow battery. I have also demonstrated major advances to the design of the first all-organic
aqueous redox flow battery.
17
18
Chapter 2
Experimental Techniques and Methodology for Analysis
of Results
2.1 Electrochemical Characterization
All experiments for the electrochemical characterization of the individual redox molecules were
conducted in a standard three-electrode cell. The cell consists of a working electrode, a platinum-wire
counter electrode for all systems, a mercury/mercurous sulfate reference electrode (E
o
= +0.65 V) for
acidic systems, and a mercury/mercury oxide reference electrode (E
0
= +0.1 V) for alkaline studies (Figure
6). For most experiments, the working electrode consisted of a rotating disk electrode made of glassy
carbon (Pine) (Figure 7). However, for studies on different electrode surfaces various other electrodes
were used, including an isomolded graphite rod (Graphtek LLC), a carbon fiber microelectrode (Pine),
carbon felt (SGL), Toray paper, or an edge-plane graphite rotating disk electrode. The quinones, in either
the fully-reduced or fully-oxidized form, were dissolved in 1 M sulfuric acid to a concentration of 1 mM.
The solutions were de-aerated and maintained under a blanket of argon gas throughout all the
experiments.
Cyclic voltammetry experiments were conducted at a scan rate of 50 mV s
-1
. Linear sweep
voltammetry experiments were conducted with the glassy carbon rotating disk electrode, at a scan rate
of 5 mv s
-1
over a range of rotation rates from 500 rpm to 3000 rpm (Figure 8). All experiments were
conducted using a Versastat 300 potentiostat and rotating electrode equipment from Pine Instruments.
19
Figure 6: Pictures of the acidic and alkaline media rotating disk electrode and cyclic voltammogram set-
up
2.1.1 Electrochemical Method of Analysis
The Rotating Disk Electrode (RDE) is a circular disk electrode embedded in teflon material
connected to a rotating shaft. The rotation rate of the shaft is controlled by an electric motor with a
speed controller (Figure 7). As the RDE rotates, a laminar flow of solution is created perpendicular to the
electrode (along the axis of rotation) which is accompanied by the removal of liquid radially across the
electrode surface. This type of steady state laminar flow across the surface of the RDE allows the
development of a boundary layer of constant thickness.
143
20
Figure 7: Rotating disk electrode set-up
This boundary layer thickness can be used in the Nernst diffusion layer model to relate the
limiting current for the oxidation and reduction processes to the boundary layer thickness (Eq. 3).
𝐼 𝑙𝑖𝑚 = 𝑛𝐹 𝐶 ∗
𝐷 𝛿 (3)
Where I lim is the limiting current density, n is the number of moles of electrons transferred per
mole of reactants, F is the Faraday constant (96485 C mole
-1
), C* is the concentration, D is the diffusion
coefficient, and δ is the diffusion layer thickness.
For a diffusion layer thickness of 50 microns, a diffusion coefficient of 3.8 x 10
-6
cm
2
/sec, and a
bulk concentration of 0.2 M, we predict from Eq. 3 a limiting current density at room temperature to be
approximately 30 mA cm
-2
. Further increase in limiting current density can be achieved by increasing the
21
concentration of reactants, reducing the diffusion layer thickness, and by increasing the diffusion
coefficient. Higher concentrations and diffusion coefficients are achieved by raising the operating
temperature, while a lower diffusion layer thickness can be achieved by increased convective mass
transport to the surface of the electrode. Based on these numbers, many of the quinone compounds
under study met the estimated limiting current density value.
The limiting current observed at various rotation rates at the rotating disk electrode was used to
obtain the diffusion coefficient by applying the Levich equation (Figure 9, Eq. 4).
I lim = 0.62 n F A D
2/3
1/2
-1/6
C* (4)
Where is the angular rotation rate in radians s
-1
, A is the area of the electrode, and is the kinematic
viscosity of the solution. Other constants are the same as defined before in Eq. 3.
I also determined the rate constant and the exchange current density for the charge transfer
process from the linear-sweep voltammetry measurements made on the rotating disk electrode for all
of the compounds investigated. However, I will present data for 4,5-dihydroxybenzene-1,3-disulfonic
acid (BQDS) in this section as an example of the calculations done for all compounds. This data is shown
in Figures 8 and 9.
When the rate of reaction at the surface of the electrode is limited by mass transport, it is
possible to determine the kinetic parameters for charge-transfer and mass transport by using the
Koutecky-Levich equation to analyze the data from the rotating disk measurement (Eq. 5).
1
𝐼 =
1
𝐼 𝑘 +
1
𝐼 𝑙𝑖𝑚 (5)
Where I is the measured current, I k is charge transfer (or kinetic) current, and I lim is mass-transport
limited (or diffusion-limited) current. Once the current was corrected for mass transport, I made a plot
of the logarithm of the kinetic current versus the electrode potential for values of overpotential greater
than 25 mV, where the Tafel approximation of the Butler-Volmer equation is applicable (Figure 9).
22
Thus, kinetic currents can be calculated from the measured current-potential curves by
correcting for mass transport effects. The kinetic current is given by
𝐼 𝑘 =
𝐼 𝑜𝑏𝑠 𝐼 𝑙𝑖𝑚 (𝐼 𝑙𝑖𝑚 −𝐼 𝑜𝑏𝑠 )
(6)
Where I obs is the measured current and I lim is the mass-transport limited current.
From the values of the kinetic current (I k) that were obtained after the mass-transport
correction (Eq. 6), the Tafel parameters were determined by plotting the overpotential, η, versus log I k
(kinetic current) by using the Tafel equation (Eq. 7). The Tafel slope can be obtained from the slope of
this line and the exchange current density (I o) can be calculated from the intercept of this line with the x
axis at η = 0 (Figure 9).
𝜂 =
2.3𝑅𝑇
𝛼𝑛𝐹 𝑙𝑜𝑔 𝐼 𝑜𝑏𝑠 −
2.3𝑅𝑇
𝛼𝑛𝐹 𝑙𝑜𝑔 𝐼 𝑘 (7)
Where α is the transfer coefficient, R is the gas constant, and T is the temperature. The slope of the
Tafel plot along with extrapolation to zero overpotential allowed us to determine the exchange current
density by combining Eq. 6 and 7.
(8)
Where E-E rev is the overpotential. The rate constant for the electrochemical reaction, k o, was calculated
from the exchange current density (Eq. 9).
𝑘 0
=
𝐼 0
𝑛𝐹𝐴 𝐶 ∗
(9)
(
𝐼 1 −
𝐼 𝐼 𝑙𝑖𝑚 ) = 𝐼 𝑜
[
𝐶 𝐶 ∗
𝑒𝑥𝑝 (−
𝛼𝑛𝐹 (𝐸 − 𝐸 𝑟𝑒𝑣 )
𝑅𝑇
)]
23
The half-wave potential, E 1/2, was determined from the electrode potential corresponding to
half the limiting current. For a reversible electrochemical reaction, E 1/2 can be approximated as the
standard reduction potential.
Figure 8: 1mM of BQDS in 1M sulfuric acid. The three-electrode set up consisted of a glassy carbon
rotating disk working electrode, a platinum counter electrode, and a MSE reference electrode. Rotation
rate is as indicated (in RPMs). Sweeps were performed from 0.1 V to 0.8 V.
Figure 9: Calculations for BQDS to determine its rate constant, exchange current density, and diffusion
coefficient. All solutions were 1mM of BQDS in 1M sulfuric acid. Linear sweeps are shown in Figure 8.
24
2.2 Quantum Mechanics-Based Calculations
To predict the E ½ values for various molecules that had not yet been synthesized, we used
density functional theory (DFT) to calculate the standard Gibbs free energy change (ΔG
o
) for the
reduction of the oxidized form of the redox couple by hydrogen. The calculations were performed in
collaboration with Dr. Fang Wang. The DFT calculations were performed at the B3LYP/6-31+G(d,p) level
of theory with thermal correction
73
and implicit consideration of water-solvation.
74
The free energy
correction for the standard state of 1 atm in the gas phase and 1 M upon solvation was applied, i.e.
ΔG solution = ΔG gas + 1.9 kcal·mol
-1
at 298 K. Considering the lower pK a value of benzenesulfonic acid
(pK a(sulfonic acid) = –2.8),
75
quinone sulfonic acid derivatives are expected to dissociate to sulfonates in 1 M
sulfonic acid aqueous solution. ΔG
o
was calculated based on the reduction of quinone derivatives with
H 2. The standard electrode reduction for the redox couple was deduced from E
o
= –ΔG
o
/nF, where n is
the number of protons involved in the reaction and F is the Faraday constant (96485 C mol
-1
).
2.3. Synthetic and Chemical Characterization of Materials
Many of these compounds were purchased from chemical supply companies, such as Sigma
Aldrich, Alfa Aesar, TCI Chemical, etc. However, some of the chemicals were not commercially available
and had to be synthesized in-house. The synthesis was carried out by Dr. Sankarganesh Krishnamoorthy
of Prof. G. K. Surya Prakash’s research group. Details of the synthetic procedures are available from his
thesis and in joint papers.
76
In general, the synthetic process aimed at preparation of sulfonated benzoquinone compounds.
Product yields ranged from 2% to 85%, depending on the compound and the synthetic procedure. For
many of the compounds, the resulting aqueous sulfonic acid solution was neutralized with a
stoichiometric amount of solid potassium carbonate. The resulting solid, at neutral pH, was filtered and
washed with water. Acetone was added to precipitate out the potassium sulfate, and the filtered
solution was concentrated by rotary evaporation and left overnight in a refrigerator for crystallization.
25
The resulting crystals were filtered and washed with cold 10% water/acetone and acetone and dried
overnight under high vacuum. The structure and purity of the synthesized compounds were always
confirmed by
1
H NMR (400 MHz with D 2O), and sometimes by
13
C-NMR and Electron-spray-ionization
high-resolution mass spectrometry (ESI-HRMS).
1
H and
13
C NMR spectra were recorded on Varian 500
MHz or 400 MHz NMR spectrometers.
1
H NMR chemical shifts were determined relative to the signal of
a residual protonated solvent, D 2O (δ 4.79 ppm).
13
C NMR chemical shifts were determined relative to
the
13
C signal of solvent, DMSO-d 6 (δ 39.52 ppm). ESI-HRMS analysis was conducted at the University of
Illinois- Urbana Champagne Mass Spectrometry Facility.
NMR samples for the crossover analysis and chemical composition were typically prepared by
diluting 150 microliters of electrolyte sample with 350 microliters of deuterated water (D 2O) and
1
H
NMR spectra were obtained on Varian 500 MHz NMR instruments with 128 scans. The samples were
obtained at various points during the cycling experiments.
2.4 Cell Set-Up
To build a flow cell, we adapted fuel cell test hardware (Electrochem Inc.) that used densified
graphite flow field plates and an electrode active area of 25 cm
2
. The reactants were circulated using
centrifugal pumps (March Pump Model #: BC-2CP-MD 12 VOLT DC) at flow rates in the range of 0.5-1.0
liter min
-1
(Figure 10).
Four different flow-field designs were tested to determine whether the “flow-by” or “flow-
through” characteristics yielded better overall cell performance. The “flow-through” configuration refers
to a fluid path that includes flow which is perpendicular and parallel to the plane of the electrode; thus,
the entire volume of the electrode has the opportunity to encounter the flow of solution. In a “flow-by”
design, the fluid flow is largely parallel to the plane of the electrode, and thus the flow resistance is
significantly lowered compared to the “flow-through” configuration. The relative utilization of the area
26
of the electrodes in these two configurations is dependent on the flow rate and the pressure drop
between the inlet and the outlet.
Figure 10: Picture of redox flow cell set-up
For some experiments, especially the initial ones, we prepared the membrane electrode
assemblies (MEAs) needed for the cell using procedures that were developed in-house, similar to those
27
used in direct methanol fuel cells.
77
Specifically, two sheets of carbon paper (Toray
030-non-teflonized,
each with an area of 25 cm
2
) were coated with an ink made from 0.1 g of carbon black (Vulcan XC-72
)
and 0.3 g of Nafion
®
with 0.5 ml of water and 0.5 ml of isopropanol. The equivalent weight of the Nafion
®
ionomer was 1100 g/mole of protons. The coated electrodes were then hot pressed on to a Nafion
®
117
or 212 membrane to form the MEA. This was then placed inside the fuel cell hardware with varying flow
field configurations. Electrodes used in these studies were prepared in a collaborative effort with
Advaith Murali.
In several of the experiments, graphite felt (SIGRACELL GFD 4.6 from SGL Group) was used as
electrodes, either as-received or after surface modification through acid treatment, heat treatment,
addition of Nafion®, or addition of carbon nanotubes. This type of felt is widely used in vanadium-redox
flow batteries.
78
To increase their wettability, the felts were immersed in a dilute solution of 10%
Nafion
®
and then dried at 140
o
C for an hour. The felts were then boiled in water for 2 hours to remove
any excess or unbound Nafion
®
and to hydrate the Nafion
®
that was retained on the felt. In another
modification, the graphite felt was coated with varying amounts of carbon black.
All full cell experiments were carried out with solutions of redox couples dissolved in 1 M sulfuric
acid at 23
o
C. The concentration and composition of the redox couples varied with the experiment. Two
glass containers served as reservoirs for the solutions of the redox couples (Figure 10). An argon flow
was maintained at all times above these solutions to avoid reaction of the reduced form of the redox
couples with oxygen. We have found that the exclusion of oxygen is critical to maintaining stable cell
capacity. The current-voltage characteristics of the cells were measured at various states of charge.
Charge/discharge studies were carried out under constant current conditions either using a potentiostat
(Versastat 300) or a battery cycler (Maccor 4200). Typical currents ranged from 0.5 A to 5 A for an
electrode area of 25 cm
2
.
28
29
Chapter 3
Kinetics of Electrochemical Reactions of Hydroquinones
3.1 Electrochemical Characterization of Reactions of Hydroquinones
Hydroquinones have the highest reduction potentials out of all quinone derivatives, consistent
with the energy of the HOMO and LUMO levels in these molecules.
53, 79-80
Therefore, I have undertaken
to characterize 23 benzoquinones in order to understand the effect of substituents on electrochemical
kinetics for use at the positive electrode in a redox flow battery. All the relevant electrochemical data,
collected for the various materials, is listed in Tables 2 and 3. I found two important factors that
determined the effectiveness of a hydroquinone molecule in a redox flow battery, besides the standard
reduction potential and solubility in water:
1) Their ability to resist chemical transformations, both by the Michael Reaction and by proto-
desulfonation
2) The rate of electron transfer, and its dependence on the intramolecular hydrogen bonding
and specific interactions with the electrode surface
Our initial screening studies aimed at determining the hydroquinone’s baseline characteristics
for achieving the required cell voltage and the current density for a redox flow battery. However, when
we subjected these molecules to charge/discharge cycling, as in practical application, we learned more
about their reactivity that led us to design the optimal benzoquinone molecule.
30
Table 2: Cyclic Voltammetry or Linear Sweep Voltammetry of Various Hydroquinones. All voltages are
against a MSE Reference Electrode (E° = +0.650 V), all quinone concentrations are 1 mM in 1 M sulfuric acid.
The working electrode was glassy carbon and the counter was a platinum wire. The scan rate was 50
mV/sec, and the rotation rate was 1500 rpm.
Name
(purchased or
synthesized in-
house by
Krishnamoorthy)
Structure CV/RDE data
Hydroquinone
(Sigma Aldrich)
2-sulfo,p-
benzoquinol
(Sigma Aldrich)
31
4,5-dihydroxy
benzene-1,3-
disulfonic acid
(Alfa Aesar)
2,5-dichloro
hydroquinone
(Sigma Aldrich)
3,5-dimethyl p-
benzoquinol (TCI
America)
HO
3
S SO
3
H
OH
OH
OH
OH
C H
3
CH
3
32
3,5-dimethyl, 6-
sulfo, p-
benzoquinol
(synthesized)
3-hydroxy, 4,6-
disulfo, o-
benzoquinol
(synthesized)
3-hydroxy, 4-
sulfo, 6-
carboxylic, o-
benzoquinol
(synthesized)
SO
3
H
OH
OH
C H
3
CH
3
33
3,5-disulfo, 4,6-
dimethyl, o-
benzoquinol
(synthesized)
3-sulfo, 4,6-
dimethyl, o-
benzoquinol
(synthesized)
OH
OH
SO
3
H
CH
3
C H
3
3-hydroxy, 4-
carboxylic, o-
benzoquinol
(Sigma Aldrich)
34
3-hydroxy, 5-
carboxylic, o-
benzoquinol
(Alfa Aesar)
3-sulfo, 5-
carboxylic, o-
benzoquinol
(synthesized)
3,5- methoxy, p-
benzoquinone
(Sigma Aldrich)
35
3,6-dimethyl p-
benzoquinol
(Sigma Aldrich)
2,5-di-tertbutyl
hydroquinone
(Sigma Aldrich)
2,6-dihydroxy,
p-benzoquinol
(Sigma Aldrich)
36
2,6-dihydroxy,
3,5-disulfo, p-
benzoquinol
(synthesized)
2,3,5,6-
carboxylic acid,
p-benzoquinone
(synthesized)
2,3,5-
trimethyl,6-
sulfo, p-
benzoquinol
(synthesized)
OH
OH
C H
3
CH
3
C H
3
SO
3
H
OH
OH
COOH
COOH HOOC
HOOC
37
Dihydroxy
benzoic acid
(Sigma Aldrich)
Dopamine
hydrochloride
(bought)
Table 3: Electrochemical Properties of Various Hydroquinones
Name Structure Molecular
Weight
(g/mol)
Redox
Potential,
E 1/2 vs NHE
Diffusion
Coefficient,
D 0, (cm
2
/sec)
Rate
Constant,
k 0, (cm/sec)
Exchange
Current
Density, I ex,
(A/cm
2
)
Hydroquinone
110 0.70 5.03E-6 2.63E-3 5.09E-5
2-sulfo,p-
benzoquinol
228 0.80 4.28E-6 5.52E-4 1.10E-5
4,5-dihydroxy
benzene-1,3-
disulfonic acid
332 1.10 3.80E-6 1.55E-4 3.00E-6
HO
3
S SO
3
H
OH
OH
38
2,5-dichloro
hydroquinone
179 0.75 4.57E-7 1.26E-4 2.42E-6
3,5-dimethyl
p-benzoquinol
138 0.61 6.71E-6 8.56E-4 2.47E-5
3-hydroxy,
4,6-disulfo, o-
benzoquinol
286 1.00 7.05E-6 4.74E-4 1.37E-5
2,6-dihydroxy,
3,5-disulfo, p-
benzoquinol
304 No clear
peaks;
broad peak
between
0.95-1.3 V
3,5-dimethyl,
6-sulfo, p-
benzoquinol
219 0.82 4.12E-6 1.80E-4 3.90E-6
4,6-dimethyl,
3-sulfo, o-
benzoquinol
OH
OH
SO
3
H
CH
3
C H
3
219 0.90 5.02E-6 9.70E-4 2.80E-5
2,3,5-
trimethyl,6-
sulfo, p-
benzoquinol
OH
OH
C H
3
CH
3
C H
3
SO
3
H
254 0.78 3.91E-6 1.04E-4
3.31E-6
3,5-disulfo,
4,6-dimethyl,
o-benzoquinol
362 1.10 2.49E-6 5.02E-4 1.60E-5
3-hydroxy, 4-
sulfo, 6-
carboxylic, o-
benzoquinol
250 0.85; 1.30 3.66E-6
1.46E-6
2.11E-5
1.04E-3
6.1E-7
3.02E-5
2,3,5,6-
carboxylic
acid, p-
benzoquinone
286 1.05 1.49E-6 8.96E-5 2.59E-6
3-hydroxy, 4-
carboxylic, o-
benzoquinol
170 1.15 5.89E-6 1.27E-3 3.70E-5
OH
OH
C H
3
CH
3
SO
3
H
OH
OH
C H
3
CH
3
OH
OH
COOH
COOH HOOC
HOOC
39
3-hydroxy, 5-
carboxylic, o-
benzoquinol
170 0.9; 1.25 2.98E-6
3.42E-7
6.02E-3;
9.71E-4
1.74E-4;
2.81E-5
3-sulfo, 5-
carboxylic, o-
benzoquinol
234 1.3 5.76E-6 6.98E-3 2.02E-4
3,5- methoxy,
p-
benzoquinone
168 0.41 4.08E-6 1.68E-3 4.85E-5
3,6-dimethyl
p-benzoquinol
138 0.65 3.64E-6 2.27E-3 6.59E-5
2,5-di-
tertbutyl
hydroquinone
222 0.60 Low
solubility,
insignificant
results
2,6-dihydroxy,
p-benzoquinol
142 0.9 2.39E-6 6.15E-5 1.78E-6
Dihydroxy
benzoic acid
154 0.82 5.93E-6 3.35E-5 9.7E-7
Dopamine
hydrochloride
189 0.83 5.70E-6
1.13E-4 3.28E-6
Quinone-based redox systems have been extensively reported in the literature
81-83
and it is well
known that these systems undergo a proton-coupled electron transfer. The rate constants for charge
transfer were generally high – at least an order of magnitude higher than those observed for the
vanadium redox couples.
61
Additionally, the value of the transfer coefficients being close to 0.5 and the
high values of rate constants suggest an “outer-sphere” process.
84-88
40
Importantly, the benzoquinone molecule must have the desired standard reduction potential.
The half-wave potential values (Table 3) are consistent with the values reported in the literature for the
various compounds tested.
33-35, 37, 39, 89, 90
In general, the addition of aromatic rings lower the standard
reduction potential. Consequently, the anthraquinones are more suitable as negative electrode
materials (as discussed in the next section), and hydroquinones were the focus of positive electrode
material studies. The addition of sulfonic acid groups to the benzoquinone ring increased the standard
reduction potential, which is consistent with the lowering of molecular orbital energies by electron-
withdrawing groups. This raising of the standard reduction potential was also seen with other electron-
withdrawing groups, such as carboxylic acid groups.
47, 91
The opposite effect was observed with electron-
donating groups, such as methyl and hydroxy groups, were added to the ring.
Not only were these effects seen experimentally, but they were confirmed by theoretical
calculations, performed by Dr. Sankarganesh Krishnamoorthy, from Prof. Surya Prakash’s group. As
more methyl groups are added and sulfonic acid groups are removed the HOMO and LUMO levels
increase in energy. As the opposite happens, the HOMO and LUMO levels decrease in energy (Figure
11). These experimentally calculated values are consistent with the values of E 1/2 from experiments
(Table 3).
41
Figure 11: Theoretical calculations for the various hydroquinone molecules under study done by Dr.
Sankarganesh Krishnamoorthy
The values of diffusion coefficients for quinones were approximately an order of magnitude
smaller in aqueous solutions than in non-aqueous solvents such as acetonitrile.
92
In aqueous solutions,
the observed extent of decrease in the values of diffusion coefficients with increase in molecular mass is
about 6x10
-9
cm
2
/sec per unit of molecular mass. This coefficient is an order of magnitude lower than
that observed in acetonitrile.
103
Thus, besides the effect of molecular mass, the molecular diameters
resulting from the solvation and the interaction of ionic groups with water through hydrogen bonding
have a significant effect on the diffusion coefficient values in water solutions.
It can be seen in Table 3 that most, if not all, of these benzoquinones have rate constants on the
order of 10
-4
cm/sec. While the rate constants for the various compounds were at least one order of
magnitude greater than that of those in the vanadium system, the diffusion coefficients were
comparable to that of vanadium on a glassy carbon surface (10
-6
cm
2
/sec),
62, 72, 93
making the quinone
42
redox couples very attractive from the standpoint of electrode kinetics compared to the vanadium redox
flow battery system. Additionally, rate constants are within the range of values found widely in the
literature for quinones.
84, 94
However, it was observed that benzoquinones with a sulfonic acid group in
an ortho-position to a hydroxyl group allows for intramolecular hydrogen bonds to form, thus stabilizing
the energy level of the transition state. This higher energy transition state ultimately reduces the rate
constant for the kinetics of proton-coupled electron transfer in the molecule.
Figure 12: Calculations of intramolecular hydrogen bonding on a benzoquinone ring with one or two
sulfonic acid group substituents
As sulfonic acid groups are added to the ring, the intra-molecular hydrogen bonding interactions
in the quinone molecules bearing sulfonic acid groups increase. This intra-molecular hydrogen bonding
plays a critical role in the rate limiting step of proton-coupled electron transfer,
95
due to the increased
stability of the transition state of the compound. Therefore, an increased activation energy is required
for concerted proton and electron transfer in order to break the intra-molecular hydrogen bond that
forms in the transition state. This stability makes the resident hydrogen atom and the incoming proton
compete in order to interact with the carbonyl oxygen. According to our calculations, hydroquinone
sulfonic acid preferentially adopts a conformation, allowing the formation of intra-molecular hydrogen
43
bonding which leads to a stabilization energy of 1.6 kcal mol
-1
. Similarly, intra-molecular hydrogen
bonding provides extra stabilization of other hydroquinone sulfonic acid derivatives (Figure 12). Thus,
the intra-molecular hydrogen bonding could explain the lowering of the rate constants observed with
the addition of sulfonic acid groups.
3.2 Sulfonic Acid Substituted Anthraquinones
I have also studied compounds that could potentially be used for the negative side of an organic
redox flow battery. In particular, I have studied a variety of sulfonic acid substituted anthraquinones.
The addition of a ring to a benzoquinone compound pushes the reduction potential of these molecules
to more negative values because of the electron-donating nature of the ring addition. Therefore,
substituted anthraquinones would be ideal for use as a negative side material.
96
The large focus of the anthraquinone study was on those molecules that had sulfonic acid
groups attached to the ring.
97
Sulfonic acid groups were essential to render anthraquinone soluble. Once
the sulfonic acid groups were added to this compound, the solubility greatly increased, especially once
these compounds were transformed into their acidic, as opposed to their salt, form (discussed in more
detail in Chapter 5). The addition of a second sulfonic acid group greatly increases the solubility of the
anthraquinone molecule; therefore, we determined that di-sulfonic acid anthraquinone molecules
(AQDS) should be used for the negative side of ORBAT.
However, di-sulfonic acid substituted anthraquinones are difficult to synthesize in pure form
because of the possibility of forming various isomers.
98
When the second sulfonic acid group is added on
to the mono-substituted anthraquinone, selectivity of placement is very low.
99
In fact, when purchasing
our main anthraquinone compound (2,7-disulfonic acid anthraquinone), it was always sold as a mixture,
with between 10 and 20% being made up of other isomers. It is difficult and expensive to separate out
these different isomers, hence the use of an isomer mixture.
44
Several disulfonic acid anthraquinones were purchased and tested to determine their
electrochemical properties (Table 4, Figure 13). There was a significant lowering of exchange current
density when the sulfonic acid groups were placed in the 1, 4, 5, or 8 positions (those positions which
could undergo hydrogen bonding with the carbonyl oxygen),
37
which was also shown to cause a similar
lowering on the benzoquinone derivatives. We observe that intramolecular hydrogen bonding would
lead to slower charge transfer kinetics.
Table 4: Electrochemical Properties of Di-sulfonic Acid Substituted Anthraquinones
Name of redox couple E 1/2 vs.
MSE, Volt
Rate Constant
(cm/sec)
Diffusion Coefficient
(cm²/sec)
Solubility
Anthraquinone Completely
Insoluble
Insoluble
Anthraquinone sulfonic acid -0.52 2.25E-4 3.71E-6 0.2 M
2,6 Anthraquinone
disulfonic acid
-0.60 1.52E-4 3.04E-6 ~0.5 M
2,7 Anthraquinone
disulfonic acid
-0.57 1.21E-4 2.82E-6 ~0.5 M
1,8 Anthraquinone
disulfonic acid
-0.70 1.01E-4 3.28E-6 ~0.5 M
1,5 Anthraquinone
disulfonic acid
Not reversible
~0.5 M
However, it was observed that all of the tested di-sulfonic acid anthraquinone compounds are
relatively reversible, as their rate constants are on the same order of magnitude as some of the most
reversible benzoquinones. These rate constants have been shown to be higher than those of the
vanadium redox couples that are employed in the more-advanced vanadium redox flow batteries.
100
The
tested compounds also have similar reduction potentials, though there is some variation especially with
those with sulfonic acid substituents in the 1, 4, 5, and 8 positions. Not only do these substituent
45
positions lower the exchange current densities and rate constants, but they have also been shown to
lower the reduction potentials of these molecules.
37
Figure 13: Cyclic voltammetry (a) and rotating disk electrode linear sweep voltammetry (b) on a glassy
carbon disk of various anthraquinone derivatives. All scans were done at 50 mV/sec and at a rotation
rate of 1500 rpm. Redox couples were at a concentration of 1 mM dissolved in 1 M sulfuric acid. A MSE
reference electrode (E° = +0.650 V) and a platinum counter electrode were used.
46
However, even if an AQDS isomer mixture is employed in the redox flow battery, it should not
affect the overall performance of the cell, as the kinetics and the electrode potentials are similar for all
compounds. Additionally, the placement of the sulfonic acid groups did not appear to affect the
solubility of the anthraquinone isomer – it appears that the addition of the sulfonic acid group affects
the solubility to a much greater degree than does its placement. This distinction is likely due to the fact
that the solvation free energy is determined by the type and number of substituent groups, as opposed
to the placement.
53
3.3 Selection of a Benzoquinone and Anthraquinone Molecule for Initial Cycling
4,5-dihydroxybenzene-1,3-disulfonic acid (BQDS) was down-selected as the best positive side
material for the initial studies of ORBAT. BQDS has high solubility (around 2 M), a high reduction
potential (0.45 V vs MSE), and a sufficiently high rate constant (1.55E-4 cm/sec) and diffusion coefficient
(3.80E-6 cm
2
/sec). For the negative side, 9,10-anthraquinone 2,6-disulfonic acid (AQDS) was selected
due to its low reduction potential (-0.58 V vs MSE) and a sufficiently high rate constant (1.52E-4 cm/sec)
and diffusion coefficient (4.04E-6 cm
2
/sec). Additionally, both compounds were readily available for
commercial purchase. These compounds were the baseline materials for many of the experiments
performed on ORBAT (see Chapter 5 for more information on full cell cycling results).
3.3.1 Initial Cycling Results and Discussion
One of the first experimental discoveries was the fact that BQDS underwent a chemical
transformation during cycling (it was later determined to be the Michael reaction). The chemical
transformation was confirmed by taking samples from the positive side of the cell drawn at different
stages of the extended cycling at 100 mA/cm
2
and analyzing them by NMR (Figure 14). We determined
that BQDS underwent a slow chemical transformation in addition to the expected electro-oxidation and
electro-reduction during charge and discharge. The two doublets at 7.34 ppm (J = 2.2 Hz, 1H) and 7.10
ppm (J =2.2 Hz, 1H) in the proton NMR at the start of the cycling corresponded to 4,5-
47
dihydroxybenzene-1,3-disulfonic acid (BQDS), with protons on the aromatic ring present in two distinct
environments. After the first charge, the
1
H NMR spectrum showed a new “aromatic” proton signal at
7.14 ppm (singlet) that was attributed to 1,2,4-trihydroxybenzene-3,5-sulfonic acid. After extended
cycling (120 cycles), the ratio of 1,2,4-trihydroxybenzene-3,5-sulfonic acid to BQDS had increased. After
400 cycles, all the “aromatic” proton signals of BQDS and other reaction products disappeared,
indicating the occurrence of yet another transformation that produced compounds with no peaks on the
proton NMR spectrum. The lack of NMR-active “aromatic” protons on the aromatic ring suggested the
formation of 1,2,4,6- tetrahydroxybenzene-3,5-sulfonic acid.
Figure 14: 1H NMR at 400 or 500 MHz of BQDS samples taken after various numbers of cycles. Solutions
were diluted with deuterated water (D2O).
Although the samples for NMR were obtained only after 120 cycles and 400 cycles, the
electrochemical data (Figure 16) indicated that changes were likely complete even in the earlier stages
of cycling. The changes to the charge-discharge curves and the NMR spectra during cycling were
consistent with the stepwise transformation of 4,5-dihydroxybenzene-1,3-disulfonic acid to 2,4,5-
trihydroxybenzene-1,3-disulfonic acid and ultimately to 1,2,4,6-tetrahydroxybenzene-3,5-disulfonic acid
48
(Figure 15). These transformations involve electrochemical oxidation of the hydroxyl group to the
quinone, followed by the addition of water to form reduced products. The addition of nucleophiles such
as water to the α,β-unsaturated carbonyl compounds in a 1,4-fashion is called the Michael reaction. In
our study, the nucleophilic addition of water is accompanied by re-aromatization and exchange of the
proton.
101
The Michael reaction necessitated additional charge input for the re-oxidation of the product
during charging.
Figure 15: a) Schematic of the transformation of BQDS during charging; b) Schematic of the Michael
reaction showing the nucleophilic addition of water followed by re-aromatization and proton exchange.
49
Figure 16: Rotating disk electrode linear sweep voltammetry at glassy carbon disk (a) AQDS and (b)
BQDS at various states of charge during cycling. The %SOC values refer to the amount of charge with
reference to the theoretical discharge capacity based on BQDS. The theoretical capacity is based on the
total charge required to transform BQDS to its “cyclable” form. This value is tantamount to 6 Faradays to
charge every mole of BQDS. All scans were done at 50 mV/sec and at a rotation rate of 1500 rpm. Redox
couples were at a concentration of 1 mM dissolved in 1 M sulfuric acid. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used.
50
The limiting current at a rotating disk electrode during a linear sweep voltammetry experiment
was used to analyze the progress of conversion of the materials from the discharged to the charged
state. The value of the limiting current indicated the concentration of the reduced and oxidized species
during charge and discharge. When AQDS was in its fully discharged form, there was a reduction current
indicating presence of the oxidized form (Figure 16a). Similarly, when the solution on the negative side
was in the fully charged state, there was no longer any reduction current ― only an oxidation current ―
indicating the full conversion of the reactants.
However, the linear sweep voltammetry studies on BQDS indicated changes to the composition
of the solution during the charging process. The persistent presence of the reduced state of the positive
side material (Figure 16b) was consistent with the chemical transformations occurring during charging.
The electrochemical studies confirmed that BQDS undergoes three redox transformations during
charging that include two Michael addition steps followed by re-aromatization with intermediates in the
“reduced” state (Figure 15). Following these transformations, the 1,2,4,6-tetrahydroxybenzene-3,5-
disulfonic acid became the positive electrode material that was repeatedly cycled against AQDS at the
negative electrode.
3.3.2 The Issues Arising from the Michael Reaction of Water with BQDS
In this cell with BQDS as the positive electrolyte and AQDS as the negative electrolyte, undesired
chemical transformations of BQDS occur in addition to the necessary redox reactions. While AQDS is
stable and does not undergo such chemical transformations, BQDS undergoes the Michael reaction with
water. There are several detrimental effects of the Michael reaction, which include:
1) Three equivalents of the negative side material (AQDS) are needed to convert every equivalent
BQDS into a stable and rechargeable positive side material. However, only one equivalent of the
AQDS is actually used during discharge and charge. This need for excess AQDS is shown in Figure
17b – the charge acceptance in the first cycle was about three times that of the discharge
51
capacity output following this charge. Therefore, the coulombic efficiency in the first cycle was
about 33%. However, in subsequent cycles the coulombic efficiency gradually converged to
100% (Figure 17a). After BQDS had reached its fully charged form, 66% of the AQDS remained
unused during the subsequent cycles. We found that stable capacity could be reached only after
about five cycles of charge and discharge (Figure 17a). However, this excess requirement of
AQDS adds to the overall cost of the system, as two-thirds of the negative side material remains
unutilized beyond the first charge.
2) The products formed after the Michael reactions were expected to reduce the cell voltage, as
additional hydroxyl substituents are added to BQDS. A reduction of about 150 mV in the charge
voltage is noted after the first few cycles (Figure 17). This reduction in the value of the electrode
potential is slightly lower than the 100 mV/ hydroxyl group expected from the study of
substituent effects.
47
3) Incomplete utilization of AQDS results in a reduced concentration of the discharged form of
AQDS, which reduces the maximum current density during subsequent charging. Placing a fully-
transformed BQDS (1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid) molecule in the cell at the
start of cycling would avoid the need for excess AQDS.
Therefore, these challenges lead to the need to design our own molecule, as opposed to one which
is simply available for purchase.
52
Figure 17: 25 cm
2
redox flow cell, 0.2 M BQDS and AQDS, charged and discharged at 2A (80 mA/cm
2
)
with graphite felt electrodes. a) Coulombic efficiency for the first 20 cycles when cycled between 1 V and
0 V (100 % depth-of-discharge); b) Charge curves over the first four cycles.
53
3.4 The Design of a Michael-Reaction Resistant Benzoquinone
To overcome the disadvantages resulting from the Michael reaction, a molecule that is resistant to
this reaction is needed. Based on the Michael reaction transformation, the version of BQDS that was
made after two complete Michael transformations, 1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid
(THBDS), was synthesized. Figure 18 shows the CV and RDE curves of THBDS. This CV and RDE are vastly
different from the CV and RDE of the starting BQDS compound. The CV does not exhibit any sharp or
well-defined peaks – there is only one large, broad peak over approximately a 400 mV range. This peak
broadness could be indicative of different oxidation reactions all occurring at potentials very near one
another. Additionally, the RDE is unable to reach a limiting current density, and the current density
continually decreases as the number of sweeps increases. These characteristics could also indicate that a
molecule is undergoing several simultaneous reactions, as well as sluggish electron transfer kinetics. This
performance ultimately indicates that this type of compound has:
1) intra-molecular hydrogen bonding occurring with the addition of the sulfonic acid groups.
Therefore, when the -para quinone tries to undergo reduction, it is unable to due to the
interaction of the quinone carbonyl group with the α-OH group.
2) competing reactions of the -ortho and -para hydroquinone reactions, and that both are
occurring at potentials close to one another.
Because both of these factors are occuring on this molecule, the kinetics are limited and the CV
shape is asymmetrical.
54
Figure 18: CV and RDE of 1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid on a glassy carbon working
electrode. A MSE reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The
scan rate was 50 mV/sec. The quinone had a concentration of 1 mM in a 1 M sulfuric acid solution.
However, when CV and RDE experiments were performed on the cycled BQDS material, the
similarities become apparent. The cycled BQDS no longer has any sharp peaks, at any potential, but
instead several smaller, broader peaks (Figure 19a). This may be indicative of the hydroxylated
hydroquinone molecules that have formed due to the Michael reaction that has already taken place; the
CV reveals that there are several different molecules in this cycled solution that are still being
transformed, as well as undergoing cycling. These peaks reveal that even after 30 cycles the BQDS has
still not fully converted to the final THBDS form. The presence of several molecules in a single solution
can be confirmed when compared to the synthesized THBDS molecules – this CV shows no peaks, but
simply one large, broad peak. The cycled BQDS eventually has the same CV as the synthesized THBDS
(Figure 19).
55
Figure 19: a) Cyclic voltammograms of 1,2,4,6-tetrahydroxybenzene -3,5-disulfonic acid (THBDS, red),
and cycled BQDS (after 30 cycles – blue) and (after 100 cycles – yellow), and b) linear sweeps of cycled
BQDS material after 30 cycles (blue), and THBDS (red). A glassy carbon rotating disk working electrode, a
mercury/mercurous sulfate reference electrode and a platinum counter electrode were used. The scan
rate was 50 mV/sec and the rotation rate was 1500 rpm. Redox materials had a concentration of 1 mM
in 1 M sulfuric acid.
56
1,2,4,6-tetrahydroxybenzene was also synthesized, before being sulfonated. This compound
demonstrated very reversible behavior, especially when compared to the disulfonated version of this
molecule (Figure 20). This differing CV shape means that the sulfonic acid group has a large hydrogen
bonding effect on the carbonyl oxygen group compared to the hydroxyl groups. It also demonstrates
that the -para quinone reaction is favored over the -ortho quinone reaction. Since there are no longer
any sulfonic acid groups in the α-positions to one of the paraquinone groups, that electron transfer is no
longer hindered.
Figure 20: For both compounds, a MSE reference electrode (E° = +0.650 V) and a platinum counter
electrode were used. The scan rate was 50 mV/sec. The quinone had a concentration of 1 mM in a 1 M
sulfuric acid solution. Blue: CV of 1,2,4,6-tetrahydroxy benzene; Green: CV of 1,2,4,6-tetrahydoxy, 3,5-
benzenedisulfonic acid.
The CVs indicated the need to design a molecule that was not based on BQDS at all – but instead
based on other substituent groups, such as methyl, hydroxy, or carboxylic acid groups, with appropriate
placement. The following sections discuss several benzoquinones that were synthesized and studied to
determine their electrochemical properties, including studies done on different electrode surfaces.
57
3.4.1 Surface Effects on CV Performance
It was also determined that the electrode surface plays a crucial role in the performance of those
hydroquinones that had non-planar substituent groups. It is important for the redox molecule to be able
to approach the electrode surface in such a way that electron transfer to the quinone oxygen groups is
facile. For those molecules that were planar – those with methyl, hydroxy, or carbonyl substituent
groups – the change in the roughness of the surface did not alter the shape of the CV or the reduction
potential. Even though methyl is not planar, since its three hydrogen atoms are in a tetrahedral
geometry around the carbon atom, the relatively small hydrogen atoms do not affect the overall
planarity of the molecule to the same extent that the bulky tetrahedral sulfonic acid group does.
Therefore, the methyl groups can be considered as maintaining the planarity of the hydroquinone
molecule. However, when bulky sulfonic acid groups were added to the quinone ring, the surface greatly
changed the shape of the cyclic voltammogram.
For example, the shape of the CV of BQDS on a graphite rod is shown in Figure 21. The CV shape
indicated that BQDS exhibits the characteristics of a reversible redox couple, which is in direct contrast
to the CV that is obtained when the working electrode is glassy carbon. It was also determined that an
electron transfer process is occurring – as opposed to an adsorption process – because the limiting
current density increased linearly with the square root of scan rate (Figure 21 – inset). The linearity of
current with the square root of scan rate confirms that the electron transfer process was a solution
species interaction and not a surface species interaction; meaning that there was no adsorption on the
electrode surface, but instead the CV demonstrated fast and reversible kinetics of the BQDS molecule
from solution. This observation is consistent with the CV of the tetra-hydroxy compound.
Once the CV was recorded on the surface of a graphite rod, the large separation between the
oxidation and reduction peaks (about 400 mV) for BQDS is no longer present. The peaks have only 60
mV of separation on the graphite rod surface – this number generally indicates a reversible two-
58
electron, two-proton redox reaction, based on the Nernst equation. Glassy carbon is known to have a
flat surface with low electron density,
102
while graphite has a much higher rougher surface rich in
electron density; this difference can play a crucial role in electron transfer reactions.
103
Recent studies
have shown that the basal plane of highly oriented pyrolytic graphite (HOPG) has substantially high
activity for facile electron transfer reactions.
104-106
In fact, it was observed that electron transfer kinetics
at a freshly-cleaved, highly-oriented pyrolytic graphite electrode are at least as fast as on platinum for
several different metal complexes, including Fe(CN)6
4−/3−
.
107
Figure 21: 1 mM BQDS in 1 M sulfuric acid, with a graphite rod working electrode. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was as indicated, in
mV/sec. Inset shows the peak height variance with the square root of scan rate
Additionally, BQDS was tested on a graphite micro-electrode (Figure 22). This electrode
demonstrates the reversible behavior of BQDS. The inherent characteristics of micro-electrodes are that
they will reach a diffusion limited current without stirring (due to their very small radius) and the ability
59
to access a growing radial diffusion field. This behavior of the micro-electrode is shown in Figure 22, and
confirms the fast kinetics of BQDS.
Figure 22: 1 mM BQDS in 1 M sulfuric acid, with a graphite fiber working micro-electrode. A MSE
reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec.
When BQDS is allowed to interact with a rougher electrode surface, the CV shows sharper
peaks, indicating a faster rate of electron transfer. Thus, we hypothesized that the larger, bulkier, and
more electronegative sulfonic acid groups on the benzene ring affected how the BQDS molecule
interacted with the electrode surface. Once those sulfonic acid groups were added to the quinone, the
molecule was no longer planar, making it difficult to closely interact with the flat electrode surfaces.
Thus, the kinetics changed dramatically on flat surfaces. This makes it difficult to use the CVs on glassy
carbon to predict the performance of the molecule in a full redox flow cell, as the flow cells employ
graphitic carbon paper or felt electrodes.
The cyclic voltammetric data on BQDS is also consistent with the performance of the
tetrahydroxy benzene molecule. Since tetrahydroxy benzene has a planar transition state, the molecule
60
can interact well with the flat surface of glassy carbon. This means that concerted and simultaneous
two-proton and two-electron transfer is easy to achieve. However, once sulfonic acid substituent groups
are added to the ring it is no longer planar; the electron transfer is inhibited on the flat glassy carbon
surface, which is consistent with the poorly defined peaks in the CV.
Figure 23: Cyclic voltammograms of 1 mM benzoquinone di-tertbutyl in 1M sulfuric acid on a graphite
rod working electrode, at increasing scan rates (as indicated in the figure, in mV). A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used.
In order to confirm the hypothesis that the electrode surface structure plays a major role in
determining the observed CV, we undertook studies on several benzoquinone variations on different
electrode surfaces to better understand those effects. We purchased an extremely bulky compound –
benzoquinone di-tertbutyl – and tested its activity on a graphite rod. The compound showed extremely
reversible behavior (Figure 23), indicating that this compound was able to exhibit reversible behavior so
long as its contact with the electrode surface was sufficient for rapid electron transfer. Additionally, the
linearity of current with the square root of scan rate confirmed that the electron transfer process was a
61
solution species interaction and not a surface species interaction. In contrast, the corresponding CV on a
glassy carbon electrode is seen in Table 2 and such reversible behavior is not observed.
For a final investigation of the difference between using a graphitic surface and a glassy carbon
surface, we performed experiments with a di-methoxy benzoquinone molecule (Figure 24). The more
reversible behavior was apparent on the graphite surface, as opposed to that on the glassy carbon
surface. Even though methoxy groups are not nearly as bulky as sulfonic acid groups, these molecules
show better reversibility on a graphite surface. The difference in current density between the two
surfaces is due to the difference in electrode surface area.
Thus, we were able to confirm that bulkier – or non-planar – substituent groups had better
reversibility on non-planar electrode surfaces, such as a graphite rod. In contrast, molecules with planar
substituent groups had better reversibility on planar electrode surfaces, such as the glassy carbon
electrode.
Figure 24: Cyclic voltammograms of 1 mM 2,6-dimethoxy, 1,4-benzoquinone in 1 M sulfuric acid on a
graphite rod (orange) and a glassy carbon (blue) working electrode. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec.
62
3.4.2 Pyrogallol Derivatives
Pyrogallol is 1,2,3-trihydroxy benzene, a compound with three hydroxy groups all in the -ortho
position to one another on a benzene ring. Several different pyrogallols were tested, with various
substituent groups (methyl, carboxylic acid, or sulfonic acid) added to the ring. It was determined that
once another hydroxyl group was placed directly next to the ortho-quinone, the mechanism changes
from a single, concerted “two-electron, two-proton” transfer pathway, to a two “one-electron, one-
proton” transfer pathway. This transition is clearly seen in the cyclic voltammograms in Figure 25.
Additionally, the reduction process appears to be hindered (if it occurs at all) as evidenced by the low
reduction current in the CVs. Additionally, the first oxidation is not affected by the occurrence of the
second oxidation; if the second oxidation peak was not a one-electron transfer, but instead some sort of
irreversible adsorption or other transformation, it could affect the reversibility of the first oxidation
process. However, when the voltage sweep was cut off before the potential for the second oxidation
was reached, the shape of the CV was unaffected (Figure 26). This behavior confirms that two one-
electron transfers are occurring, and furthermore, that they are both irreversible.
This change observed in the electron transfer pathway is most likely due to the hydrogen
bonding effects when the quinone group is surrounded by hydroxyl groups. The strong hydrogen
bonding effect from α-hydroxy groups to the quinone oxygens has been discussed in the literature.
90, 108-
110
The electrochemical behavior of these types of quinones is more irreversible because the electron
addition involves a hydrogen bonded carbonyl group.
111-112
It has also been discussed in the literature
that the hydrogen bonding of the acidic proton on the α-hydroxy group stabilizes the intermediate anion
radical. This stabilization is what leads to the two peaks forming.
110, 113
63
Figure 25: Cyclic voltammograms of various substituted pyrogallol molecules. The quinone
concentrations were 1 mM in 1 M sulfuric acid on a glassy carbon working electrode. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec.
Figure 26: Cyclic voltammograms of molecules at different voltage cut-offs. The quinone concentration
was 1 mM in 1 M sulfuric acid on a glassy carbon working electrode. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec.
64
The change in the mechanism occurs regardless of the electrode surface employed (Figure 27).
This effect is also seen in all of the different pyrogallol compounds, regardless of the arrangement of the
substituent groups or their type. The effect may vary in strength based on the substituent group and
position; however, it is always present. The variations in CV shape are due to the differences in the
strength of the intra-molecular hydrogen bonding for the specified molecule – as the number of
potential hydrogen bonds increases, so does the presence of two distinct peaks. For example, those
pyrogallols with sulfonic acid groups have two very distinct peaks because of the hydrogen bonding that
is possible with those groups, as well as the ever-present α-OH group. However, when the sulfonic acid
groups are removed, the only possible hydrogen bonding that can occur is with the α-OH group, and
thus the presence of two peaks decreases. As the amount of hydrogen bonding decreases, the electron
transfer mechanism goes from a two one-electron transfer pathway to a one two-electron transfer
pathway.
Figure 27: 3-hydroxy, 4,6-disulfo, o-benzoquinol cyclic voltammograms on various electrode surfaces
(labeled). The quinone concentration was 1 mM in 1 M sulfuric acid. A MSE reference electrode (E° =
+0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec.
65
Further confirmation of this hypothesis is shown in Figure 28. Simply by removing the third
hydroxy group, the CV shifts from two, irreversible oxidation peaks to one, reversible oxidation-
reduction peak. These CVs show that the third hydroxy group plays an important role in the
intramolecular hydrogen bonding effects on pyrogallol compounds (Figure 29).
Figure 28: Cyclic voltammograms of 3-hydroxy, 4,6-disulfo, o-benzoquinol (red) and 3-sulfo, 5-carboxyl
acid, o-benzoquinol (blue) on a glassy carbon working electrode. The quinone concentration was 1 mM in
1 M sulfuric acid. A MSE reference electrode (E° = +0.650) and a platinum counter electrode were used.
The scan rate was 50 mV/sec.
Figure 29: Intra-molecular hydrogen bonding effects after first proton-electron transfer for two different
pyrogallol molecules
66
The only apparent exception is the tetra-hydroxy compound. Even though this compound is also
derived from pyrogallol, it does not exhibit the same sort of CV shape and kinetic parameters as other
pyrogallol derivatives, because of the positioning of the fourth hydroxy group. The position of this group
turns the pyrogallol from an ortho-quinone to a para-quinone, which is the more reversible quinone.
Therefore, it is clear that it is important to design a molecule that is a paraquinone, as this sort of
configuration lends itself to more reversible behavior.
3.4.3 Methyl Substituted Benzoquinones
Finally, the last type of benzoquinone that was tried were methyl-substituted benzoquinones
(either the para- or ortho- quinone). Methyl groups were employed to fill the open positions on the ring
next to the quinone groups; there was no chance of hydrogen bonding occurring with these compounds.
These groups also ensured that the benzoquinone would stay a planar molecule. As shown in Table 5,
the methyl substituted benzoquinones have reversible kinetics and a relatively symmetrical CV shape.
Table 5: Kinetic Parameters and CVs of Methyl-Substituted Hydroquinones in Acid Media. All voltages are against a
MSE Reference Electrode (E° = +0.650 V), all quinone concentrations are 1 mM in 1 M sulfuric acid. The working
electrode was glassy carbon and the counter was a platinum wire. The scan rate was 50 mV/sec.
Name Structure CV Diffusion
Coefficient,
D 0, (cm
2
/s)
Rate
Constant,
k 0, (cm/sec)
4,6-
dimethyl o-
benzoquinol
OH
OH
SO
3
H
CH
3
C H
3
5.49E-6
3.30E-3
67
3,5-
dimethyl p-
benzoquinol
6.71E-6 8.56E-4
3,6-
dimethyl p-
benzoquinol
3.64E-6 2.27E-3
3,5-
dimethyl, 6-
sulfo, p-
benzoquinol
2.82E-6 1.3E-4
3,5-disulfo,
4,6-
dimethyl, o-
benzoquinol
2.49E-6 5.02E-4
OH
OH
C H
3
CH
3
SO
3
H
OH
OH
C H
3
CH
3
68
2,3,5-
trimethyl,6-
sulfo, p-
benzoquinol
OH
OH
C H
3
CH
3
C H
3
SO
3
H
3.91E-6 1.04E-4
To reduce the susceptibility of the molecule to the Michael addition reaction it was essential to
design it with the fewest possible open positions on the hydroquinone. Sulfonic acid groups were still
desirable substituents, as they increased the reduction potential and the solubility of the molecule in
water. To limit the number of open positions on the phenyl ring, methyl groups were used as
substituents. Though the disulfonic acid molecule was desirable, adding a second sulfonic acid group to
DHDMBS proved to be a challenge for synthesis. Furthermore, we speculated that the one open position
on the DHDMBS in the quinone form is less prone to nucleophilic attack by water compared to BQDS as
the two methyl groups, being electron-donating, lower the nucleophilicity of the open position.
3.4.3.1 Proto-desulfonation
Even though the addition of methyl groups prevented the Michael addition reaction, it also
increased the chance of proto-desulfonation to occur. Proto-desulfonation is a process by which the
substituted sulfonic acid group is removed from the ring and replaced by a proton (Figure 30). Unlike
most other electrophilic substitution reactions, sulfonation is reversible. Desulfonation is the opposite of
the sulfonation process by which the sulfonic acid was originally made. Since the reverse reaction
involves a protonation step, this forward reaction is called proto-desulfonation. The methyl group
substituents on the ring increase the likelihood of this process because of the electron donating nature.
The thermodynamic reaction curve shown in Figure 31 shows that with the addition of sulfuric acid, the
equilibrium shifts from a sulfonated aryl group to a protonated group, due to the increased presence of
69
protons that are available for proto-desulfonation to occur. The effects of this process are discussed in
more detail in Chapter 5.
Figure 30: Mechanism for Proto-desulfonation
Figure 31: Thermodynamic diagram of proto-desulfonation, based on pH. Adapted from reference 114
114
Likewise, it is harder to add a second sulfonic acid group onto the ring because of this same
process. In general, electron withdrawing groups on the phenyl ring lower reactivity towards
electrophiles – in this case, a proton – because the intermediate carbocation formed during protonation
is destabilized by the electron-withdrawing groups such as SO 3H. Calculation of HOMO energies, as
70
performed by Dr. Sankarganesh Krishnamoorthy, show that the HOMO of the di-sulfonated version of
DHDMBS is 10.6 kcal/mol lower in energy, and therefore it is harder to remove an electron from that
molecule compared to DHDMBS. In other words, it should be harder to protonate the di-sulfonated
molecule compared to the mono-sulfonated molecule. We were able to determine that the proto-
desulfonation occurred more rapidly at higher concentrations of acid (See Table 6). The study,
conducted by Archith Nirmalchandar, was done by stirring a solution of DHDMBS (1 M) with different
sulfuric acid concentrations at different temperatures. The time taken to observe the precipitated
desulfonated material was recorded. The higher concentration of acid aided proto-desulfonation
because of a higher concentration of protons. Therefore, it is necessary to work with lower acid
concentrations in order to avoid proto-desulfonation.
Table 6: Effect of Temperature and Added Acid Concentration on Proto-desulfonation Rate
Concentration
of H 2SO 4 (M)
Time for visual
observation of proto-
desulfonation at 23 °C
Time for visual
observation of proto-
desulfonation at 40 °C
Time for visual
observation of proto-
desulfonation at 50 °C
Time for visual
observation of proto-
desulfonation at 60 °C
1 - - - Overnight
2 - 17 days 10 days Overnight
3 - 7 days 5 days Overnight
4 21 days 2 days Overnight Overnight
Therefore, even though this disulfonic acid molecule is harder to make, it is also much more
stable and less susceptible to any sort of chemical degradation, both by the Michael Reaction or proto-
desulfonation. The synthesis of a di-methyl di-sulfonic acid benzoquinone needs to be undertaken.
Additionally, even though the hydroquinones with methyl groups have faster rate constants and are less
susceptible to the Michael reaction, they are more susceptible to proto-desulfonation because of the
electron-donation by the methyl groups.
3.5 Summary
It is concluded that the substituent plays a substantial role in the electrochemical characteristics
of benzoquinone redox chemistry. Substituents not only modify the solubility and the reduction
71
potential of the molecule, but they have a large effect on the kinetics of the electron transfer reaction
due to the intramolecular hydrogen bonding effects. We found that the effects of hydrogen bonding for
hydroxy or sulfonic acid groups in the both the alpha or beta positions can lower the rate constant. This
lowering is due to the transfer of the proton in the hydrogen bond that occurs when the molecule is
accepting a proton during reduction.
90
Because of these effects, it was essential to design a benzoquinone molecule with methyl
groups on the ring, so as to limit the hydrogen bonding effects and the subsequent inhibition of the
electron transfer. However, substituent groups that raise the reduction potential and solubility of the
benzoquinone molecule, i.e. sulfonic acid groups, are just as necessary. Therefore, for the positive side
redox couple for ORBAT, it is necessary to design a benzoquinone molecule that will meet the required
high reduction potential and solubility limits, while not compromising on electron transfer rate.
It is also crucial that the molecule does not undergo any chemical transformations, i.e. the
Michael reaction or proto-desulfonation. There are chemical (as well as some electrolyte engineering)
solutions that can be made to limit these detrimental chemical transformations. Molecular design and
synthesis is one method that was employed to limit these transformations, although changing the acid
concentration in the electrolyte can help limit proto-desulfonation as well.
Finally, the surface on which the electron transfer occurs is essential in obtaining accurate
performance results for the molecule. It was discovered that using a surface with a larger number of
defects and higher local electron density can enhance the electron transfer for non-planar
benzoquinone molecules, though it does not seem to have the same effect on planar derivatives of
hydroquinone.
It is also clear that sulfonic acid substituted anthraquinones are excellent candidate molecules
for the negative side of an acid-based redox flow battery. They are highly stable at acidic pH and can
72
achieve sufficient solubility in order to ensure economic viability. Anthraquinones have been cycled
against several different positive side redox couples – in many different cell configurations – and we
have yet to observe any sort of chemical degradation or transformation of this molecule.
73
74
Chapter 4
Alkaline Studies
4.1 Introduction and Advantages of an Alkaline RFB System
Even though it was found that many of the hydroquinones are good candidate materials for
positive side electrodes in acidic systems, and can easily undergo concerted proton-electron transfer,
there are several disadvantages with such a system. Those disadvantages include the chemical
transformations that they can undergo – namely, the Michael addition reaction and proto-
desulfonation. While the Michael addition reaction will actually be enhanced in alkaline media, proto-
desulfonation will not be favored and so some chemical transformations may be avoided in alkaline
media. However, there are several other advantages to switching to an alkaline system, besides limiting
proto-desulfonation:
1) the higher pH values of the electrolyte will lead to decreased corrosion of metal parts, resulting in
lower materials cost for the stack
2) potentially higher solubility values will lead to higher energy density and power density
3) potentially lower resistances due to the higher degree of ionization will lead to improved efficiencies
4) ability to move away from using more expensive Nafion® membranes to hydrocarbon membranes,
such as the Tokuyama membranes employed in some alkaline fuel cells.
Because of all of these potential advantages, it was important to investigate molecules that
could be used in alkaline redox flow batteries.
4.2 Napthoquinones
The naphthoquinones were studied only in alkaline media because their two-ring configuration
made them unlikely candidates for the acidic system – their reduction potentials would have been in
75
between the benzoquinones and anthraquinones, thus rendering them non-ideal candidates for either
the positive or negative side of the acid-based cell. However, they were tested for their utility in alkaline
media because their second ring limits the Michael addition reaction from occurring, thus making them
more stable compounds for use in an alkaline redox flow battery.
Four naphthoquinone derivatives were tested (Table 7), and these compounds showed
relatively reversible behavior in alkaline media. The only compound that did not show reversible
behavior is the naphthoquinone with two hydroxy groups as substituents – 2,3-dichloro-5,8-dihydroxy –
1,4-naphthoquinone. Once all of the hydroxy groups of the compound transformed into carbonyl groups
in the oxidized form of the compound, the solubility was greatly lowered, thus making it difficult for the
molecule to return to its original reduced form. Additionally, there was the possibility of hydrogen
bonding occurring between the two hydroxy groups once the naphthoquinone was in the reduced form,
which could limit its ability to be re-oxidized. However, for all of the other naphthoquinone derivatives,
the performance in alkaline media yielded reversible CV shapes and high solubility (~1 M). The reduction
potentials of these molecules were too negative for these molecules to be used as positive side
materials in an alkaline RFB. However, they have the potential to be used as negative side materials.
76
Table 7: Cyclic Voltammograms of Naphthoquinoe Derivatives in Alkaline Media. All voltages are against
a MMO Reference Electrode (E° = +0.140 V), all quinone concentrations are 1 mM in 1 M potassium
hydroxide. The working electrode was glassy carbon and the counter was a platinum wire. The scan rate
was 50 mV/sec.
Compound Name Structure Performance in 1M NaOH (vs. MMO)
2,3-dichloro-5,8-
dihydroxy-1,4-
naphthoquinone
2,3-dichloro-1,4-
naphthoquinone
2-amino-3-chloro-
1,4-
naphthoquinone
77
2-hydroxy-1,4-
naphthoquinone
4.3 Hydroxy Substituted Anthraquinones
Hydroxy substituted anthraquinones were also studied for their performance only in alkaline
media due to their limited solubility in acidic solutions. Like the sulfonic acid substituted
anthraquinones, the substituent position did not greatly affect the standard reduction potential of these
molecules – all of the dihydroxy substituted compounds were within 100 mV of each other (Figure 32).
These consistent reduction potential values confirm that it is the type of substituent – not the
placement – that most affects the reduction potential. When a third hydroxy group was added to the
molecule, the reduction potential does become more negative, which is consistent with what we know
about electron-withdrawing groups.
47
These compounds appear to be candidates for negative side
materials for an alkaline RFB due to their very negative reduction potentials (Figure 32, Table 8).
78
Figure 32: Linear sweeps of various hydroxy-substituted anthraquinones on a glassy carbon working
electrode (see labels). The quinone concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec, and the rotation rate was 1500 rpm.
Table 8: Structures and Bi-Functional Potentials of Hydroxy-Substituted Anthraquinone Derivatives in
Alkaline Media
Compound Structure E 1/2 (Volts) vs MMO Bi-functional Material
Potential
1,2,4-trihydroxy
anthraquinone
O
O
OH
OH
OH
-0.84 V Yes, 840 mV difference
2,6-dihydroxy
anthraquinone
O
O
OH
O H
-0.8 V No
79
1,5-dihydroxy
anthraquinone
O
O
OH
O H
-0.65 V No
1,4-dihydroxy
anthraquinone
O
O
OH
OH
-0.63 V Yes, 880 mV difference
1,2-dihydroxy
anthraquinone
O
O
OH
OH
-0.73 V Yes, 1 V difference
1,2-dihydroxy, 3-
sulfonic acid
anthraquinone
O
O
OH
OH
SO
3
H
-0.75 V Yes, 930 mV difference
As can be seen in Figure 33, some of the hydroxy anthraquinone compounds have two separate
redox reactions at very different potentials. This is due to the placement of the hydroxy substituent
groups. If there are at least two hydroxy groups on one of the rings of the anthraquinone molecule (as
seen on 1,2-dihydroxy-, 1,2-dihydroxy, 3- sulfonic acid-, 1,2,4-trihydroxy-, and 1,4-dihydroxy-
anthraquinone), those hydroxy groups have similar behavior to a hydroquinone molecule; i.e. they will
have an oxidation current at a positive electrode potential. Therefore, the carbonyl oxygens in the 9,10
positions behave like the parent anthraquinone molecule and have a very negative reduction potential,
while the two hydroxy groups on another ring behave like a hydroquinone molecule and have a very
80
positive reduction potential, which shifts depending on the other substituents on this “hydroquinone”
ring.
Figure 33: Linear sweeps of various hydroxy-substituted anthraquinones on a glassy carbon working
electrode (see labels). The quinone concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec, and the rotation rate was 1500 rpm.
These characteristics of the hydroxy anthraquinone compounds open up the potential for a bi-
functional material. We define a bi-functional redox couple as a molecule that can theoretically perform
the role of both a “positive” side material and a “negative” side material; i.e. this single molecule has
two separate redox couples that occur at very different potentials. Therefore, this kind of molecule can
be placed on both sides of a redox flow cell, since the molecule can perform both the roles of a positive
or negative redox couple. This kind of cell is called a “symmetric” cell since the same molecule will be
found on both sides. A symmetric cell would make crossover a non-issue for redox flow cells, which is a
problem which has constantly plagued vanadium redox flow cells. The difference between the two
standard reduction potentials of this single molecule will determine the cell voltage, similar to an
asymmetric cell. Therefore, it would be important for the difference between the molecule’s two
81
reduction potentials to be as large as possible. The hydroxy anthraquinone derivatives give a large
potential difference (840 - 1000 mV) which would be relatively suitable for use in a commercial redox
flow battery.
Unfortunately, none of the “hydroquinones” moieties in the hydroxy anthraquinone molecule
exhibit reversible behavior – there is a clear oxidation current, but no reduction current. This
irreversibility is attributed to the lowered solubility of the molecule once it has been transformed into its
fully oxidized form with four carbonyl oxygen groups, as opposed to its reduced form with four hydroxy
groups. But if the molecule can be substituted with sulfonic acid to make the “hydroquinone” moieties
more reversible and more soluble, this new molecule would have the potential to be a bi-functional
material. A few other groups have investigated the idea of a bi-functional molecule,
115-116
but none of
these molecules have proven to be effective, yet. Hydroxy anthraquinones could be a potential pathway
towards making a fully-functional, symmetric redox flow cell with a truly bi-functional molecule.
4.4 Quinoxalines
Quinoxaline derivatives were tested for their chemical and electrochemical properties in both
acidic and basic solutions. These kind of molecules were selected for study because of their small size,
potential for two-electron reductions, and high solubility in acidic media (7 M). The two-electron
reduction is shown below, in Figure 34.
N
N
2 e
-
+ 2H
+ N
H
N
H
Figure 34: Scheme of quinoxaline reduction
82
It was determined that in acidic media quinoxalines undergo two one-electron reactions, as
opposed to the single two-electron reaction that quinones undergo (Figure 35). Protons accompany
these electron transfers.
Figure 35: Linear sweeps of quinoxaline (yellow) and 5-methyl quinoxaline (red) on a glassy carbon
working electrode. The quinoxaline concentrations were 1 mM in 1 M sulfuric acid. A MSE reference
electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan rate was 50 mV/sec, and
the rotation rate was 1500 rpm.
This behavior is consistent with what has been reported in the literature – in acidic solutions
(under pH 4), quinoxalines go through a radical intermediate before being transformed into their fully
reduced form (Figure 36). This formation of the radical leads to dimerization through the radical that is
formed on the ring. When a radical semi-quinone molecule interacts with the starting quinoxaline
molecule, 2,2'-bisquinoxaline is formed.
117
Once the quinoxaline molecule was placed in a cell for full
scale cycling, dimerization was rapid – within a few days of cycling, the solution became so viscous that
the pumps were no longer able to provide the required level of circulation. Therefore, these molecules
were deemed unsuitable for acid-based RFBs.
83
N
N
N
H
N
H
N
H
N
pH< 4
2e
-
+ 2H
+
pH> 4
e
-
+ H
+
e
-
+ H
+
Figure 36: Scheme of quinoxaline reduction/oxidation at different pH. Adapted from reference 118
118
However, in alkaline media (with pH values above 4) the two proton and two electron transfer
happens simultaneously, thus bypassing radical formation and the subsequent polymerization. This
made quinoxalines better suited for use in an alkaline RFB. Figure 37 shows the CVs of several
quinoxaline derivatives – the peaks are all symmetrical and all of their reduction potentials are quite
negative (below 870 mV vs MMO).
As was discussed in the previous section, it is important to avoid crossover in a redox flow
battery system. Crossover leads to significantly lower cycle life, as well as eventually lower energy
densities and current densities. Therefore, it was decided that the only quinoxaline that would be
suitable for use in an alkaline RFB was the 6-carboxylate quinoxaline, as the other two derivatives will
cross over through the membrane, because they are not ionized. The 6-carboxylate quinoxaline will stay
in its anion form, thus being rejected by a cation-exchange membrane such as Nafion®.
84
Figure 37: Cyclic voltammograms of quinoxaline (red), 5-methyl quinoxaline (blue), and 6-carboxlyate
quinoxaline (green) on a glassy carbon working electrode. The quinoxaline concentrations were 1 mM in
1 M potassium hydroxide. A MMO reference electrode (E° = +0.140 V) and a platinum counter electrode
were used. The scan rate was 50 mV/sec.
4.5 Riboflavin
Not only are quinoxaline and naphthoquinone molecules good options for the negative side of
an alkaline RFB, but derivatives of riboflavin were determined to be good candidates as well (Figure 39).
Riboflavin, also known as Vitamin B2, is safe and non-toxic (as it is used as a vitamin supplement) and is
also quite cheap and easily accessible.
119
In fact, in 2008, the total world riboflavin production capacity
was around 10,000 tons (estimated to be worth $150 million) whereas annual demand is around 6,000
tons.
120
Though riboflavin has been known for its nutritional and medical benefits for decades now,
121-122
it has the potential to be used as a redox molecule in electrochemical systems, including batteries.
123
It
is relatively insoluble in aqueous systems; however, by adding a phosphate group the molecule’s
85
solubility is increased 10-fold (up to 1.5 M), making it a viable option for a negative side compound in
alkaline media (Figure 38).
Figure 38: Cyclic voltammograms of riboflavin (blue), riboflavin phosphate (red) on a glassy carbon
working electrode. The redox couple concentrations were 1 mM in 1 M potassium hydroxide. A MMO
reference electrode (E° = +0.140 V) and a platinum counter electrode were used. The scan rate was 50
mV/sec.
Figure 39: Riboflavin Molecule
86
Riboflavin has, however, been shown to undergo chemical transformations with time. Riboflavin
is also known to have photo-sensitivity; when irradiated by visible light in the presence of oxygen,
riboflavin undergoes electron transfer that can produce superoxide radicals, singlet oxygen, hydroxy
radicals, and hydrogen peroxide.
124-125
In its as-purchased and reduced form riboflavin dimerizes and
thus its electrochemical properties change over time when in alkaline media, independent of any
electrochemistry done on the molecule.
123
The dimer has a lower standard reduction potential than the
monomer.
This degradation was seen in some of our cycling of full redox flow cells. Figure 40 shows that a
cell was set up with riboflavin phosphate as the negative side material, and potassium ferrocyanide as a
positive side material (which has been shown to be an adequate positive side material for alkaline-based
RFBs).
52
The first cycle had approximately only 40% utilization, and then the capacity faded quickly. This
fast fade rate was due to the riboflavin phosphate dimerization process that was taking place. The dimer
of the riboflavin phosphate had a lower standard reduction potential than the monomer; therefore, the
original cell voltage cutoff at 1.2 V was too low to fully utilize the dimer. Because the dimer’s reduction
potential was not reached during this cycling process, the capacity faded as the original monomer
solution was converted to the dimer of riboflavin phosphate.
87
Figure 40: 50 mM Riboflavin phosphate as the positive side material and 100 mM potassium
ferrocyanide as the negative side material in 1 M KOH. Cycling was done at a current density of 2
mA/cm
2
.
4.6 Hydroquinones
Even though there were several options for the negative side of the alkaline RFB, it was difficult
to find compounds that were stable and also had the required high reduction potentials for use on the
positive side of the cell.
The hydroquinones show remarkably different behavior once the pH is raised from zero (1M
sulfuric acid) to 14 (1M potassium or sodium hydroxide). It appears that partially substituted
benzoquinones are unstable in basic media due to oxidation – many of the CVs in Table 9 show that the
hydroquinones can easily undergo oxidation but do not have any sort of reduction current. There are
two reasons for the poor reversibility of hydroquinones in alkaline media: (1) it is known that the
Michael transformation is more facile under basic conditions, (2) it is also known that in basic conditions
88
that two one-electron reactions are more likely than a single two-electron reaction; the formation of a
semiquinone radical can lead to dimerization of the hydroquinone, which could limit the reversibility of
the compound.
Table 9: Hydroquinone Molecules Tested in Alkaline Media. All voltages are against a MMO Reference
Electrode (E° = +0.140 V), all quinone concentrations are 1 mM in 1 M potassium hydroxide. The working
electrode was glassy carbon and the counter was a platinum wire. The scan rate was 50 mV/sec.
Name Structure CV/RDE Data
Hydroquinone
2-sulfo, p-
benzoquinol
4,5-dihydroxy
benzene-1,3-
disulfonic acid
HO
3
S SO
3
H
OH
OH
89
3,5-dimethyl, 6-
sulfo, p-
benzoquinol
3-hydroxy, 4,6-
disulfo, o-
benzoquinol
3-sulfo, 5-
carboxylic, o-
benzoquinol
2,3,5,6-carboxylic
acid, p-
benzoquinone
SO
3
H
OH
OH
C H
3
CH
3
OH
OH
COOH
COOH HOOC
HOOC
90
2,5-dihydroxy, 1,3-
benzene disulfonic
acid
OH
O H
SO
3
H
SO
3
H
duroquinone
Tri-methyl, mono
sulfonic acid
hydroquinone
C H
3
CH
3
C H
3
SO
3
H
OH
OH
Dihydroxybenzoic
acid
91
It is known the Michael addition reaction is actually more likely to occur in alkaline solutions –
the nucleophilic addition of hydroxide to an open position on the ring is aided by the higher pH.
Therefore, partially-substituted hydroquinones are unstable in alkaline media, due to the chemical
transformations that they are likely to undergo. Additionally, dimerization of hydroquinone molecules
can occur through semi-quinone radical formation for benzoquinone at high pH values. The fully
substituted hydroquinones did not appear to have any sort of chemical transformations following their
electrochemical oxidation – they showed varying small amounts of reduction current. Therefore, fully
substituted hydroquinones are the only viable candidates for alkaline redox flow batteries.
However, the fully substituted hydroquinones either were difficult to synthesize or had low
values of reduction potentials. For example, the tetra-carboxylic acid para-benzoquinone was
synthesized by Dr. Sankarganesh Krishnamoorthy, and it showed excellent electrochemical
characteristics. However, this molecule was very difficult to synthesize – the yield was around 2% after
several synthetic steps. Therefore, even though it was important to keep the fully substituted
benzoquinone, it was also important to discover a molecule that is easy to synthesize.
The best hydroquinone candidate is a tri-methyl mono-sulfonic acid hydroquinone. When tested
in alkaline media, it appeared stable and had a relatively high reduction potential. One of the future
steps will be to place it in a cell and cycle it against one of the promising negative side materials. As can
be seen in Figure 41, the separation between reduction potentials is approximately 600 mV, a significant
cell voltage for a commercially viable RFB.
92
Figure 41: RDE of 2,6-dihydroxy, 9-10-anthraquinone and 2,3,6-trimethyl, 1,4-dihydroxy benzenesulfonic
acid and the potential difference between them. Rotation rate of 1500 rpm. 1mM concentration of
quinone in 100 mL of 1M sodium hydroxide. A MMO reference electrode (E° = +0.140 V) and a platinum
counter electrode were used. The scan rate was 50 mV/sec, and the rotation rate was 1500 rpm.
4.7 Summary
Though there are several reasons to have an alkaline based system for an organic RFB, the most
suitable combination of redox couples for the positive and negative side compounds have yet to be
discovered. However, there are several promising options for molecules that are soluble and stable in
alkaline media, while giving the requisite reduction potentials. These promising combinations include
using quinoxaline or hydroxy-anthraquinones as negative side materials and fully-substituted
hydroquinones as positive side materials. Quinoxaline derivatives are excellent candidates for the
negative side of an alkaline RFB, based on their reduction potentials, reversibility, and solubility.
Partially substituted hydroquinones undergo chemical transformations which make them generally
unsuitable for use in an alkaline RFB. However, 2,3,6-trimethyl, 1,4-dihydroxy benzenesulfonic acid is a
93
promising fully-substituted hydroquinone candidate, which needs to be investigated further for its use
as a positive side material.
Finally, the hydroxy-substituted anthraquinones are good candidate molecules for the negative
side of an alkaline-based RFB, as well as opening up the possibility of a bi-functional molecule – one that
could be used on both sides of the battery. Bi-functional molecules expand the range of possibilities for
redox flow batteries and are a promising approach to overcome the problem of crossover.
56
94
95
Chapter 5
Full Cell Cycling of an Acid-Based Redox Flow Battery
5.1 AQDS/BQDS Full Cell Studies and Results
For many of our cycling experiments in full cells, the positive and negative electrolyte materials
were BQDS and AQDS, respectively. This combination was used because of their relatively high
difference in standard reduction potential values, their high solubility, fast electrochemical kinetics, and
their ease of procurement. These materials were helpful in studying and optimizing the conditions for
operation of the full cell including flow field configuration, electrode structure, optimal flow rates, and
choice of membranes. In this chapter, we discuss the effect of operating conditions as well as the
optimized performance characteristics, such as power density and efficiency measurements.
5.1.1 Solubility of the Acid Form and Sodium Salt of the Quinones
High values of solubility of the redox couples is necessary for achieving high current densities
and high efficiencies in ORBAT. We had used the “as received” sodium sulfonate salts of BQDS and AQDS
(Sigma-Aldrich) in all our preliminary experiments. AQDS and BQDS had solubility values of 0.5 to 1.7 M,
respectively, in 1 M sulfuric acid at 23
o
C. When the sodium sulfonate salts of the quinones were
transformed to the free acid form (sulfonic acid) by passing the solution through an ion-exchange
column (Amberlyst 15(H) ion exchange resin), the solution concentrations of AQDS and BQDS could be
increased to as high as 1.5 M and 4 M, respectively. For fully-dissociated electrolytes resulting in
solvated ions in solution, solubility is governed by the difference between the crystal lattice energy and
the solvation energy.
126
The higher solubility observed for the acid form compared to the sodium salt is
consistent with the solvation energy of a sodium ion being 162.7 kcal/mole less than that of the
formation of a hydronium ion.
127
96
To test the effect of solubility and concentration on cell performance, current-voltage
measurements were made when the redox couples were in their fully charged form in the cell. In this
measurement, the current was controlled while voltage was measured. We observed a five-fold increase
in current density with the free sulfonic acid form over the sodium salt form (Figure 42); this increase in
performance can only be attributed to the increase in concentration of the reactants and improved ionic
conductivity of the membrane. The higher solubility of the acid form avoided the abrupt drop in cell
voltage resulting from concentration polarization caused by inadequate mass transport, especially at
high current densities and low states-of-charge. The impedance of the cell at 10 kHz with 1 M sulfuric
acid was 0.30 Ohm cm
2
, and the addition of 0.5 M of the sodium salt of AQDS raised this value of cell
resistance to 0.58 Ohm cm
2
.Thus, the absence of sodium ions in the electrolyte increased the ionic
conductivity of the Nafion
®
membrane, as protons have a higher mobility in the membrane compared to
the sodium ions.
101, 128
Figure 42: Current-voltage curves for the flow cell: 1M of BQDS/AQDS in acid form (labeled A) and 0.2 M
of BQDS/AQDS in sodium salt form (labeled B). Both cells employed Toray® paper electrodes and a “flow-
by” flow field design.
97
5.1.2 Effect of Flow Field Modifications
Employing the proper type of flow field in a redox flow battery is important in utilizing the
electrode surface area. Four different flow-field designs were tested to determine whether “flow-by” or
“flow-through” characteristics yielded better overall cell performance. The different designs had varying
percentages of both of these characteristics. The “flow-through” configuration refers to a fluid path that
includes flow that is perpendicular and parallel to the plane of the electrode; thus, the entire volume of
the electrode has the opportunity to encounter the flow of solution. In a “flow-by” design, the fluid flow
is largely parallel to the plane of the electrode, and flow resistance is lowered. The relative utilization of
the area of the electrodes in these two configurations is dependent on the flow rate and the pressure
drop between the inlet and the outlet. It is also dependent on the kind of electrode structure that is
employed (see next section).
To determine the impact of various types of flow fields on cell performance, we made current-
voltage measurements with BQDS and AQDS. Changes to the flow-field and electrode structure resulted
in at least a three-fold increase in power density (Figure 44). In these tests, an adequate amount of
AQDS was used at the start of cycling tests to ensure complete electrochemical conversion of BQDS by
sequential Michael addition reactions. The flow field was modified from a flow-by configuration to an
interdigitated design that added “flow-through” characteristics to the original flow field. We still had a
small component (~ 20%) of flow-by due to features that allowed some by-passing of the flow
interdigitated structure (Figure 43). With this improved design, the solutions were forced to move
through the felt electrode structure, instead of simply flowing past the electrode. Several studies
confirm the benefits of the flow-through design for redox flow batteries.
16, 129-132
By utilizing more of the
surface area of the electrode and by improving mass transport to the electrode surface, an increase in
the cell performance was achieved. We note that the benefits of the flow-through design tended to
increase with the flow rate. However, the flow-through design caused an increase in pressure drop for
98
the flow of the solutions. The pressure drop was high when the carbon-coated Toray
paper electrodes
were used with a simple interdigitated flow field. With an open electrode structure, such as the SGL
graphite felt, and with a combination of flow-through and flow-by designs, we could minimize the
pressure drop and still maintain the high level of cell performance. We will continue pursuing research
on optimization of the flow-field design with open electrode structures for improving cell performance.
Figure 43: Diagram of improved flow field arrangement with modified interdigitated flow field with
about 20% of “flow-by” properties.
Inlet
outlet
Custom channel blocks
Flow-through
Flow-by
Transverse Flow-by
Transverse
Flow-by
Flow-through
Graphite Felt
Graphite Base
Column/pin structures
99
Figure 44: Polarization Curves (top figure) and power density curves (bottom figure) at 100% state-of-
charge for A) 1M AQDS/BQDS in 1M Sulfuric acid, with a “flow-through” flow field and carbon-coated
carbon felt electrodes and B) 1M AQDS/BQDS in 1M sulfuric acid, with a “flow-by” flow field and carbon-
coated Toray® paper electrodes.
100
5.1.3 Electrode Design
Efforts on improving the electrode design focused on increasing the accessible surface area of
the electrodes. The initial experiments utilized Toray paper electrodes. However, it was determined that
using (modified) carbon felt electrodes can greatly increase surface area.
133
Additionally, it is known that
attaching high surface area carbon to electrode surfaces will improve performance in redox flow
batteries.
134-137
This work was done in collaboration with my colleague, Advaith Murali. Double-layer
capacitance measurements were used to determine the electrochemically active surface area. While the
graphite felt treated with 2.5 wt % of Nafion
®
appeared more wettable than the un-treated felt, both
these samples had nearly the same double-layer capacitance (Figure 45, Table 10). We expected that
when the surface became more wettable more of the electrochemical active area would increase and,
consequently, the double layer capacitance would increase. While Nafion® increased the apparent
wettability, it did not affect the accessible active area of the electrode to any significant degree. This
observation suggests that the inaccessible area was in the micropores of the graphite fibers that were
not affected by the addition of Nafion®.
Table 10: Double Layer Capacitance and Electrochemically Active Surface Area of Various Types of SGL
Carbon Felt (Geometric Area of 25 cm
2
) Determined from the Data in Figure 45.
Electrode Material
(geometric area of 25 cm
2
)
Capacitance (mF)
Electrochemically Active
Surface Area (cm
2
)
Unmodified Graphite Felt 2032.5 81.3
2.5 % Nafion®-Coated Graphite
Felt
2364.1 94.6
10% Carbon-Coated Graphite Felt 16026 641.0
30% Carbon-Coated Graphite Felt 28249 1129.9
101
Figure 45: The imaginary component of the impedance plotted against the reciprocal of the angular
frequency of excitation. The equations are for the line fits for the plotted data.
The addition of Vulcan XC-72
carbon black into the graphite felt material greatly increased the
double-layer capacitance (Table 10). The incorporation of 30 weight % of the Vulcan XC-72 and 2.5% of
Nafion® in the graphite felt yielded the highest capacitance. Thus, the surface area of the electrode was
substantially increased without any noticeable impact on the resistance to flow through the electrode
structure. The graphite felt treated with 2.5 wt% of Nafion® and 30% carbon black was employed in
further optimization of ORBAT performance.
5.1.4 Power Density
After selecting the electrode with the highest surface area (30 wt% Vulcan XC-72) and a flow-by
flow field for use in the redox flow cell, a power density as high as 100 mW/cm
2
could be achieved with
the high concentration of reactants in the acid form (Figure 44). Continuous cycling at a power density
102
of 55 mW/cm
2
at about 23
o
C could be sustained at a round-trip energy efficiency of 70% for at least 100
cycles (Figure 46). We expect to achieve higher power density by replacing BQDS with a redox couple
with a higher standard reduction potential that will avoid Michael reaction transformations. With
further optimization of the flow field structure, reduction of the thickness of the electrodes, and small
increases in operating temperature, we expect that a power density of 300 mW/cm
2
– similar to the
standard operating power density of commercial vanadium redox flow batteries – can be achieved.
138-139
5.1.5 Coulombic and Energy Efficiency
Maintaining 100% coulombic efficiency at all depths of discharge is essential for large-scale
energy storage applications. The charge-discharge efficiency of 100% was maintained for at least 100
cycles after the transformations in the first 4 to 5 cycles (Figure 46). Capacities as large as 10 Ah were
realized from the cells. While the initial charge-discharge efficiency was low due to the aforementioned
chemical transformations of BQDS, the charge-discharge efficiency stayed close to 100% once the
transformations were complete. The value of 100% coulombic efficiency over the 100 cycles is
consistent with the absence of crossover as confirmed by the electrochemical measurements, NMR
data, and independent ex-situ studies.
The round-trip energy efficiency of ORBAT when charged and discharged at 100 mA/cm
2
was
about 70% over 100 cycles. The cells were cycled between 1.0 V and 0 V with full utilization of the
solution at the positive electrode, since it is has been confirmed that the solution at the positive
electrode limits the capacity of the cell. The electrolyte path length in the graphite felt electrodes
(thickness of 2.5- 3.0 mm) contributed significantly to the ohmic resistance of the cell. This resistance
was considerably less with Toray
paper electrodes. Thus, the value of energy efficiency was determined
from the area under the charge and discharge curves after correcting the cell voltage for losses from the
ohmic resistance in the cell. The energy efficiency can be improved further by reducing the mass
transport losses.
103
These mass transport losses can be mitigated by decreasing the diffusion layer thickness and
increasing the contact area of the redox-active materials. The diffusion layer thickness can be decreased
by increasing the flow rate; a faster flow rate does not allow for the build-up of a diffusion layer, as the
solution is quickly forced past the electrode surface. More contact area can be achieved by modifying
the flow field and electrode structure. Of course, an open electrode surface with high surface area
increases contact area. However, the flow field ensures that the flow of the liquid is able to access this
open, available surface area. A flow field that forces the liquid to move through the electrode (“flow-
through”, instead of past it (“flow-by”) can lead to an increase in surface area, thus leading to lower
mass transport losses.
Figure 46: Coulombic efficiency and energy efficiency over 100 cycles for a 25 cm
2
flow cell with 1 M
AQDS/BQDS with a modified interdigitated flow field and carbon-coated graphite felt electrodes (flow
rate of 1 L/min). The cell was charged and discharged between 0 V and 1 V at 100 mA/cm
2
. Energy
efficiency was calculated based on voltage-time curves that were corrected for the internal resistance.
104
5.2 3,6-dihydroxy-2,4-dimethylbenzenesulfonic acid (DHDMBS) Studies in a
Redox Flow Cell
We studied 3,6-dihydroxy-2,4-dimethylbenzenesulfonic acid (DHDMBS) as a positive side
material, aiming to mitigate the Michael reaction that was encountered with BQDS.
140-141
To reduce the
susceptibility of the molecule to the Michael reaction it was essential to design a molecule with the
minimum number of unsubstituted positions on the benzoquinone. Sulfonic acid groups were still
desirable substituents, as they increased the reduction potential and the solubility of the molecule in
water. To limit the number of open positions on the benzene ring, methyl groups were used as
substituents; further discussion on the selection of this molecule can be found in Chapter 3.
Therefore, we combined 1 M DHDMBS in 1 M aqueous sulfuric acid on the positive side of the
flow cell with 1 M anthraquinone-2,7-disulfonic acid (2,7-AQDS) in 1 M aqueous sulfuric acid on the
negative side. We used 2,7-AQDS instead of 2,6-AQDS, because of inadequate supply and the high cost
of obtaining 2,6-AQDS. Earlier, we had tested the electrochemical and kinetic parameters of both 2,6-
AQDS and 2,7-AQDS and found that all properties were comparable (see Table 4 in Chapter 3 for
details).
5.2.1. Improved Performance
In the flow cell experiments, we ensured that the same number of equivalents of DHDMBS and
2,7-AQDS were used on both sides. The charge and discharge curves at 100 mA/cm
2
showed a single
plateau (Figure 47), which is markedly different from that of the BQDS where multiple plateaus were
observed from the very first cycle. Furthermore, with BQDS the first discharge only returned one-third of
the first charge capacity, consistent with the transformations caused by the Michael reaction, needing
three times the amount of charge to reach the final charged state. When DHDMBS was used, the first
charge and discharge capacities were almost equal, proving that there was no Michael transformation
occurring with DHDMBS.
105
Figure 47: Charge and discharge curves for 1 M DHDMBS as the positive electrolyte and 1 M 2,7-AQDS as
the negative electrolyte, each dissolved in 100 mL of 1 M sulfuric acid. The cell used interdigitated flow
fields and graphite felt electrodes. The membrane was Nafion® 117. Charge and discharge currents were
2.5 Amperes, and the cut-off voltage was 1.0 V during charge, and 0.005 during discharge. The dashed
lines represent the IR-corrected voltages; the solid lines the non-corrected voltage values.
Upon continuing the cycling of the cell with DHDMBS and 2,7-AQDS, we were able to achieve
100% coulombic efficiency repeatedly during charge and discharge cycles (Figure 48a), indicating that all
the charged DHDMBS was being returned to the discharge state. The 100% coulombic efficiency was yet
further proof that the Michael reaction was not occurring with the DHDMBS electrolyte.
During the 25
th
cycle the current-voltage curve was obtained in the fully charged state. The cell
was able to achieve 500 mA/cm
2
of current density (Figure 48b). The flatness of the curve at high
current densities suggested that the mass-transport limitations were not significant even at such high
current densities. However, the ohmic resistance arising from the graphite felt electrodes with 1 M
sulfuric acid contributed to 0.020 Ohm (as measured by impedance spectroscopy). At a discharge
current of 12.5 A (500 mA/cm
2
), the voltage drop due to this series ohmic resistance was 0.25 V, a very
significant number. This ohmic resistance can be lowered with the choice of thinner electrodes
106
combined with modifications to the contact area of the flow field plate.
142
To depict the impact of this
ohmic loss, we have plotted the corrected value of cell voltage along with the measured value in Figure
47.
Figure 48: Cycling studies of 1 M DHDMBS and 1 M 2,7-AQDS when charged and discharged at 100
mA/cm
2
. (a) Coulombic efficiency of flow cell; (b) Current-voltage curve at the end of charge in the 25th
cycle.
107
The relatively low cell voltage of the system could lead to significantly greater costs during scale-
up of this system as more cell units would be necessary; more cell units entail the use of more flow field
plates, electrodes, and membranes, which are expensive. However, the lower cost of the electrolyte and
active material, especially when compared to vanadium, still makes this system worthy of consideration
for a grid-storage RFB. Therefore, cell and electrode design optimization would be an essential part of
the scale up of this system. Further, we can design molecules with electron-withdrawing substituent to
achieve a higher reduction potential that would raise the cell voltage, preserving the property of being
resistant to the Michael reaction, just as DHDMBS is.
5.2.2 Confirmation of No Michael Reaction
After 25 cycles, even though the cell maintained close to 100% coulombic efficiency, the
discharge capacity showed about a 0.05% decrease every cycle (Figure 49). We removed 500 microliters
from each side of the cell at the end of the 25
th
discharge to determine the state of the materials using
electrochemical and spectroscopic techniques. This sample was compared with the sample prior to the
commencement of cycling.
1
H NMR analysis confirmed that DHDMBS was still in its original state and
had not undergone any significant chemical transformation (Figure 50). The continued presence of the
aromatic proton signal after 25 cycles showed that the Michael reaction had not occurred on DHDMBS,
unlike on BQDS. The minor peaks in the aliphatic region of the
1
H NMR corresponded to the oxidized
form of DHDMBS present even in the discharged state. Therefore, it was clear from the NMR results that
the DHDMBS had remained unchanged during the 25 cycles.
108
Figure 49: Capacity of flow cell operating with 1 M DHDMBS and 1 M 2,7-AQDS when charged and
discharged at 100 mA/cm
2
.
Further confirmation of the absence of any chemical transformation of DHDMBS was evident
from electrochemical tests. Linear-sweep voltammograms of the DHDMBS before cycling and after 25
cycles showed no change in the onset potential and no noticeable change in the limiting current for both
samples (Figure 51). These observations further confirmed that there had been no chemical
transformation of the molecule over the 25 cycles. Thus, the electrochemical and spectroscopic analysis
confirmed that the capacity decrease with cycling was not due to any chemical transformations. These
results attested to the robustness of DHDMBS to the Michael reaction with water, and confirmed the
suitability of using DHDMBS as a positive electrolyte. Thus, DHDMBS had overcome the principal issue
encountered with BQDS.
109
Figure 50: 1H- NMR studies on samples of DHDMBS before and after 25 cycles. Imidazole was added
deliberately as an internal standard for estimating concentration. Electrolyte solutions were diluted with
deuterated-water (D2O).
110
Figure 51: Linear-sweep voltammogram of DHDMBS samples before and after 25 cycles. Both samples
were at a 1 mM concentration. The scan rate was 50 mV/sec and the rotation rate was 1500 rpm. The
working electrode was a glassy carbon rotating disk electrode, the counter electrode was a platinum
wire, and the reference electrode was mercury sulfate (MSE, E° = +0.65 V).
5.2.3 Crossover through the Nafion
Membrane
While investigating the slow capacity decrease with cycling, we found that the crossover of
DHDMBS from the positive side of the cell to the negative side of the cell could be a possibility. We had
expected that the sulfonic acid molecules would be fully ionized and the anionic molecular species
would therefore be excluded by the Nafion
membrane. However, when we determined the pK a of
DHDMBS to be 2.46, we concluded that in a 1 M solution of sulfuric acid (at a pH close to zero) much of
the DHDMBS would be in the unionized form as the sulfonic acid. Thus, permeation and electro-osmotic
drag could allow DHDMBS to be transported across the membrane.
By analyzing the negative electrolyte compartment by
1
H NMR, it was determined that small
amounts of DHDMBS had indeed crossed over through the Nafion
117 membrane (Figure 52). The
111
concentration of DHDMBS in the AQDS electrolyte had slowly increased with time. Since the cells were
being cycled at 100 mA/cm
2
, the impact of concentration changes on the capacity was even more
significant than it would have been at lower current densities. From studies of the limiting current on
the rotating disk electrode, we determined that the concentration of DHDMBS had reduced by about
10% through the 25 cycles, consistent with the decrease in capacity observed. Similar investigations of
the positive electrolyte (Figure 50) did not show any AQDS crossover into the positive electrolyte even
after 25 cycles. It was surprising that, although AQDS is also not expected to be fully ionized, we did not
see its crossover to the positive side. We concluded that in addition to the state of ionization, the
molecular size has a significant effect on the permeability of the molecule through the membrane.
Thus, the slow decrease in capacity of the cell could be attributed largely to the slow fall in
concentration of DHDMBS in the positive electrolyte due to crossover. It was, therefore, important to
deter the transport of neutral molecules across the membrane, since incomplete ionization leads to
crossover. With the determination that crossover is now the main cause of a capacity fade in the RFB, it
was important to investigate methods to mitigate this phenomenon.
112
Figure 52:
1
H- NMR studies on samples of 2,7-AQDS before and after 25 cycles. Imidazole was added
deliberately as an internal standard for estimating concentration. Electrolyte solutions were diluted with
deuterated-water (D2O).
5.3 Mixed Electrolyte Studies and Potential for Bi-functional Redox Molecule
5.3.1. New Membrane Testing
Crossover of redox materials is detrimental to the lifetime of the batter; therefore, it was of
utmost importance to engineer a solution to this problem. One solution is to install a membrane that
would limit the crossover of redox materials. Consequently, two new membranes purchased from
Fumatech were tested out for the crossover properties – E750 and F1850. Their structures, thickness,
proton conductivity, and water uptake are shown in Table 11. Both of these membranes were tested
because of their lower percentage of water uptake – this property would limit the mobility of the
DHDMBS from crossover
Before cycling
After 25 cycles
}
AQDS
}
AQDS
Imidazole
Imidazole
Imidazole
Imidazole
113
quinone molecules because of the decrease in the size of water domains through which the transport of
the molecules would occur. Therefore, even though they were thinner than the Nafion® membranes,
the crossover of the quinone molecules could be expected to be lower. Additionally, the use of thinner
membranes would help lower the ohmic resistances in the cell, even though their conductivity values
are lower.
Table 11: Redox Flow Battery Membrane Physical Characteristics
Designation Source Structure Thickness,
micron
Proton
Conductivity, S
cm
-1
Water Uptake,
%
Nafion® 117 Du Pont 1100 EW 175 0.1 38
Nafion® 324 Du Pont/
W Grot
Layered composite,
1500 EW/1100 EW
25/130 ~0.1 ~35
F1850 Fumatech 1850-2200 EW
perflourinated
45-55 0.025 5-10
E750 Fumatech 750 EW H-PEEK
membrane
50 0.016 Not available
It was shown through systematic studies that E750 showed a lower degree of crossover in 100
cycles. Figure 53 shows that the capacity fade rate (% per cycle) of the Fumatech membranes is an order
of magnitude lower than Nafion® 117. Even when the fade rates were normalized for the different
thicknesses of the membranes, the fade rate of the Fumatech membranes was half that of the Nafion®
membrane. It was determined that it was not just a function of thickness, but also the size of the
pathways through the membranes that determine crossover rate. Fumatech membranes have more
molecules per unit area – they are packed tighter and provide a denser network that reduces the
passages for molecular travel. Also, these membranes have lower water swelling that ultimately limits
the chain spacing and prevents crossover pathways.
The NMR spectra and the RDE tests confirmed that crossover of the DHDMBS molecule to the
negative side of the cell was considerably reduced with the alternate membranes; the amount of the
114
crossed over material is consistent with the capacity fade rate. Therefore, simply by using a different
membrane the crossover rate of DHDMBS could be reduced.
Figure 53: Capacity and cycle life of various flow cells operating with 1 M DHDMBS and 1 M 2,7-AQDS
when charged and discharged at 100 mA/cm
2
. Various membranes were employed. Capacity fade rate
was normalized for thickness of membrane.
5.3.2. Mixed Electrolyte Studies
Another solution to the problem of crossover is to mix the positive and negative electrolytes and
use this mixed material on both sides, termed here as a symmetric cell. This solution to the crossover
issue, while requiring twice the amount of active materials to be placed in the tanks, could be used as an
interim measure in place of designing a new membrane that completely avoids crossover. Using a
mixture of the electrolytes on both sides is an approach that is applicable regardless of the actual
molecules that we would be using in the final version of the system. Additionally, the membrane
115
solution is likely to be unique to the specific redox molecules in use. It has been shown in the literature
that mixed electrolytes in RFBs have worked in a variety of situations.
56
We have set up cells that contain the mixed electrolytes and have charged and discharged them
to their full capacity. It is to be expected that the mixed electrolyte cells will also exhibit a slow fade in
capacity. However, if the solutions in the positive and negative reservoirs are mixed and sent back to the
tanks, the capacities can be recovered. This procedure, called “mix and split,” is an expedient to mitigate
the sustained effect of crossover in a conventional cell. This capacity recovery is shown in Figure 54 –
every spike in capacity was after the solutions from both tanks were removed, mixed together, and then
split in two and placed on either side of the battery once again. This “mix and split” protocol was done
after the starting capacity had reduced by 25%. The total fade rate over more than 450 cycles was 0.3%.
Figure 54: AQDS/DHDMBS mixed electrolyte cell in 1 M sulfuric acid solutions. Solutions were mixed and
split after capacity had faded 25%. A Nafion® 117 membrane was used.
While this “mix and split” protocol is relatively easy to employ and requires little constant
maintenance, there are some drawbacks, including the fact the solutions are exposed to oxygen during
116
the mixing. These molecules are very sensitive to oxygen and can be easily oxidized; therefore, this
method could end up limiting the full capacity return of the solutions. To address these issues, we have
developed a protocol called “lead switching.” In lead switching – or lead reversal – the direction of
charge is reversed at the end of the discharge by switching the connections of the positive and negative
electrodes. Since a mixed electrolyte cell is a symmetric cell, at the end of discharge both sides of the
cell will have the same composition and thus the “positive” or “negative” sides are equivalent.
Therefore, even if one of the redox couples (for example, DHDMBS) crosses over, lead switching can
basically reverse the detrimental effects of crossover.
This method especially proves its usefulness and practicality when the leads are reversed at the
end of every cycle. This kind of cycling proves that it is possible to obtain the full capacity from the
battery in every single cycle. This switching protocol ensures that any crossed over material would be
able to be fully utilized again immediately. This protocol also proves that this concept would be the most
practical approach in a real-world situation. Being able to use the battery at 100% of its capacity
whenever the demand is needed is crucial for commercial applications, and this protocol proves that
this feature is feasible to attain.
Figure 55 shows that the capacity for the mixed cell is still fading, even though at a much lower
rate than with any other membrane or cycling protocol method. While lead switching can mitigate the
effects of crossover, we discovered yet another process that could be causing the capacity loss. This
process is proto-desulfonation of the DHDMBS (see Chapter 3 for more details). Proto-desulfonation of
DHDMBS produces a compound significantly less soluble than the original DHDMBS, leading to
precipitation of the compound. We have found insoluble materials in the pumps and in the graphite felt
electrodes. Proto-desulfonation could be occurring more rapidly in a mixed electrolyte cell because of
the higher free acid concentration that is used in mixed cells; in addition to the sulfonic acid group on
the DHDMBS, there are the two sulfonic acid groups on the AQDS and the acidity from the 1 M sulfuric
117
acid added to the electrolyte. This higher free acid concentration greatly increases the propensity of
proto-desulfonation in these mixed cells. Ex-situ experiments were done to determine the extent to
which this phenomenon was occurring. Figure 56 shows how the DHDMBS concentration decreased
over the course of cycling, but did not reappear on the negative side of the cell. This was due to the loss
of material dissolved in solution, which could no longer be detected with RDE studies. NMR data was
also taken at various points during several different cycling experiments, readily confirming the
existence of the proto-desulfonated molecule (the peak that appears at around 1.3 ppm (Figure 57)).
Figure 55: Capacity and cycle life of various flow cells operating a mixed (symmetric) cell with 0.5 M
DHDMBS and 0.5 M 2,7-AQDS on both sides. Charge and discharge current density was 100 mA/cm
2
.
Various membranes were employed.
118
Figure 56: Linear-sweep voltammogram of DHDMBS/AQDS mixed cell samples before and after 216
cycles. Both samples were at a 1 mM concentration. The scan rate was 50 mV/sec and the rotation rate
was 1500 rpm. The working electrode was a glassy carbon rotating disk electrode, the counter electrode
was a platinum wire, and the reference electrode was mercury sulfate (MSE, E° = +0.65 V).
Figure 57:
1
H- NMR studies on DHDMBS/2,7-AQDS mixed cell samples before and after 216 cycles.
Imidazole was added deliberately as an internal standard for estimating concentration. Electrolyte
solutions were diluted with deuterated-water (D2O).
119
5.3.2.1 Potential for Bifunctional Molecule
Yet another approach to the mitigation of crossover is to use a single type of molecule on both
sides of the cell. This molecule will have to possess “bi-functional” properties, i.e. it should be able to
function as both a positive electrolyte and a negative electrolyte material. Such a molecule would be
called a bi-functional redox couple. In fact, a bi-functional molecule builds on the concept of a mixed
electrolyte cell and improves upon it, since only one type of molecule is needed. In the case of the
previous mixed cell tests, the solubility and the energy density were halved. As long as the molecule
were not an adduct of two separate redox molecules, this energy density decrease could be avoided.
However, it would still have all of the advantages the mixed cell has, i.e. the non-issue of crossover.
Additionally, the logistics of implementing “lead-switching” or “mix and split” can be avoided. These
advantages of the bifunctional molecule are quite attractive from an economic and operational
viewpoint. This concept was introduced in the previous chapter – hydroxy anthraquinones have the
potential to be such a molecule. However, we have also demonstrated a proof-of-concept for this idea
by synthesizing an anthraquinone tethered to a hydroquinone sulfonic acid. Before sulfonation of this
compound, the molecule was practically insoluble in acidic solution and an inconclusive CV was
obtained. However, once sulfonated, two clear peaks emerged, corresponding to the anthraquinone and
benzoquinone reactions (Figure 59).
O
O
OH
OH
SO
3
H
SO
3
Figure 58: Sulfonated bifunctional molecule
120
The difference between the two reduction potentials is 850 mV, which is a desired cell voltage
for a RFB. Further work needs to be done to scale up this molecule for cycling studies, as the synthesis
for this molecule is tedious.
Figure 59: Cyclic voltammogram (red) and RDE (blue) of the sulfonated bi-functional molecule on a
glassy carbon working electrode. The bi-functional molecule concentration was 1 mM in 1 M sulfuric
acid. A MSE reference electrode (E° = +0.650 V) and a platinum counter electrode were used. The scan
rate was 50 mV/sec for both, and the RDE rotation rate was 1500 rpm.
5.4 Summary
In conclusion, we have shown that there are several parameters that determine the optimal
cycling conditions. First, the redox couples themselves – it is of the utmost importance that they do not
undergo any sort of chemical transformations. Chemical transformations are detrimental to the long-
term cycle life of an RFB. This problem can be avoided by choosing and designing molecules that are
resistant to chemical transformations. An example of such a compound was demonstrated in DHDMBS.
Secondly, flow field, electrode structure, and membrane selection are important engineering
parameters to optimize the RFB. It was determined that a flow field with more flow-through
characteristics used with carbon felt electrodes was an optimal design for the flow cell. Mitigation of
121
crossover is important to increase the cycle life of the battery. Fumatech membranes could limit the
transport of the hydroquinone molecules because of their lower water content. Consequently, these
membranes gave significantly lower crossover rates as opposed to the Nafion® membranes that had
been previously employed. The capacity fade rate could be reduced by this approach.
Finally, we have shown that a cell with mixed electrolytes has the potential to mitigate the
effects of crossover. Both “mix and split” and “lead-switching” protocols were found to be effective.
Finding a proper protocol for regaining lost capacity over time would be the main question of interest
during scale-up of a symmetric redox flow battery.
We have also shown that a single type of bifunctional molecule is viable for use as both the
positive and negative electrolytes in ORBAT. Such a bifunctional material, which could be placed on both
sides of the battery, can also mitigate the effects of crossover. This approach opens up new avenues for
future research.
122
123
Chapter 6
Conclusions
My work has focused on all aspects of the design of an all-organic aqueous redox flow battery.
We invented this type of battery at USC; it has the potential to be implemented on a grid-scale level, as
it utilizes environmentally-friendly and inexpensive materials. As more renewable energy power sources
(such as wind turbines and solar panels) are added to grids across the world, it is of paramount
importance to design energy storage that is efficient, long-lasting, robust, cheap, and environmentally-
sustainable. An organic redox flow battery (ORBAT), like the one discussed in this thesis, could be the
solution as it has the potential to fulfill all of these requirements. Another advantage of ORBAT is that
the energy and power densities are de-coupled, making it easily scalable for the required application.
My work concerns many aspects of ORBAT, from the study of the properties of the redox-active
molecules to cell design and lifetime studies. I have studied the effect that molecular architecture, use
of acidic or alkaline electrolytes, electrode and cell design, operating conditions, and membrane type
have on the overall performance of ORBAT cells. Specifically, I have understood the effect that
molecular structure has on electrochemical kinetics and long-term cycleability, the effect of cell design
and electrode structure on power densities and efficiency, the effect of membranes on crossover and
cycle life, and the effect of pH on the reactivity of the molecules.
For the acid-based ORBAT system, it was found that hydroquinone molecules could provide the
required reduction potentials and solubility values for use as the positive side electrolyte. However, the
chosen substituents play a substantial role in the electrochemical characteristics of hydroquinone redox
chemistry. Substituents not only modify the solubility and the reduction potential of the molecule, but
they have a large effect on the kinetics of the electron transfer reaction due to the intramolecular
124
hydrogen bonding effects. It was found that the hydrogen bonding effects for hydroxyl groups or
sulfonic acid groups can lower the rate constant of electron transfer.
Therefore, we designed a hydroquinone molecule with methyl groups on the ring, so as to limit
the hydrogen bonding effects and the subsequent inhibition of the electron transfer rates. However, we
had to ensure that substituent groups are used which raise the reduction potential and solubility of the
hydroquinone molecule, i.e. sulfonic acid groups. Therefore, for the positive side redox couple for
ORBAT, we designed a benzoquinone molecule that will meet the required high reduction potential and
solubility limits, while not compromising on electron transfer rate.
We found that it is also crucial that the redox-active molecules do not undergo any chemical
transformations, i.e. the Michael reaction or proto-desulfonation. There are chemical and electrolyte
engineering solutions to limit these detrimental chemical transformations. Molecular design and
synthesis is one method that was employed to limit these transformations, while changing the acid
concentration in the electrolyte helped limit proto-desulfonation as well.
Sulfonic acid substituted anthraquinones are excellent candidate molecules for the negative
electrolyte of an acid-based RFB. They are highly stable over a wide pH range, and can achieve sufficient
solubility in order to maintain economic viability. Substituted anthraquinones have been cycled against
several different positive side redox couples, in a variety of different cell configurations; we have yet to
see any sort of chemical degradation or transformation of this molecule.
While we have used the acid electrolytes for our baseline ORBAT studies, there are a few
advantages that an alkaline system offers. The higher pH values of the electrolyte will lead to decreased
corrosion of metal parts, resulting in lower material costs for the stack. Additionally, the potentially
higher solubility values will lead to better energy densities. There is also the potential to have lower
resistances due to the higher degree of ionization, which will lead to improved efficiencies. Finally,
125
alkaline based RFBs can employ a hydrocarbon-based membrane, instead of using expensive Nafion®
membranes.
Though there are several cost-effective reasons to have an alkaline based system for an organic
RFB, viable redox couples have yet to be discovered; however, we have found some potential
candidates. Quinoxaline derivatives and riboflavin are excellent candidates for the negative side of an
alkaline RFB based on their reduction potentials, reversibility, and solubility. The hydroxy-substituted
anthraquinones were also found to be good candidate molecules for the negative side of an alkaline-
based RFB. These anthraquinones can also be used as bi-functional molecules that could be placed on
both sides of the flow battery.
In addition to studying molecular architecture and the effect of pH on the reactivity of the
molecules, I also investigated the effect that electrode structure, cell design, and membrane choice have
on ORBAT performance. It was determined that a flow field with more flow-through characteristics used
with carbon felt electrodes was an optimal design for achieving high power densities. This improved
performance was partially due to the increased electrochemically-active surface area of the modified
carbon felt electrodes. However, the flow field design was also crucial for accessing this increased
surface area. A flow field with a large amount of “flow-though” characteristics ensured that the liquid
was forced through the electrode structure and was able to access all available surface area.
Mitigation of crossover is essential to increase the cycle life of the battery. Fumatech
membranes limit the transport of the hydroquinone molecules because of their lower water content
compared to Nafion® type membranes. Consequently, these membranes gave significantly lower
crossover rates compared to the Nafion® membranes that are widely employed in redox flow cells. The
capacity fade rate could be reduced by this approach.
We have also shown that a cell with mixed electrolytes (a symmetric cell) has the potential to
mitigate the effects of crossover as well. Both “mix and split” and “lead-switching” cycling protocols
126
were found to be effective in reducing the capacity fade rate. Designing a flow battery system that had
these sorts of protocols built into it to alleviate the detrimental effects of crossover on cycle life would
be the principle engineering challenge for scale-up of symmetric flow cells with mixed electrolytes.
We have also shown that a single type of bifunctional molecule is viable for use as both positive
and negative electrolytes in ORBAT. Such a bifunctional material, which could be placed on both sides of
the battery, can also mitigate the effects of crossover. This novel approach leads to new avenues for
future research.
127
PUBLICATIONS, PATENTS, AND PRESENTATIONS
Publications:
1. L. Hoober-Burkhardt, S. Krishnamoorthy, B. Yang, A. Murali, A. Nirmalchandar, G. K. Surya Prakash, S.
R. Narayanan, J. Electrochem. Soc. 2017, volume 164, issue 4, A600.
2. B. Yang, L. Hoober-Burkhardt, S. Krishnamoorthy, G. K. Surya Prakash, S. R. Narayanan, J.
Electrochem. Soc. 2016 volume 163, issue 7, A1442
3. B. Yang, L. Hoober-Burkhardt, F. Wang, G. K. Surya Prakash, S. R. Narayanan, J. Electrochem. Soc.
2014 volume 161, issue 9, A1371
4. F. M. Toma, A. Sartorel, M. Iurlo, M. Carraro, S. Rapino, L. Hoober-Burkhardt, T. Da Ros, M.
Marcaccio, G. Scorrano, F. Paolucci, M. Bonchio, M. Prato, ChemSusChem 2011, issue 4, 1447
Patents:
14/307030: INEXPENSIVE METAL-FREE ORGANIC REDOX FLOW BATTERY (ORBAT) FOR GRID-SCALE
STORAGE, Narayan, Sri; Prakash, Surya G. K.; Yang, Bo; Hoober-Burkhardt, Lena; Krishnamoorthy,
Sankarganesh, 6/17/2014
Oral Presentations:
1. L. Hoober-Burkhardt, B. Yang, S. Krishnamoorthy, A. Murali, G. K. Surya Prakash, and S. R. Narayanan,
Organic Redox Flow Batteries for Large-Scale Energy Storage, Honolulu, Hi., Pacific Rim Meeting on
Electrochemical and Solid State Science, October 2-7, 2016
128
2. B. Yang, L. Hoober-Burkhardt, G. K. Surya Prakash, and S. R. Narayanan, Studies on Aqueous Redox
Flow Batteries Based on Water-Soluble Quinone Redox Couples, Chicago, Il., Electrochemical Society
Meeting, May 24-28, 2015
3. B. Yang, L. Hoober-Burkhardt, G. K. Surya Prakash, and S. R. Narayanan, Understanding the
Performance of Aqueous Organic Redox Flow Batteries, Orlando, Fla., Electrochemical Society Meeting,
May 11-16, 2014
Poster Presentations:
1. L. Hoober-Burkhardt, B. Yang, G. K. Surya Prakash, and S. R. Narayanan, “Properties of Redox Couples
for Use in Organic Redox Flow Batteries,” Chicago, Il., Electrochemical Society Meeting, May 24-28, 2015
2. L. Hoober-Burkhardt, B. Yang, G. K. Surya Prakash, and S. R. Narayanan, “Applications of Quinone-
Based Redox Chemistry for Aqueous Flow Batteries,” Orlando, Fla., Electrochemical Society Meeting,
May 11-16, 2014
3. L. Hoober-Burkhardt, G. K. Surya Prakash, S. R. Narayanan, “Inexpensive Flow Batteries Based on
Organic Redox Couples”, San Francisco, Ca., Electrochemical Society Meeting, October 27 -Nov 1, 2013
4. L. Hoober-Burkhardt, G. K. Surya Prakash, and S. R. Narayanan, “An Inexpensive Metal-Free Organic
Redox Flow Battery (ORB) for Grid-Scale Storage”, San Diego, Ca., ARPA-E Grid Scale Energy Storage
Meeting, March 27-28, 2013
5. L. Hoober-Burkhardt, G. K. Surya Prakash, and S. R. Narayanan, “An Inexpensive Metal-Free Organic
Redox Flow Battery for Grid-Scale Storage,” Newport Beach, Ca., Engineering Conferences International,
June 2013
129
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Abstract (if available)
Abstract
This dissertation is focused on all aspects of the design of an all-organic, aqueous, redox flow battery (ORBAT). This kind of battery has the potential to be implemented on a grid-scale level, as it utilizes environmentally-friendly and inexpensive materials. As more renewable energy power sources, such as wind turbines and solar panels, are added to electricity grids across the world, it is of paramount importance to design energy storage that is efficient, long-lasting, robust, cheap, and environmentally-friendly. An organic redox flow battery like the one discussed in this thesis could be the solution to the need for robust and long-lasting energy storage, as it has the potential to fulfill all of these requirements. Another advantage of ORBAT is that the energy and power densities are de-coupled, thus making it easily scalable for the required application. ❧ This work covers many aspects of ORBAT, from the study of the properties of the redox-active molecules to cell design. Knowing the physiochemical properties, understanding the mechanisms of the electrochemical and chemical reactions, and measuring the interactions of the electrode surfaces with the quinones is of paramount importance to effectively utilize these quinone molecules in large scale energy storage systems. These properties are dependent on a number of factors, including solvent system, geometry and molecular structure of the quinone, concentration, temperature, pH, electrode material, and structure of the electrode surface. I have focused primarily on studying the effects of ring substituents, pH, and electrode materials—I have tested 40 different quinone and quinone-related structures for their reduction potentials, solubility, and electron and proton transfer kinetics. As a result, I have been able to select the most suitable redox couples and have validated their performance in a redox flow battery. I have also demonstrated major advances to the design of an all-organic aqueous redox flow battery, which is the first of its kind.
Linked assets
University of Southern California Dissertations and Theses
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Asset Metadata
Creator
Hoober-Burkhardt, Lena
(author)
Core Title
Small organic molecules in all-organic redox flow batteries for grid-scale energy storage
School
College of Letters, Arts and Sciences
Degree
Doctor of Philosophy
Degree Program
Chemistry
Publication Date
04/19/2019
Defense Date
03/23/2017
Publisher
University of Southern California
(original),
University of Southern California. Libraries
(digital)
Tag
battery,energy storage,OAI-PMH Harvest,quinone,redox flow battery
Language
English
Contributor
Electronically uploaded by the author
(provenance)
Advisor
Narayan, Sri (
committee chair
), Prakash, Surya (
committee member
), Shing, Katherine (
committee member
)
Creator Email
hooberbu@usc.edu,lenahoober@gmail.com
Permanent Link (DOI)
https://doi.org/10.25549/usctheses-c40-361711
Unique identifier
UC11256082
Identifier
etd-HooberBurk-5242.pdf (filename),usctheses-c40-361711 (legacy record id)
Legacy Identifier
etd-HooberBurk-5242.pdf
Dmrecord
361711
Document Type
Dissertation
Rights
Hoober-Burkhardt, Lena
Type
texts
Source
University of Southern California
(contributing entity),
University of Southern California Dissertations and Theses
(collection)
Access Conditions
The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the a...
Repository Name
University of Southern California Digital Library
Repository Location
USC Digital Library, University of Southern California, University Park Campus MC 2810, 3434 South Grand Avenue, 2nd Floor, Los Angeles, California 90089-2810, USA
Tags
battery
energy storage
quinone
redox flow battery