University of Southern California Dissertations and Theses

Advancing lithium batteries and related electrochemical technologies for a sustainable future

(USC Thesis Other)

Advancing lithium batteries and related electrochemical technologies for a sustainable future

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Content Copyright 2022 Billal Zayat

ADVANCING LITHIUM BATTERIES
AND RELATED ELECTROCHEMICAL TECHNOLOGIES
FOR A SUSTAINABLE FUTURE

by

Billal Zayat

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)

December 2022

ii
DEDICATION

إﻟ ﻰ أ ﻣ
ّ ﻲ

iii
ACKNOWLEDGMENTS

The work in chapters 3, 4, and 5 was done as part of the SCALAR Energy Frontier
Research Center. Oxide films discussed in chapter 3 were fabricated by Daniel Robertson at the
University of California, Los Angeles. The polymers discussed in chapters 4 and 5 were
synthesized by Pratyusha Das (Barry Thompson group) at the University of Southern California
(USC). Most of the full cell testing shown in chapters 3 and 4 was performed by Rodrigo Elizalde-
Segovia (Narayan group) at USC. Grazing Incidence Wide Angle X-ray Scattering (GIWAXS)
data and polymer swelling data was collected by Charlene Salamat (Tolbert group) at UCLA.
The work in chapter 6 was performed in collaboration with Dr. Jayakanth Ravichandran,
Dr. Brent Melot, and Dr. Shaama Sharada. Shantanu Singh (Ravichandran group) synthesized the
sulfides, and Dr. Ahamed Irshad (Narayan group) performed the hydrogen evolution reaction
experiments.
The work in chapter 7 was performed in collaboration with Dr. Debanjan Mitra, a former
member of the Narayan Research Group. Mainly, Dr. Mitra performed some of the physical
characterization and collected some of the electrochemical data in section 7.2.
The work in chapter 8 was done in collaboration with Dr. Antonina Nazarova and Dr.
Valery Fokin at USC. Dr. Nazarova synthesized all the compounds and collected all NMR data
that supplemented this work.

iv
TABLE OF CONTENTS

DEDICATION ....................................................................................................................................... ii
ACKNOWLEDGMENTS ....................................................................................................................... iii
LIST OF TABLES ................................................................................................................................... vi
LIST OF FIGURES ................................................................................................................................. vii
LIST OF PUBLICATIONS ..................................................................................................................... xvi
ABSTRACT ....................................................................................................................................... xvii
CHAPTER 1 BACKGROUND AND MOTIVATION ....................................................................... 1
1.1 The Challenge of Climate Change ................................................................................ 1
1.2 Electrochemistry and Sustainability .............................................................................. 4
1.3 Lithium Batteries and Renewable Energy ..................................................................... 6
1.4 Water Electrolysis for Clean Hydrogen Production ...................................................... 8
1.5 Other Applications of Electrochemistry ........................................................................ 9

CHAPTER 2. EXPERIMENTAL TECHNIQUES .............................................................................. 10
2.1 Cycling and Measuring the Conductivity of Polymer Binders ..................................... 10
2.2 Coin Cell Fabrication and Cycling ................................................................................ 12
2.3 Voltammetry at Rotating Disk Electrode ...................................................................... 16
2.4 Electrode Fabrication and Testing in Alkaline Conditions ........................................... 17
2.5 Physical Characterization Methods ............................................................................... 22
2.6 Cyclic Voltammetric Kinetic Studies ............................................................................ 23

CHAPTER 3. DEVELOPING AN IN-SITU CONDUCTIVITY MEASUREMENT TECHNIQUE
FOR POLYMER BINDERS AND ELECTRODES ..................................................... 26
3.1 Introduction ................................................................................................................... 26
3.2 Electrode Geometry and Methodology ......................................................................... 33
3.3 Method Validation with p-Dopable Conducting Polymers ........................................... 34
3.4 Applying Method to n-Dopable Polymers .................................................................... 45
3.5 Expanding the Technique to Oxides ............................................................................. 47

CHAPTER 4. ENHANCING LITHIUM-ION BATTERIES USING BIFUNCTIONAL
CONDUCTIVE BINDERS ........................................................................................... 52
4.1 PProDOT-Hx 2 as a Conductive Binder for Lithium-ion Batteries ................................ 52
4.2 Increasing PProDOT Ionic Conductivity with Oligoether Side Chains ........................ 69
4.3 Introducing Conjugation-Break Spacers to Improve PProDOT-Hx 2 Binding .............. 82

v
CHAPTER 5. IMPROVING THE PERFORMANCE OF LITHIUM-SULFUR BATTERIES
USING FUNCTIONALIZED CONJUGATED POLYMERS ..................................... 97
5.1 Introduction ................................................................................................................... 97
5.2 Electrochemical Cycling of Polymer Thin Films .......................................................... 99
5.3 Electronic and Ionic Conductivity of Polymer Thin Films ........................................... 101
5.4 Physical Characterization of Binders ............................................................................ 104
5.5 Polysulfide Shuttle Current Measurement .................................................................... 107
5.6 Performance of Lithium-Sulfur Full Cells .................................................................... 110

CHAPTER 6. PROBING TRANSITION METAL PEROVSKITE CHALCOGENIDES AS
PROMISING ELECTROCATALYSTS ....................................................................... 118
6.1 Introduction ................................................................................................................... 118
6.2 Synthesis and Physical Characterization ....................................................................... 119
6.3 Hydrogen Evolution Reaction ....................................................................................... 122
6.4 Oxygen Evolution Reaction .......................................................................................... 125

CHAPTER 7. DEVELOPING AN EFFICIENT AND INEXPENSIVE ALL-IRON WATER
ELECTROLYSIS DEVICE .......................................................................................... 129
7.1 Introduction ................................................................................................................... 129
7.2 Steel-Based Electrodes for Oxygen Evolution .............................................................. 131
7.3 Oxygen Electrodes for All-iron Water Electrolyzer ...................................................... 143
7.4 Steel-Based Electrodes for Hydrogen Evolution ........................................................... 150
7.5 Fabrication and Characterization of All Iron Electrolyzer ............................................ 159

CHAPTER 8. ELECTROCHEMICAL STUDIES OF THE CYCLOADDITION ACTIVITY OF
BISMUTH(III)ACETYLIDES TOWARDS ORGANIC AZIDES UNDER
COPPER(I)-CATALYZED CONDITIONS ................................................................. 165
8.1 Introduction to OrganoBismuth(III) Compounds in Drug Design ................................ 165
8.2 Synthesis and Characterization of Diphenyl Sulfone Bismuth(III) Acetylides ............ 167
8.3 Kinetic Cyclic Voltammetry for Mechanism Elucidation ............................................. 168
8.4 Kinetic Model Parameter Estimation from Experimental Data .................................... 176
REFERENCES ....................................................................................................................................... 186

vi
LIST OF TABLES

Table Page
3.1 Comparison of Various Conductivity Measurement Methods ......................................... 32
3.2 Fitting Parameters for Nyquist Plots in Figures 3.13 and 3.14 ......................................... 43
5.1 Sulfur and Fluorine Content on Lithium Electrodes Using EDS Analysis ....................... 116
6.1 HER Mechanism in Acidic Conditions ............................................................................. 124
7.1 Mass Gain, Overpotentials, and Tafel Slopes of CS Oxygen Electrodes ......................... 131
7.2 The Kobussen Pathway of Oxygen Evolution .................................................................. 133
7.3 Mass Gain, Overpotentials, and Tafel Slopes of NS Oxygen Electrodes ......................... 136
7.4 HER Mechanism in Alkaline Conditions ......................................................................... 156
7.5 Comparison with the Performance of Various Industrial Electrolyzers ........................... 164
8.1 Rate Parameters Derived from the Cyclic Voltammogram Kinetic Studies ..................... 183

vii
LIST OF FIGURES

Figure Page
1.1 (a) Global CO2 emissions between 1880 and 2020.
1
(b) Global temperature deviation
from 1880-1900 averages. Data from Goddard Institute for Space Studies. .................... 1
1.2 Annual global temperature anomaly (℃) at 2 m height in 2020 vs. 1981-2010 annual
average. Computed with KNMI Climate Explorer. ........................................................... 2
1.3 Global mean sea level change vs. 1870-1890. Data from NOAA Laboratory for Satellite
Altimetry. ........................................................................................................................... 3
1.4 Historic and projected world temperature change with respect to 1880-1900 averages
from The Intergovernmental Panel on Climate Change. Computed with KNMI
Climate Explorer ............................................................................................................... 4
1.5 Global Emissions by Sector in 2010. Data from Fifth Assessment Report of the
Intergovernmental Panel on Climate Change. ................................................................. 5
1.6 A schematic of a lithium-sulfur battery. ........................................................................... 7
1.7 Comparison of specific capacity and specific energy between Li-S and LIB. ................. 8

2.1 A representative ink sample and the resulting prepared electrode. .................................. 17
2.2 Experimental setup of RDE experiments for OER. .......................................................... 17
2.3 Schematic of the three-electrode cell used for electrochemical testing in alkaline
conditions. ......................................................................................................................... 19
2.4 Reaction scheme of the copper(I) complex formation. ..................................................... 24
2.5 Reaction scheme of bismuth(III) triazolide[X] formation. ............................................... 25

3.1 Schematic of a typical battery electrode. .......................................................................... 27
3.2 Electrode geometries used in impedance measurements for (a) low resistivity and (b)
high resistivity materials. .................................................................................................. 30

viii
3.3 SEM image of (a) spin-coated P3HT on the gold interdigitated electrode, (b) P3HT film,
and (c) and (d) electrode gold digits and the coated polymer film. .................................. 33
3.4 The gold interdigitated electrode in the (a) 2-electrode and (b) 3-electrode
configuration. .................................................................................................................... 34
3.5 Structures of p-dopable polymers used in this study. ....................................................... 35
3.6 CV plots of (a)P3HT, (b) P3HT/PEO, and(c) PEDOT:PSS, at 10 and 50 mV s
−1
. .......... 35
3.7 Electrochemical doping and 3-electrode impedance measurement protocol for
determining ionic conductivity. ........................................................................................ 36
3.8 A schematic of the polymer electrode in the 3-electrode cell. .......................................... 37
3.9 (a) Equivalent circuit of the transmission line model, (b) expected impedance response,
and (c) experimental impedance response. ....................................................................... 38
3.10 3-electrode impedance data of P3HT film. (a) Bode plots as a function of
electrochemical doping and (b) Nyquist plots at various potentials. ................................ 39
3.11 Ionic conductivities of P3HT, PEDOT:PSS, and P3HT/PEO. ......................................... 40
3.12 Potential and current profiles during electrochemical doping preceding the EIS
measurement protocol to determine electronic conductivity. ........................................... 41
3.13 A representative Nyquist impedance plot of P3HT obtained at 3.3 V vs Li
+
/Li using the
two-electrode configuration and the fitting circuit (inset) used to obtain the electronic
conductivity. ...................................................................................................................... 42
3.14 2-electrode Nyquist plots of (a) P3HT/PEO and (b) PEDOT:PSS at 3.3 V vs Li
+
/Li. ..... 43
3.15 Electronic conductivities of P3HT, PEDOT:PSS, and P3HT/PEO. ................................. 44
3.16 Bode plots of 2-electrode EIS measurement of P3HT film (a) between 2.8 and 3.3 V
and (b) between 3.3 and 4 V. ............................................................................................ 44
3.17 Structure of P(NDI2OD-T2). ............................................................................................ 45
3.18 Cyclic voltammetry plot of P(NDI2OD-T2) at the various scan rates indicated. ............. 46
3.19 The electronic and ionic conductivities of P(NDI2OD-T2) as a function of electrode
potential. ............................................................................................................................ 47
3.20 (a) Photograph of the oxide electrode. (b) and (c) SEM images of the electrode. ............ 48
3.21 A schematic of the oxide micro-electrode. ....................................................................... 48
3.22 CV of TiNb2O7 at various scan rates. ................................................................................ 49

ix
3.23 CV data of TiNb2O7 at 20 mV s
−1
. .................................................................................... 49
3.24 2-electrode EIS Nyquist plot of TiNb2O7 at 1.6 V. ........................................................... 50
3.25 Electronic conductivity of TiNb2O7 as a function of potential. ........................................ 51

4.1 Schematic of PProDOT-Hx2 synthesis using DArP. ........................................................ 56
4.2 Initial cycling of PProDOT-Hx2. ...................................................................................... 57
4.3 CV data for PProDOT-Hx2 thin film at various potential intervals at 10 mV s
−1
. ............ 59
4.4 CV data for PProDOT-Hx2 as a function of various scan rates from 10 to 100 mV s
−1
. .. 60
4.5 Log of the peak current (i) vs. log of the scan rate (v) for data shown in Figure 4.4. ...... 61
4.6 Long-term CV curves of PProDOT-Hx2 between 3.0 and 4.0 V at 10 mV s
−1
. ............... 61
4.7 (a) The 3-electrode configuration used to electrochemically dope the polymer and
determine its ionic conductivity. (b) A representative Nyquist impedance plot obtained
at 4 V vs Li/Li
+
using the 3-electrode configuration to obtain the ionic conductivity. .... 62
4.8 Ionic conductivities of P3HT and PProDOT-Hx2. ............................................................ 63
4.9 (a) The 2-electrode configuration used to measure electronic conductivity. (b) A
representative Nyquist impedance plot obtained at 4 V vs Li/Li
+
using the 2-electrode
configuration and the fitting circuit (inset) used to obtain the electronic conductivity. ... 64
4.10 The electronic conductivities of P3HT and PProDOT-Hx2. ............................................. 65
4.11 TEM image of (a) 96-4% NCA-PProDOT-Hx2 and (b) 90-3-3-4% NCA-SP-CNT-
PProDOT-Hx2 electrodes. ................................................................................................. 66
4.12 Rate capability of the NCA-PProDOT-Hx2 and NCA-PVDF. ......................................... 67
4.13 The galvanostatic charge-discharge curves of the (a) NCA-PProDOT-Hx2 and (b)NCA-
PVDF. ............................................................................................................................... 68
4.14 (a) Long-term cycling for NCA-PProDOT-Hx2 and NCA-PVDF at a rate of 2C. The
corresponding galvanostatic charge-discharge curves of the (b) NCA-PProDOT-Hx2
and (c) NCA-PVDF at different cycles. ............................................................................ 69
4.15 Synthesis scheme of the (Hex:OE) PProDOT copolymer family. .................................... 70
4.16 Electrochemical performance of PProDOT-Hx2. (a) Cycling Voltammograms at 10 mV
s
−1
in various potential windows up to 4.5 V, (b) CV scans at different scan rates, and
(c) long-term cycling at 10 mV s
−1
. .................................................................................. 71

x
4.17 CV scans of (Hex:OE) PProDOTs at 5 mV s
−1
. ............................................................... 72
4.18 CV data at 10 mV s
−1
for (a) (65:35) PProDOT, (b) (75:25) PProDOT, (c) (85:15)
PProDOT, and (d) (95:05) PProDOT at various potential intervals. ................................ 73
4.19 CV data at varying scan rates between 100 and 20 mV s
−1
for (a) (65:35) PProDOT, (b)
(75:25) PProDOT, (c) (85:15) PProDOT, and (d) (95:05) PProDOT. ............................. 74
4.20 Long-term CV curves for (a) (65:35) PProDOT, (b) (75:25) PProDOT, (c) (85:15)
PProDOT, and (d) (95:05) PProDOT thin films at 10 mV s
−1
. ......................................... 75
4.21 Capacity retention of (Hex:OE) PProDOT copolymer family after 100 cycles at 10 mV
s
−1
as a function of oligoether content. ............................................................................. 76
4.22 CV curves for (a) oligoether homopolymer, (b) (15:85) PProDOT, (c) (75:25)
PProDOT, and (d) (50:50) PProDOT thin films at 100 mV s
−1
. ....................................... 77
4.23 Electronic conductivities of the PProDOT copolymer series. .......................................... 78
4.24 Ionic conductivities of the PProDOT copolymer series. .................................................. 79
4.25 Swelling study of the (Hex:OE) PProDOTs using propylene carbonate electrolyte. ....... 80
4.26 Specific capacity as a function of cycle number for the Li-NCA-(Hex:OE) PProDOT
cells. .................................................................................................................................. 81
4.27 Schematic of CBS co-polymer family synthesis using DArP. ......................................... 83
4.28 Initial CV data for T6T (a) 5%, (b) 10%, and (c) 20%. .................................................... 83
4.29 Initial CV data for T8T (a) 5%, (b) 10%, and (c) 20%. .................................................... 84
4.30 Initial CV data for T10T (a) 5%, (b) 10%, and (c) 20%. .................................................. 85
4.31 CV data for T6T (a) 5%, (b) 10%, and (c) 20% at various potential windows at 10 mV
s
−1
. ..................................................................................................................................... 86
4.32 CV data for T8T (a) 5%, (b) 10%, and (c) 20% at various potential windows at 10 mV
s
−1
. ..................................................................................................................................... 86
4.33 CV data for T10T (a) 5%, and (b) 10% at various potential windows at 10 mV s
−1
. ....... 87
4.34 CV data for T6T (a) 5%, (b) 10%, and (c) 20% at various scan rates. ............................. 87
4.35 CV data for T8T (a) 5%, (b) 10%, and (c) 20% at various scan rates. ............................. 88
4.36 CV data for T10T (a) 5% and (b) 20% at various scan rates. ........................................... 88
4.37 Long-term CV curves for T6T (a) 5%, (b) 10%, and (c) 20% thin films at 10 mV s
−1
. ... 89

xi
4.38 Long-term CV curves for T8T (a) 5, (b) 10, and (c) 20% thin films at 10 mV s
−1
. .......... 89
4.39 Long-term CV curves for T10T (a) 5%, and (b) 10% thin films at 10 mV s
−1
. ................ 90
4.40 Electronic conductivities of the T6T polymers. ................................................................ 91
4.41 Electronic conductivities of the T8T polymers. ................................................................ 91
4.42 Electronic conductivities of the T10T polymers. .............................................................. 92
4.43 Comparison of electronic conductivity of polymers with 5% CBS content. .................... 92
4.44 (a) Electronic and (b) ionic conductivities of the CBS polymers. .................................... 93
4.45 GIWAXS data of (a) T6T polymers and (b) CBS polymers with 5% CBS content. ........ 94
4.46 Full cell testing of CBS polymers. (a) Specific capacity as a function of cycle number
and (b) rate capability test. ................................................................................................ 96

5.1 Polymers investigated for binders in Lithium-Sulfur (Li-S) Batteries. ............................ 99
5.2 Cyclic voltammograms of initial cycling for N2200 and N2200-OE. .............................. 100
5.3 CV data of N2200 and N2200-OE as a function of scan rate. .......................................... 101
5.4 CV data for long term cycling of N2200 and N2200-OE thin films. ............................... 101
5.5 Cyclic voltammograms of N2200 and N2200-OE at 5 mV s
−1
. ....................................... 102
5.6 (a) Electronic and (b) ionic conductivity of N2200 and N2200-OE as a function of
potential. ............................................................................................................................ 102
5.7 Integration of the GIWAXS diffractograms of (a) the full integration of N2200 and the
inset is of the in-plane integration of the polymer, and (b) the full integration of N2200-
OE and the inset is of the in-plane integration of the polymer. ........................................ 105
5.8 UV-Vis spectra of LiPS exposed to (a) PVDF, (b) N2200, and (c) N2200-OE. (d)
Absorbance at 470 nm of LiPS solution after 3 hours of polymer exposure. ................... 106
5.9 Photographs of LiPS (a) before exposure to polymers and after 3 hours of exposure to
(b) PVDF, (c) N2200, and (d) N2200-OE. ....................................................................... 107
5.10 (a) Potentiostatic intermittent shuttle current measurement curve of a Li-Sulfur-N2200
cell. (b) Voltage profile prior to and during shuttle current measurements on a Li-Sulfur-
N2200 cell. ........................................................................................................................ 108

xii
5.11 Polysulfide shuttle current measurement. a) Potentiostatic observed current as a
function of time at approximately 2.3 V vs Li/Li
+
after an OCV relaxation step and b)
steady-state current or shuttle current as a function of potential during discharge for
sulfur-N2200), sulfur-N2200-OE and sulfur-PVDF electrodes. ....................................... 109
5.12 Steady-state current or shuttle current as a function of potential after 35 cycles during
discharge for a Li-Sulfur-N2200 and Li-Sulfur-PVDF cells. ........................................... 109
5.13 (a) GCD profiles for Li-sulfur-N2200, Li-sulfur-N2200-OE, and Li-sulfur-PVDF cells
at C/20. Potentiostatic impedance response was measured at OCV as a function of SOC
during charge and discharge (denoted with yellow circles). (b) Real part of the
impedance response against frequency for the Li-sulfur-N2200, Li-sulfur-N2200-OE,
and Li-sulfur-PVDF cells at approximately 2.28 V, 2.15 V, and 2.1 V vs Li
+
/Li. ........... 111
5.14 Rate Capability for Li-sulfur-n-dopable polymer cells. Specific capacity as a function
of cycle number with increments in the discharge rate and a constant charging rate at
C/10 for sulfur electrode with (a) moderate loadings and (b) high loadings. ................... 112
5.15 Specific capacity as a function of cycle number at a discharge/charge rate of C/2 for
sulfur-N2200, sulfur-N2200-OE and Sulfur-PVDF. ........................................................ 113
5.16 Impedance response as function of cycle number at approximately 2.2 V vs Li
+
/Li for
the same (a) Li-sulfur-N2200, (b) Li-sulfur-N2200-OE, and (c) Li-sulfur-PVDF cells
for the 15
th
, 200
th
, and 495
th
cycles. .................................................................................. 114
5.17 SEM images of (a) pristine lithium and lithium from cycled (b) (b) PVDF, (c) N2200,
and (d) N2200-OE cells along with photographic images of the electrodes (insets). ....... 115
5.18 EDS spectra of (a) pristine lithium, and lithium from cycled (b) PVDF, (c) N2200, and
(d) N2200-OE cells. .......................................................................................................... 116

6.1 Powder XRD spectra of (a) LaMnO3, (b) LaFeO3, (c) LaCoO3, and (d) LaNiO3. ........... 120
6.2 CS2 annealing setup for LaMS3 sulfurization. .................................................................. 121
6.3 Powder XRD patterns for LaMS3 with representative reference pattern for LaNiS3. ....... 121
6.4 Powder XRD patterns for LaMnS3 annealed for (a) four hours and (b) 28 hours. ........... 122
6.5 Polarization curves for (a) BaTiS3, (b) BaZS3, and (c) BaVS3. ........................................ 123
6.6 Tafel slopes of (a) BaTiS3, (b) BaZS3, and (c) BaVS3. ..................................................... 123
6.7 Electrocatalytic activity of various transition metal perovskite sulfides LaMS3 towards
hydrogen evolution (a) polarization curves and (b) the corresponding Tafel slopes. ....... 125
6.8 OER Polarization curves for barium transition metal chalcogenides. .............................. 126

xiii
6.9 Electrocatalytic activity of LaMS3 materials for oxygen evolution, (a) OER polarization
(b) the corresponding Tafel slopes. ................................................................................... 127
6.10 Polarization curves for oxygen evolution reaction on (a) LaMnS3 and LaMnO3, and (b)
LaCoS3 and LaCoO3. ........................................................................................................ 128

7.1 Steady-state polarization plots of NS oxygen electrodes. ................................................. 132
7.2 The Tafel plots of NS oxygen electrodes. ......................................................................... 133
7.3 SEM images of NS oxygen electrodes. ............................................................................. 134
7.4 XRD data of NS oxygen electrodes. ................................................................................. 135
7.5 SEM images of CS oxygen electrodes. ............................................................................. 137
7.6 XRD data of CS oxygen electrodes. ................................................................................. 138
7.7 Polarization curved of NS-200 and CS-200 oxygen electrodes. ....................................... 139
7.8 XPS data of Ni 2p3/2 for NS-200 and Co 2p3/2 for CS-200 oxygen electrodes. ................ 139
7.9 Electrochemical pre-treatment protocol for the nickel-treated steel electrodes. ............... 140
7.10 Polarization curves of the NS-200 oxygen electrodes with(purple) and without (pink)
pre-treatment. .................................................................................................................... 141
7.11 Electrode capacitance as a function of electrochemical pre-treatment cycles. ................. 142
7.12 SEM images of (a) untreated steel mesh substrate, pre-treated steel mesh substrate after
(b) 1 cycle, (c) 10 cycles, and (d) 20 cycles. ..................................................................... 143
7.13 SEM image of the (a) as-prepared oxygen electrode and (b) the oxygen electrode after
electrochemical testing. ..................................................................................................... 144
7.14 EDS spectrum of the as-prepared oxygen electrode. ........................................................ 145
7.15 XRD spectra of the oxygen electrode before and after electrochemical testing. .............. 146
7.16 Binding energy of Ni(2p) for the as-prepared oxygen electrode. ..................................... 146
7.17 (a) Polarization data of the oxygen electrode and (b) the corresponding Tafel slope. ..... 147
7.18 Electrochemical impedance data of the oxygen electrode (a) as a function of
overpotential and (b) the corresponding EIS data fitting. ................................................. 148
7.19 Log(1/RCT) vs overpotential for the oxygen electrode. ..................................................... 149

xiv
7.20 Chronopotentiometry of the oxygen electrode at 10 mA cm
−2
. ........................................ 149
7.21 SEM images of the hydrogen electrode. (a) and (b) Cross-section and top-view of the
as-prepared Ni-Mo co-sputtered surface. (c) and (d) Cross-section and top-view of the
Ni-Mo co-sputtered surface after electrochemical testing. ............................................... 151
7.22 Reflectivity profile of NiMo co-sputtered layer on glass slide. ........................................ 152
7.23 EDS spectrum of as-prepared hydrogen electrode. ........................................................... 153
7.24 XRD spectra of the hydrogen electrode before and after electrochemical testing. .......... 154
7.25 Binding energy of (a) Ni (2p) and (b) Mo (3d) for the as-prepared hydrogen electrode. . 155
7.26 (a) Polarization data of the HER electrode and (b) the corresponding Tafel slope. ......... 156
7.27 EIS Nyquist plots of the hydrogen electrode as a function of overpotential and (b) the
EIS data fitting. ................................................................................................................. 157
7.28 Log(1/RCT) vs overpotential for the hydrogen electrode. .................................................. 158
7.29 Chronopotentiometry of the hydrogen electrode at 10 mA cm
−2
. ..................................... 159
7.30 Schematic of an all-iron alkaline water electrolyzer: (1) Nickel-molybdenum coated
iron electrode, (2) Zirfon PERL separator, (3) graphite plate with columnar flow field,
(4) heating pad, (5) nickel-coated plate with columnar flow field, (6) gold-coated
current collector, and (7) K-type thermocouple. ............................................................... 160
7.31 (a) Galvanostatic data of the electrolyzer at different temperatures and (b) the
corresponding slopes. ........................................................................................................ 161
7.32 Tafel slopes of the electrolyzer at 30 °C. .......................................................................... 162
7.33 (a) Impedance spectroscopy data of the electrolyzer as a function of overpotential and
(b) EIS data fitting at 35 mA cm
−2
. ................................................................................... 163
7.34 Constant current operation of the electrolyzer at 1 A cm
−2
at room temperature. ............ 163

8.1 Schematic of the general molecular structure and reaction. ............................................. 166
8.2 (a) Synthetic route for the synthesis of diphenyl sulfone bismuth(III) acetylides. (b)
Diphenyl sulfone bismuth(III) acetylides. (c) 1-ethynyl-4-methylbenzene bismuth(III)
acetylides with functionalized diphenyl sulfone ligands. ................................................. 167
8.3 Proposed mechanistic model of azide-bismuth(III) acetylide copper(I)- catalyzed
cycloaddition reaction. ...................................................................................................... 169

xv
8.4 Reaction scheme of kinetic studies of copper(I)-catalyzed reaction of 5-bismuth(III)
triazolides formation with cyclic voltammetry technique. ............................................... 170
8.5 Cyclic voltammogram of the copper(I) triflate catalyst at 25°C in dry DMSO recorded
at 100 mV sec
−1
. ................................................................................................................ 172
8.6 Kinetic profiles of the cyclic voltammograms obtained after the addition of acetylide. .. 173
8.7 Cyclic voltammogram of the copper(I) triflate catalyst at 25°C in dry DMSO before
(blue) and after (orange) the addition of acetylide. ........................................................... 174
8.8 Kinetic profiles of the cyclic voltammograms obtained after the addition of azide. ........ 175
8.9 CV of A[1] before and after the addition of azide. ........................................................... 176
8.10 Reaction scheme of the first step, the π-intermediate complex formation of the copper(I)
catalyst and bismuth(III) acetylide. ................................................................................... 177
8.11 Current change after the addition of (a) bismuth(III) acetylide[1] and (b) the azide. ...... 178
8.12 Current change after the addition of (a) bismuth(III) acetylide[2] and (b) the azide. ...... 178
8.13 Current change after the addition of (a) bismuth(III) acetylide[3] and (b) the azide. ...... 178
8.14 Current change after the addition of (a) bismuth(III) acetylide[4] and (b) the azide. ...... 179
8.15 Current change after the addition of (a) bismuth(III) acetylide[5] and (b) the azide. ...... 179
8.16 Current change after the addition of (a) bismuth(III) acetylide[6] and (b) the azide. ...... 179
8.17 Concentration changes of the Cu
+
/Co
0
redox pair for step 1. ........................................... 182
8.18 CV of [Cu] after the addition of (a) bismuth(III) acetylide and (b) azide. (c) A
comparison of the [Cu] CV before (blue) and after (orange) the addition of the
bismuth(III) acetylide complex and (purple) the addition of azide. ................................. 185

xvi
LIST OF PUBLICATIONS

Zayat, B.; Elizalde-Segovia, R.; Das, P.; Salamat, C.Z.; Irshad,

A.; Tolbert, S.H.; Thompson,

B.C.;
Narayanan S.R.; The Role of Functionalized Conducting Polymer Binders in Improving Power
Density and Cycle Life of Lithium-Sulfur Batteries. Journal of the Electrochemical Society
2022. (Submitted)
Das, P.; Elizalde-Segovia, R.; Zayat, B.; Salamat, C. Z.; Pace, G.; Zhai, K.; Vincent, R. C.; Dunn,
B. S.; Segalman, R. A.; Tolbert, S. H.; Narayan, S. R.; Thompson, B. C., Enhancing the Ionic
Conductivity of Poly(3,4- Propylenedioxythiophenes) with Oligoether Side Chains for Use as
Conductive Cathode Binders in Lithium-Ion Batteries. Chemistry of Materials 2022.
Nazarova, A. L.; Zayat, B.; Fokin, V. V.; Narayan, S. R., Electrochemical Studies of the
Cycloaddition Activity of Bismuth(III) Acetylides Towards Organic Azides under Copper(I)-
Catalyzed Conditions. Frontiers in Chemistry 2022, 10.
Elizalde-Segovia R.; Das P.; Zayat B.; Irshad A.; Narayan S.R.; Thompson B. C., Understanding
the Role of π-Conjugated Polymers as Binders in Enabling Designs for High-Energy/High-Rate
Lithium Metal Batteries. Journal of the Electrochemical Society 2021, 168 (11) , 110541.
Zayat, B.; Das, P.; Thompson, B. C.; Narayan, S. R., In Situ Measurement of Ionic and Electronic
Conductivities of Conductive Polymers as a Function of Electrochemical Doping in Battery
Electrolytes. Journal of Physical Chemistry C 2021, 125 (14), 7533-7541.
Zayat, B.; Mitra, D.; Irshad, A.; Rajan, A. S.; Narayanan, S. R., Inexpensive and Robust Iron-
based Electrode Substrates for Water Electrolysis and Energy Storage. Current Opinion in
Electrochemistry 2021, 25.
Elizalde-Segovia, R.; Irshad, A.; Zayat, B.; Narayanan, S. R., Solid-State Lithium-Sulfur Battery
Based on Composite Electrode and Bi-layer Solid Electrolyte Operable at Room Temperature.
Journal of the Electrochemical Society 2020, 167 (14), 140529.
Zayat, B.; Das, P.; Wei, Q.; Salamat, C. Z.; Magdău, I.-B.; Elizalde-Segovia, R.; Rawlings, D.;
Lee, D.; Pace, G.; Irshad, A.; Ye, L.; Schmitt, A.; Segalman, R. A.; Miller, T. F.; Tolbert, S.
H.; Dunn, B. S.; Narayan, S. R.; Thompson, B. C., Dihexyl-Substituted Poly(3,4-
Propylenedioxythiophene) as a Dual Ionic and Electronic Conductive Cathode Binder for
Lithium-Ion Batteries. Chemistry of Materials 2020, 32 (21), 9176-9189.
Zayat, B.; Mitra, D.; Narayanan, S. R., Inexpensive and Efficient Alkaline Water Electrolyzer
with Robust Steel-Based Electrodes. Journal of the Electrochemical Society 2020, 167, 114513.

xvii
ABSTRACT

Climate change is one of the most difficult challenges we have faced as humans.
Addressing this issue requires a series of structural changes to our economy and way of life.
Electrochemistry is one of the tools that will help us achieve zero emissions due to its ability to
replace many greenhouse emitting processes with sustainable alternatives. For example, electric
motors powered by lithium batteries are replacing many internal combustion engines that rely on
fossil fuels. Hydrogen is increasingly being adopted as a clean fuel replacement as well.
One obstacle to lithium-ion battery (LIB) adoption is long charge times and low cycle life,
an issue that many try to address by using conductive polymers instead of insulating ones as
binders to improve electron transport. We have developed an in-situ technique that allows for the
electrochemical characterization and electronic and ionic conductivity measurement of these
polymers as a function of potential (Chapter 3). In Chapter 4, we demonstrate the use of
bifunctional conducting polymers as binders for LIBs to greatly improve rate capability and cycle
life. In Chapter 5, we utilize n-dopable conducting polymers as binders for the sulfur electrodes in
lithium-sulfur batteries to limit the polysulfide shuttling and significantly enhance cycle life.
In Chapter 6, we probe a new class of transition metal perovskite chalcogenides as
promising electrocatalysts for the splitting of water into hydrogen and oxygen. We demonstrate
the advantages of using perovskite sulfides compared to the more commonly used perovskite
oxides. In Chapter 7, we develop robust and inexpensive iron-based catalysts for the hydrogen and
oxygen evolution reactions. We also demonstrate an efficient all-iron alkaline electrolyzer that
operates at 20 mA cm
−2
. Finally, we use electrochemistry to perform kinetic studies on bismuth
complexes for drug delivery applications in Chapter 8.

1
CHAPTER 1.

BACKGROUND AND MOTIVATION

1.1 The Challenge of Climate Change
Global carbon dioxide emissions have been continuously increasing since the start of the
industrial revolution in the 18
th
century. In fact, the rate of emissions significantly increased
starting in the 1950s due to a dramatic surge in fossil fuel consumption for energy production. This
resulted in almost six times increase in the CO2 emissions between 1950 and 2020 (Figure 1.1a).
1-
3
The massive climb in emissions has been accompanied by an increase in the average global
temperature.
4-7
As Figure 1.1b shows, we observe a phase of intense global warming that started
in the 1950s and significantly increased in the 1980s.

Figure 1.1. (a) Global CO2 emissions between 1880 and 2020.
1
(b) Global temperature deviation
from 1880-1900 averages. Data from Goddard Institute for Space Studies.

To emphasize the increasing rate of global warming, a heat map that shows the annual
global temperature difference between 2020 and the annual global temperature average between
1981-2010 is shown in Figure 1.2. We can see that temperature increases are universal and

2
occurring across the globe but with significant increases of over 2 ℃ in many regions of the
arctic circle. The increased global temperature leads to measurable physical effects such as
increased heat content of the oceans, decreased marine and terrestrial ice and snow surface
coverage, increased droughts, increased frequency of heatwaves and intense rains, and extreme
weather events.
8-15

Figure 1.2. Annual global temperature anomaly (℃) at 2 m height in 2020 vs. 1981-2010 annual
average. Computed with KNMI Climate Explorer.

The warmer temperatures result in the melting of ice caps at the north and south poles
which in turn causes the sea levels to rise. As Figure 1.3 shows, there has been a significant
increase in the global mean sea level since the 1800s. Furthermore, the rate of sea level rise
increased to over 3 mm per year since 1992. Since most of the world population lives near oceans
and seas, a rising sea level would be detrimental to many societies and regional economies

3
resulting in the migration of tens of millions of people and the loss of billions of dollars.
16-19
Global
warming and rising sea levels do not only affect human populations, but can be devastating to
aquatic life as well.
20-22

Figure 1.3. Global mean sea level change vs. 1870-1890. Data from NOAA Laboratory for
Satellite Altimetry.

Unfortunately, not enough is being done to mitigate the negative impacts of global
warming and to reduce greenhouse gas emissions. Figure 1.4 shows the projection of the global
temperature increase up to the year 2100. If we continue our current path (scenario RCP8.5), the
earth will witness a temperature increase of about 5 ℃ compared to 1880-1900 levels. Such
increase will have everlasting and devastating effects on our lives on this planet. However, if we
take immediate actions today to quickly limit greenhouse gases (scenario RCP4.5), we may be
able to reach an emission peak by 2040 followed by a strong decline until 2080. This will result
in less than 3 ℃ gain in global temperature compared to 1880-1900 levels, allowing us to avoid
the most extreme consequences of global warming. As humans, we have no option but to
aggressively limit our carbon dioxide emissions for the sake of our survival on this planet.

4

Figure 1.4. Historic and projected world temperature change with respect to 1880-1900 averages
from The Intergovernmental Panel on Climate Change. Computed with KNMI Climate Explorer.

1.2 Electrochemistry and Sustainability
When we break down greenhouse gas emissions by economic activities that lead to their
production (Figure 1.5), we observe that the burning of coal, natural gas, and oil for electricity and
heat is the largest single source of global greenhouse gas emissions at 25%.
23
Agriculture in the
form of cultivations of crops and livestock contributes to approximately 24% of emissions. In
addition, there are significant emissions (21%) from industry due to chemical, metallurgical, and
mineral transformation processes, and waste management activities.
24
Greenhouse gas emissions
due to the burning of fossil fuels for road, rail, air, and marine transportation contribute to 14% of
the global emissions. Thus, energy production and transportation have an enormous contribution
to global warming, and their negative effect on climate change needs to be mitigated. However,
greenhouse gas emissions are not restricted to energy production and transportation only.
Agriculture and manufacturing are sizeable contributors to global warming and they, in turn, need
to be radically transformed to reign in greenhouse gas emissions. Due to the multiple sources of

5
greenhouse gases and the multitude of sectors and processes that contribute to their emissions,
there is no single simple solution to our climate problem.

Figure 1.5. Global Emissions by Sector in 2010. Data from Fifth Assessment Report of the
Intergovernmental Panel on Climate Change.

Electrochemistry, nonetheless, can provide us with the tools to address climate change
challenges on multiple fronts. The broad application of electrochemistry is because many
processes such as energy storage, energy conversion, and industrial production of chemicals
involves the transfer of electrons. Thus, electrochemistry is an ideal tool to move away from
traditional high-emission processes towards sustainable substitutes.
For example, significant work has been done on capturing carbon dioxide from the
atmosphere and converting it into fuels and value-added products as opposed to extracting non-
renewable fossil fuels for energy production.
25-29
In addition, solar energy is replacing more
traditional energy sources such as coal at a large scale.
30, 31
Electrochemical techniques are also
being used to replace traditional industrial manufacturing techniques. The electrochemical
reduction of nitrogen to ammonia is being investigated as a more sustainable approach to
producing ammonia instead of the more traditional Haber-Bosch process.
32-35
Hydrogen can be
electrochemically produced by the splitting of water in a process that only generates oxygen as a

6
byproduct.
36, 37
The electrochemical pathway is a significantly more sustainable option compared
to the more common steam-methane reforming approach.
38-40
Thus, water electrolysis is a
sustainable pathway towards energy storage and production through the hydrogen economy.
41-43

Furthermore, fuel cells are being increasingly used to replace internal combustion engines, where
the chemical energy of fuels such as hydrogen are electrochemically converted into energy.
44-46

1.3 Lithium Batteries and Renewable Energy
As we move away from fossil fuels and towards renewable energy sources, the need to
store the produced energy becomes more important due to the mismatch between the rates of
production and consumption of electricity. For example, peak production of electricity using solar
energy is around noon, while peak consumption of electricity is in the evening hours. As a result,
it is essential that we have the ability to store and extract energy in accordance with our
consumption levels. Therefore, there has been renewed interest in developing new battery
technology and energy storage systems to meet our demand.
47
One of the most common batteries
today is the lithium-ion battery (LIB) which dominates in the portable devices market which
includes smart phones, tablets, power tools and laptop computers. LIBs are also used in powering
electric vehicles because of their high energy density. In a typical LIB, a transition metal oxide is
used as the cathode, where the transition metal can be cobalt, nickel, or manganese, along with
small amounts of aluminum. The anode, however, is composed of graphite and the electrolyte is
typically a lithium salt such as lithium bis-(trifluoromethanesulfonyl)imide (LiTFSI) or LiPF6 in a
non-aqueous alkyl carbonate solvent. In Chapters 3 and 4, we focus on synthesizing conductive
polymers as binders for LIB cathodes. In this strategy, we attempt to significantly improve the
discharge and charge rate capability of the cells, an area of great interest. Furthermore, we develop

7
a technique that allows for the in-situ electrochemical characterization of these polymers which
results in fast screening of possible polymer binder candidates.
Even with continuous improvements to lithium-ion batteries, they are approaching their
theoretical maximum capacity. For example, the LiCoO2 cathode has an approximate theoretical
capacity of 272 mAh g
−1
. A promising replacement to LIB that has garnered significant attention
is the lithium-sulfur battery that has the potential of storing two to three times the energy of a
LIB.
48, 49
In a lithium-sulfur battery, the cathode is composed of elemental sulfur while the anode
is composed of metallic lithium. 1M LiTFSI in a 1:1 volume ratio of 1,3 dioxolane and 1,2
dimethoxyethane is the electrolyte most commonly used with lithium nitrate as an additive. During
discharge, lithium ions migrate to the cathode where sulfur is reduced to Li2S in a complex reaction
that involves the formation of several polysulfide species (Figure 1.6).

Figure 1.6. A schematic of a lithium-sulfur battery.
The lithium sulfur (Li-S) battery is a promising next generation energy storage technology
to replace the current lithium-ion battery due to the high theoretical specific capacity of the sulfur
electrode at 1672 mAh g
−1
and the high gravimetric energy density of the Li-S cell at 2500 Wh

8
kg
−1
(Figure 1.7).
50, 51
In addition, sulfur is substantially more abundant and less expensive than
the active materials in LiBs such as nickel and cobalt.
52-54

Figure 1.7. Comparison of specific capacity and specific energy between Li-S and LIB.
However, commercial adoption of Li-S batteries have been hindered by several issues
including poor discharge rate capability, low coulombic efficiency, and poor cycle life.
55, 56
These
performance limitations are mainly due to the inherently insulating nature of sulfur, the formation
of soluble polysulfides, and the large volume expansion of the sulfur electrode.
57
To address those
shortcomings, several strategies have been deployed with varying degrees of success in the last
decade aiming at improving the performance of Li-S batteries .
58-62
In Chapter 5, we discuss the
work we have done to limit the polysulfide shuttling and improve the cycle life and rate capability
by using n-dopable conducting polymers as binders for the sulfur electrode.
1.4 Water Electrolysis for Clean Hydrogen Production
The use of transition metals and other expensive materials in lithium batteries hinders their
full-scale adoption for large scale energy storage and heavy transportation applications such as
cargo ships and long-range trains. Hydrogen, on the other hand, is an attractive alternative as a fuel
due to its high specific energy of 33.6 kWh kg
−1
. Hydrogen production by the electrolysis of water

9
using renewable energy is a sustainable approach to produce clean hydrogen instead of methane
steam reforming.
36, 63-65
In Chapter 6, we probe transition metal perovskite chalcogenides as
promising materials for electrochemical water splitting for their performance as electrocatalysts
for both hydrogen evolution reaction (HER) and oxygen evolution reaction (OER).
Water electrolysis may be carried out under either acidic or alkaline conditions. Alkaline
water electrolysis allows the use of less expensive catalysts for HER and OER.
66, 67
In addition,
alkaline systems have been in operation for over 100 years and their performance can rival that of
the acidic electrolyzers.
68
For large-scale adoption, the water electrolyzers must use low-cost and
abundantly-available materials, so that the approach is economically competitive and sustainable.
However, over 50% of capital costs for alkaline water electrolysis systems is attributed to the
electrolysis stacks.
69
Further, at least 50% of the stack costs are governed by the cost of electrode
materials, chiefly made of nickel. While there has been much focus on improving the energy
efficiency of water electrolysis through improved electrocatalysts for the hydrogen and oxygen
evolution reaction,
70-73
replacing expensive nickel-based electrode substrates has received very
little attention.
74-76
Thus, going beyond preliminary studies, we focus on developing a scalable
alkaline electrolyzer that is both robust and inexpensive in Chapter 7.
1.5 Other Applications of Electrochemistry
Going beyond energy storage and energy conversion, electrochemistry was used to perform
kinetics studies on the cycloaddition activity of bismuth(III)acetylides towards organic azides
under copper(I)-catalyzed conditions. In that context, bismuth-doped systems have become
important in the area of near-infrared (NIR)-emitters and drug-delivery materials. In Chapter 8,
we discuss the use of cyclic voltammetry to derive kinetic parameters of various synthesized
ligands which exploit bismuth’s coordination chemistry.

10
CHAPTER 2.

EXPERIMENTAL TECHNIQUES

2.1 Cycling and Measuring the Conductivity of Polymer Binders
2.1.1 Electrode Preparation
Interdigitated microelectrodes (IDM) were purchased from Metrohm Dropsens (DRP-G-
IDEAU5-U20) and were rinsed with isopropanol and dried under argon before use. The electrode
is composed of two interdigitated gold electrodes with two connection tracks on a glass substrate
(L 22.8 × W 7.6 × H 0.7 mm). Each microelectrode is composed of 250 pairs of digits with a digit
length of 6760 µm, a height of 200 nm, and a gap of 5 µm between the digits.
Poly(3-hexylthiophene) (P3HT) (Mw =85-100 kDa) and poly(ethylene oxide) PEO
(Mw=100 kDa) were purchased in powder form from Sigma-Aldrich. poly(3,4-
ethylenedioxythiophene) polystyrene sulfonate (PEDOT:PSS) (Clevios PH 1000) aqueous
solution was purchased from Heraeus and used as received. Poly[N,N′-bis(2-octyldodecyl)-
naphthalene-1,4,5,8-bis(dicarboximide)-2,6-diyl]-alt-5,5′-(2,2′-bithiophene) (P(NDI2OD-T2))
was synthesized as described in-house. Polymer solutions (20 mg mL
-1
) of P3HT, PEO, and
P(NDI2OD-T2) were prepared by dissolving the polymer in 1,2-dichlorobenzene (99%, Sigma-
Aldrich). The solution was then heated at 40°C under argon for two hours to ensure complete
dissolution of the polymer in the solvent. To prepare the P3HT/PEO sample, equal amounts of the
solution of P3HT and PEO were mixed for 24 hours under argon.
5 μL of the prepared solutions were spin-coated on the gold IDM at 1000 RPM for 30
seconds to produce a 50 nm polymer film. The prepared electrodes were then annealed under

11
vacuum at 110℃ for two hours then transferred to a nitrogen glovebox. At least three IDMs were
prepared for each polymer sample.
2.1.2 Electrochemical Doping
All electrochemical tests were performed in an argon glovebox at room temperature. A 3-
electrode cell was assembled in an argon glovebox to electrochemically dope the polymer and
determine its ionic conductivity. The two terminals of the interdigitated gold electrodes were
shorted to form the working electrode. Pieces of lithium foil were used as the counter and reference
electrodes. 1 M solution of bis(trifluoromethane)sulfonimide lithium salt (LiTFSI) in a mixture of
1:1 by volume of ethylene carbonate and dimethyl carbonate (EC/DMC) was used as the
electrolyte. To electrochemically dope the polymer, the potential of the working electrode was
held for 300 seconds. Following this, we allowed adequate time for the electrode potential to relax
to a stable open circuit value. We attribute this relaxation of the open circuit potential to the
equilibration of ion-concentration gradients within the polymer following doping. Any de-doping
that occurs may also be part of this relaxation process. Such an equilibration allows us to achieve
a uniformly doped film. Under these conditions, we could associate the level of doping to the stable
open circuit potential. Following impedance measurements, the applied potential for doping was
then stepped up to the next value. The potential range of the doping varied with the polymer
studied.
2.1.3 Cyclic Voltammetry
Cyclic voltammetry (CV) scans of the polymer films were performed to identify the
oxidation and reduction processes, and the potential range associated with electrochemical doping.
The IDM was held in the electrochemical doping configuration, and the electrode potential was
scanned repeatedly in a cyclic fashion at a preset scan rate in the range of 1 to 100 mV s
−1
over a

12
chosen window of electrode potential based on the polymer type. The current response was
recorded as a function of the electrode potential, and the potentials corresponding to the current
peaks were noted.
2.1.4 Electrochemical Impedance Measurements
To determine ionic conductivity, the impedance of the electrode was measured as a
function of frequency (100 kHz to 100 mHz) at a sinusoidal excitation of ±10 mV peak-to-peak
for each value of electrode potential reached after progressive doping. The potentiostatic EIS
measurement was repeated at different equilibrated values of electrode potential to determine the
change in ionic conductivity as a function of degree of doping.
To measure the electronic conductivity accurately, the impedance measurement was
carried out between the terminals of the two interdigitated gold electrodes. Electrochemical doping
of the polymers was performed in the 3-electrode configuration as before and the cell was allowed
to relax for 100 seconds to reach an equilibrium value. Following the doping process, the electrode
connections were switched from the 3-electrode configuration to a 2-electrode configuration. EIS
measurement was then performed at open circuit potential between 100 mHz and 100 kHz at a
sinusoidal excitation of ± 10 mV peak-to-peak.
2.2 Coin Cell Fabrication and Cycling
2.2.1 Lithium-ion Cells
To prepare the cathode slurries for Li-ion batteries, 90 mg of polyvinylidene fluoride
(PVDF, Sigma-Aldrich) was slowly added to 1 mL of N-methyl-2-pyrrolidone (NMP, Sigma-
Aldrich), and the solution was stirred for 30 minutes until complete dissolution of the binder was
achieved. 2.03g of LiNi0.8Co0.15Al0.05O2 (NCA, Quallion Corp., Sylmar, CA), 67.5 mg of Super P
carbon black (Sigma-Aldrich), and 67.5 mg of multiwalled carbon nanotubes (CNTs, OD×ID×L:

13
10 nm × 4.5 nm × 3-6 μm) were combined and manually mixed with a mortar and pestle. The
powder mixture was then added to the binder solution and 1 mL of NMP was then added. The
solution was allowed to stir for two hours. The resulting slurry was poured on aluminum foil, and
a doctor blade was used at a thickness setting of 8 μm to form an electrode coating. The coating
was then dried for 12 hours at 70 ℃ after which electrode disks 14 mm in diameter were punched
out of the coating and weighed individually. Electrodes were then dried in a vacuum oven at 110
℃ for 12 hours before being loaded into an argon glovebox.
The weight ratio of 90:3:3:4 of NCA:Super P:CNT:PVDF was used unless otherwise
stated. To prepare electrodes with conducting binders which are not soluble in NMP, 20 mg mL
−1

solution of the polymer in o-dichlorobenzene (ODCB) was prepared before mixing with NCA and
conducting carbon. To accurately determine the weight of the electrode, five 14-mm disks of
aluminum foil were weighed and their average mass was subtracted from the total weight of the
electrode.
CR2032 coin-type cells were fabricated with NCA electrodes as the working electrode,
metallic lithium (MTI, D×T :16 mm × 0,6 mm) as counter, and Celgard 2325 (PP/PE/PP) as
separator. Conventional carbonate electrolyte 1.0 M Lithium bis(trifluoromethanesulfonyl)imide
in 1:1 in ethylene carbonate and diethyl carbonate, (LiTFSI EC/DEC =50/50 (v/v), Sigma-Aldrich)
was utilized. The approximate electrolyte amount was ≈ 50 µL. Cells were crimped with a pressure
controlled electric crimper machine (MSK-160E, MTI) between 0.9-1.0 metric tons. Cell assembly
was performed in an argon-filled glove box (VAC system 60387, NexGen 2P, Vacuum
Atmospheres Company) with less than 0.5 ppm of moisture and 0.2 ppm of oxygen.
Battery test station BCS-815 series potentiostat/galvanostat with electrochemical
impedance spectroscopy (EIS) (BioLogic) was employed for the galvanostatic charge-discharge

14
(GCD) cycling, rate capability tests, and EIS measurements of all coin cells. 4 hours equilibration
time at open circuit voltage were imposed to all cells before testing. C-rate was defined based on
the reversible capacity of NCA at 1C that corresponds to 160 mA g
−1
.Two formation cycles at
C/10 charge-discharge were carried before cycling for all cells. For the true discharge rate
capability tests, the cells were charged at C/5 and discharged in increments of C/5, C/2, 1C, 2C,
4C, 6C, and C/5 for five cycles at each rate. For the symmetric rate capability tests, the cells were
charged-discharged for 5 cycles at the rates of C/5, C/2, 1C, 2C, 4C, 6C, and C/5. Cycle-life cells
were charged-discharged at a constant rate of 1C unless otherwise mentioned. The charging
potential cut-off for all cells was imposed at 4.2 V vs Li/Li
+
while the discharge cut-off was at 2.7
V. The EIS response was measured at open circuit voltage in the frequency range of 0.1 Hz to 100
kHz using a sinusoidal excitation amplitude of ± 5 mV (peak to peak). Specific capacity (mAh g
−1
)
was based on the weight of the active material (NCA). All testing was done at room temperature
(21 °C).
2.2.2 Lithium-Sulfur Cells
Sulfur electrodes with PVDF binder where prepared by dissolving PVDF (MTI) in N-
methyl-2-pyrrolidone (Sigma-Aldrich) (0.04 g mL
−1
) then mixing elemental sulfur (Aldrich,
99.5% purity), Ketjen Black (KB, Akzo Nobel EC-600JD), Super P (MTI), and the polymer
solution in a weight ratio of 70:15:10:5. First, sulfur was melt infused into KB at 155 °C for 12
hours. Then, Super P was added and the powders were gently mixed in a mortar and pestle. The
resulting slurry was coated onto aluminum foil utilizing the doctor blade technique, vacuum-dried
at 70°C overnight, and then punched into 14 mm diameter disks.
N2200 and N2200-OE containing sulfur electrodes were prepared in the same weight ratio
as sulfur-PVDF cells. However, the sulfur, Ketjen Black, Super P mixture was added to the ODCB

15
polymer solution (40 g L
−1
) to form a slurry The approximate sulfur areal loading of moderate and
high loading electrodes was 1.8 ± 0.3 and 3.4 ± 0.3 mg cm
−2
, respectively.
2032-type coin-cells were assembled in argon-filled glove box (VAC system 60387,
NexGen 2P, Vacuum Atmospheres Company) with less than 0.5 ppm of moisture and 0.2 ppm of
oxygen. Metallic lithium (MTI, D × T:16 mm × 0.6 mm) was utilized as the counter electrode and
Celgard 2325 (PP/PE/PP) as the separator. 1M LiTFSI in DOL/DME was utilized in an electrolyte
to sulfur ratio of 12 µL mg
−1
. The electrolyte did not contain any electrolyte additive. Cells were
crimped with a pressure-controlled electric crimping machine (MSK-160E, MTI) between 0.9–1.0
metric tons.
Galvanostatic charge/ discharge (GCD) cycling, rate capability testing, polysulfide shuttle
current, and EIS measurements were carried in a battery test station BCS-815 series
potentiostat/galvanostat with EIS (BioLogic) at room temperature (21 °C). An equilibration time
of 24 hours at OCV was imposed on all the cells before cycling. Two “formation” cycles were
carried at C/15. C-rate definition is based on the theoretical capacity of sulfur where 1C = 1675
mA g
−1
. The charging and discharging potential cut-off were set at 2.8 V and 1.55 V vs Li/Li
+
,
respectively. Shuttle current measurements were performed by first discharging the cell at C/20
for 1 hour and then allowing it to rest at open circuit voltage (OCV) with a voltage change cutoff
of 0.1 mV h
−1
. Then, the cell was held at the measured potential of OCV and the current applied
to the cell was observed for 1 hour during which the current reached a steady-state value. This
steady-state current was recorded as the shuttle current and was measured throughout the potential
window. The EIS response of the full cell was measured at open circuit voltage in the frequency
range of 0.1 Hz to 100 kHz using a sinusoidal excitation amplitude of ±4 mV (peak to peak). True-
discharge rate capability testing, for the moderate loading electrodes, was carried out by charging

16
the cells at 0.3 mA cm
−2
(C/10) and discharging them in increments of 0.2, 0.3, 0.6, 1.5, 3, and 6
mA cm
−2
that correspond to C/15, C/10, C/5, C/2, 1C and 2C, respectively. High loading electrodes
were charged at 0.57 mA cm
−2
(C/10) and discharged in increments of 0.29, 0.57, 1.14, 2.85, 5.7,
and 11.4 mA cm
−2
that correspond to C/15, C/10, C/5, C/2, 1C and 2C respectively. Cycle-life
characterization was carried at a constant charge/discharge rate of C/2 for 500 cycles. Specific
capacity was calculated based on the weight of sulfur since the capacity contribution from the
conductive polymer is negligible.
2.3 Voltammetry at Rotating Disk Electrodes
Rotating disk electrode (RDE) experiments were used to electrochemically characterize
powders as electrocatalysts for the oxygen evolution reaction (OER) and the hydrogen evolution
reaction (HER). First, the materials were ground into a fine powder using a mortar and pestle. 10
mg of the powder catalyst were then added to 1 mL solution of dimethyl sulfoxide (DMSO) along
with 10 μL of 5% by weight Nafion binder solution. For HER experiments, carbon black was
added as a conductive additive to improve the kinetics of the reaction. The ink was then sonicated
for three hours in 30-minute increments. To prepare the electrode, 10 μL of the ink was drop-cast
on either a polished gold electrode for OER experiments or polished platinum electrode for HER
experiments. The electrode was then placed in a furnace and heated at 80 ℃ for 30 minutes in air.
This coating process was repeated two more times for a total catalyst loading of 1.0 mg cm
−2
. It is
important to note that adequate loading levels are essential to this experiment: loadings should be
high enough to completely cover the surface (Figure 2.1) but not excessively high so not to affect
conductivity (especially if no conductive additive is used) and the electrochemical performance of
the materials.

17

Figure 2.1. A representative ink sample and the resulting prepared electrode.
To test the electrochemical performance of the catalyst materials, the electrode was then
placed in an RDE setup (Pine Research) where it was rotated at 1000 RPM. Electrochemical tests
were performed in a three-electrode setup (Figure 2.2). For OER, experiments were done in 1 M
potassium hydroxide solution with a nickel counter electrode and a mercury/mercuric oxide
(MMO) reference electrode filled with 1 M KOH solution. HER experiments were performed in
0.5 M sulfuric acid using platinum wire as counter electrode and mercury/mercuric sulfate
(MSE) as reference electrode.

Figure 2.2. Experimental setup of RDE experiments for OER.
2.4 Electrode Fabrication and Testing in Alkaline Conditions
2.4.1 Low-Carbon Steel Mesh as a Substrate
Electrodes were prepared from commercial low-carbon steel mesh (Mcmaster-Carr, wire
diameter of 0.2 mm, open area of 31%, opening size of 0.23 mm, and a mesh size of 60´60).

18
Each of these electrodes were cut to be 5 cm ´ 5 cm. The cut mesh was rinsed with acetone to
remove any grease and then dipped in 0.1M sulfuric acid for one minute at room temperature to
remove surface oxides. The mesh was sonicated in acetone for 15 minutes and then dried under
vacuum at 55 °C for 15 minutes.
2.4.2 Oxygen Electrode Pre-treatment
In certain experiments, the steel mesh substrate was electrochemically pre-treated before
applying the OER electrocatalyst layer. To increase the surface area of the substrate, the steel mesh
was polarized in a solution of 30% potassium hydroxide (5.35 M) in the presence of 0.5 g sodium
sulfide per liter of solution per gram of steel. At first, the steel mesh was polarized for 2 hours at
a constant reduction current of 10 mA g
−1
of steel. Under these conditions, the potential of the
electrode was -1.2 V vs. MMO, while hydrogen evolved at the electrode. Following the cathodic
polarization, the mesh was polarized anodically at a constant current of 1 mA g
−1
of steel during
which the potential of the electrode was about -0.75 V vs. MMO until a cutoff voltage of -0.7 V
vs. MMO was reached. Under these conditions, the conversion of iron to iron (II) hydroxide
occurred.
2.4.3 Formation of the Oxygen Electrode Catalytic Layer
To form the electrocatalyst layer of the oxygen electrode, the steel substrate was placed in
either a 0.08 M solution of nickel nitrate or cobalt nitrate for 15 minutes accompanied by
sonication, following which the substrate was heated to 150
o
C to form an evaporated layer of
nickel nitrate. The electrode was dried and then annealed at the target temperature in air for 30
minutes to decompose the coating of the precursor.
2.4.4 Fabrication of the Hydrogen Electrode
The steel substrate for the hydrogen electrode was subjected to the same cleaning procedure

19
used for the oxygen electrode. To prepare the hydrogen electrode, nickel and molybdenum were
co-sputtered on the steel mesh substrate using direct current (DC) magnetron sputter deposition
(Denton Vacuum Explorer 20) with a pre-sputtering time of 5 minutes to clean the target surface
and a sputtering time of 40 minutes. DC co-sputtering was carried out under 6.5 mTorr of argon
with a nickel cathode power of 80 W and molybdenum cathode power of 15 W resulting in a total
metal loading of approximately 0.6 mg cm
−2
.
2.4.5 Electrochemical Characterization of Electrodes
Electrochemical characterization of individual electrodes was performed in three-electrode
cells (Figure 2.3). The working electrode was a 5 cm ´ 5 cm catalyzed steel electrode with nickel
mesh counter electrodes. A mercury-mercuric oxide electrode in 30% potassium hydroxide was
used as the reference electrode in all the experiments. The potential of the reversible hydrogen
electrode (RHE) vs. MMO was calculated according to Equation 2.1. The area used in the current
density calculations was the geometric area of both sides of the working electrode after subtracting
the open area of the mesh.

Figure 2.3. Schematic of the three-electrode cell used for electrochemical testing in
alkaline conditions.

20
Electrode potentials measured against MMO were converted to potentials versus the
Reversible Hydrogen Electrode (RHE) using equation 2.1 where ERHE refers to the potential of the
electrode vs. the reference hydrogen electrode, EMMO is the potential vs. the Mercury/Mercuric
oxide reference electrode (MMO) and aOH- is the activity of hydroxide ions in the electrolyte.
77
In
30% KOH, RHE potential vs MMO was calculated to be −0.932 V.
!
!"#
=!
$$%
+0.098 V+0.828V+0.059log(0
%"
!) (2.1)
The electrocatalytic activity of the oxygen and hydrogen electrodes was determined by steady-
state potentiostatic polarization experiments at room temperature with a 15-minute hold at each
potential using a multi-channel potentiostat (AMETEK Scientific Instruments, VersaSTAT 4).
Potentiostatic data for the hydrogen electrode was collected between -0.8 V and -1.4 V vs. MMO
in 50 mV increments while potentiostatic data for the oxygen electrode was collected between 0.3
V and 1.0 V vs. MMO in 50 mV increments. The measured potential was corrected for the drop
across the uncompensated solution resistance using the solution resistance value obtained from the
electrochemical impedance measured at 10 kHz. The Tafel slope was obtained by plotting the
potentiostatic data from -1.1 V to -1.4 V (vs. MMO) for the hydrogen electrode and from 0.5 V
to 0.8 V (vs. MMO) for the oxygen electrode. These potential windows corresponded to the current
density regime where the formation of gas bubbles did not impact the measurement. Furthermore,
electrochemical impedance at various values of electrode potential was measured in the frequency
range of 10 kHz to 100 mHz with a sinusoidal voltage excitation of 2 mV peak-to-peak. To
determine the durability of the electrode, three sets of potentiostatic experiments were performed
(each run conducted over 8 hours) followed by 24 hours of open circuit in 30% potassium
hydroxide between each potentiostatic run. The electrode was then placed in 30% potassium
hydroxide for one week. Finally, the electrochemical stability of the electrode was tested by

21
passing a constant current at 10 mA cm
−2
for 100 hours and observing the changes in electrode
potential.
To determine the double layer capacitance of the hydrogen and oxygen electrodes,
electrochemical impedance spectroscopy was collected at open circuit and the resulting Nyquist
plot was fitted to a modified Randles-type equivalent circuit that includes the constant phase
element to represent the distributed capacitance arising from the high-surface area of the electrodes
as shown in section 7.4. The fitted constant phase element (CPE) was then used to calculate the
double layer capacitance using equation 2.1. CDL is the double-layer capacitance in farad, CPE is
the constant phase element, and n is the unit-less exponent (0 < n < 1). Rs and Rct refer to solution
resistance and charge-transfer resistance, respectively.
3
&'
=
⎝
⎛
36!
7
1
8
(
+
1
8
)*
9
+,-
⎠
⎞
+
-
(2.1)

2.4.6 Electrolyzer Assembly
An electrolyzer was assembled using the 5 cm x 5 cm steel mesh-based hydrogen and
oxygen electrodes. A commercially available special porous polymeric material stable in alkali
with low gas permeability (ZIRFON PERL, Agfa Specialty Products) was used as the separator.
This separator material consisted of an open mesh polyphenylene sulfide fabric that is coated with
a mixture of a polysulfone and zirconium oxide. A solution of 30% potassium hydroxide was used
as the electrolyte. A nickel plate with a columnar flow field was used as the current collector on
the oxygen electrode, and a graphite plate with a columnar flow field was used on the side of the
hydrogen electrode. Centrifugal pumps (24W BC-2CP-MD, March Pumps) were used to circulate
the electrolyte through the electrolyzer at a rate of 1L minute
−1
. The cell was heated using two 40-
watt heating pads (2 in x 2 in) attached on to the current collector plates. The temperature of the

22
cell and electrolyte was controlled using a temperature controller and a K-type thermocouple
(92000-00, Cole-Parmer Instrument Co.) placed at the cell.
2.4.7 Electrochemical Characterization of Electrolyzer
All electrochemical tests on the assembled electrolyzer were performed using a VersaSTAT
single-channel potentiostat (AMETEK Scientific Instruments) with a single channel power
booster (Princeton Applied Research, 20 Amp max). The electrochemical performance of the
electrolyzer was determined by galvanostatic polarization studies at room temperature between
100 mA cm
−2
and 1 A cm
−2
in 100 mA cm
−2
increments with a 15-minute hold at each current
density value. These galvanostatic measurements were repeated at 45 °C, 60 °C, and 70 °C. In
addition, the electrochemical impedance of the electrolyzer was measured under the passage of
direct current. The direct current was held for 15 minutes over the range of 15 to 40 mA cm
−2

with the current stepped up in 5 mA cm
−2
increments until a steady state was reached at each
current density. Then, the electrochemical impedance was measured at each steady-state cell
voltage under the passage of direct current. Finally, the electrochemical stability of the cell was
examined under galvanostatic polarization at 1 A cm
−2
for 100 hours at room temperature.
2.5 Physical Characterization Methods
A Nova NanoSEM 450 was used to obtain Scanning electron microscopy (SEM) images of
substrates and electrodes to examine their morphology. Energy dispersive X-ray Spectroscopy
(EDS) was performed using JSM 7001F to determine the atomic content and ratios on the surface
of electrodes. X-ray diffraction (XRD) spectroscopy was performed with a Rigaku Ultima IV
diffractometer using a Cu K-alpha source to investigate the crystallinity and phase composition of
the substrates and catalyst. The surface oxidation states were examined using X-ray photoelectron
spectroscopy (Kratos Axis Ultra DLD) using an aluminum K-alpha source (1486.6 eV). X-ray
reflectivity measurement using a Rigaku Ultima IV diffractometer was utilized to determine the
thickness of the co-sputtered nickel-molybdenum layer for the hydrogen electrode in Chapter 7
using a glass slide with the catalytic thin layer deposited under the same sputtering conditions as

23
the electrodes.
2.6 Cyclic Voltammetric Kinetic Studies
2.6.1 General Experimental Details
Anhydrous DMSO was purchased from MilliporeSigma and used without further
purification. Copper(I) trifluoromethanesulfonate toluene complex (≥99.7%) was purchased from
Sigma-Aldrich. All glassware was oven-dried (140°C) and purged by vacuum-N2 cycles in the
antechamber of the glovebox before use. All electrochemical measurements were performed inside
a glovebox under a nitrogen atmosphere. Anhydrous DMSO served as the solvent. A 3-electrode
glass cell with a glassy carbon working electrode (BASi, MF-2012), copper foil as the counter and
reference electrodes, and copper(I) triflate in anhydrous DMSO) was used for the measurements.
Electrochemical data was collected on a multi-channel potentiostat (AMETEK Scientific
Instruments, VersaSTAT 4). Voltammogram at a scan rate of 100 mV s
−1
was performed on the
copper compound between −0.3 V and 0.8 V vs. Cu
+
/Cu
0
. Kinetic cyclic voltammetry
measurements were processed as current versus time dependences. Graphical interpretation of
processed experimental data was performed in Origin 9.0 data analysis and graphing software
package (OriginLab Corporation) or Microsoft Excel 2016. Two distinct reaction steps were
examined. The first step involves the formation of the bismuth(III)-acetylide – copper(I) complex,
and the second step involves the formation of the bismuth(III) triazolide complex.
2.6.2 Determination of the Rate Constant of Bismuth(III)-Acetylide – Copper(I) Complex
Formation
As shown in Figure 2.4, the first part of the reaction mechanistic study involved preparing
stock solutions of the corresponding bismuth(III) acetylides (0.075 M), and copper(I) triflate
(0.016 M) in anhydrous DMSO. 0.5 mL of the copper(I) triflate toluene complex stock solution

24
and 6.0 mL of DMSO were transferred into the 3-electrode cell, agitated, and used to performed
cyclic voltammetry studies of three redox cycles (to ensure reproducibility of the data and the
stability of copper(I) triflate complex). Six acetlyde compounds were examined and labeled A[1]
through A[6]. Details about the examined compounds can be found in section 8.2.

Figure 2.4. Reaction scheme of the copper(I) complex formation.
Kinetic CV measurements were recorded in the same 3-electrode cell with a glassy carbon
working electrode, copper foils as counter, and reference electrodes. Bismuth(III) acetylide[X]
stock solution in anhydrous DMSO (0.5 mL) was introduced into the electrochemical cell and
agitated for some time before the start of the electrochemistry data acquisition. Continuous
electrochemical data collection at 100 mV s
−1
was performed until no further changes in the redox
cycles were observed.
2.6.3 Determination of the Apparent Rate Constant of the Bismuth(III) Triazolide[X] Formation
Figure 2.5 shows the second step of the overall reaction in which the bismuth(III)
triazolide[X] compound is formed. A solution of the (2-azidoethyl)benzene (0.135M) in anhydrous
DMSO (0.5 mL) was introduced to an electrochemical cell with already preformed π-intermediate
(section 2.6.2). The mixture was agitated for some time before the start of the electrochemistry
data acquisition.

25

Figure 2.5. Reaction scheme of bismuth(III) triazolide[X] formation.
Kinetic CV experiments were recorded in a 3-electrode cell with a glassy carbon working
electrode, copper foils as counter and reference electrodes. Continuous electrochemical
measurements were recorded at 100 mV s
−1
between −0.3 V and 0.8 V vs. Cu
+
/Cu until no further
changes in the redox cycles were observed. An aliquot sample was taken and analyzed by
1
H NMR
to confirm the conversion rate of the reaction.

26
CHAPTER 3.

DEVELOPING AN IN-SITU CONDUCTIVITY MEASUREMENT
TECHNIQUE FOR POLYMER BINDERS AND ELECTRODES
Reprinted with permission from The Journal of Physical Chemistry C 2021 125 (14), 7533-7541.
Copyright 2021 American Chemical Society.

3.1 Introduction
3.1.1 Conductive Polymers as Multifunctional Binders
Lithium-ion batteries have become the dominant energy storage technology for portable
devices and electric vehicles due to the relatively high specific energy and power density of these
batteries
78-80
. The electrodes in these batteries employ traditionally, mechanical binders such as
polyvinylidene fluoride (PVDF) that are electrical insulators. Due to their insulating nature, such
binders tend to hinder charge transport in the electrodes leading to a reduced performance.
81-83
A
conducting polymer that is in the doped (or conducting) state over the operating potential range of
the cathode or anode can confer the following benefits: (1) a thin film of coating of the conducting
polymer on the cathode or anode particles will increase the electrical inter-connectivity between
the electrode particles and thus increase the utilization of the material at high rates of discharge
and charge, (2) the amount of carbon additive used in the electrodes can be reduced to realize a
higher energy density, (3) if the conducting polymers can be made elastic, then volume changes in
the anode and cathode can be accommodated without loss of electrical interconnectivity. A
schematic of a battery electrode is shown in Figure 3.1.

27

Figure 3.1. Schematic of a typical battery electrode.
While planar, rigid aromatic units of conjugated polymers are essential for charge
transport, their soft nature is responsible for mixed electronic and ionic conductivity, and swelling
in solvating environments, advantageous to battery electrodes. Consequently, conductive polymers
have been extensively explored as binders for both low and high-voltage electrodes in lithium-ion
batteries (LIBs) due to their superior ion and electron transport properties.
84-101
Due to their
significant impact on performance, conducting polymers have also been investigated in lithium-
sulfur batteries, supercapacitors, and pseudocapacitors.
102-106
Furthermore, nanostructured
conducting polymers have garnered much attention as active materials in organic electrodes.
107

Thin films of neutral conjugated polymers such as poly(3-hexylthiophene) (P3HT) and
poly[N,N′-bis(2-octyldodecyl)-naphthalene-1,4,5,8-bis(dicarboximide)-2,6-diyl]-alt-5,5′-(2,2′-
bithiophene) (P(NDI2OD-T2)) are rendered electronically conductive when electro-oxidation or
electro-reduction generates mobile charge carriers, typically described as polarons and bipolarons.
Simultaneously, counter ions are introduced into the polymer structure for ensuring charge
compensation.
108, 109
This process by which the polymer becomes electrically conducting is often

28
referred to as “electrochemical doping”. Electro-oxidation and electro-reduction produce “p-
doped” polymers or “n-doped” polymers, respectively. The doping process results in remarkable
changes in electronic conductivity spanning several orders of magnitude. For example,
electrochemical doping increases the electronic conductivity of poly(3-hexylthiophene)-
poly(ethylene oxide) (P3HT-PEO) block copolymer from 10
−8
to 10
−2
S cm
−1
.
110
Similarly, the
structural and morphological changes induced by electrochemical doping can be expected to alter
the ionic conductivity in these polymers. Electrochemical doping also introduces anions/cations
and solvent that can induce swelling of the polymer, leading to substantial changes to the ionic
conductivity.
111
When the conductive polymer is part of a battery electrode, the changes in
electronic and ionic conductivity occur simultaneously with the charging and discharging of the
electrode. The degree of doping and the number of charge carriers is controlled by holding the
electrode potential at various values at which electro-oxidation or electro-reduction occur. We
achieve optimal battery performance when the polymer remains in an electrically conductive state
and is chemically stable over the entire operating potential window of the battery electrode.
Although changes to the ionic conductivity are also expected, the effect of electrochemical
doping is relatively unexplored. Such ionic conductivity changes are also very likely to influence
the observed battery performance. Thus, measuring the electronic and ionic conductivity of the
polymer thin films in the relevant battery electrolyte as a function of potential (i.e. electrochemical
doping) will not only help explain the observed changes in battery electrode performance,
112
but
will also provide fundamental insights for the rational design of the next generation of conducting
polymer binders and additives for various electrochemical devices.
110
The focus of the present
study is on developing and demonstrating a reliable experimental method for such in situ
measurements that combine uniform electrochemical doping in battery electrolyte and practical

29
considerations of using thin polymer films with electronic and ionic conductivities that vary over
several orders of magnitude.
3.1.2 Considerations for Measuring Electronic and Ionic Conductivity
There are several considerations for measuring the electrical conductivity of conductive
polymers for battery applications: (1) The carbon additive usually present in battery electrodes
masks the electronic conductivity contribution from the polymer binder necessitating a carbon-
free electrode for studying just the polymer’s properties. (2) We must be able to vary the electrode
potential over the specific operating potential window of the battery electrode. For example, this
window is between 3 to 4.5 V vs Li
+
/Li for lithium-ion intercalating cathodes and 0 to 1.5 V for
lithium metal and lithium-ion intercalating anodes. (3) Measurements must be made in the relevant
solvent and electrolyte so as to include the effect of swelling by the solvent, well known to alter
the polymer’s properties.
113, 114
and (4) uniform electrochemical doping relies on the diffusion and
migration of dopant ions through a dense film. Thus, a dense film produced by spin-coating must
be usually a few tens of nanometers thick for rapid and uniform electrochemical doping. However,
with electrochemical deposited layers, a porous film may be produced in which ion-transport can
be facilitated even at a thickness of a micron. A routine four-probe conductivity measurement on
relatively thick films on an insulating planar substrate (an arrangement often used in electrical
property measurements of thermoelectric materials) is unsuitable because of the inability to
perform rapid and uniform in situ electrochemical doping of the entire volume of a dense polymer
film that is coated on an insulating substrate.
115, 116

Additional practical considerations are: (1) The electronic conductivity of polymers of
interest to battery applications must have good electronic conductivity. It is desirable that in the
doped state the conductivity values are as high 1 S cm
−1
. These high conductivity values lead to a

30
very small resistance across a thin film. Consequently, the voltage drop across these films will be
too small to be reliably measured without using large currents. (2) The technique should have the
dynamic measurement range to accommodate a change of several orders of magnitude in the
electronic conductivity of the polymers when transitioning from an un-doped to a doped polymer.
Thus, we need an electrode arrangement that combines the ability to perform in situ
electrochemical doping on a thin polymer film and have a wide dynamic range of sensitivity.
In addition to electronic conductivity, the values of ionic conductivity are also expected to
change over a wide range with doping. The ionic conductivity values are usually three to four
orders of magnitude lower than that of electronic conductivity, yet important in achieving the
required battery performance. Thus, the measurement of both ionic and electronic conductivity
using a single experimental setup presents a challenge. Specifically, we require the ability to switch
between two separate geometries to enable measurement of very large and very small resistance
values as described in Figure 3.2.

Figure 3.2. Electrode geometries used in impedance measurements for (a) low resistivity and (b)
high resistivity materials.

A large thickness and small area of cross-section are needed for measuring low values of
resistivity, while for measuring high values of resistivity, a thin layer with a large area of cross-

31
section is desirable. Thus, meeting these conflicting requirements of electrode geometry is also a
challenge.
3.1.3 Limitations of Currently Used Methods
In reviewing the literature we found several reports of the measurement of ionic and
electronic conductivity of polymers.
117-121
However, many of these reported methods are not suited
for in situ measurement and separation of electronic and ionic conductivity of polymer thin films
as a function of electrochemical doping in a relevant electrolyte. For example, the electronic
conductivity of a conductive block copolymer has been reported as a function of electrochemical
doping by using electrochemical impedance spectroscopy (EIS) but the measurement was
performed on a thick polymer film (over 100 μm) in a solid-state system, free of liquid electrolyte,
and without the ability to measure ionic conductivity.
110
On the other hand, ionic and electronic
conductivity of polymer thin films in LIB electrolyte has been reported simultaneously but without
chemical or electrochemical doping.
122
EIS has also been successfully used to determine the ionic
conductivity of polythiophene derivatives, but the salts were simply mixed with the polymer and
then cast as a film, without any electrochemical doping.
123
Karlsson et al
124-126
have used direct
current measurements and analysis of chronoamperometry to determine in situ the electronic
conductance and ion-diffusion coefficient. This type of measurement is suitable for the direct
measurement of electronic conductivity, however ionic conductivity is calculated using the
diffusion coefficient values obtained from transient measurements. In Table 3.1 we summarize the
strengths and limitations of previously reported methods. Briefly, for the various reasons listed in
Table3.1, the previously used techniques are limited in their ability to measure both electronic and
ionic conductivity accurately as a function of progressive uniform electrochemical doping of the
polymer in relevant electrolyte environments for battery applications.

32
Table 3.1. Comparison of Various Conductivity Measurement Methods
Technique Brief Description/ Example Advantages Disadvantages
Four-Probe on insulating
substrates
4 electrodes attached to film on an insulating
substrate. Used for various applications such as
OTFT’s
127
and free-standing films
Geometry ideal for measuring electronic
conductivity of thin films

Electrode arrangement inappropriate for electrochemical
doping because of using an insulating substrate.
Performed typically on dry film not in contact with liquid
electrolyte. Not suited for measurement of ionic conductivity
EIS with ion-blocking
electrodes (Huggins
Approach)

The Huggins approach of measuring the impedance
of a material with two ion-blocking electrodes yields
ionic and electronic conductivity.
128

P3HT-PEO block copolymer conductivity as a
function of Li-salt mixing and ratio of PEO/P3HT
129

Can simultaneously obtain ionic and electronic
conductivity as a function of chemical doping
Electrode geometry inappropriate for electrochemical
doping.
Thick polymer film needed if conductivity is high.
Both electronic and ionic conductivities can only be obtained
when they are similar in magnitude.
EIS + Electrochemical
Doping

Huggins approach of measuring impedance with
two ion-blocking electrodes is used (see above).
P3HT-PEO block copolymer electronic
conductivity as a function of electrochemical doping
using a 3 electrode cell
110

Can obtain electronic conductivity as a function
of electrochemical doping
Thick polymer films needed (>100 microns) for ensuring
measurable resistance values for films with high
conductivity values.
Typically used without liquid electrolyte
Ionic conductivity only obtained when polymer is undoped
otherwise shorted by electronic conductivity
EIS + Interdigitated
Microelectrode (IDM)

EIS to obtain ionic conductivity of polymer thin
films on IDM
123

IDM enhances signal.
A thin film can be used
Conductivity can be obtained as a function of
chemical doping
Geometry inappropriate for electrochemical doping
When either conductivity dominates, the other conductivity
value cannot be determined
No liquid electrolyte can be used
Direct Current Transient
Method with
interdigitated electrode
DC current between two working electrodes on the
polymer film is used to measure electronic
conductance. Analysis of current transient resulting
from potential step is used to separate ion-diffusion
from electron conduction.
124-126

Uses the geometry of the interdigitated
electrode array. Can measure the electronic and
ionic conductance by switching between two
measurement modes.
Analysis of transients require correction for capacitive
charging current although these errors can be minimized by
using small potential steps. The assumption of faradaic
currents being equal at the two working electrode potentials
must be verified for the specific polymer under test.
This Work
Uses polymer thin film on IDM
Allows for electrochemical doping in LIB
electrolyte
2-probe measurement to determine electronic
conductivity in LIB electrolyte at various potentials
3-electrode measurement to determine ionic
conductivity at any given potential
Electronic conductivities up to 10 S cm
−1
and
ionic conductivities up to 10
−3
S cm
−1
can be
measured (based on electrode geometry)
The difference in orders of magnitude between
ionic and electronic conductivities does not
affect measurement
Can be used with different LIB electrolytes
Upper limit of conductivity measurement is dictated by
electrode dimensions.
Polymer should be insoluble in liquid electrolyte

33
3.2 Electrode Geometry and Methodology
To allow for concurrently evaluating electronic and ionic conductivity as a function of
electrochemical doping in liquid electrolyte, we have combined the merits of using the planar and
interdigitated electrode geometry with EIS measurements. The test electrode is an interdigitated
gold electrode with two terminals. The electrode is coated with a thin layer of the polymer film
(Figure 3.3).

Figure 3.3. SEM image of (a) spin-coated P3HT on the gold interdigitated electrode, (b) P3HT
film, and (c) and (d) electrode gold digits and the coated polymer film.

The two gold-electrode terminals are electrically shorted and the electrochemical doping
is carried out by polarizing this test electrode in a three-electrode cell configuration in a solution
of lithium-ion battery electrolyte at various potentials (Figure 3.4a). Cyclic voltammetry (CV)
scans on the interdigitated electrode were also conducted in the three-electrode configuration to

34
track the electrochemical doping process and identify the electrode potentials at which the doping
occurs. Ionic conductivity of the film was measured in this three-electrode configuration using
EIS. The electrode geometry under these conditions is that of a planar electrode of large area and
a small thickness, well suited for measuring large resistivity values (Figure 3.2b). After measuring
the ionic conductivity, in a separate experiment on the same film, the electronic conductivity is
measured between the terminals of the two gold interdigitated electrodes (Figure 3.4b).

Figure 3.4. The gold interdigitated electrode in the (a) 2-electrode and (b) 3-electrode
configuration.

In this two-electrode geometry, the area of cross-section is small as governed by the
thickness of the film, while the path length for current is long as determined by the distance
between the digits of the gold electrodes (Figure 3.2a). The latter geometry was well-suited for
measuring very low values of resistivity.
3.3 Method Validation with p-Dopable Conducting Polymers
To validate the technique, we used P3HT, 1:1 P3HT/PEO mixture, and PEDOT:PSS as
model systems for p-dopable polymers (structures shown in Figure 3.5).

35

Figure 3.5. Structures of p-dopable polymers used in this study.
3.3.1 Cyclic Voltammetry
Figure 3.6 shows cyclic voltammograms of the polymer thin films. Cyclic voltammetry of
P3HT films showed two current peaks at 3.45 V and 3.85 V vs. Li
+
/Li, attributed to the electro-
oxidation and anion-doping of P3HT (Figure 3.6a).
130
The CV scans of P3HT/PEO (Figure 3.6b)
showed similar oxidation peaks to those of P3HT while the CV scans of PEDOT:PSS (Figure 3.6c)
showed oxidation and reduction processes occurred over a wide potential window.

Figure 3.6. CV plots of (a)P3HT, (b) P3HT/PEO, and(c) PEDOT:PSS, at 10 and 50 mV s
−1
.
3.3.2 Ionic Conductivity Measurements
The impedance of the polymer thin film electrodes was measured in the three-electrode
configuration following electrochemical doping to different levels as discussed in section 2.1. The
electrochemical doping and 3-electrode impedance measurement protocol to determine ionic
conductivity is shown in Figure 3.7.

3.0 3.5 4.0
-1
0
1
2
3.0 3.5 4.0
-1
0
1
2
1 2 3 4
-0.5
0.0
0.5
(c)
(a)
10 mV/s
50 mV/s

(µA (mVs
-1
)
-1
)
Rate-Normalized Current
Potential (V vs Li
+
/Li)
(b)
(µA (mVs
-1
)
-1
)
Rate-Normalized Current
Potential (V vs Li
+
/Li)
10 mV/s
50 mV/s
(µA (mVs
-1
)
-1
)
Rate-Normalized Current
Potential (V vs Li
+
/Li)
10 mV/s
50 mV/s

36

Figure 3.7. Electrochemical doping and 3-electrode impedance measurement protocol for
determining ionic conductivity.

The sinusoidal potential perturbation is accompanied by double-layer charging, charge-
transfer, diffusion and migration of ions, and the transport of electrons. The impedance response
arising from these processes distributed across the thickness of the porous polymer film has been
analyzed rigorously by Garcia-Belmonte et al.
131-133
A schematic of the thin film polymer electrode
in the 3-electrode cell is shown in Figure 3.8.

37

Figure 3.8. A schematic of the polymer electrode in the 3-electrode cell.
Under these conditions the electrode impedance can be represented by a “finite length”
transmission line model with a reflective boundary.
134, 135
This model with distributed circuit
elements has been applied successfully to study the diffusion of charge carriers in
electrochemically-doped thin polymer films.
136
For a 100 nm thick polymer film, the sinusoidal
response at low frequencies (10 Hz to 0.1 Hz) arises from the entire volume and finite thickness
of the porous film. Specifically, the ionic resistance of the polymer phase (Rion), the electronic
resistance of polymer phase (Re), and the interfacial faradaic impedance (Zf) are distributed circuit
elements of a generalized transmission line equivalent circuit along with a geometric capacitance
CG of the entire thin film (Figure 3.9a).

38

Figure 3.9. (a) Equivalent circuit of the transmission line model, (b) expected impedance
response, and (c) experimental impedance response.

For such a model, assuming rapid interfacial charge-transfer, it was shown by Albery et
al
134
that at low frequencies, the real component of the impedance Zreal tends to attain a constant
value given by Equation 3.1.
!
!"#$
=
$
%&'
+ $
"
3
+ $
(
(3.1)
Here Rs is the uncompensated solution resistance obtained from the high frequency intercept. The
value of Zreal was obtained by extrapolation of the low-frequency line to meet the real axis of the
Nyquist plot (Figure 3.9c). With the electronic resistance Re obtained separately (from the
electronic conductivity measurement described in the next section), Rion was obtained using
Equation 3.1. The polymer film sits between the gold digits and its electronic conductivity is
greater than the ionic conductivity. Consequently, the current lines for ionic flow into the

39
electrolyte are perpendicular to the surface of the polymer film. Thus, the ionic conductivity of the
polymer film, σion was calculated using Equation 3.2, based on the dimensions of the film and the
inter-digitated electrode.
+
%&'
=
1
$
%&'
×
ℎ
.×(/−1)×1
(3.2)
In Equation 3.2, h is the thickness of the polymer film from the base of the substrate and between
the gold digits, l and N are the length and the number of the digits of the gold electrode,
respectively, and d is the distance between the gold electrodes. Electrochemical impedance was
measured at various electrochemical doping levels (potential) as shown in Figure 3.10.

Figure 3.10. 3-electrode impedance data of P3HT film. (a) Bode plots as a function of
electrochemical doping and (b) Nyquist plots at various potentials.

To validate the measurement method and analysis, we studied various conjugated polymer
films. For films of P3HT, at 3.0 V, prior to any significant doping, the ionic conductivity was 1 ×
10
−9
S cm
−1
(Figure 3.11). No ionic conductivity data of un-doped P3HT films could be found in
the literature. This low value of ionic conductivity is not surprising as the un-doped P3HT film is
semi-crystalline and lacks the ability to solvate Li
+
or anions. Therefore, prior to doping we could
expect the ionic conductivity to arise from the porosity and solvent-induced swelling of the film.

40
Upon electrochemical doping at a potential of 3.4 V vs Li
+
/Li, anions were introduced into the
polymer film, and the ionic conductivity of P3HT increased by an order of magnitude to 3 × 10
−8

S cm
−1
.

Figure 3.11. Ionic conductivities of P3HT, PEDOT:PSS, and P3HT/PEO.
The large decrease in impedance is evident in the Bode and Nyquist plots of the P3HT film
(Figure 3.10). The ionic conductivity continued to increase with further doping and reached a value
of 9 × 10
−8
S cm
−1
at 4.0 V vs Li
+
/Li. On the other hand, P3HT/PEO showed a constant ionic
conductivity of 4 × 10
−7
S cm
−1
across the examined potential window (Figure 3.11). Thus, we
may conclude that most of the ionic transport is supported by the flexible and ion-solvating
ethylene oxide groups of PEO consistent with expectations from previous reports.
137
PEDOT:PSS
showed an ionic conductivity of 4 × 10
−5
to 6 × 10
−5
S cm
−1
in the potential window of 3.2 to 4.0
V vs Li
+
/Li. This value of ionic conductivity agrees closely with previous reports in the literature
for PEDOT:PSS.
122

41
3.3.3 Electronic Conductivity Measurements
The detailed protocol for electronic conductivity measurement is discussed in Section 2.1.
A representative 2-electrode impedance measurement protocol to determine electronic
conductivity is shown in Figure 3.12.

Figure 3.12. Potential and current profiles during electrochemical doping preceding the EIS
measurement protocol to determine electronic conductivity.

The Nyquist plot in the two-electrode configuration (Figure 3.13) showed two semicircular
arcs. The impedance data was analyzed by the method of Huggins
128
for mixed conductors.
Specifically, the data was fitted to the Debye equivalent circuit model for mixed conduction
through the film (Figure 3.13 inset).
138
This equivalent circuit embodies the mixed electronic and
ionic conductivity including the capacitance properties of the film.

42

Figure 3.13. A representative Nyquist impedance plot of P3HT obtained at 3.3 V vs Li
+
/Li using
the two-electrode configuration and the fitting circuit (inset) used to obtain the electronic
conductivity.

At high frequencies, the current flow occurs through the charge and discharge across the
capacitive elements. At low frequencies, the electrodes block the ionic current, while the electronic
current can still flow between the two terminals of the interdigitated gold electrodes. Such a circuit
yields an impedance response with two semi-circles (Figure 3.13). In fitting the experimental data,
the constant phase element was used in the place of the capacitor element.
Representative Nyquist plots for PEDOT:PSS and P3HT/PEO at 3.3 V are shown in Figure
3.14. The fitting parameters used for the Nyquist plots and the associated errors are shown in Table
3.2. The diameter of the semicircle observed at the higher frequencies corresponded to a parallel
combination of the ionic and electronic resistances, while the intersection of the second semicircle
with the real axis at lower frequencies corresponded to the electronic resistance of the conducting
polymer film, Re. The electronic conductivity of the polymer film, σe, was then calculated based on
the two-electrode geometry (Equation 3.3) that is different from that applied to the ionic
conductivity measurement.

43
+
"
=
1
$
"
×
1
.×(/−1)×ℎ
(3.3)

Figure 3.14. 2-electrode Nyquist plots of (a) P3HT/PEO and (b) PEDOT:PSS at 3.3 V vs Li
+
/Li.

Table 3.2. Fitting Parameters for Nyquist Plots in Figures 3.13 and 3.14
Polymer
Rs
(Ω)
Rs error
(%)
CPE
(S sec
n
)
CPE error
(%)
n
n error
(%)
Re
(Ω)
Re error
(%)
CPE
(S sec
n
)
CPE error
(%)
n
n error
(%)
P3HT 43.27 0.83 3.79E−07 16.39 1.00 12.72 11.38 3.84 3.16E−04 8.63 0.58 17.04
P3HT/PEO 38.04 0.15 9.35E−07 35.15 0.91 2.91 9.95 0.58 6.44E−05 7.97 0.74 1.04
PEDOT:PSS 38.95 0.82 6.99E−06 24.25 0.97 18.48 1.98 1.61 1.79E−05 18.67 0.86 11.70

The electronic conductivity of P3HT in the un-doped state was determined to be 8 × 10
−6
S cm
−1
at 3.0 V vs Li
+
/Li (Figure 3.15). This value is in agreement with previously reported values
from four-point probe measurements.
139
The electronic conductivity of P3HT, however, increased
by several orders of magnitude upon electrochemical doping and reached a maximum of 1 × 10
−1
S cm
−1
at 3.5 V and remained constant until 3.9 V. The electronic conductivity then decreased
slightly upon further electrochemical doping to 7 × 10
−2
S cm
−1
at 4.0 V vs Li
+
/Li. The significant
change in P3HT film impedance as a function of electrochemical doping is reflected in the Bode
plots (Figure 3.16).

44

Figure 3.15. Electronic conductivities of P3HT, PEDOT:PSS, and P3HT/PEO.

Figure 3.16. Bode plots of 2-electrode EIS measurement of P3HT film (a) between 2.8 and 3.3 V
and (b) between 3.3 and 4 V.

The electronic conductivity of P3HT/PEO, as expected, was similar to that of P3HT when
measured as a function of electrode potential. This trend showed that electron transport in P3HT
was not negatively affected upon mixing with PEO. The electronic conductivity of PEDOT:PSS
was determined to be 2 S cm
−1
. This value is in agreement with the literature reports using a four-
point probe method.
122
PEDOT:PSS as synthesized is in the doped or oxidized state since PSS

45
functions as a polyanionic dopant. Thus, the open-circuit potential of this polymer is at 3.2 V vs.
Li
+
/Li prior to any CV scans. The electronic conductivity remained stable over the potential
window investigated with only a slight decrease to 1 S cm
−1
above 3.5 V vs Li
+
/Li (Figure 3.16).
De-doping of PEDOT:PSS was found to occur at about 1.5 V vs. Li+/Li. However, de-doping of
p-doped polymers below 2.5 V is not relevant to the study of lithium-ion cathode binders.
3.4 Applying Method to n-Dopable Polymers
To validate the technique with n-dopable polymers, P(NDI2OD-T2) was used as a model
n-dopable polymer with its structure shown in Figure 3.17.

Figure 3.17. Structure of P(NDI2OD-T2).
Electrochemical doping of the P(NDI2OD-T2) film (Figure 3.18) was observed in the
potential range of 1.7 V to 3.0 V vs Li
+
/Li as shown in Figure 3.18 with electrochemical n-doping
occurring at 2.2 V vs Li
+
/Li. De-doping (oxidation) starts at 2.2 V and the main oxidation peak is
observed at 2.55 V vs Li
+
/Li.

46

Figure 3.18. Cyclic voltammetry plot of P(NDI2OD-T2) at the various scan rates indicated.
The ionic conductivity of un-doped P(NDI2OD-T2) at 2.4 V was 3 × 10
−10
S cm
−1
(Figure
3.19). The ionic conductivity increased with electrochemical n-doping to reach 6 × 10
−9
S cm
−1
as
the electrode was held at more negative potentials up to 1.6 V vs. Li
+
/Li (Figure 3.19). Thus, an
increase in ionic conductivity of at least two orders of magnitude was observed with both p-doped
and n-doped polymers upon electrochemical doping. Thus, the effect of electrochemical doping
was to increase the ionic conductivity.

47

Figure 3.19. The electronic and ionic conductivities of P(NDI2OD-T2) as a function of electrode
potential.

3.5 Expanding the Technique to Oxides
After validating our technique for both p-dopable and n-dopable polymers, we attempted
to apply it to oxide materials that are usually the active materials in lithium-ion batteries.
Challenges arise, however, in thin-film fabrication. Unlike conducting polymers, oxides are not
easily spin-coated into uniform and dense films. As a result, alternate fabrication methods on
interdigitated micro electrodes had to be researched. For example, controlled sedimentation from
hexane solvent onto a substrate followed by annealing at 400 ℃ yielded a dense and uniform oxide
film as shown in Figure 3.20.

48

Figure 3.20. (a) Photograph of the oxide electrode. (b) and (c) SEM images of the electrode.
Due to the oxide materials requiring to be annealed at elevated temperatures, gold micro-
electrodes with glass substrates are not suitable materials for oxide electrodes. As a result, Oxide
thin films were fabricated on silicon substrates and gold digits were printed on the surface of the
oxides to form the gold micro-electrodes, as shown in Figure 3.21. This geometry is slightly
different than the one shown in the previous section but still plays a similar role. d is the distance
between the gold digits (200 μm) and h is the thickness of the oxide film (150 nm).

Figure 3.21. A schematic of the oxide micro-electrode.
As with polymers, we are able to cycle the oxide thin-film at relatively high scan rates as
can be seen in the TiNb2O7 cyclic voltammograms in Figure 3.22.

49

Figure 3.22. CV of TiNb2O7 at various scan rates.
Furthermore, we can reliably and continuously cycle the oxide thin film even though there
are no binders or conductive additives. Figure 3.23 shows multiple CV cycles at 20 mV s
−1
in
which no noticeable change is observed between cycles. This indicates that the mechanical
integrity of the film is maintained, and the oxide layer is adhered well to the substrate.

Figure 3.23. CV data of TiNb2O7 at 20 mV s
−1
.
Electrochemical impedance data was collected as before to obtain the electronic
conductivity as a function of potential. The electrode was held potentiostatically at a specific
potential in the 3-electrode setup for 300 seconds. The electrode was then allowed to equilibrate

50
at open circuit. Finally, the geometry was changed into two electrodes, and 2-electrode
potentiostatic EIS was collected as shown in Figure 3.24.

Figure 3.24. 2-electrode EIS Nyquist plot of TiNb2O7 at 1.6 V.
The impedance data was analyzed as before using Huggin’s approach for mixed conductors
in which the Nyquist plot was fitted to the Debye equivalent circuit model for mixed conduction
through the film. The extracted electronic resistivity values, Re, was used to obtain the electronic
conductivity, se, of the oxide film using Equation 3.4:
s
"
=
1
$
"
×
1
.×(/−1)×ℎ
(3.4)
where d is the distance between digits, l is the digit length (0.57 cm), N is the number of digits (8),
and h is the film thickness. Obtaining EIS data at various potentials then extracting electronic
resistivity values allowed us to determine the electronic conductivity of the oxide film as a function
of potential as shown in Figure 3.25.

51

Figure 3.25. Electronic conductivity of TiNb2O7 as a function of potential.
Thus, we have demonstrated that we are not only able to measure the conductivity of
polymers, but we are also able to determine the conductivity of oxides which form the active
materials in LiB cathodes. Therefore, this technique is robust and can provide invaluable insight
into properties of battery electrode materials.

52
CHAPTER 4.

ENHANCING LITHIUM-ION BATTERIES USING
BIFUNCTIONAL CONDUCTIVE BINDERS
Reprinted with permission from Chemistry of Materials 2020 32 (21), 9176-9189, and Chemistry
of Materials 2022 34 (6), 2672-2686 Copyright © 2022 American Chemical Society.

4.1 PProDOT-Hx2 as a Conductive Binder for Lithium-ion Batteries
4.1.1 Introduction
Lithium-ion batteries (LIBs) are the dominant energy storage technology used today in
electronics and electric vehicles because of the need for high specific energy and power density.
78-
80
Traditional polymeric binders, such as polyvinylidene fluoride (PVDF), are used in these energy
storage devices to hold the active electrode material together with conductive carbon additives,
and to bind both materials to the metallic current collector. PVDF and related binders are chosen
for their electrochemical stability, but they suffer from low electronic and Li
+
ion conductivities,
thereby increasing electrode impedance by interfering with charge transport pathways.
81-83

Conjugated polymers that combine the advantages of organic conductors and traditional polymers,
serve as excellent candidates for multi-functional binders, offering mixed electronic and ionic
conduction along with mechanical integrity, thereby improving electrochemical performance.
88, 97,
140-148
Besides the promise of providing robust mechanical, electrochemical, and thermal stability
along with good wetting/coating properties,
84, 146, 149
it is their properties of mixed conduction that
lets them function both as a binder and as a conductive additive. Because of this, conjugated
polymers have been explored for use as both cathode and anode binders in LIBs.
84-91, 150-152

53
Neutral (un-doped) conjugated polymers usually exhibit low electronic conductivity in the
range of 10
−10
to 10
−5
S cm
−1
, but these values can be increased significantly by electrochemical
or chemical doping.
153-156
While high electrical conductivity in conjugated polymers is based on
achieving both sufficient levels of doping and sufficient levels of crystallinity, ionic conductivity
is dictated by solvated ion motion through the polymer matrix and favored by a lower degree of
crystallinity and higher porosity.
157, 158
While many doped conjugated polymers show excellent
electronic conductivity, in order to augment performance, an important focus is on improving the
ionic conductivity, which is often significantly lower.
97
A broader understanding of the critical
structure-function relationships in conductive polymer binders is desired in order to leverage their
use to increase rate capability, battery cycle life,
85, 101, 159, 160
and specific capacity.
84, 101, 149, 161

Regioregular poly(3-alkylthiophenes) (P3ATs) are a broadly-studied family of conjugated
polymers possessing relatively narrow band gaps, good processability, and high hole mobility.
89,
162, 163
Their solubility in many organic solvents enables formation of uniform thin film coatings
on electrodes as opposed to the colloidal nature of doped polymer suspensions like poly(3,4-
ethylenedioxythiophene)/poly(4-styrenesulfonic acid) (PEDOT:PSS).
164
Upon doping, poly(3-
hexylthiophene) (P3HT) exhibits enhanced electronic conductivity (up to 15 S cm
−1
) within the
potential window of a majority of LIB cathodes.
165, 166
Recent work from our team on the use of
P3HT as a conductive binder for LiNi0.8Co0.15Al0.05O2 (NCA) cathodes showed high power density
and excellent cycling stability compared to PVDF-NCA electrodes.
112
While P3HT enhances
battery performance relative to PVDF, it still lacks Li
+
solubilizing functionality, limiting ionic
conductivity. The semi-crystalline morphology of P3HT poses a further hindrance to Li
+
transport
due to the dense chain packing. While the hexyl side chains of P3HT are not conducive to the
diffusion of Li
+
, other side chains such as oligoethylene-glycol may be more beneficial for

54
promoting Li
+
hopping as explored in recent work.
123, 167
Reichmanis and co-workers
160, 168
have
also investigated blends of polythiophenes with polyethylene glycol (PEG) as a multi-component
binder for anodes in LIBs in order to target mixed conductivity.
Unfortunately, polythiophenes also have another drawback as binders for LIB electrodes.
Namely, they do not possess the necessary chemical stability required for repeated cycling at high
potentials as cathode binders or low potentials as anode binders.
169, 170
An alternative to P3ATs
and a soluble alternative to PEDOT-PSS is poly(3,4-propylenedioxythiophenes) (PProDOTs),
which are known to exhibit remarkable stability to repeated cycling over a broad potential window
that is well suited for cathode materials commonly used in LIBs.
171
By functionalizing the 3- and
4- positions of thiophene with an alkylenedioxy bridge, the resulting polymers show a significantly
lower onset for oxidation due to the electron donating oxygens.
172
Further, the propylene bridge in
ProDOTs allows symmetrical disubstitution and hence good polymer solubility in organic
solvents.
152

PProDOTs first garnered interest due to their excellent electrochromic properties,
reversible electrochemical doping and fast switching times.
173
In fact, PProDOTs have excellent
electrochemical stability when p-doped (oxidized) as well as faster counter-ion influx/efflux upon
doping/dedoping than polythiophenes due to their open morphology, making them promising
candidates for batteries and capacitors.
174-176
. Reynolds and coworkers have recently reported a
ProDOT-based copolymer displaying redox stability over thousands of cycles along with
reasonable charge-storage capacity.
177
As a general class, dioxythiophenes have been established
for long-term cycling stability when used as supercapacitors
178, 179
demonstrating capacity
retention over more than 400,000 cycles.
175

55
As the most well-known dioxythiophene, PEDOT has shown excellent stability when used
as a protective coating for LiFePO4 and other cathodes.
180-182
resulting in enhancement on capacity
and charge/discharge behavior. Such conductive coatings are also known to lead to an overall
increase in electrical conductivity. In an attempt to provide mixed conductivity in three
dimensional electrodes, Hammond and coworkers investigated PEDOT:PSS blended with PEO.
122

The favorable interaction of PEO with the PSS matrix suppressed crystallinity and elevated the
ionic conductivity, resulting in mixed conductivity at an optimum PEO:PEDOT ratio. However,
PEDOT:PSS is intrinsically insoluble, and the colloidal suspension does not provide the same
uniform thin film coating abilities offered by soluble polymers.
In order to address the current shortcomings of polymer binders for cathode in LIBs and to
explore the utility of soluble dioxythiophene polymers, the focus of the current work is the
investigation of the mixed electronic and ionic conductivity of dihexyl-substituted poly(3,4-
propylenedioxythiophene) (PProDOT-Hx2). While the propylenedioxythiophene backbone
accounts for enhanced electrochemical stability in comparison to polythiophenes, the hexyl side
chains enhance the solubility of the polymer in non-polar processing solvents improving the
processability of the electrode. An open, partly disordered morphology compared to highly
crystalline polymers like P3HT, allows efficient ion insertion/de-insertion, facilitating ionic
conductivity. Although the electrochemical and optical characterization of PProDOT-Hx2 has been
reported,
183, 184
its properties relevant to cathode binders for LIBs has not been investigated.
Therefore, here we present a combined theoretical and experimental study to establish the
fundamental mixed ionic and electronic conductivity of PProDOT-Hx2, and the comparative
performance of PProDOT-Hx2 as a conductive binder for the NCA cathode relative to PVDF.

56
4.1.2 Synthesis of PProDOT-Hx2
PProDOT-Hx2 has previously been synthesized by oxidative polymerization,
electropolymerization,
171, 185
and by Grignard metathesis polymerization.
36,186
Recently, Reynolds
and coworkers have also demonstrated the synthesis of dioxythiophene polymers using Direct
Arylation Polymerization (DArP).
176, 177
In an attempt to avoid the toxic, unstable, and hard to
purify metallated monomers typically used in traditional cross-coupling reactions, DArP has been
developed as a greener and more atom economical approach.
187, 188
DArP is a C-H activation
method whereby one aromatic C-H bond couples with an aryl halide C-X bond forming a C-C
bond, thus, eliminating the need for additional synthetic steps. In the current study, PProDOT-Hx2
was synthesized on the gram scale using novel conditions for DArP with a number-average
molecular weight (Mn) of 19.1 kDa and a dispersity of 1.6 (Figure 4.1).

Figure 4.1. Schematic of PProDOT-Hx2 synthesis using DArP.
4.1.3 Electrochemical Properties of PProDOT-Hx2
To be an effective conductive binder, it is necessary that PProDOT-Hx2 be conducting over
the potential window of operation for the battery cathode material. Also, the electrochemical
doping should have fast kinetics to allow for rapid doping and de-doping during cycling. Thus, the
electrochemical doping (oxidation) behavior of the PProDOT-Hx2 was first investigated in thin
film format in a three-electrode cell with 1 M LiTFSI dissolved in 1:1 (by volume) ethylene
carbonate (EC)/dimethyl carbonate (DMC) as the electrolyte and lithium foil as the counter and
reference electrodes. The cyclic voltammetry (CV) curves between 3 and 4.2 V (vs Li/Li
+
) at a

57
scan rate of 10 mV s
−1
indicated that the electrochemical doping process was highly reversible
(Figure 4.2). In the first cycle, a shoulder is observed at ~3.3 V which is known to correspond with
polaron (radical cation) formation. The first major oxidation peak (oxidation 1) at 3.36 V
corresponds to bipolaron formation and is similarly established and supported.
152, 173, 184, 189-192

Figure 4.2. Initial cycling of PProDOT-Hx2.
Upon cycling to higher potentials, a second major oxidation (oxidation 2) is observed at
~3.8 V. This higher oxidation process is not generally reported in dioxythiophene polymers, which
are rarely cycled to this high a potential. However, model studies in solution on well-defined
ProDOT-Hx2 oligomers (n = 4-12) indicate that this higher oxidation is due to the generation of
additional charge carriers, where the higher doping levels are enabled by the stabilizing effect of
the oxygen atoms along the backbone.
193
While the exact nature of the higher oxidation event is
not fully established, the oligomer work suggests a more significant role of cation character on the
oxygens in the ether bridge in the reversibly generated charge carrier species. Upon subsequent
scans, the polaron pre-peak is no longer observed at this scan rate and two pairs of reversible redox

58
peaks were observed at 3.22 to 3.16 V and 3.76 to 3.68 V (vs. Li/Li
+
) with the anodic peak listed
first in each pair. The reversibility of both oxidation peaks is key to the use of these polymers as
binders for cathode materials. The shift of the oxidation peaks to lower potentials between the first
and following cycles is attributed to the enhanced conductivity of the polymer after
electrochemical doping, which reduced the overpotential. Furthermore, we have calculated the
coulombic efficiency for the first cycle. Based on integrating the CV curves (Figure 4.2), the initial
coulombic efficiency (ICE) of PProDOT-Hx2 is 38.4%. The low value of ICE is likely due to side
reactions such as the oxidation of the electrolyte and/or the formation of a solid electrolyte
interphase (SEI) on the film. The contributions from side reactions are significant since the films
are only 50 nm thick.
For use as binders in LIBs, stability over a broad range of potentials is key, and so we
further expanded the potential window to higher potential to investigate the stable operational
range of the polymer. CV curves were obtained with gradually increasing potential ranges (Figure
4.3). The shapes of the redox peaks were retained up to a voltage cutoff of 4.5 V. At potentials
beyond 4.5 V, however, the redox peak features began to change, indicating decomposition of the
electrolyte and polymer. The stable and wide operation potential range (3.2 – 4.5 V) of PProDOT-
Hx2 makes it an ideal candidate for use as a binder for a number of cathode materials, including
LiCoO2, LiMn2O4 and LiMO2 (M=Ni, Co, Mn, Al).

59

Figure 4.3. CV data for PProDOT-Hx2 thin film at various potential intervals at 10 mV s
−1
.
To ensure that electrochemical doping/de-doping of the PProDOT-Hx2 binder does not
limit cathode cycling, we also examined the kinetics of the electrochemical doping at high rates.
Kinetics were analyzed based on a series of CV measurements at various scan rates (Figure 4.4)
from 10 to 100 mV s
−1
. The minimal shifts of the redox peaks with increasing scan rates indicated
rapid reaction rates.

60

Figure 4.4. CV data for PProDOT-Hx2 as a function of various scan rates from 10 to 100 mV s
−1
.
To further quantify the kinetics of polymer doping, we examined the relation between the
measured current (i) and scan rate (v). According to Equation 4.1.
194

4 = 56
)
, (4.1)
where b can be determined by the slope of a plot of log (i) vs. log (v) plots for each redox peak. A
value of b equal to 0.5 indicates a process controlled by semi-infinite diffusion, while a b value
close to 1 indicates a non-diffusion controlled or a surface-controlled charge-storage process. From
the plot of log (i) vs. log (v) (Figure 4.5), the b values of the oxidation 1 anodic and cathodic peaks
are 0.96 and 0.89 respectively. For the oxidation 2 anodic and cathodic peaks the b values were
0.88 and 0.94, respectively. These b values indicate rapid redox processes in the polymer thin
films. The fast kinetics for the electrochemical p-doping of PProDOT-Hx2 are expected to facilitate
rapid electron transport when used as a conductive binder in cathode composites.

61

Figure 4.5. Log of the peak current (i) vs. log of the scan rate (v) for data shown in Figure 4.4.
To verify the redox stability of the polymer, extended CV scans were performed on
PProDOT-Hx2 thin films (Figure 4.6) showing a capacity retention of 72% over 500 cycles when
cycling from 3.0 to 4.0 V vs. Li/Li
+
at 10 mV s
−1
.

Figure 4.6. Long-term CV curves of PProDOT-Hx2 between 3.0 and 4.0 V at 10 mV s
−1
.

62
CV data like that shown previously also allows us to calculate the specific capacity of the
polymer which is 18.5 mAh g
−1
at 10 mV s
−1
. While this capacity is not important for battery
performance, as binders are used at just a few percent levels, it does allow us to calculate the
doping level, which is important for electrical conductivity. In this case, the value corresponds to
one dopant per 4.5 monomer units
4.1.4 Ionic Conductivity of PProDOT-Hx2
Our conductivity technique that we have developed and discussed in Chapter 3 was used
to examine the conductivity of PProDOT-Hx2 and compare it to P3HT, which is a common
conducting polymer that has been previously used as a binder in LIBs. For determining ionic
conductivity, the electrochemical impedance of the PProDOT-Hx2 thin-film was measured in the
three-electrode configuration (Figure 4.7a) following doping to different levels. The impedance
data (Figure 4.7b) was analyzed using a transmission line equivalent circuit model used for
describing porous electrode structures.
136

Figure 4.7. (a) The 3-electrode configuration used to electrochemically dope the polymer and
determine its ionic conductivity. (b) A representative Nyquist impedance plot obtained at 4 V vs
Li/Li
+
using the 3-electrode configuration to obtain the ionic conductivity.

63

The ionic conductivity of the polymer film was then calculated from the measured ionic
resistance based on the dimensions of the electrode and the film thickness. The ionic conductivity
of P3HT is 10
−9
S cm
−1
up to 3.2 V (vs Li/Li
+
) and increased by an order of magnitude to 10
−8
S
cm
−1
at potentials above 3.2 V vs Li/Li
+
(Figure 4.8). On the other hand, PProDOT-Hx2, upon
doping showed an ionic conductivity of 10
−7
S cm
−1
which is at least an order of magnitude higher
than that of P3HT. Despite a slight decrease beyond 3.2 V, the ionic conductivity of PProDOT-
Hx2 remained relatively constant with increasing levels of electrochemical doping. The higher
ionic conductivity of PProDOT-Hx2 can be attributed to the Li
+
solvation abilities of the oxygens
present in PProDOT-Hx2 but absent in P3HT.

Figure 4.8. Ionic conductivities of P3HT and PProDOT-Hx2.
4.1.5 Electronic Conductivity of PProDOT-Hx2
Upon electrochemical doping of PProDOT-Hx2 in the 3-electrode configuration (Figure
4.7a) the electronic resistance of the polymer thin film was determined by analysis of the
electrochemical impedance in the 2-electrode configuration (Figure 4.9a). We used the Huggins

64
approach
128
and fit the impedance data to the Debye equivalent circuit model for mixed
conduction through the film (Figure 4.9b).
138

Figure 4.9. (a) The 2-electrode configuration used to measure electronic conductivity. (b) A
representative Nyquist impedance plot obtained at 4 V vs Li/Li
+
using the 2-electrode
configuration and the fitting circuit (inset) used to obtain the electronic conductivity.

The electronic conductivity was then calculated by accounting for the electrode geometry,
and the trend in conductivity relative to potential is shown in Figure 4.10. For comparison, the
electronic conductivity of both PProDOT-Hx2 and P3HT in their neutral, un-doped states were
approximately 10
−5
S cm
−1
(Figure 4.10). The electronic conductivity of both polymers increased
by several orders of magnitude with doping, reaching a maximum of 1x10
−1
S cm
−1
at around 3.6
V corresponding to complete conversion of polarons to bipolarons. The electronic conductivity
then slightly decreased for PProDOT-Hx2 to 5 × 10
−2
S cm
−1
above 3.6 V. The slight decrease
could be due to potential changes in the nature of the charge carriers associated with oxidation
2,
193
or to observed changes in morphology at higher doping levels. The electronic conductivity of
P3HT reached a maximum of 1 × 10
−1
S cm
−1
at 3.55 V and remained stable above 3.55 V with a
slight decrease in electronic conductivity above 3.9 V.

65

Figure 4.10. The electronic conductivities of P3HT and PProDOT-Hx2.
4.1.6 Performance as a Cathode Binder
To examine the performance of PProDOT-Hx2 as a polymer binder in Li-ion batteries,
LiNi0.8Co0.15Al0.05O2 (NCA) electrodes incorporating PProDOT-Hx2, Super P carbon black, and
multiwalled carbon nanotubes (CNT) with a weight ratio of 90-3-3-4% (NCA-SP-CNT-PProDOT-
Hx2) were prepared. As can be observed from the TEM image of just NCA and PProDOT-Hx2 (96-
4%) (Figure 4.11a), PProDOT-Hx2 is homogeneously distributed and acts as a binder to the NCA
particles, leading to a well-connected network of active materials. Figure 4.11b shows the TEM
image of the full composite electrode.

66

Figure 4.11. TEM image of (a) 96-4% NCA-PProDOT-Hx2 and (b) 90-3-3-4% NCA-SP-CNT-
PProDOT-Hx2 electrodes.

To characterize the rate capability of the NCA-PProDOT-Hx2 electrodes, galvanostatic
charge−discharge (GCD) tests were performed in coin cells. The coin cells were assembled using
lithium as the counter electrode and 1 M LiTFSI in EC/DMC as the electrolyte. The rate capability
of the NCA-PProDOT-Hx2 electrodes were compared to NCA-PVDF control electrodes, and NCA
electrodes without the addition of any binder for a mass loading of ~ 6 mg cm
−2
to demonstrate
the benefits of PProDOT-Hx2 (Figures 4.12).

67

Figure 4.12. Rate capability of the NCA-PProDOT-Hx2 and NCA-PVDF.
Different C-rates were utilized based on 1 C = 160 mA g
−1
as previously reported.
112
At a
rate of C/5, the NCA-PProDOT-Hx2 and NCA-PVDF cathodes delivered nearly identical
discharge capacities of 168 and 166 mAh g
−1
, respectively. At slow rates below 1C, they continued
to deliver comparable specific capacities. Nevertheless, at higher rates (above 2C), the NCA-
PProDOT-Hx2 cells exhibited enhanced rate capability compared to the NCA-PVDF cells and
binder-free cells. At 6C, the NCA-PProDOT-Hx2 delivered a capacity of 111 mAh g
−1
, while
NCA-PVDF only provides 20 mAh g
−1
(Figure 4.12). We can attribute these improved results with
PProDOT-Hx2 to the enhanced electronic and ionic conductivity of the doped PProDOT-Hx2
compared to the insulating PVDF binder. The NCA-PProDOT-Hx2 cells also showed significantly
reduced polarization at all rates above 1C compared to the NCA-PVDF cells and the binder-free
NCA cells (Figure 4.13).

68

Figure 4.13. The galvanostatic charge-discharge curves of the (a) NCA-PProDOT-Hx2 and
(b)NCA-PVDF.

While high rate is important for many applications, stability is also key, and PVDF is
designed for its chemical inertness. Surprisingly, a long-term cycling comparison at 2C between
the PProDOT-Hx2 and PVDF cells showed that in the first 120 cycles, the capacities are
comparable, but after that point, a decay is observed in the NCA-PVDF cells that is not observed
in those made with PProDOT-Hx2 (Figure 4.14a). Thus, the NCA-PProDOT-Hx2 cathode showed
better cycling stability with less decay of the potential drop in the GCD curves compared to NCA-
PVDF electrodes. Overall, this indicates that the PProDOT-Hx2 polymer was more stable during
long-term cycling (Figure 4.14b and 4.14c).

69

Figure 4.14. (a) Long-term cycling for NCA-PProDOT-Hx2 and NCA-PVDF at a rate of 2C. The
corresponding galvanostatic charge-discharge curves of the (b) NCA-PProDOT-Hx2 and (c)
NCA-PVDF at different cycles.

4.2 Increasing PProDOT Ionic Conductivity with Oligoether Side Chains
4.2.1 Introduction
While PProDOT-Hx2 was a step closer to achieving a single component multifunctional
polymeric binder system, there still is a lot of room for performance improvement, specifically in
the improvement of Li
+
ion conduction. Ionic conductivity of conjugated polymers is generally
several orders of magnitude lower than the electronic conductivity.
195, 196
At the same time, ionic
diffusion distances required are generally much shorter than the distances required for effective
electrode based electronic conductivity, so the ideal compromise is not well defined. Elevating
ionic conductivity in order to enhance mixed conduction could significantly improve the rate
capability and capacity retention of LIBs, but a chemically tunable system is required to optimize
both. One of the ways of achieving this is by appending ion-transporting, polar oligoether side
chains to the conjugated backbone,
177, 197-199
where, similar to crown ethers, the oxygen atoms
along the side chains coordinate with Li
+
ion. For example, Reynolds et al. have reported an
oligoether functionalized propylenedioxythiophene (ProDOT)-based copolymer with redox
stability for thousands of redox cycles along with excellent charge-storage capacity.
199
It has also
been well reported that incorporation of oligoether side chains endows the conducting polymer

70
with a larger surface area, decreased π-stacking distance
123, 200
and a larger volumetric
capacitance.
201, 202
Thus, we directed our efforts towards incorporating oligoether side chains into
our previously reported PProDOT-Hx2 system.
In this section, we investigated a series of random copolymers based on dihexyl-substituted
propylenedioxythiophenes, whereby the hexyl side chains of the homopolymer (PProDOT-Hx2)
have been replaced to varying extents with oligoether side chains, (Hex:OE) PProDOTs. We have
presented a detailed experimental study for a comprehensive understanding of the balance between
electronic and ionic conduction and their ultimate usage as a mixed conductive binder for NCA
cathodes.
4.2.2 Synthesis of( Hex:OE) PProDOT Random Copolymers

Figure 4.15. Synthesis scheme of the (Hex:OE) PProDOT copolymer family.
The (Hex:OE) PProDOT random copolymer series was synthesized using direct arylation
polymerization as shown in Figure 4.15. The hexyl (Hex) versus oligoether (OE) side chain ratio
(Hex:OE) was tuned from (65:35) to (95:5) to generate a family of four random copolymers, which
were subsequently used for further characterization and ultimate investigation for cathode binder
applications in LIBs with NCA cathodes. Incorporation of 50 percent or more of OE relative to
hexyl side chains was also examined but resulted in dissolution of the polymer in battery
electrolyte as will be discussed in the following section. Therefore, high OE content over 50% is

71
unsuitable for polymer binder applications. The selected four (Hex:OE) random copolymers
(Figure 4.15), along with PProDOT-Hx2, were found insoluble in battery electrolyte and were
synthesized on the gram scale and analyzed relative to PVDF.
4.2.3 Electrochemical Properties of (Hex:OE) Random Copolymers
To determine the suitability of these polymers for use in LIB cathodes, we examined the
electrochemical doping (oxidation), potential window of stability, kinetics, and long-term cycling
of the polymer films in battery electrolyte. For reference we performed the same electrochemical
testing on PProDOT-Hx2. The electrochemical data for PProDOT-Hx2 is shown in Figure 4.16.

Figure 4.16. Electrochemical performance of PProDOT-Hx2. (a) Cycling Voltammograms at 10
mV s
−1
in various potential windows up to 4.5 V, (b) CV scans at different scan rates, and (c)
long-term cycling at 10 mV s
−1
.

The cyclic voltammograms at a slow scan rate of 5 mV s
−1
for the polymers (Figure 4.17)
showed two redox peaks. A main oxidation peak appears at 3.21 V vs Li/Li
+
. A second oxidation
peak appears at 3.77 V for PProDOT-Hx2. The second oxidation peak appears to shift to lower
potentials as the OE content increases to reach 3.69 V vs Li/Li
+
for (65:35) PProDOT. Similarly,
two main reduction peaks appear for all polymers. One is observed at 3.12 V while the second is
observed at 3.61 V and appears to slightly shift with changing OE content.

72

Figure 4.17. CV scans of (Hex:OE) PProDOTs at 5 mV s
−1
.
Furthermore, when the potential window limit was expanded gradually from 4.0 V to 4.5V
vs Li/Li
+
(Figure 4.18), no change in redox peak shape was observed, indicating that there is no
degradation upon polarization to 4.5 V vs Li/Li
+
. The wide operating window of 2.8 to 4.5 V makes
the (Hex:OE) PProDOT copolymer family ideally suited to work with various commercially
relevant LIB cathodes such as NCA.

73

Figure 4.18. CV data at 10 mV s
−1
for (a) (65:35) PProDOT, (b) (75:25) PProDOT, (c) (85:15)
PProDOT, and (d) (95:05) PProDOT at various potential intervals.

The CV scans collected at scan rates from 20 to 100 mV s
−1
between 2.8 V and 4.1 V vs
Li/Li
+
(Figure 4.19) showed that the electrochemical doping process was highly reversible up to
100 mV s
−1
. In addition, the minimal shift in the redox peaks with increasing scan rates indicated
that the electrochemical doping process was rapid. Furthermore, b-value analysis for the
PProDOT-OE polymers yielded a b-value close to 1 that indicated a non-diffusion controlled
process. A b-value of 0.95 and 0.93 for the (85:15) PProDOT was obtained for the first and second
oxidation peaks while a b-value of 0.96 and 0.98 was obtained for first and second reduction peaks,

74
respectively. This b-value analysis indicated that the kinetics of doping are rapid, and the polymers
can facilitate rapid electron and ion transport.

Figure 4.19. CV data at varying scan rates between 100 and 20 mV s
−1
for (a) (65:35) PProDOT,
(b) (75:25) PProDOT, (c) (85:15) PProDOT, and (d) (95:05) PProDOT.

The long-term electrochemical stability of the polymers was studied by repeated cycling
over the potential range 2.8 V to 4.1 V at 10 mV s
−1
for 100 cycles. As can be seen from Figure
4.20, a significant fraction of the capacity was retained after 100 cycles for all PProDOT:OE
polymers. For example, 78% and 71% capacity retention was observed for (95:05) and (85:15)
PProDots, respectively. Furthermore, no major peak shifts were observed indicating stable
electrochemistry for these polymers.

75

Figure 4.20. Long-term CV curves for (a) (65:35) PProDOT, (b) (75:25) PProDOT, (c) (85:15)
PProDOT, and (d) (95:05) PProDOT thin films at 10 mV s
−1
.

However, we observe a significant drop in capacity retention when we increase the OE
content above 35%. As we can see from Figure 4.21, the capacity retention drops to 33% for
(65:35).

76

Figure 4.21. Capacity retention of (Hex:OE) PProDOT copolymer family after 100 cycles at 10
mV s
−1
as a function of oligoether content.

Increasing the OE content above 35% leads to high-rate polymer film dissolution and loss
in capacity during extended cycling. This is confirmed by the CVs of the PProDOT-OE polymers
with high OE content (Figure 4.22) and visual inspection of the polymer films after cycling.

77

Figure 4.22. CV curves for (a) oligoether homopolymer, (b) (15:85) PProDOT, (c) (75:25)
PProDOT, and (d) (50:50) PProDOT thin films at 100 mV s
−1
.

4.2.4 Electronic and Ionic Conductivity of (Hex:OE) PProDOT Random Copolymers
To determine the effect of oligoether side chains of the (Hex:OE) PProDOTs on the
electron transport within the polymer structure, electronic conductivity of the polymer thin films
was measured using electrochemical impedance spectroscopy (EIS) in 1 M LiTFSI in EC/DMC
as a function of electrochemical doping (electrode potential) using the method discussed in Chapter
3. Electronic conductivities of the (Hex:OE) PProDOT copolymer family is shown in Figure 4.23.

78

Figure 4.23. Electronic conductivities of the PProDOT copolymer series.
At low electrochemical doping levels (below 3.2 V vs Li/Li
+
) all (Hex:OE) PProDOTs had
a similar electronic conductivity of ~5 × 10
−6
S cm
−1
at 2.8 V vs Li/Li
+
. The electronic conductivity
of the polymer films increased steadily to reach 10
−2
S cm
−1
at 3.2 V vs Li/Li
+
which aligned with
the first oxidation peak. The conducting polymer films reached their peak electronic conductivity
beyond 3.2 V, where the values started to diverge, with (65:35) PProDOT having the lowest
maximum electronic conductivity of 1.4 × 10
−1
S cm
−1
at 3.4 V vs Li/Li
+
followed by (75:25)
PProDOT which has an electronic conductivity of 2.4 × 10
−1
S cm
−1
at 3.4 V vs Li/Li
+
. The trend
continued with (85:15) PProDOT and (95:05) PProDOT with a maximum electronic conductivity
of 6.1 × 10
−1
S cm
−1
and 1.1

S cm
−1
, respectively.
Interestingly, the electronic conductivity of all (Hex:OE) PProDOTs dropped at potentials
higher than 3.4 V vs Li/Li
+
to reach a plateau of ~2 × 10
−2
S cm
−1
at high doping levels. Although
a similar overall trend was observed with the PProDOT-Hx2 homopolymer, a higher electronic
conductivity was observed at low doping levels (below 3.2 V vs Li/Li
+
) and high doping levels

79
(above 3.4 V vs Li/Li
+
). However, PProDOT-Hx2 reaches a maximum conductivity of 1 S cm
−1
,
similar to that of (95:5) PProDOT between 3.3 and 3.5 V vs Li/Li
+
. Thus, we can conclude that
introducing oligoether side chains into PProDOT-Hx2 lowered its electronic conductivity, and
there is an inverse relationship between the oligoether content and the maximum electronic
conductivity that is attained between 3.2 and 3.5 V vs Li/Li
+
.
EIS was also used to measure the ionic conductivity of polymer thin films in 1 M LiTFSI
in EC/DMC as a function of electrochemical doping (Figure 4.24). Upon doping, the ionic
conductivity of PProDOT-Hx2 reached 1 × 10
−7
S cm
−1
. However, upon introduction of OE side
chains, the ionic conductivity of (Hex:OE) PProDOTs increased by as much as a factor of four
with (65:35) PProDOT reaching 4 × 10
−7
S cm
−1
. The incorporation of only 5% oligoether side
chains doubled the ionic conductivity compared to PProDOT-Hx2.

Figure 4.24. Ionic conductivities of the PProDOT copolymer series.
We note that while the ionic conductivity is improved with OE incorporation, it is still
orders of magnitude lower than the electronic conductivity. This trend in ionic conductivity is well

80
supported by swelling studies in propylene carbonate (Figure 4.25), as we observe that ion
conduction is proportional to the amount of Li
+
coordinating OE, which in turn is proportional to
the amount of electrolyte uptake. We observe that there is an almost linear increase in swelling
with increasing OE content, culmination in a near 2x mass increase in the polymers with the
highest OE content.

Figure 4.25. Swelling study of the (Hex:OE) PProDOTs using propylene carbonate electrolyte.

Electrolyte uptake of this magnitude can easily generate ion-transport pathways, improving
ionic conductivity. Unfortunately, swelling simultaneously disrupts π-conjugation and thereby,
electron-transport pathways, which explains the inverse trend in both electronic and ionic
conductivity with increase in OE content within the random copolymer family. Importantly, this
result suggests that tuning electrolyte swelling may be the primary method for optimizing the
interplay between electronic and ionic conductivity in Li
+
battery systems, where polar organic
solvents will always have some propensity to swell polymer binders.
4.2.5 (Hex:OE) PProDOTs as Cathode Binders
To test the function and efficacy of these polymers in a practical battery electrode, NCA
cathodes employing (Hex:OE) PProDOTs, PProDOT-Hx2, and PVDF polymers as binders were

81
repeatedly cycled at a 1C rate for 200 cycles with two formation cycles at C/10 with a cathode
mass composition of 90% NCA, 4% binder, 6% carbon and an areal loading of 3.1 ± 0.4 mg cm
−2
.
All the electrodes with conducting polymer (CP) binders, barring (65:35) PProDOT, retained 17
to 40% higher capacity over the cells with PVDF binder (Figure 4.26). While all the (Hex:OE)
PProDOTs are found to be beneficial in retaining a higher capacity compared to PVDF, (75:25)
PProDOT retained the highest capacity. We attribute the improved cycling stability to the ability
of these polymers to maintain inter-particle electronic and ionic connectivity.

Figure 4.26. Specific capacity as a function of cycle number for the Li-NCA-(Hex:OE)
PProDOT cells.

It appears that solvent-induced and doping-induced swelling of oligoether containing
conducting polymers is beneficial in maintaining good electrical interconnectivity and binding of
the electrode particles during the repeated expansion and contraction during charge/discharge.
However, the results in Figure 4.26 indicate that the benefits require an optimal level of oligoether
group incorporation. Excessive swelling is clearly undesirable and leads to loss of binding action
as seen in the steep decay of cycle life with (65:35) PProDOT. The rapid decay in capacity of

82
electrodes with (65:35) PProDOT is consistent with the results of cyclic voltammetry studies on
the polymer film that suggested partial dissolution of the polymer during extended cycling.
4.3. Introducing Conjugation Break Spacers to Improve PProDOT-Hx2 Binding
4.3.1 Introduction
In the previous section, we have showed that introducing oligoether side chains and
controlling the OE:hexyl content ratio alters the electron and ion transport properties in the hex:OE
polymer family. The tuning of these properties has a direct effect on the performance of these
polymers as binders in battery electrodes. However, the physical properties of a polymer, and not
just its electronic properties, play a crucial role in the performance of the binder. A significant
amount of research has been published aiming to understand and improve the mechanical
properties of polymer binders in various LIB anodes and cathodes.
203-206

A common method to decrease the rigidity and alter the mechanical properties of
conjugated polymers is the introduction of carbon chains of varying lengths and content ratio in
the polymer backbone to break the conjugation.
207
These carbon chains, known as conjugation
break spacers (CBS), can have a large impact on both the mechanical properties and the processing
of the conjugated polymers.
208, 209
Furthermore, the mechanical properties can be carefully tuned
by systematically introducing CBS into the conjugated polymers.
210-212
In this section, we perform
a systematic study on the effects of introducing conjugation-break spacers (CBS) on the
mechanical and electronic properties of PProDOT-Hx2.
4.3.2 Synthesis of CBS PProDOT Polymers
Direct arylation polymerization was used as described in previous sections to synthesize a
family of nine PProDOT:CBS random copolymers with varying lengths of carbon chain (6-carbon,

83
8-carbon, and 10-carbon), referred to as T6T, T8T, and T10T, and CBS percent content (5%, 10%,
and 20%). Figure 4.27 showed a schematic of the synthesis procedure.

Figure 4.27. Schematic of CBS co-polymer family synthesis using DArP.
4.3.3 Electrochemical Performance of CBS:PProDOT Co-Polymers
To examine the effectiveness of the synthesized co-polymers as potential binders in LIB
cathodes, their electrochemical behavior in thin film format was examined in a three-electrode cell
with 1M LiTFSI in EC/DMC as the electrolyte and Li foil as the reference and counter electrodes.
All measurements were performed inside an argon glovebox. Initial CV curves were collected at
a scan rate of 50 mV s
−1
between 2.8 and 4.1 V vs Li/Li
+
. The initial CV curves of the T6T
copolymers are shown in Figure 4.28.

Figure 4.28. Initial CV data for T6T (a) 5%, (b) 10%, and (c) 20%.
In the first cycle, and oxidation peak appears at 3.51 V vs Li/Li
+
for T6T 5%, 3.47 V for
T6T 10%, and 3.59 V for T6T 20% (Figure 4.28). In subsequent cycles, the oxidation peak shifts
to 3.35 V, 3.31 V, and 3.46 V for T6T 5%, 10%, and 20%, respectively. A second oxidation peak
appears at 3.85 V and 3.87 V for T6T 5% and 10% while a first reduction peak appear at 3.17 V

84
and a second reduction peak appears at 3.70 and 3.74 V for T6T 5% and 10%, respectively.
Interestingly, only a very broad oxidation peak appears for T6T 20% and 2 broad overlapping
reduction peaks appear at 3.31 and 3.48 V vs Li/Li
+
. This suggests that the electrochemical doping
process is highly reversible for T6T 5 and 10% and much less reversible for T6T 20%.
Figure 4.29 shows the initial cycling for the T8T polymers. Similar to what was observed
with T6T polymers, an oxidation peak in the first cycle appears at 3.45 V vs Li/Li
+
for T8T 5%,
3.63 V for T6T 10%, and 3.41 V for T8T 20%. In subsequent cycles, the oxidation peak shifts to
3.32 V and 3.47 V for T8T 5% and 10%, respectively. A second oxidation peak appears at 3.85 V
and 3.88 V for T8T 5% and 10% while a first reduction peak appear at 3.14 V and 3.18 V. A
second reduction peak appears at 3.68 and 3.54 V for T8T 5% and 10%, respectively. On the other
hand, a very broad oxidation and reduction peak appears for T8T 20% after the first cycle at around
3.5 V vs Li/Li
+
. This suggests that although the electrochemical doping process is highly reversible
for T8T 5 and 10%, increasing the CBS content to 20% in T8T polymers decreases the reversibility
of the electrochemical doping process.

Figure 4.29. Initial CV data for T8T (a) 5%, (b) 10%, and (c) 20%.
Figure 4.30 shows the initial cycling of the T10T polymer family. The first cycle shows
several oxidation peaks starting at 3.45 V. With subsequent cycles, two oxidation peaks appear at
3.33 V and 3.84 V, and two reduction peaks appear at 3.15 V and 3.84 V vs Li/Li
+
, respectively.

85
A broad oxidation peak at 3.35 V appears in the initial cycle for T10T 10%. Two oxidation peaks
appear in subsequent cycling at 3.27 and 3.86 V, and two reduction peaks appear at 3.15 and 3.67
V vs Li/Li
+
. Although some redox activity is observed in the initial cycle for T10T 20%, there is
a significant loss in capacity in subsequent cycles suggesting that T10T 20% doping is not
electrochemically reversible.

Figure 4.30. Initial CV data for T10T (a) 5%, (b) 10%, and (c) 20%.
We expanded the potential window to higher potential to investigate the stable operational
range of the CBS polymers. CV curves were obtained with gradually increasing potential ranges
at 10 mV s
−1
. Figure 4.31 shows the obtained CV curves for the CBS T6T polymers. The shapes
of the redox peaks were maintained for the 5% polymer when going to 4.5 V suggesting it is
electrochemically stable in the wide potential range of 2.8 V to 4.5 V vs Li/Li
+
. Similarly, the
shapes of the peaks in T6T 10% were largely maintained with minimal shifts in the first oxidation
and reduction peaks. However, large shifts in the redox peaks were observed for T6T 20%
suggesting it is not stable in the wide potential window.

86

Figure 4.31. CV data for T6T (a) 5%, (b) 10%, and (c) 20% at various potential windows at 10
mV s
−1
.

Similarly, the shapes of the redox peaks for T8T 5% were retained up to a voltage cutoff
of 4.5 V (Figure 4.32). Small shifts in the redox peaks, however, were observed when expanding
the potential window while cycling T8T 10% suggesting it is less electrochemically stable in the
2.8-4.5 V window range. CV data for T8T 20% showed significant shifts in the redox peaks
suggesting it is not electrochemically stable.

Figure 4.32. CV data for T8T (a) 5%, (b) 10%, and (c) 20% at various potential windows at 10
mV s
−1
.

Figure 4.33 shows the CV data for T10T 5% and T10T 10% at 10 mV s
−1
at expanding
potential windows up to 4.5 V vs Li/Li
+
. Small shifts in the first oxidation peak are observed for
T10T 5% and significant shifts are observed for the reduction peaks. Peak shifts are also observed
for T10T 20% suggesting that T10T is not electrochemically stable when expanding the potential
window to 4.5 V vs Li/Li
+
. The stable and wide operation potential range of 3.2 to 4.5 V vs Li/Li
+

87
for T6T 5%, T6T 10%, and T8T 5% suggests that these three polymers are the most promising
CBS polymers for use as binder for cathode materials, including LiCoO2, LiMn2O4 and LiMO2
(M=Ni, Co, Mn, Al).

Figure 4.33. CV data for T10T (a) 5%, and (b) 10% at various potential windows at 10 mV s
−1
.

To ensure that electrochemical doping and de-doping of the CBS polymer binders does not
limit cathode cycling, we also examined the kinetics of the electrochemical doping at high rates.
Kinetics were analyzed based on a series of CV measurements at various scan rates from 10 to 100
mV s
−1
(Figures 4.24-4.26).

Figure 4.34. CV data for T6T (a) 5%, (b) 10%, and (c) 20% at various scan rates.
To quantify the kinetics of polymer doping, we examined the relation between the
measured current and scan rate as previously discussed in Section 4.1 where we calculated a b
value for each redox peak. A value of b equal to 0.5 indicates a process controlled by semi-infinite
diffusion, while a b value close to 1 indicates a non-diffusion controlled or a surface-controlled

88
charge-storage process. Figure 4.34 shows that the b values for all redox peaks, with the exception
of the oxidation peak for T6T 20%, are above 0.9 for T6T 5% and T6T 10%. These b values
indicate rapid redox processes in the polymer thin films.

Figure 4.35. CV data for T8T (a) 5%, (b) 10%, and (c) 20% at various scan rates.
Similarly, Figure 4.35 shows that the T8T polymer films can be electrochemically doped
at high scan rates. b values close to 1 for all redox peaks confirm rapid redox processes. Figure
4.36 also shows rapid doping kinetics for T10T 5% and T10T 10% polymers. Thus, the fast
kinetics of the electrochemical p-doping of the CBS polymers are expected to facilitate rapid
electron transport when used as conductive binders in cathode composites.

Figure 4.36. CV data for T10T (a) 5% and (b) 20% at various scan rates.
To examine the long-term electrochemical stability of CBS polymers, they were
continuously cycled between 2.8 V and 4.1 V vs Li/Li
+
at a scan rate of 10 mV s
−1
for 100 cycles.
Figure 4.37 shows that a significant fraction of the capacity is retained for T6T 5% and T6T 10%

89
after 100 cycles (64 and 59%, respectively). In addition, no major peak shifts were observed
indicating stable electrochemistry. However, we observe a significant drop in capacity for T6T
20% to just 21% after 100 cycles.

Figure 4.37. Long-term CV curves for T6T (a) 5%, (b) 10%, and (c) 20% thin films at 10 mV
s
−1
.
The CV data in Figure 4.38 shows that T8T 5% retains the highest capacity of 61% after
100 cycles. T8T 10% retains 54% of its initial capacity after 100 cycles, and we observe peak
shifts after 100 cycles suggesting that it may not be electrochemically stable in the long term. T8T
20% loses most of its capacity and is not suitable for long-term cycling.

Figure 4.38. Long-term CV curves for T8T (a) 5, (b) 10, and (c) 20% thin films at 10 mV s
−1
.
The CV data for T10T in Figure 4.39 shows that T10T 5% has good capacity retention
after 100 cycles with no significant change in redox peak shapes or positions. T10T 10%, however,
loses most of its initial capacity suggesting it is not suitable for long-term cycling.

90

Figure 4.39. Long-term CV curves for T10T (a) 5%, and (b) 10% thin films at 10 mV s
−1
.
4.3.4 Conductivity of CBS:PProDOT Co-Polymers
To determine the effect of conjugation break spacers in PProDOT on electron transport
within the polymer structure, electronic conductivity of the polymer thin films was measured using
electrochemical impedance spectroscopy (EIS) in 1 M LiTFSI in EC/DMC as a function of
electrochemical doping using the method discussed in Chapter 3.
Figure 4.40 shows the electronic conductivity of T6T polymers as a function of electrode
potential. At 2.9 V vs Li/Li
+
, the electronic conductivity of T6T 5% is 1×10
−5
S cm
−1
. This is half
an order of magnitude lower that PProDOT-Hx2. Upon doping, the electronic conductivity of T6T
5% increases to a maximum of 5×10
−2
S cm
−1
at 3.3 V. The electronic conductivity slightly
decreases reaching 9×10
−3
S cm
−1
at 3.7 V. The electronic conductivity is relatively constant above
3.7 V vs Li/Li
+
. Although a similar trend is observed with PProDOT-Hx2, its electronic
conductivity reaching 1 S cm
−1
, one-and-a-half orders of magnitude higher than T6T 5%. Upon
increasing the CBS content percent even further, the electronic conductivity drops further to
2×10
−6
S cm
−1
and 1×10
−6
S cm
−1
for T6T 10% and T6T 20%, respectively, when undoped. T6T
10% reaches a maximum conductivity upon doping of 1.9×10
−4
S cm
−1
at 3.3 V, approximately
four orders of magnitude lower than that of PProDOT-Hx2. The maximum electronic conductivity
of T6T 20% is another magnitude lower reaching 2×10
−5
S cm
−1
at 3.3 V vs Li/Li
+
.

91

Figure 4.40. Electronic conductivities of the T6T polymers.
A similar trend is observed when introducing the longer chain conjugation break spacers
(Figure 4.41). The electronic conductivity of T8T 5% reaches a maximum value of 2×10
−3
S cm
−1

at 3.3 V while the electronic conductivity of the T8T 10% polymer reaches a maximum value of
3×10
−5
S cm
−1
at 3.3 V vs Li/Li
+
. Those values are approximately three and five orders of
magnitude lower when compared to PProDOT-Hx2 for T8T 5% and T8T 10%, respectively.

Figure 4.41. Electronic conductivities of the T8T polymers.

92

Figure 4.42 Electronic conductivities of the T10T polymers.
The trend in electronic conductivity extends to the T10T polymers (Figure 4.42) where the
maximum electronic conductivity of T10T 5% is 3×10
−4
S cm
−1
at 3.3 V. The electronic
conductivity decreases further upon incorporating a higher CBS content % where the maximum
electronic conductivity of T10T 20% is 3×10
−6
S cm
−1
at 3.3 V vs Li/Li
+
.

Figure 4.43. Comparison of electronic conductivity of polymers with 5% CBS content.

93
To examine the effect of the conjugation break spacer length on the electron transport of
the CBS polymers, the electronic conductivities of all 5% CBS polymers were compared as a
function of potential (Figure 4.43). Upon introducing 5% of the 6-carbon break spacer, the
maximum electronic conductivity drops by one and a half orders of magnitude compared to
PProDOT-Hx2. Increasing the chain length to eight carbons reduces the maximum electronic
conductivity by another order-and-a-half to a total of three orders of magnitude compared to
PProDOT-Hx2. The 10-carbon conjugation break spacer further reduces the maximum electronic
conductivity by another order of magnitude. As a result, the maximum electronic conductivity of
T10T 5% is approximately four orders of magnitude lower than PProDOT-Hx2.

Figure 4.44. (a) Electronic and (b) ionic conductivities of the CBS polymers.
Although the introduction of conjugation break spacers has a significant effect on the
electronic transport properties of the conjugated polymers, the examination of ionic conductivity

94
as a function of electrode potential (Figure 4.44b) revealed that conjugation break spacers do not
have any significant effect on the ionic transport within the polymer.
To understand the observed trends in electronic conductivity for the various CBS, we
investigated the polymer structure of the T6T polymer family at the first and second oxidation
peak (Figure 4.45a) using grazing incidence wide-angle X-ray scattering (GIWAXS). This
experiment provides information about total scattering intensity (which correlates with
crystallinity), lattice constants (which change upon doping), and orientation of the polymer chains
with respect to the substrate. All polymers are fairly disordered, with only two semicrystalline
peaks observed: a lamellar (100) peak at ∼0.3−0.4 Å, which corresponds to the side-chain spacing
between the polymer chains, and a π-stacking (010) peak at ∼1.4 Å, which corresponds to the
distance between polymer chains along the lattice vector closest to the π-stacking direction.

Figure 4.45. GIWAXS data of (a) T6T polymers and (b) CBS polymers with 5% CBS content.
Figure 4.45a shows that the lamellar peak is the highest for T6T 5% followed by T6T 10%
then T6T 20% at the first oxidation peak. Similar, but much less pronounced, trend is observed for
the π-stacking peak as well. This suggests that increasing the CBS % content leads to more disorder
in the polymer structure which results in lower electronic conductivity. Increasing the potential to
the second oxidation peak reveals that the lamellar peak height is enhanced further compared to
the first peak suggesting further enhancement in the crystalline order of the T6T polymers.

95
However, this enhancement is the highest for T6T 5% which decreases with increasing CBS
content.
To understand the effect of conjugation break spacer chain length on the morphology of
the polymers, we compared the structure of the 5% CBS polymers using GIWAXS at the bipolaron
peak (Figure 4.45b). When increasing the chain length from six carbons to eight carbons, we
observe a significant reduction in the lamellar peak. Going to a ten-carbon chain reduces the
lamellar peak intensity even further. A small reduction in the π-stacking (010) peak is observed
for the T10T 5% polymer as well. This observation suggests that there is a direct correlation
between the crystallinity of the polymer films and the length of the CBS chain.
Based on electrochemical testing and electronic conductivity determination of the CBS
polymers, it was determined that CBS T6T 5%, T6T 10%, and T8T 5% would provide good
physical as well as electrochemical properties as conductive binders in LIB cathodes. Thus, NCA-
CBS and NCA-PVDF cathodes were prepared as described in Chapter 2. The cathodes with a
composition of 90% NCA, 3% binder, 7% carbon and an areal loading of 3.4 ± 0.4 mg cm
−2
were
cycled in coin cells with a lithium foil anode and 1M LiTFSI in EC/DMC. Figure 4.46a shows the
cycle life data for the NCA cells while Figure 4.46b shows the rate capability of the NCA
electrodes.

96

Figure 4.46. Full cell testing of CBS polymers. (a) Specific capacity as a function of cycle
number and (b) rate capability test.

As Figure 4.46a shows, all NCA-CBS cells retained significantly more capacity after 300
cycles compared to the NCA-PVDF cells. Furthermore, NCA electrodes prepared with T6T 10%
and T8T 5% retained higher capacity compared to PProDOT-Hx2. Thus, the introduction of
conjugation break spacers improves binder performance in NCA electrodes. Rate capability
studies (Figure 4.46b) revealed that cells with PProDOT-Hx2 and PProDOT-CBS significantly
improve the rate capability when compared to NCA-PVDF cells. The CBS cells were tested to 8C.
Specifically, CBS electrodes can retain up to 80 mAh g
−1
at 6C. Therefore, the electrochemical
data as well as the full cell testing shows that we were able to tune the composition of the multi-
functional conductive binders with conjugation break spacers of different chain sizes and CBS %
content to improve the rate capability and cycle life of lithium-ion cells.

97
CHAPTER 5.

IMPROVING THE PERFORMANCE OF LITHIUM-SULFUR BATTERIES
USING FUNCTIONALIZED CONJUGATED POLYMERS

5.1 Introduction
Mesoporous carbon has been reported as a strategy to decrease the polysulfide shuttle due
to the small pore diameter.
213, 214
However, the large surface area of the carbon often allows the
polysulfides to escape into solution. To address this problem, composites of the carbon materials
with polymers have also been widely investigated.
215
Unlike the commonly used polyvinylidene
fluoride (PVDF) binder, polymer materials such as polyvinylpyrrolidone can serve not only as a
binder but also as a trap for polysulfides.
216
If the chemical groups in the polymer can interact
strongly with the polysulfides, the shuttling can be reduced. In this regard, carbonyl-containing
functional groups such as esters and amides have been found to sequester the polysulfides through
a Li-O interaction.
217
Thus, polymers with good mechanical properties such as catechol-conjugated
chitosan sulfate and gum arabic that have an affinity for polysulfides, have been used as binders
leading to enhanced cycle life in Li-S batteries.
218, 219

Although such polymers are effective as binders and can trap the polysulfides, their non-
conductive nature impedes electron transport through the composite structure. Thus, the addition
of such non-conductive polymer coatings, although beneficial with respect to trapping the
polysulfide, can also increase the resistance of the carbon composites and reduce sulfur utilization
and rate capability.
216
Consequently, significant research has focused on using electronically
conductive polymers as binders for the sulfur electrode. To this end, several p-dopable polymers

98
(i.e., polymers capable of producing free charge carriers upon oxidation) such as polythiophene
and its derivatives, PEDOT, polyaniline, and polypyrrole have been utilized with the goal of
simultaneously enhancing electron transport in the sulfur electrode and reducing the polysulfide
shuttling.
220-226
Furthermore, there were attempts to develop dual ionic and electronically
conducting polymer binder for sulfur electrodes by copolymerizing poly(3,4-
ethylenedioxythiophene) (PEDOT) with polyethylene glycol (PEG).
227
While these polymers can
offer electronic conductivity and stability on p-doping in the potential range of 3.1 V to 4.3 V vs.
Li
+
/Li, they remain insulating below 3.1 V vs. Li/Li
+
.
196, 228
This voltage window of electronic
conductivity and stability of these p-dopable conducting polymers are well matched with the
cycling window of cathode materials such as nickel cobalt oxide.
110, 112, 196, 229
However, as the
voltage window for the cycling of the sulfur electrode is between 1.7 and 3.0 V vs. Li/Li
+
, it is
important to choose a polymer that is an electronic conductor in this potential range. Although p-
dopable conductive polymers are more common, there is no known p-dopable conjugated polymer
that will remain in the p-doped state below ~2.3 V vs. Li/Li
+
.
230-232
Any p-dopable polymer with a
lower oxidation potential would be remarkably unstable in air and profoundly difficult to
synthesize and handle. Therefore, n-dopable polymers were selected as the primary candidates in
this study.
There are several n-dopable polymers that will exist in the conducting state between 2.6 –
1.6 V vs. Li/Li
+
. A well-known example is N2200 (or P(NDI2OD-T2)); poly{[N,N′-bis(2-
octyldodecyl)-naphthalene-1,4,5,8-bis(dicarboximide)-2,6-diyl]-alt-5,5′-(2,2′-bithiophene)}
which has an onset of reduction at ~2.6 V vs. Li/Li
+
and remains conductive down to potentials
below 1.6 V.
233, 234
In addition, the carbonyl groups in the polymer backbone should increase the
interaction of the polymer with the polysulfides. Furthermore, the ion transport properties were

99
enhanced by generating a copolymer of N2200 containing oligoether side-chains
235
to form
N2200-OE (or PNDI (OD:OE)) with OD:OE ratio of 83:17.
229
The structures of these polymers
are shown in Figure 5.1.

Figure 5.1. Polymers investigated for binders in Lithium-Sulfur (Li-S) Batteries.
The structural, electron-transport and ion-transport properties of these polymers were
characterized and the interaction of the polymers with the polysulfides were measured. By
comparing the cycling performance and rate capability of lithium-sulfur cells with N2200, N2200-
OE, and PVDF, we could verify the multiple anticipated benefits of these functionalized
electronically conducting polymers as binders in the sulfur electrode.
5.2 Electrochemical Cycling of Polymer Thin Films
Thin films of N2200 and N2200-OE were electrochemically characterized in a three-
electrode cell with Li-S battery electrolyte in an argon glovebox between 3.0 and 1.7 V vs Li/Li
+
,
a potential range relevant for the sulfur electrode. Initial thin film cycling was performed at 100
mV s
−1
(Figure 5.2). In the first cycle, the cyclic voltammogram of N2200 showed a reduction
peak at 2.08 V and two oxidation peaks at 2.33 and 2.55 V vs Li/Li
+
. The capacity of N2200
increased significantly with subsequent cycles from 5.4 nAh in the first cycle to 92 nAh after 20
cycles. The cyclic voltammogram on the 25th cycle showed that the reduction peak shifted slightly

100
to 1.97 V and the oxidation peaks shifted to 2.39 and 2.53 V vs Li/Li
+
. The CV data was stable
after the 20th cycle with no peak shifts or further change in capacity. The slow increase in capacity
is likely associated with solvent swelling of the polymer network, which facilitates ion transport
through the film. The CV data for N2200-OE, on the other hand, showed a small reduction peak
at 2.55 V and a major reduction peak at 2.15 V vs Li/Li
+
(Figure 5.2b). A small oxidation peak
appears at 2.37 V and a major oxidation peak appears at 2.54 V vs Li/Li
+
. No major changes in the
CV data were observed in the subsequent cycles. The slight shifts to slightly more positive
potentials in the major cathodic peak for both polymers in subsequent cycles is attributed to the
enhanced conductivity of the polymers after electrochemical doping that led to lowering of the
overpotential for reduction. Unlike N2200, the capacity in the first cycle for N2200-OE was 298
nAh and slightly decreased to 287 and 283 nAh in the second and fifth cycles, respectively. This
is likely due to the ability of N2200-OE to swell significantly more compared to N2200, allowing
electrolyte to enter the structure and electrochemically dope the polymer film much faster than
with N2200.

Figure 5.2. Cyclic voltammograms of initial cycling for N2200 and N2200-OE.
To ensure that the electrochemical doping/de-doping of the polymer binders does not limit
cathode cycling, we also examined the kinetics of the electrochemical doping based on a series of

101
CV measurements at various scan rates (Figure 5.3) from 20 to 100 mV s
−1
. The minimal shifts of
the redox peaks with increasing scan rates indicate rapid reaction rates.

Figure 5.3. CV data of N2200 and N2200-OE as a function of scan rate.
Finally, we examined the long-term electrochemical stability of the polymer thin films by
cycling them between 1.7 and 3.0 V vs Li/Li
+
at 10 mV s
−1
for 100 cycles (Figure 5.4). Both
polymers exhibited stable redox properties in the potential range of 1.7 to 3.0 V vs Li/Li
+
over
multiple cyclic voltammogram scans

Figure 5.4. CV data for long term cycling of N2200 and N2200-OE thin films.
5.3 Electronic and Ionic Conductivity of Polymer Thin Films
To probe the electron and ion transport properties of N2200 and N2200-OE, we measured
the electronic and ionic conductivities of the polymer thin films in 1 M LiTFSI in DOL/DME as a

102
function potential of electrochemical doping. Cycling the polymers at a slow scan rate of 5 mV s
−1

(Figure 5.5) revealed a cathodic peak at 2.16 V and two anodic peaks at 2.31 V and 2.50 V vs
Li/Li
+
for N2200. In the cathodic region for N2200-OE, a small shoulder appears at 2.64 V and a
main peak appears at 2.30 V. For the anodic region, we observed a shoulder at 2.36 V and a main
oxidation peak at 2.51 V vs Li/Li
+
.

Figure 5.5. Cyclic voltammograms of N2200 and N2200-OE at 5 mV s
−1
.
Electronic conductivity data (Figure 5.6a) shows that the electronic conductivity of undoped
N2200 is 2×10
−6
S cm
−1
above 2.5 V vs Li/Li
+
, consistent with the absence of any significant
electrochemical doping inferred from the cyclic voltammograms (Figure 5.5). The electronic
conductivity of N2200 starts to increase below 2.5 V vs Li/Li
+
which aligns with the onset of
electrochemical n-doping of the polymer (Figure 5.5) and levels off at 2×10
−4
S cm
−1
around 2.2
V. Thus, we observed two orders of magnitude increase in electronic conductivity upon doping of
N2200. N2200-OE, on the other hand, has a slightly lower electronic conductivity of 8×10
−7
S
cm
−1
when undoped at potentials greater than 2.65 V vs Li/Li
+
. Its electronic conductivity starts to
increase at potential below 2.65 V, the potential aligning with the onset of n-doping as observed

103
in Figure 5.5, to reach a maximum of 1×10
−4
S cm
−1
just below 2.4 V. The electronic conductivity
remains relatively constant throughout the potential window tested. Thus, the electronic
conductivity of N2200-OE reaches a lower maximum conductivity compared to N2200, but the
onset of the conductivity increase occurs at a higher potential compared to N2200. The
introduction of oligoether side chains seems to lower electronic conductivity which has been
observed in our previous work.
229

Figure 5.6. (a) Electronic and (b) ionic conductivity of N2200 and N2200-OE as a function of
potential.

The ionic conductivity data (Figure 5.6b) revealed that N2200 has ionic conductivity of
4×10
−10
S cm
−1
when undoped. Upon doping, the ionic conductivity increases to 7×10
−9
S cm
−1
at
1.7 V vs Li/Li
+
. Figure 5.6b shows that N2200-OE has an ionic conductivity of 6×10
−9
S cm
−1
at
2.8 V (when undoped) and 2×10
−7
S cm
−1
at 1.7 V vs Li/Li
+
. Thus, the addition of the oligoether
side chains appears to contribute to over one-order of magnitude improvement in ionic
conductivity across the tested potential range. This increase in ionic conductivity is consistent with
the structural changes that occur during n-doping of the polymer, as observed in the GIWAXS
studies, discussed below. Furthermore, the increased ionic conductivity of N2200-OE compared
to N2200 is well supported by the swelling behavior of these polymers. Almost all organic polymer
swell to some degree when placed in organic solvents like the DOL:DME electrolyte used here,
and the OE side chains significantly enhance that process. The addition of the oligoether sidechains

104
increases the % mass on swelling in the DOL:DME solvent. We observed that neutral N2200 had
a 100% mass uptake, while N2200-OE had a 220% mass uptake. The solvent swollen binder has
many pathways for ionic motion through the polymer network and thus should show enhanced
ionic conductivity. Solvent swelling upon electrochemical doping is likely also responsible for the
increase in ionic conductivity upon doping.
5.4 Physical Characterization of Binders
GIWAXS was used to understand how the structure of N2200 and N2200-OE evolve
during the cycling process. GIWAXS provides information about lattice constants, scattering
intensity, and the orientation of the polymer with respect to the substrate. For the pristine N2200
polymer, shown in Figure 5.7a in black, the lamellar (100) peak, predominantly in-plane of the
substrate, is observed at 0.25 Å
−1
, corresponding to a d-spacing of 2.6 nm. The (020) pi-stacking
peak, out of plane of the substrate, is at 1.6 Å
−1
, corresponding to a d-spacing of 0.39 nm. Pristine
N2200-OE, shown in Figure 5.7b in black, has similar d-spacing. The lamellar (100) peak is
located at 0.27 Å
−1
and the pi-stacking (020) peak at 1.6 Å
−1
, corresponding to a d-spacing of 2.3
nm, and 0.39 nm, respectively. The smaller lamellar spacing for N2200-OE is due to the curling
of the oligoether sidechains.
229, 236, 237

In the configuration that these polymers are used, they are n-type polymers, where the
polymers are reduced to induce electronic conductivity. Thus, 2.3 V vs. Li/Li
+
corresponds to
partial reduction and 2.1 V vs. Li/Li
+
to higher doping. When the polymers are reduced, an n-
polaron, which carries a negative charge is delocalized along the polymer chain and Li
+
is the
counter-ion. For both polymers, electrochemical doping does not change the orientation of the
polymers, but it does change the crystallinity, and the extent of change of crystallinity is different.
While the crystallinity of the N2200-OE polymer decreases somewhat upon doping, the overall

105
polymer structure remains relatively unchanged (Figure 5.7b). Aside from the decrease in intensity
of the principal (100) lamellar peak upon doping, all other peaks retain similar intensities and the
peak positions are constant, regardless of the doping potential. This indicates that the small Li
+

cation is likely incorporated in the polar oligoether side chains and does not significantly affect
the overall backbone stacking. On the other hand, the crystallinity of N2200 (Figure 5.7a) is more
significantly reduced by electrochemical doping, particularly at lower potentials. This is likely
because there is no low-energy place for the Li
+
to sit in the lamellar side chain region, and so the
Li
+
associates more closely with the polar polymer backbone.

Figure 5.7. Integration of the GIWAXS diffractograms of (a) the full integration of N2200 and
the inset is of the in-plane integration of the polymer, and (b) the full integration of N2200-OE
and the inset is of the in-plane integration of the polymer.

To examine the interaction between the conducting polymer binders and the polysulfides,
UV-Vis spectra of LiPS solutions were collected after exposing the solutions to polymer

106
electrodes. Figure 5.8 shows the UV-Vis spectra of LiPS solutions after being exposed to PVDF,
N2200, and N2200-OE over a period of three hours. No significant change in the LiPS spectrum
is observed when it is exposed to the PVDF electrode. However, a significant reduction in the UV-
Vis absorption is observed when LiPS solution is exposed to either N2200 or N2200-OE. This was
visually confirmed by taking photographs of the LiPS solutions (Figure 5.9).

Figure 5.8. UV-Vis spectra of LiPS exposed to (a) PVDF, (b) N2200, and (c) N2200-OE. (d)
Absorbance at 470 nm of LiPS solution after 3 hours of polymer exposure.

Figure 5.8d shows that LiPS absorption at 470 nm, which corresponds to the formation of
Li2S3,
238
rapidly decreased when exposed to either N2200 and N2200-OE electrodes, although a
slightly lower absorption was observed for the LiPS solution exposed to N2200-OE. This suggests

107
that both conducting polymers are effectively and continuously binding to the polysulfides
suggesting that they may reduce the polysulfide shuttle in LiS batteries.

Figure 5.9. Photographs of LiPS (a) before exposure to polymers and after 3 hours of exposure to
(b) PVDF, (c) N2200, and (d) N2200-OE.

5.5 Polysulfide Shuttle Current Measurement
The polysulfide shuttle current measurement is a methodology previously reported by our
team
239
that quantifies the degree of polysulfide shuttling occurring between the electrodes in a Li-
S cell. Briefly, the measurement is carried after “cell formation” where the cell is discharged to a
target state-of-charge (SOC), followed by an open circuit voltage step that leads the voltage to
relax, then, when voltage reaches a steady-state value it is held constant. The applied current that
is required to maintain the potentiostatic conditions is known as the “polysulfide shuttle current”
and changes as a function of time exhibiting a transient region at the beginning followed by a
steady-state region or plateau (Figure 5.10).
239

108

Figure 5.10. (a) Potentiostatic intermittent shuttle current measurement curve of a Li-
Sulfur-N2200 cell. (b) Voltage profile prior to and during shuttle current measurements on a Li-
Sulfur-N2200 cell.

In the potential range where the soluble polysulfides are present, the observed steady-state
current is attributed to the polysulfide shuttle. 70 wt% sulfur Li-S cells were assembled without
additives to determine the shuttle current. At approximately 2.3 V vs Li/Li
+
, the Li-S-N2200 and
Li-S-N2200-OE cells exhibited four times lower shuttle current than the Li-S-PVDF cells (Figure
5.11a). Figure 5.11b shows the observed shuttle current as function of potential in the polysulfide
soluble potential region that goes from 2.4 to 2.15 V vs Li/Li
+
, where the conducting polymer (CP)
cells had a lower shuttle current than those with PVDF suggesting that N2200 and N2200-OE
interact with the soluble polysulfides diminishing the shuttling. This is consistent with the
observed interactions in the UV-Vis measurements (Figure 5.8).

109

Figure 5.11. Polysulfide shuttle current measurement. a) Potentiostatic observed current as a
function of time at approximately 2.3 V vs Li/Li
+
after an OCV relaxation step and b) steady-
state current or shuttle current as a function of potential during discharge for sulfur-N2200),
sulfur-N2200-OE and sulfur-PVDF electrodes.

Furthermore, to determine the benefit of the polymer during cycling we measured the
shuttle current in the 35
th
cycle (Figure 5.12). The shuttle current remains significantly lower in
the Li-S-N2200 cell compared to the Li-S-PVDF cell showing an interaction between the CP
polymers and LiPS even after continuous cycling.

Figure 5.12. Steady-state current or shuttle current as a function of potential after 35 cycles
during discharge for a Li-Sulfur-N2200 and Li-Sulfur-PVDF cells.

110
5.6 Performance of Lithium-Sulfur Full Cells
5.6.1 Electrochemical Impedance Spectroscopy
Electrochemical Impedance Spectroscopy (EIS) was measured as function of SOC at
various points on the galvanostatic charge-discharge (GCD) curves for the Li-S-N2200, Li-S-
N2200-OE, and Li-S-PVDF cells (Figure 5.13a). While the difference between the cell voltage
following current interruption and the discharge/charge voltage is indicative of the polarization
losses, greater insights and resolution of the polarization contributions are gained through the
analysis of the impedance of the cell as function of frequency.
The real part of the impedance response (Zre) normalized to the electrode active mass
loading (ohm mg
−1
) as a function of frequency was analyzed at 2.28, 2.15, and 2.1 V vs Li/Li
+

(Figure 5.13b). For all SOCs, at high frequencies, Zre had the almost the same value indicating that
the ohmic resistance of the electrolyte is similar for all cells. However, at mid and low frequencies,
the cells with the conducting polymers (CPs) had approximately 4 to 5 times lower impedance
than the PVDF cells. The reduction in impedance of the CP cells is attributed to the enhanced
charge-transfer kinetics and lower mass transport resistance arising from the increased electronic
and ionic conductivity of electrodes with conducting polymer binders that leads to improved
utilization of the sulfur electrode.

111

Figure 5.13. (a) GCD profiles for Li-sulfur-N2200, Li-sulfur-N2200-OE, and Li-sulfur-PVDF
cells at C/20. Potentiostatic impedance response was measured at OCV as a function of SOC
during charge and discharge (denoted with yellow circles). (b) Real part of the impedance
response against frequency for the Li-sulfur-N2200, Li-sulfur-N2200-OE, and Li-sulfur-PVDF
cells at approximately 2.28 V, 2.15 V, and 2.1 V vs Li
+
/Li.

5.6.2 Full Cell Electrochemical Cycling
Rate capability of Li-S-N2200, Li-S-N2200-OE, and Li-S-PVDF cells was tested at
moderate (1.79 mg cm
−2
) and high sulfur electrode loadings (3.4 mg cm
−2
). These tests were
conducted following 24 hours of equilibration after assembly to allow adequate time for polymer
swelling. Following equilibration, we carried out two “formation” cycles at C/15 rate before
conducting the rate capability tests.
At moderate loadings, the CP cells showed a higher rate capability than the PVDF cells
(Figure 5.14a) consistent with the results of EIS measurements (Figure 5.13b), confirming the
benefit of the CPs in reducing the impedance to electron and ion transport. At 2C rate, the Li-S-
N2200-OE cells had a remarkable specific capacity of more than 400 mAh g
−1
attributed to the
polymer’s enhanced electronic and ionic conductivity that leads to a higher accessible surface area

112
for reaction and ion transport in the sulfur electrode. Even more significant, for the higher loading
electrodes, the N2200-OE and N2200 cells showed even higher rate capability compared to the
PVDF cells, indicating that the higher ionic and electronic conductivity of the CP binders is even
more important in thicker electrodes, where Li
+
transport is hampered due to the higher electrode
tortuosity and lower accessible surface area (Figure 5.14b).
240

Figure 5.14. Rate Capability for Li-sulfur-n-dopable polymer cells. Specific capacity as a
function of cycle number with increments in the discharge rate and a constant charging rate at
C/10 for sulfur electrode with (a) moderate loadings and (b) high loadings.

GCD cycling and EIS as a function of cycle number was measured to test the effectiveness
of the CPs in enhancing the cycle-life of the cells. After 500 cycles, the Li-S-N2200 and Li-S-
N2200-OE cells had an impressive capacity retention of approximately 67 and 82%, respectively
(Figure 5.15). In contrast, the Li-S-PVDF cell showed significant capacity decay in the first 100
cycles with a capacity retention of only 18% after 200 cycles.

113

Figure 5.15. Specific capacity as a function of cycle number at a discharge/charge rate of C/2 for
sulfur-N2200, sulfur-N2200-OE and Sulfur-PVDF.

Figure 5.16 shows the Nyquist plots at the 15
th
, 200
th
, and 495
th
cycles. Compared to the
CP cells after 200
th
cycles the PVDF cells showed 5 and 4 times higher resistance at mid-low and
high frequencies respectively, indicating an increase in the faradaic, mass-transport, and ohmic
resistance after cycling, attributed to the effects of polysulfide shuttling in the cell. Thus, n-dopable
CPs can lower the cell’s overall impedance and reduce the polysulfide shuttling to extend the
battery cycle-life. Furthermore, in comparing the CP containing cells after the 495
th
cycle, the
N2200-OE cell had approximately 3 times lower impedance than the N2200 cell, suggesting higher
effectiveness of the N2200-OE binder (Figure 5.16b). We expect the mechanical elasticity induced
by the increased degree of swelling in the N2200-OE binder to benefit the expansion and
contraction of the sulfur electrode.

114

Figure 5.16. Impedance response as function of cycle number at approximately 2.2 V vs Li
+
/Li
for the same (a) Li-sulfur-N2200, (b) Li-sulfur-N2200-OE, and (c) Li-sulfur-PVDF cells for the
15
th
, 200
th
, and 495
th
cycles.

5.6.3 Post-Cycling Analysis
The lithium electrodes were extracted from the coin cells after 200 cycles for the Li-S-
PVDF cells and 500 cycles for the Li-S-N2200 and Li-S-N2200-OE cells. SEM images of the
lithium electrodes were compared to SEM images of pristine lithium (Figure 5.17). The lithium
electrode from the cycled Li-S-PVDF cells showed a significant color change to yellowish brown
unlike the lithium electrodes from the Li-S-N2200 and Li-S-N2200-OE cells. Furthermore, the
SEM images revealed that the lithium electrode from the cycled PVDF cell had a non-
homogeneous surface with two distinct regions: one with increased surface roughness compared
to the pristine lithium, and one with a smooth surface resembling that of the pristine lithium.
Although a rough surface is expected due to the continuous stripping and plating of lithium upon
cycling, the areas of smooth surface are likely due to surface passivation due to the deposition of
the insoluble sulfides resulting from polysulfide shuttling. The lithium electrode surface of the
N2200 and N200-OE cells, on the other hand showed a more homogeneous roughened surface
indicative of uniform stripping and plating of lithium without any noticeable surface passivation.

115

Figure 5.17. SEM images of (a) pristine lithium and lithium from cycled (b) (b) PVDF, (c)
N2200, and (d) N2200-OE cells along with photographic images of the electrodes (insets).

To understand the impact of polysulfide shuttling in the full cells, EDS analysis was performed
the lithium electrodes from the cycled cells (Figure 5.18).
241
We analyzed specifically for sulfur,
as we were interested in the formation of insoluble sulfides due to polysulfide shuttling. Since
sulfur can also arise from the decomposition of the lithium TFSI salt, we also analyzed for the
fluoride content.

116

Figure 5.18. EDS spectra of (a) pristine lithium, and lithium from cycled (b) PVDF, (c) N2200,
and (d) N2200-OE cells.

The weight ratio of sulfur to fluorine (S/F) on the surface of the lithium anodes of all the
cycled cells was then calculated and compared to that of LiTFSI. An increase in S/F ratio is
indicative of an increase in sulfur content on the electrode surface from the formation of insoluble
sulfides. Table 5.1 shows the weight % of sulfur and fluorine along with the S/F weight ratio.
Table 5.1. Sulfur and Fluorine Content on Lithium Electrodes Using EDS Analysis
Sample Sulfur Weight % Fluorine Weight % S/F Ratio
LiTFSI 25 53 0.47
PVDF 36 8 4.50
N2200 15 45 0.33
N2200-OE 10 32 0.31

The S/F weight ratio on the lithium electrode of the PVDF cells of 4.50 was significantly
higher than that of LiTFSI (0.47) indicative of a substantial content of sulfur from insoluble
sulfides formed by the higher level of polysulfide shuttling process in these cells. On the other
hand, the S/F ratio for N2200 and N2200-OE cells were 0.33 and 0.31, respectively. These values
of S/F ratio were comparable to that of plain lithium TFSI salt. Thus, we concluded that the surface

117
of lithium anode in cells with N2200-based polymer binders at the cathode had almost no insoluble
lithium sulfides formed at the anode. These results are consistent with the substantially reduced
value of polysulfide shuttle current and explain the significantly longer cycle life observed in the
N2200-polymer cells.
Thus, we have demonstrated that n-dopable conducting polymers N2200 and N2200-OE
when used as binders at the sulfur electrode can significantly enhance the capacity output and cycle
life of lithium-sulfur batteries compared to standard insulating binders such as PVDF. The benefit
of these conducting polymer binders resulted from the electronic and ionic conductivity of the
polymers over the range of potential relevant to the operation of the sulfur electrode. Furthermore,
the strong affinity of the polysulfides to the functional groups on the polymer curtailed the degree
of polysulfide shuttling in the electrolyte leading to markedly enhanced capacity retention as high
as 82% after 500 cycles. The analysis of the lithium anode after cycling indicated the absence of
insoluble sulfides and the significantly increased cycle life confirmed the suppression of the
polysulfide shuttling with the use of N2200 polymer binder at the sulfur cathode. The modification
of the N2200 polymer with oligoether side chains increases the swelling of the polymer and results
in enhanced ionic conductivity. The electronic and ionic conductivity of the conducting polymer
binders led to a significant reduction in impedance of the lithium-sulfur cell compared to the
baseline PVDF cells, and was manifested as a notable increase in rate capability. These benefits
are observed both at moderate and high loadings of the sulfur electrode. These results emphasize
the advantages of using conducting polymer binders that can support electron and ion transport in
the potential range of electrode operation, and the presence of functional groups in the polymer
that interact strongly with the soluble polysulfides in addressing the limitations of lithium-sulfur
cells.

118
CHAPTER 6.

PROBING TRANSITION METAL PEROVSKITE CHALCOGENIDES
AS PROMISING ELECTROCATALYSTS

6.1 Introduction
Transition metal oxides have been extensively studied for their activity as electrocatalysts
for various reactions including carbon dioxide reduction, oxygen reduction, the oxygen evolution
reaction (OER), and the hydrogen evolution reaction (HER).
242-248
In addition, perovskite oxides,
with the chemical formula ABO3, where A is an alkali, alkaline or rare earth metal and B is an
octahedrally coordinated transition metal, have been also examined as electrocatalysts for water
electrolysis.
249-254
However, the large energy differences between the transition metal d-orbital
conduction band and the O 2p orbital valence band lead to very large energy gaps of over 3 eV,
resulting in relatively lower electronic conductivity and no photo-activity.
Transition metal perovskite chalcogenides of the formula AB(S, Se)3, on the other hand,
possess excellent photo-activity and higher electrical conductivity due to smaller band gap and
higher electronic mobility of charge carriers but are significantly less explored as
electrocatalysts.
255-259
In this chapter, we examine a variety of transition metal chalcogenides such
as BaMS3 (M= Ti, Zr, V) and BaTiSe3, as well as lanthanum transition metal sulfides (LaMS3)
where M = Mn, Fe, Co, Ni as electrocatalysts for water electrolysis. Specifically, we characterize
the materials as electrocatalysts for HER in 0.5 M sulfuric acid and for OER in 1 M potassium
hydroxide.

119
6.2 Synthesis and Physical Characterization
BaTiS3, BaTiSe3, and BaZS3 were synthesized as described before
260-262
using
conventional solid-state reaction of barium sulfide or barium selenide, elemental sulfur, and
titanium or zirconium with iodine as the catalyst. Lanthanum metal sulfides were synthesized using
sol-gel method in which the LaMO3 precursor was first synthesized by mixing lanthanum nitrate
and the transition metal nitrate with citric acid in 1:1:2 molar ratio in deionized water then dried
at 120 ℃. The Fe and Co mixtures were then heated for 48 hours at 1300 ℃ while the Ni and Mn
mixtures were heated for four hours at 700 ℃. Powder X-ray diffraction data for the synthesized
transition metal oxides (Figure 6.1) confirms the successful synthesis of the oxides. The
synthesized materials were supplied for electrochemical characterization by Prof. Jayakanth
Ravichandran and his graduate students.

120

Figure 6.1. Powder XRD spectra of (a) LaMnO3, (b) LaFeO3, (c) LaCoO3, and (d) LaNiO3.
To prepare the corresponding sulfides, the metal oxides were sulfurized using CS2 gas. A
CS2 annealing setup was used to convert LaMO3 to LaMS3 as shown in Figure 6.2.

121

Figure 6.2. CS2 annealing setup for LaMS3 sulfurization.
LaFeO3 was annealed at 975 ℃ for twenty hours, while LaCoO3 and LaNiO3 were annealed
at 1100 ℃ for five hours. The corresponding powder x-ray diffraction patterns for the isostructural
LaMS3 powders are shown in Figure 6.3.

Figure 6.3. Powder XRD patterns for LaMS3 with representative reference pattern for LaNiS3.

122
It is important to note that obtaining phase pure materials proved to be challenging. In fact,
the synthesis of LaMnS3 did not yield a product with good purity as can be seen from the powder
XRD pattern in Figure 6.4. Nonetheless, the obtained material was still tested for electrocatalytic
properties.

Figure 6.4. Powder XRD patterns for LaMnS3 annealed for (a) four hours and (b) 28 hours.
6.3 Hydrogen Evolution Reaction
The electrocatalytic activity of the perovskite sulfides for the hydrogen evolution reaction
was examined as described in Chapter 2 in 0.5 M sulfuric acid. Initially, catalyst inks were
prepared with 20% by mass Super P and without Super P as conducting carbon additive. As Figure
6.5 shows, the addition of conducting carbon significantly reduced the overpotential for all BaMS3
samples by almost 200 mV at 10 mA cm
-2
. Thus, conducting carbon was added to all subsequent
LaMS3 inks tested.

123

Figure 6.5. Polarization curves for (a) BaTiS3, (b) BaZS3, and (c) BaVS3.
All three BaMS3 materials were shown to be active towards the hydrogen evolution
reaction with BaZrS3 having the lowest overpotential of 546 mV at 10 mA cm
−2
. BaVS3 had the
second lowest overpotential of 556 mV at 10 mA cm
−2
while BaTiS3 had the highest overpotential
of 656 mV at 10 mA cm
−2
. BaTiSe3 did not show significant activity towards HER.
To further examine the mechanism by which the barium transition metal sulfides catalyze
the hydrogen evolution reaction, Tafel slopes were plotted as shown in Figure 6.6. The
corresponding Tafel slopes obtained from the polarization curved were 165, 134, and 130 mV
decade
−1
for BaTiS3, BaZS3, and BaVS3, respectively. The addition of Super P did not result in a
significant change in the Tafel slope except in the case of BaZrS3 at high current densities.

Figure 6.6. Tafel slopes of (a) BaTiS3, (b) BaZS3, and (c) BaVS3.
The Tafel slope values can be used to determine the rate limiting step in the hydrogen
evolution reaction. As Table 6.1 shows, the first step in the hydrogen evolution reaction in acidic

124
media is the adsorption of the hydronium ion in the Volmer step. The adsorbed hydrogen atoms
then form hydrogen gas by the Heyrovsky pathway or the Tafel pathway. The obtained Tafel slopes
are closest to 120 mV decade
−1
indicating that the rate limiting step is most likely the Volmer step.
Table 6.1 HER Mechanism in Acidic Conditions
Step Reaction Tafel Slope (mV decade
−1
)
Volmer
Surface+H
!
O
"
+e
#
⇌H
$%&
+H
'
O 120
Heyrovsky
H
$%&
+H
!
O
"
+e
#
⇌H
'())
+H
'
O 40
Tafel
2H
$%&
⇌H
'())
30

Lanthanum transition metal sulfides were also examined for the electrocatalytic activity
for the hydrogen evolution reaction. As shown in Figure 6.7a, all the LaMS3 samples examined
had a significantly higher electrocatalytic activity compared to barium transition metal sulfides.
Furthermore, a trend is observed as we substitute the transition metal. LaMnS3 had the lowest HER
activity with an overpotential of 590 mV at 10 mA cm
−2
. Replacing Mn with Fe, however,
significantly improved HER activity to achieve an overpotential of 360 mV at 10 mA cm
−2
.
LaCoS3 achieved similar performance to iron with an overpotential of 364 mV at 10 mA cm
−2
.
LaNiS3 showed the highest activity for HER with an overpotential of 327 mV at 10 mA cm
−2
. The
corresponding Tafel slopes shown in Figure 6.7b suggest that LaMnS3, LaFeS3, and LaCoS3 have
the Volmer step as the rate-determining step while LaNiS3 may have the Heyrovsky or Tafel step
as the rate determining step.

125

Figure 6.7. Electrocatalytic activity of various transition metal perovskite sulfides LaMS3
towards hydrogen evolution (a) polarization curves and (b) the corresponding Tafel slopes.

6.4 Oxygen Evolution Reaction
The electrocatalytic performance of BaTiS3, BaZrS3, and BaTiSe3 in the oxygen evolution
reaction was examined in 1 M potassium hydroxide as described in Chapter 2. Unlike the hydrogen
evolution reaction, all examined barium transition metal chalcogenides showed no activity towards
the oxygen evolution reaction. In fact, the absence of a significant current in the polarization curves
of Figure 6.8 suggest that the electrocatalytic surface is passivated upon oxidative polarization,
rendering it electrochemically inactive. Thus, it was concluded that the tested materials were only
suitable as HER catalysts and could not be used as OER catalysts in alkaline conditions.

126

Figure 6.8. OER Polarization curves for barium transition metal chalcogenides.
Lanthanum transition metal sulfides, on the other hand, showed mostly good activity
towards OER in alkaline conditions (Figure 6.9). LaCoS3 showed the highest activity with an
overpotential of 370 mV at 10 mA cm
−2
. LaNiS3 and LaFeS3 also showed good catalytic activity
with overpotentials of 387 and 389 mV at 10 mA cm
−2
, respectively. Interestingly, LaMnS3
showed poor activity towards OER with a Tafel slope of 128 mV decade
−1
. The Tafel slopes for
the other lanthanum transition metal sulfides LaCoS3, LaFeS3, and LaNiS3, were 55.6, 62.2, and
78.6 mV decade
−1
, respectively, considerably lower than the corresponding nickel containing
sulfide.

127

Figure 6.9. Electrocatalytic activity of LaMS3 materials for oxygen evolution, (a) OER
polarization (b) the corresponding Tafel slopes.

Although LaMnS3 did not show good catalytic activity towards OER, this result was not
surprising as the oxide counterpart is known to be a poor OER electrocatalyst.
263
As can be seen
in Figure 6.10a, there was no significant difference in the polarization curved between LaMnO3
and LaMnS3. When comparing the OER activity of the rest of the lanthanum transition metal
sulfides to their oxide counterparts, however, we observed a significant enhancement in OER
catalytic activity for the sulfides. For example, the overpotential of LaCoS3 was almost 100 mV
lower than that of LaCoO3 at 10 mA cm
−2
as can be seen in Figure 6.10b. This enhanced activity
over the oxide counterparts was also observed for the other lanthanum metal sulfides where LaFeS3
showed an overpotential that is over 120 mV lower than that of LaFeO3.
264-266
Thus, we can
conclude that the lanthanum transition metal sulfides could be significantly better catalysts for the
oxygen evolution reaction in alkaline conditions compared to their lanthanum transition metal
oxide analogs.

128

Figure. 6.10. Polarization curves for oxygen evolution reaction on (a) LaMnS3 and LaMnO3, and
(b) LaCoS3 and LaCoO3.

In summary, all tested BaMS3 compounds showed activity towards HER with BaZrS3
showing the highest activity with an overpotential of 546 mv at 10 mA cm
−2
. LaMS3 compounds
showed significantly higher activity towards HER compared to the barium transition metal
sulfides, with LaNiS3 showing the highest activity with an overpotential of 327 mV at 10 mA cm
−2
.
On the other hand, barium transition metal sulfides showed no activity towards OER while
lanthanum transition metal sulfides showed improved OER activity especially when compared to
their lanthanum transition metal oxide analogs.

129
CHAPTER 7.

DEVELOPING AN EFFICIENT AND INEXPENSIVE ALL-IRON WATER
ELECTROLYSIS DEVICE
This dissertation chapter is based on a recent publication: Zayat, B.; Mitra, D.; Narayanan, S. R.
Inexpensive and Efficient Alkaline Water Electrolyzer with Robust Steel-Based Electrodes. J.
Electrochem. Soc. 2020, 167, 114513. © IOP Publishing. Reproduced with permission. All rights
reserved.

7.1 Introduction
Almost 95% of the hydrogen produced in the U.S. today is by steam-reforming of methane
in which methane gas reacts with steam at high pressure and temperature to produce hydrogen and
carbon dioxide, contributing to 3% of global carbon dioxide emissions.
267
Water electrolysis, on
the other hand, is a promising alternative for clean hydrogen production where water is converted
to hydrogen and oxygen in a single step using electricity from renewable sources. Yet, water
electrolysis only accounts for less than 5% of hydrogen production.
268
Water electrolysis may be
carried out under either acidic or alkaline conditions. Alkaline water electrolysis allows the use of
less expensive catalysts for the hydrogen evolution reaction (HER) and the oxygen evolution
reaction(OER).
66, 67
In addition, alkaline systems have been in operation for over 100 years and
their performance can rival that of the acidic electrolyzers.
68
For large-scale adoption, the water
electrolyzers must use low-cost and abundantly-available materials, so that the approach is
economically competitive and sustainable. Thus, our studies have focused on enabling the

130
economic competitiveness of alkaline water electrolysis through designing electrodes with low-
cost materials.
Nickel is widely used as an electrode material in alkaline media due to its high stability
under cathodic and anodic conditions. Nickel also exhibits high electrocatalytic activity compared
to many other transition metals.
269-271
Furthermore, nickel alloys such as nickel-molybdenum
exhibit higher activity for hydrogen evolution than nickel itself, with overpotentials around 150
mV at 10 mA cm
−2
.
272-280
Transition metal oxides that combine nickel, iron and cobalt are
commonly used for OER with overpotentials of approximately 300 mV at 10 mA cm
−2
.
76, 281-283

However, such electrodes are usually prepared by applying a thin coating of the electrocatalyst on
a high-surface area nickel foam electrode substrate. The nickel substrate constitutes more than
95% of the mass of the electrode and adds a significant cost to the electrolysis system. Although
the core of the earth is believed to be made of molten nickel, the abundance of nickel in the earth’s
crust is about 0.01%. Consequently, the large-scale adoption of nickel as an electrode substrate
material for water electrolysis would be a challenge from the viewpoint of cost and sustainability.
Alternate materials must not only be inexpensive and abundant but also support current densities
in excess of 100 mA cm
−2
and be durable for several years. To this end, we have focused on
replacing nickel substrates with iron-based substrates, specifically those made with low-carbon
steel.
Iron is 600 times more abundant than nickel in the earth’s crust. Iron is also 50 times
cheaper than nickel. The cost savings from the use of iron is further amplified when the cost of
industrial steel mesh is compared with specialized nickel foam. Unmodified iron, however, suffers
from chemical instability under anodic potentials in alkali solutions. The insulating passive film
formed on the surface of iron in alkali breaks down under anodic conditions resulting in the

131
dissolution of iron. In previously published work from our group,
284, 285
iron was rendered stable
and active towards oxygen evolution by surface modification with a thin layer of high-surface area
α-nickel hydroxide. Such an electrode was stable for over 1000 hours of operation with
performance comparable to nickel-based catalysts.
Building on that work, we used low carbon steel mesh as a substrate to further reduce the
overall electrode costs. We examined surface modifications using nickel nitrate and cobalt nitrate
at various temperatures in ambient atmosphere to produce robust and inexpensive oxygen
electrodes.
286
In addition, we examined direct current magnetron sputter deposition of catalytically
active metals on low carbon steel to produce equally robust and inexpensive hydrogen electrodes.
Finally, we utilized the developed oxygen and hydrogen electrodes to fabricate an all-iron alkaline
water electrolyzer device that is capable of operating at high current densities of up to 1 A cm
−2
.
287

7.2 Steel-Based Electrodes for Oxygen Evolution
7.2.1 Effect of Annealing Temperature on Nickel-treated Electrodes
To examine the effect of the annealing temperature on the catalytic performance of the
oxygen electrode, oxygen electrodes were prepared by following the procedure in section 2.4
where the electrodes were coated with 0.08 M nickel nitrate precursor. The electrodes were then
annealed at temperatures between 200 and 800℃ in air with a heating rate of 10 °C min
−1
in each
case, producing nickel-modified steel (NS) oxygen electrodes.
Table 7.1. Mass Gain, Overpotentials, and Tafel Slopes of CS Oxygen Electrodes
Electrode
Annealing
Temperature (°C)
Mass Gain after
Annealing (mg)
Overpotential at
10 mA cm
−2
(mV)
Tafel Slope
(mV decade
−1
)
NS-200 200 5 257 43
NS-400 400 7 324 43
NS-600 600 42.2 563 140
NS-800 800 641.6 553 113

132
Figure 7.1 shows the polarization data of NS electrodes annealed at four different
temperatures. The electrode prepared at the lowest temperature (NS-200) showed the lowest
overpotential (257 mV) at 10 mA cm
−2
. Increasing the annealing temperature resulted in an
increase in the overpotential as summarized in Table 7.1. The electrode annealed at the highest
temperature of 800℃ (NS-800) showed the highest overpotential of 553 mV at 10 mA cm
−2
.

Figure 7.1. Steady-state polarization plots of NS oxygen electrodes.
The Tafel plots (Figure 7.2) for NS-200 and NS-400 were the lowest at 43 mV/decade.
This Tafel slope under high surface coverage of hydroxyl ions is consistent with the Kobussen
Pathway for OER (Table 7.2) with the second step being rate determining.
288
These results were
consistent with previously published work from our group on surface-modified iron electrodes for
oxygen evolution.
284, 285

133
Table 7.2 The Kobussen Pathway of Oxygen Evolution
Rate Determining Step Tafel Slope
Step 1
M+OH
$+
#
+e
#
→MOH+e
#
-
Step 2 MOH+OH
#
→MO+H
'
O+e
#
2RT/3F
Step 3 MO+OH
#
→MO
'
H
#
∞
Step 4 MO
'
H
#
+OH
#
→MO
'
#
+H
'
O+e
#
2RT/F
Step 5 MO
'
#
→M+O
'
+e
#
∞

When increasing the annealing temperature to 600℃, the Tafel slope increases to 140
mV/decade. Similarly, annealing at 800℃ results in an oxygen electrode with a high Tafel slope
of 113 mV/decade. This suggests that the oxygen evolution reaction proceeds through a similar
mechanism for NS-200 and NS-400. The large change in Tafel slope for NS-600 and NS-800
suggests that the reaction proceeds through a different mechanism compared to that of the
electrodes annealed at the lower temperatures.

Figure 7.2. The Tafel plots of NS oxygen electrodes.
To help explain the observed trend in the electrochemical performance of the NS oxygen
electrodes, SEM images (Figure 7.3) of the four electrodes were collected. SEM image of NS-200
and NS-400 showed a high surface area porous morphology with 100 to 200 nm pore size. SEM

134
images of NS-600 and NS-800, however, revealed flake like structures throughout the surfaces
which block the pores and the effective surface area.

Figure 7.3. SEM images of NS oxygen electrodes.
XRD analysis (Figure 7.4) was performed to gain insights into the crystal-phases in the
electrodes. Only α-iron phase was observed in the NS-200 electrode which we attributed to the
steel mesh substrate indicating that the catalyst layer is amorphous. From previous studies in our
group on sintered iron electrodes,
285
this phase was identified as α-nickel hydroxide (100 nm thick)
that gets converted into highly active γ-nickel oxyhydroxide upon anodic polarization. However,
increasing the annealing temperature to 400 °C results in the growth of magnetite, which is not as
active and conductive as γ-nickel oxyhydroxide. In addition, non-conductive iron (III) oxide
phases were observed in NS-600 and NS-800. Therefore, the lack of active phase and the growth

135
of non-conductive phases on the substrate led to a gradual decrease in the OER catalytic activity
for electrodes that were prepared at higher temperatures. The growth of oxide layers on the
substrate during the annealing process may explain the increased mass gain for the NS electrodes
prepared at higher temperatures (Table 7.1).

Figure 7.4. XRD data of NS oxygen electrodes.
7.2.2 Examining Cobalt-treated Electrodes for Oxygen Evolution
We substituted the nickel treatment with cobalt treatment to understand the effect of
catalyst composition on OER. Four cobalt-modified steel (CS) electrodes were prepared using 0.08
M cobalt nitrate precursor as describes in section 2.4. The electrodes were then physically and
electrochemically characterized as previously described. The electrodes’ designations, annealing
temperatures, mass gains after annealing, and electrochemical activities are listed in Table 7.3.
Like NS oxygen electrodes, we observed an increased mass gain after annealing and a higher Tafel

136
slope when going to higher annealing temperatures. This was accompanied by an increase in
overpotential indicating lower electrochemical activity when annealing at higher temperatures.
Table 7.3. Mass Gain, Overpotentials, and Tafel Slopes of NS Oxygen Electrodes
Electrode
Annealing
Temperature (°C)
Mass Gain after
Annealing (mg)
Overpotential at
10 mA cm
−2
(mV)
Tafel Slope
(mV decade
−1
)
CS-200 200 4.3 337 43
CS-400 400 5.9 362 43
CS-600 600 32.1 446 56
CS-800 800 621.6 553 113

SEM images of CS oxygen electrodes (Figure 7.5) showed a similar trend to those of NS
oxygen electrodes (Figure 7.3). CS-200 and CS-400 showed a porous morphology while CS-600
and CS-800 showed a growth of structures blocking the pores. The reduction in available surface
area for OER is therefore one of the main reasons for the reduced activity for the CS electrodes
prepared at higher temperatures.

137

Figure 7.5. SEM images of CS oxygen electrodes.
The XRD data of the CS oxygen electrodes (Figure 7.6) also revealed the growth of non-
conductive and less active iron oxides on the electrode substrates when increasing the annealing
temperature. This further explains the reduction in activity and the larger increase in mass after
annealing when preparing the electrodes at higher temperature.

138

Figure 7.6 XRD data of CS oxygen electrodes.
7.2.3 Comparing Ni-treated and Co-treated Oxygen Electrodes
To help understand the difference between Co-treated and Ni-treated oxygen electrodes,
the best two performing electrodes (NS-200 and CS-200) were selected for further
characterization. Figure 7.7 shows that NS-200 is more active towards OER than CS-200 with an
overpotential of 257 mV at 10 mA cm
−2
compared to 337 mV for CS-200. However, both oxygen
electrodes have a Tafel slope of 43 mV/decade suggesting that the reaction proceeds through the
same mechanism for both electrodes.
The SEM images of these electrodes (Figures 7.3and 7.5) indicated that NS-200 was more
porous than CS-200, indicating a higher surface area which may partially explain the increased
activity for NS-200.

139

Figure 7.7. Polarization curved of NS-200 and CS-200 oxygen electrodes.
However, XPS data of the two electrodes (Figure 7.8) sheds light on the origin of the higher
intrinsic activity for NS-200. A more hydrated phase of nickel hydroxide is present on the surface
of NS-200 as can be seen in Figure 7.8.
289, 290
while cobalt (II) oxide and cobalt hydroxide were
found in 1:1.2 ratio, respectively, on the CS-200 electrode surface.
289
The presence of cobalt (II)
oxide in CS-200 decreases the oxygen evolution activity since a hydrated phase like nickel
hydroxide is more active towards oxygen evolution.

Figure 7.8. XPS data of Ni 2p3/2 for NS-200 and Co 2p3/2 for CS-200 oxygen electrodes.
1.3 1.4 1.5 1.6
0
20
40
60
Current Density (mA cm
-2
)
Potential vs RHE (V)
Ni-treated
Co-treated

140
Cobalt oxide is formed from the thermal decomposition of cobalt hydroxide during the
annealing process at 200 °C which is within the known decomposition temperature for cobalt
hydroxide to the oxide (180-230 °C).
291
. However, the thermal decomposition temperature for
nickel hydroxide is around 300 °C, well below the NS-200 annealing temperature, explaining the
presence of only the active α-nickel hydroxide phase.
292

7.2.3 Electrochemical Pre-treatment of Oxygen Electrodes in Sulfide Solutions
To further increase the current density of the nickel-treated steel mesh electrode, we
attempted to increase the surface area of the substrate before applying the catalytic layer by
electrochemically pre-treating the surface as described in section 2.4.2. We borrowed from the
alkaline battery industry in which the surface area of iron electrodes is enhanced by using sulfur
additives to transform iron into high-surface area iron (II) hydroxide as shown in reaction 7.1.
293,
294
The presence of the sulfides de-passivates the surface, allowing for the conversion to occur in
the electrochemical pre-treatment protocol shown in Figure 7.9.
Fe+ 2OH
*
→ Fe(OH)
+
+ 2e
*
E
,
= −0.877 V (7.1)

Figure 7.9. Electrochemical pre-treatment protocol for the nickel-treated steel electrodes.
0 50 100 150
-1.2
-1.1
-1.0
-0.9
-0.8

Fe(OH)
2
Fe
Discharge (1mA/g)
Potential vs MMO (V)
Time (minutes)
Charge (10mA/g)

141
Steady-state polarization data (Figure 7.10) reveal that pre-treating the oxygen electrode before
applying the catalytic layer results in a lower overpotential compared to the non-treated electrode.
By performing electrochemical pre-treatment, the overpotential at 10 mA cm
−2
decreased from
257 to 235 mV, approximately a 9% reduction in overpotential. Furthermore, the pre-treated
electrode was able to sustain significantly higher current densities compared to the un-treated
electrode.

Figure 7.10. Polarization curves of the NS-200 oxygen electrodes with(purple) and without
(pink) pre-treatment.

To examine the effect of pre-treatment on the surface area of the electrode, double-layer
capacitance was measured as described in section 2.4 and plotted as a function of pre-treatment
cycles in Figure 7.11. After one cycle of pre-treatment, capacitance increased over 16 times that
of the un-treated electrode, suggesting a significant change to the surface area of the electrode.
Double-layer capacitance, however, continuously decreases after the first pre-treatment cycle to
less than 6 times that of the un-treated electrode by the 20
th
cycle.
1.2 1.4 1.6
0
50
100
150
200
Current Density (mA/cm
2
)
Potential (V vs RHE)
Not Pre-treated
Pre-treated

142

Figure 7.11. Electrode capacitance as a function of electrochemical pre-treatment cycles.
To examine the morphology of the steel substrates after electrochemical pre-treatment,
SEM images were collected before and after various pre-treatment cycles as shown in Figure
(7.12). After only one cycle of pre-treatment, the surface is transformed from relatively smooth to
highly porous. However, the increased cycling leads to a reduction in the surface area that is
apparent in Figure 7.12d after 20 cycles.
0 5 10 15 20
0
2
4
6
8

Capacitance (F)
Cycle Number

143

Figure 7.12. SEM images of (a) untreated steel mesh substrate, pre-treated steel mesh substrate
after (b) 1 cycle, (c) 10 cycles, and (d) 20 cycles.

7.3 Oxygen Electrode for All-iron Alkaline Water Electrolyzer
Based on our extensive work to develop a robust and highly active oxygen electrode, we
concluded that a nickel-treated oxygen electrode at 200 ℃ provides the lowest overpotential
towards OER. Electrochemically pre-treating the steel substrate further improved the activity by
increasing the surface area of the electrode. Thus, a steel substrate was electrochemically treated
for one cycle in the presence of sodium sulfide then placed in a nickel nitrate solution and annealed
at 200 ℃ to form an oxygen electrode to be used in an all-iron alkaline water electrolyzer.
7.3.1 Physical Characterization of the Oxygen Electrode
Scanning electron microscopy was used to examine the morphology of the as-prepared
oxygen electrode (Figure 7.13a) and the oxygen electrode after electrochemical testing (Figure

144
7.13b). A rough surface with a relatively high surface area was observed. The “coral-like structure”
was maintained after 100-hour stability tests. This stability is further supported by the absence of
significant change in the double layer capacitance values at 538 mV of 6.06 F and 6.11 F, before
the test and after the test, respectively.

Figure 7.13. SEM image of the (a) as-prepared oxygen electrode and (b) the oxygen electrode
after electrochemical testing.

Energy dispersive X-ray Spectroscopy (EDS, JSM 7001F) was used to determine the
atomic content on the surface of the oxygen electrode. The EDS spectrum of the as-prepared OER
electrode (Figure 7.14) indicated the presence of nickel and oxygen on the surface of the electrode
along with sulfur that most likely was incorporated during the sulfide pre-treatment process. Iron
was also observed in the EDS spectrum due to the electron beam penetration below the thin
catalytic layer.
(a)
(b)
1 μm 1 μm

145

Figure 7.14. EDS spectrum of the as-prepared oxygen electrode.
The XRD pattern of the as-prepared oxygen electrode (Figure 7.15) showed two main
peaks corresponding to Fe(110) and Fe(200) and two small peaks at 2-theta values of 23° and 34°
corresponding to Fe2O3(012) and Fe2O3(104), respectively. The XRD pattern of the OER electrode
after electrochemical tests showed the two main elemental iron peaks indicating that the bulk of
the substrate had remained unchanged after the tests.

146

Figure 7.15. XRD spectra of the oxygen electrode before and after electrochemical testing.
The XPS investigation of the surface of the as-prepared OER electrode at the Ni 2p binding
energy (Figure 7.16) revealed a peak at 855.6 eV corresponding to Ni
2+
consistent with earlier
studies from our laboratory that showed nickel(II) hydroxide as the catalytic layer on the surface
of the oxygen electrode.

Figure 7.16. Binding energy of Ni(2p) for the as-prepared oxygen electrode.

147
7.3.2 Electrochemical Characterization of the Oxygen Electrode
Steady-state electrochemical testing was performed as described in section 2.4. The oxygen
electrode showed high electrocatalytic activity with a low overpotential of 235 mV at 10 mA cm
−2

(Figure 7.17a). The Tafel slope (figure 7.17b) was 46.9 mV/decade over the range of 1-100 mA
cm
−2
. As discussed in section 7.2, this Tafel slope under high surface coverage of hydroxyl ions
was consistent with the Kobussen Pathway for OER with the second step being rate determining.

Figure 7.17. (a) Polarization data of the oxygen electrode and (b) the corresponding Tafel slope.
The Nyquist plots of the oxygen electrode (Figure 7.18) showed two semicircles with the
diameter of the second semicircle (appearing at lower frequency) decreasing with increasing
overpotential. The EIS data fitting (Figure 7.18b) showed a series resistance appearing at high
frequency, which corresponded to the ohmic resistance of the potassium hydroxide solution. The
resistance, R1, derived from the first semicircle at high frequency corresponded to a thin conductive
oxide film on the substrate that is present under ambient conditions. The assignment of R1 to such
a resistive layer is consistent with the observation that this resistance value does not change with
overpotential distinguishing it from a charge-transfer resistance. The diameter of the second
semicircle at low frequency corresponded to R2 the charge-transfer resistance for the faradaic

148
process of oxygen evolution. This assignment is consistent with the decrease of the value of R2
with increasing overpotential.

Figure 7.18. Electrochemical impedance data of the oxygen electrode (a) as a function of
overpotential and (b) the corresponding EIS data fitting.

To confirm that R2 corresponded to the charge-transfer resistance of the oxygen electrode,
we used the relationship between charge-transfer resistance and overpotential (Equations 7.2 and
7.3). RCT corresponds to the charge-transfer resistance while ηCT and I represent the charge-transfer
overpotential and the current density, respectively. B and io are the Tafel slope and exchange current
density, respectively.
$
-.
=
1C
-.
14
(7.2)
.DE (
1
$
-.
) = .DE(
4
&
F
)+
C
-.
F
(7.3)
Thus, according to Equation 7.3, a plot of log(1/RCT) vs overpotential would be a straight
line and the inverse of the slope would correspond to the Tafel Slope. Plotting log(1/R2) vs
overpotential (Figure 7.19) yielded a Tafel slope of 52.9 mV decade
−1
, which is close to the value
of Tafel slope value of 46.9 mV decade
−1
obtained from the potentiostatic polarization

149
measurements. Thus, R2 was attributed to the charge-transfer resistance of the oxygen electrode.

Figure 7.19. Log(1/RCT) vs overpotential for the oxygen electrode.
Finally, the electrochemical stability of the oxygen electrode was examined over 100
continuous hours of polarization at 10 mA cm
−2
. The electrode potential showed a marginal
increase of overpotential corresponding to 1.2 μV hour
−1
(Figure 7.20) indicating excellent
stability characteristics for the oxygen electrode.

Figure 7.20. Chronopotentiometry of the oxygen electrode at 10 mA cm
−2
.
200 220 240 260 280
-0.5
0.0
0.5
1.0
log(1/Rct)
Overpotential (mV)
y=0.0189x-4.178

150
7.4 Steel-Based Electrodes for Hydrogen Evolution
To produce an inexpensive and robust hydrogen electrode, we attempted to use the same
inexpensive steel mesh substrate that we have used for the oxygen electrode. Since iron is stable
under cathodic conditions, simply modifying the surface sputtering a catalytic layer is a simple
and scalable process to produce a highly active hydrogen electrode. Therefore, nickel and
molybdenum were co-sputtered on a low-carbon steel mesh as described in section 2.4 to produce
a robust hydrogen electrode with low overpotential.
7.4.1 Physical Characterization
The SEM images of the as-prepared HER catalyst (Figure 7.21) showed that co-sputtering of
nickel and molybdenum for 40 minutes fully covered the surface of the low-carbon steel mesh
and protected it from being exposed to the alkaline electrolyte. The cross-sectional image of the
HER electrode (Figure 7.21b) showed the typical columnar growth observed in DC sputtering
and an approximate thickness of around 400 nm. The cross-sectional image of the hydrogen
electrode after electrochemical tests (Figure 7.21c) showed that there was no reduction in the
thickness of the sputtered catalytic layer. Furthermore, there was no observed change in the
surface morphology or composition after the electrochemical tests (Figure 7.21d).

151

Figure 7.21. SEM images of the hydrogen electrode. (a) and (b) Cross-section and top-view of
the as-prepared Ni-Mo co-sputtered surface. (c) and (d) Cross-section and top-view of the Ni-Mo
co-sputtered surface after electrochemical testing.

To confirm the thickness of the hydrogen electrode, X-ray reflectivity measurement was
performed as described in section 2.5. The measurement confirmed a coating thickness of 400 nm
and a roughness of 6.6 nm as shown in Figure 7.22.

152

Figure 7.22. Reflectivity profile of NiMo co-sputtered layer on glass slide.
EDS measurements showed that nickel and molybdenum are homogenously distributed
with an atomic ratio of nickel to molybdenum of 9:1 (Figure 7.23). The observation of iron in the
spectrum is due to the penetration of the electron beam in the EDS measurement past the coating
into the carbon steel mesh.

153

Figure 7.23. EDS spectrum of as-prepared hydrogen electrode.
The XRD pattern of the as prepared hydrogen electrode (Figure 7.24) showed 2 main peaks
corresponding to Fe(110) and Fe(200) indicating that the bulk of the electrode was elemental iron.
The XRD pattern of the HER electrode after electrochemical testing indicated that there was no
discernible change to the electrode substrate as a result of the tests.

154

Figure 7.24. XRD spectra of the hydrogen electrode before and after electrochemical testing.
X-ray photoelectron spectroscopy was used to characterize the electrode surfaces, as-
prepared. For the HER electrode, the Ni 2p binding energy values at 852.7 eV and 856.0 eV
(Figures 7.25a) indicated the presence of Ni
0
and Ni
2+
, respectively. The Mo 3d binding energy
peaks at 228.1 eV and 232.3 eV (Figure 7.25b) were attributed to Mo
0
and Mo
6+
, respectively.
Using the areas under the curves, it was estimated that the molar ratio of Ni
2+
to Ni
0
was 3.4 to 1
while the molar ratio of Mo
6+
to Mo
0
was 6.1:1. Although elemental Ni and Mo were sputtered on
the steel electrode, Ni
2+
and Mo
6+
were spontaneously formed due to the exposure of the electrode
to ambient atmosphere.

155

Figure 7.25. Binding energy of (a) Ni (2p) and (b) Mo (3d) for the as-prepared hydrogen
electrode.

7.4.2 Electrochemical Characterization
The steady-state polarization data of the hydrogen electrode (Figure 7.26a) showed that the
catalytic layer was highly active towards HER with an overpotential of 166.8 mV at 10 mA cm
−2
,
while the Tafel slope (Figure 7.26b) was calculated to be 115.4 mV decade
−1
. The performance of
this electrode is comparable to hydrogen electrodes that are nickel based.
270, 273
It is well known in
the literature that the first step in the hydrogen evolution reaction is the surface dissociation of
water in the Volmer step (Table 7.4) resulting in adsorbed hydrogen atoms on the catalyst surface.
The adsorbed hydrogen then combines to form hydrogen gas by the Heyrovsky pathway or the
Tafel pathway.
295
The Tafel slope values can be used to determine the rate limiting step. The Tafel
slope of 115.4 mV decade
−1
indicated that the Volmer step was the rate determining step. This
observation is consistent with reports in the literature on Ni-Mo-based catalysts.
274, 296

156

Figure 7.26. (a) Polarization data of the HER electrode and (b) the corresponding Tafel slope.
Table 7.4 HER Mechanism in Alkaline Conditions
Step Reaction Tafel Slope (mV decade
−1
)
Volmer
Surface+H
'
O+e
#
⇌H
$%&
+OH
#
120
Heyrovsky
H
$%&
+H
'
O+e
#
⇌H
'())
+OH
#
40
Tafel
2H
$%&
⇌H
'())
30

The electrochemical impedance measurements of the hydrogen electrode yielded a Nyquist
plot with two semicircles (figure 7.27a). The first semicircle appearing at high frequency remained
relatively unchanged with changes in electrode potential while the diameter of the second
semicircle appearing at lower frequency decreased with increasing overpotential. Just as in the
case of the oxygen electrode, the impedance data was fitted to an equivalent circuit model (Figure
7.27b) with Rs corresponding to the resistance of the potassium hydroxide solution, R1
corresponding to the resistance due to the compact oxide layer formed while fabricating the
electrode, and R2 corresponding to the charge-transfer resistance of the HER electrode.

157

Figure 7.27. EIS Nyquist plots of the hydrogen electrode as a function of overpotential and (b)
the EIS data fitting.

To confirm that R2 corresponded to the charge-transfer resistance in the HER electrode,
log(1/R2) was plotted vs overpotential (Figure 7.28) yielding a Tafel slope of 116.3mV decade
−1

which is close to the value of 115.4 mV decade
−1
obtained from potentiostatic polarization
experiments. This type of cross verification established that R2 was indeed the charge-transfer
resistance of the hydrogen electrode.

158

Figure 7.28. Log(1/RCT) vs overpotential for the hydrogen electrode.
The stability test performed on the hydrogen electrode (Figure 7.29) showed that the
electrode was stable under continuous polarization at 10 mA cm
−2
for at least 100 hours whereupon
the overpotential increased at 3.4 μV hour
−1
during 100 hours of continuous operation. This
miniscule increase of overpotential is a preliminary indication of good stability for extended
operation.
100 150 200 250
-0.2
0.0
0.2
0.4
0.6
0.8
1.0
log(1/Rct)
Overpotential (mV)
y=0.0086x-1.084

159

Figure 7.29. Chronopotentiometry of the hydrogen electrode at 10 mA cm
−2
.
7.5 Fabrication and Characterization of All Iron Electrolyzer
7.5.1 Electrolyzer Assembly
An all-iron alkaline electrolyzer was assembled using the hydrogen and oxygen electrodes
fabricated with the modified steel substrates, as previously described. The electrodes with a
geometric are of 25 cm
2
were assembled with a commercially available Zirfon separator (Agfa),
and 30% potassium hydroxide as the electrolyte. A schematic of the electrolysis device is shown
in Figure 7.30.

160

Figure 7.30. Schematic of an all-iron alkaline water electrolyzer: (1) Nickel-molybdenum coated
iron electrode, (2) Zirfon PERL separator, (3) graphite plate with columnar flow field, (4) heating
pad, (5) nickel-coated plate with columnar flow field, (6) gold-coated current collector, and (7)
K-type thermocouple.

7.5.2 Electrolyzer Electrochemical Characterization
The electrochemical performance of the electrolyzer was measured using steady-state
galvanostatic polarization up to 1 A cm
−2
and up to 70 °C. The electrolyzer, at 100 mA cm
−2
,
operated at voltages of 1.83 V and 1.71 V, at room temperature and 70 °C, respectively (Figure
7.31a).

161

Figure 7.31. (a) Galvanostatic data of the electrolyzer at different temperatures and (b) the
corresponding slopes.

The Tafel slopes obtained from the IR-corrected voltage-current data (Figure 7.31a) are
not constant, showing increasing values of slope with increasing overpotential (Figure 7.32). We
ascribe the increasing slopes at high current densities in Figure 7.31b to the excessive generation
of gas that leads to shielding of the electrode surface with bubbles. This issue is central to
electrolyzers operating at high current densities and often special expedients are devised in
industrial electrolyzers to remove the bubbles from blocking the electroactive surface.
Furthermore, the Tafel slope does not vary as a function of temperature indicating that the reaction
mechanism of water splitting is not changing with temperature.
1.8 2.0 2.2 2.4
200
400
600
800
1000
(b) (a)

Current Density (mA/cm
2
)
IR-Corrected Cell Voltage (V)
30 °C
45 °C
60 °C
70 °C
2.0 2.5 3.0
600
800
1000

Overpotential (mV)
log j (mA/cm
2
)

162

Figure 7.32. Tafel slopes of the electrolyzer at 30 °C.
Electrochemical impedance measurements were collected between 15 and 35 mA cm
−2
(1.62 V to 1.69 V). The Nyquist plots at all cell voltage values showed two semicircles with their
diameters decreasing with increasing current density suggesting that these semicircles
corresponded to charge-transfer processes (Figure 7.33a). The two semicircles in the Nyquist plot
at 35 mA cm
−2
were fitted with two charge-transfer resistances, R1 and R2 (Figure 7.33b). Using
equations 7.2 and 7.3 to determine the resulting Tafel slopes from each charge-transfer resistance,
it was concluded that R1 corresponded to the hydrogen evolution reaction while R2 corresponded
to the oxygen evolution reaction.
2.0 2.5 3.0
600
700
800
900
1000
1100
827.4 mV/decade
577.0 mV/decade
Overpotential (mV)
log j (mA/cm
2
)
30
o
C
306.7 mV/decade

163

Figure 7.33. (a) Impedance spectroscopy data of the electrolyzer as a function of overpotential
and (b) EIS data fitting at 35 mA cm
−2
.

To test for durability, the cell was operated at 1 A cm
−2
for 100 hours at room temperature.
The cell voltage of 2.13 V was maintained for the entire period with no measurable change in this
value (Figure 7.34).

Figure 7.34. Constant current operation of the electrolyzer at 1 A cm
−2
at room temperature.

164
The all-iron electrolyzer demonstrated here was found to have performance comparable to
industrial electrolyzers, based on the literature data
297-300
in Table 7.5. Most importantly, we could
achieve this performance with steel mesh-based electrodes that could potentially reduce the cost
of the electrolyzer significantly as 50% of the cost of electrolyzer stacks is attributed to the nickel-
based electrodes.
66, 301, 302

Table 7.5. Comparison with the Performance of Various Industrial Electrolyzers
288, 298-300

Company Temperature (°C) Pressure Cell V oltage at 300 mA cm
−2

De Nora 80 Ambient 1.85 V
Electrolyzer Corporation 70 Ambient 1.9 V
Teledyne Energy Systems 82 2 bar 1.9 V
General Electric 80 4 bar 1.7 V
This work 70 ambient 1.84 V

165
CHAPTER 8.

ELECTROCHEMICAL STUDIES OF THE CYCLOADDITION ACTIVITY OF
BISMUTH(III) ACETYLIDES TOWARDS ORGANIC AZIDES UNDER
COPPER(I)-CATALYZED CONDITIONS
This dissertation chapter is based on a recent publication: Nazarova, A. L.; Zayat, B.; Fokin, V.
V.; Narayan, S. R., Electrochemical Studies of the Cycloaddition Activity of Bismuth(III)
Acetylides Towards Organic Azides Under Copper(I)-Catalyzed Conditions. Frontiers in
Chemistry 2022, 10.

8.1 Introduction to OrganoBismuth(III) Compounds in Drug Design
With an increased interest in the area of novel, non-toxic, and biocompatible nanomaterials,
bismuth (Bi)-doped systems have become important in the area of near-infrared (NIR)-emitters
and drug-delivery materials.
303-306
Bismuth has a negligibly low level of toxicity and
carcinogenicity as compared to its highly-abundant neighbors in the periodic table (tin, lead,
antimony, arsenic).
307-309
However, biochemical and industrial applications involving bismuth-
containing organic compounds remain limited.
310, 311
Presently, commercially-available
bismuth(III) compounds are typically inorganic salts used for instance as Lewis acid catalysts in
organometallic reactions or those used for the preparation of nanoparticles.
312-315
Recently the
coordination with electron donors such as sulfur and nitrogen was reported to increase the thermal,
air, and hydrolytic stability of bismuth in the oxidation states of (III) and (V), creating new
opportunities for its applications.
316-318
Also, recently bismuth(III) and (V) have been reported to
show redox activity when participating in bimolecular interactions in solution.
319
The investigation

166
of the redox characteristics as well as the role of transition metals involved in catalytic processes
through reaction kinetic profiling remains well established in organometallic chemistry.
320-322

To exploit bismuth’s coordination chemistry by modifying its electronic effects and
geometry, we have focused our study on sulfonyl-merged bis-anionic aryl-tethered scaffolds.
315,
323, 324
By altering the electron-withdrawing nature of the sulfone moiety and the ability of the
oxygen atom to intramolecularly coordinate to the bismuth center, a variety of sulfone-type ligands
that influence the bismuth(III) acetylides activity under copper(I)-catalyzed conditions were
studied (Figure 8.1).

Figure 8.1. Schematic of the general molecular structure and reaction.
We performed a mechanistic study of the reactivity of para-phenyl substituted bismuth(III)
acetylides in the copper(I)-catalyzed cycloaddition reactions with organic azides.
325
Cyclic
voltammetry and nuclear magnetic resonance were used to derive kinetic parameters and the rate
law of the transformations of the various ligands, although this work will focus primarily on the
CV studies.

167
8.2 Synthesis and Characterization of Diphenyl Sulfone Bismuth(III) Acetylides
The rational design of the bismuth(III) acetylides was a crucial factor for gaining insights
into the catalytic cycle. Two types of derivatization strategies were applied in the synthesis of the
organobismuth acetylides: (1) diphenyl sulfone ligand derivatization with different functional
groups; and (2) para-position phenylacetylene substitution with electron-deficient, neutral, and
electron-rich functional groups.
315

Figure 8.2. (a) Synthetic route for the synthesis of diphenyl sulfone bismuth(III)
acetylides. (b) Diphenyl sulfone bismuth(III) acetylides. (c) 1-ethynyl-4-methylbenzene
bismuth(III) acetylides with functionalized diphenyl sulfone ligands.

168
To study factors influencing the electronic environment of the bismuth(III) center,
acetylides A[1] through A[6] were prepared using a synthetic protocol adapted from a previously
reported procedure for acetylide A[3] containing a diphenyl sulfone scaffold as the tethered ligand
(Figure 8.2b).
326
Acetylides A[7] through A[10] were synthesized by coupling thiols with aryl
iodides and subsequent oxidation of the diphenyl sulfides (Figure 8.2c).
327, 328
Thus, copper-
catalyzed C(aryl)-S bond formation between aryl iodides and aryl sulfides, followed by the
oxidation of the sulfide with m-CPBA, metalation with n-butyllithium, and subsequent treatment
with in situ generated Bi(Br)2Ph yielded triphenyl bismuth derivatives (Figure 8.2a). Iodination of
the triphenyl bismuth species followed by C(sp)-Bi coupling provided the bismuth(III) acetylides
A[7-10]. While these substrates have not been used for the in situ kinetic studies, they were used
for comparisons of the solid-state structures.
8.3 Kinetic Cyclic Voltammetry for Mechanism Elucidation
The electronic functionality of triarylbismuth(V) ligands was reported to directly influence
their reactivity.
329-332
With this in mind, we analyzed the reactivity trends of differently
functionalized bismuth(III) acetylides towards organic azides in the copper(I) catalyzed
cycloaddition reaction. Being eager to study substrate-dependent reactivity features in the
mechanism, we made an effort for stepwise kinetic studies using cyclic voltammetry (CV).
333-336

While commonly used as an electroanalytical method, cyclic voltammetry allowed the direct
investigation of the intermolecular transannular effect on bismuth(III) acetylide reactivity towards
triazolide formation as well as the redox activity of the catalytic intermediates. Using in situ
electrochemistry as an electroanalytical tool provided vital information on the formation of
intermediates, which is critical for understanding the catalytic transformation, as well as the ability

169
to dissect the reaction rate constants and parameters from in situ experimental data. The proposed
mechanistic model of the cycloaddition reaction is shown in Figure 8.3.

Figure 8.3. Proposed mechanistic model of azide-bismuth(III) acetylide copper(I)-
catalyzed cycloaddition reaction.

A substantial drawback of the copper(I) system is its air sensitivity; therefore, all
electrochemical experiments were performed in the nitrogen atmosphere of a glovebox. Following
our developed CV kinetic analysis protocol, we were able to stepwise monitor the redox activity
of copper(I) after adding first the acetylide and then the organic azide reactant to access
quantitative kinetic parameters of the catalytic cycle (Figure 8.4). Each electrochemical kinetic
experiment was initiated by redox studies of a DMSO solution containing only the copper(I)

170
trifluoromethanesulfonate toluene complex catalyst. The mechanism of the copper (I) catalyzed
azide bismuth(III) acetylide cycloaddition has been postulated to initiate with a π-intermediate
complex formation of copper(I) and the bismuth(III) acetylide.
325
In a subsequent step, the azide
ligation/ migratory insertion occurs and reductive elimination of copper(I) yielding the 5-
bismuth(III)-triazolide.
325, 337
The kinetics of the two independent steps: (1) π-intermediate
complex of copper(I) and the bismuth(III) acetylide [A[X]·cat] formation, (2) product
formation/regeneration of [Cu
I
], were monitored by the subsequent addition of the reactants. First,
a bismuth acetylide was added and the kinetics of the π-complex formation was monitored by
continuously recording cyclic voltammograms until no further changes in the CV responses could
be detected. Subsequently, the organic azide was introduced to the mixture and the reaction
progress was again monitored by continuously recording cyclic voltammograms.

Figure 8.4. Reaction scheme of kinetic studies of copper(I)-catalyzed reaction of 5-bismuth(III)
triazolides formation with cyclic voltammetry technique.

All bismuth(III) acetylides used for the kinetic studies carried the same sterically demanding
diphenyl sulfone ligand. Therefore, no significant differences in the steric environment of the
copper species in the copper(I)-bismuth(III)acetylide complexes are expected. As for other
(metal)acetylides, coordination of the azide to copper(I) in the copper(I)-bismuth(III)acetylide

171
complex is expected to primarily depend on the electronic nature of the copper(I) species.
338-342

Therefore, the formation of the first σ-bond between the β-carbon (C(2)) of the acetylide and the
terminal nitrogen (N(3)) of the azide should determine the rate-limiting step of the overall
transformation.
325

8.3.1 Bismuth(III)-Acetylide – Copper(I) Complex Formation
To develop a quantitative model of such a complex catalytic reaction, rate constants, and
substituent effects need to be determined. To verify that the CV response is chemically reversible
and thus the catalyst and the solvent are suitable for the CV timescale, each kinetic electrochemical
experiment was first initiated with cyclic voltammetry studies of a solution of the copper(I)
trifluoromethanesulfonate toluene complex catalyst. The cyclic voltammograms of copper(I)
complex (1.25 mM, 8.12 μmol) dissolved in dry DMSO (6.5 mL) recorded at 100 mV s
-1
revealed
that the reduction of Cu(I) to Cu(0) occurred at around −0.2 V versus the Cu
+
/Cu
0
reference
electrode (Figure 8.5).

172

Figure 8.5. Cyclic voltammogram of the copper(I) triflate catalyst at 25°C in dry DMSO
recorded at 100 mV sec
−1
.

To collect the first set of quantitative kinetic data, we injected an almost five-fold excess
of bismuth(III) acetylide, A[X], into a 1.25mM catalyst solution. The final concentrations of the
bismuth(III) acetylide substrate and the catalyst solution were 5.35 mM and 1.16 mM,
respectively. The formation of the [A[X]·cat] complex was monitored by continuously recording
cyclic voltammograms until no further changes in the CV responses were detected, as shown in
Figure 8.6.

173

Figure 8.6. Kinetic profiles of the cyclic voltammograms obtained after the addition of acetylide.
Similar cyclic voltammetry studies of Cu(I) complexes with various triazolylamine ligands
for the copper(I) catalyzed azide-alkyne cycloaddition reaction have been reported while studying

174
reversible couple characteristics of Cu
I/II
.
343
Upon binding to the π-system of the bismuth(III)
acetylide, the electrode potentials during the reduction of Cu(I) were found to slightly shift towards
the higher potentials for the bismuth(III) substrates. On the contrary, the electrode potentials during
the oxidation of Cu(0) to Cu(I) remained identical and stable throughout the reaction. The final
CV response after the formation of the [A[X]·cat] complex results in a lowering of the Cu(I)
concentration and a consequent decrease in the observed current for the reduction of Cu(I) as
shown in Figure 8.7. The copper(I)-bismuth(III) acetylide complex itself is not easily reduced.

Figure 8.7. Cyclic voltammogram of the copper(I) triflate catalyst at 25°C in dry DMSO before
(blue) and after (orange) the addition of acetylide.

8.3.2 Bismuth(III) Triazolide[X] Formation
When the redox cycle indicated the complete formation of the [A[X]·cat] complex, the
azide reactant was injected and the processes on Cu(I) were monitored by continuously recording
cyclic voltammograms (Figure 8.8). Immediately after injection, the characteristic yellow color of
the [A[X]·cat] complex began to fade and the recorded current increased due to a sudden release

175
of free copper-catalyst. The electrochemical monitoring was continued until no further changes in
current were recorded and the current stabilized at a certain value.

Figure 8.8. Kinetic profiles of the cyclic voltammograms obtained after the addition of azide.

176
The cyclic voltammograms of copper(I), after the formation of the copper-bismuth(III)
acetylide complex, and after the azide addition/product formation for [A]1 are shown in Figure
8.9.

Figure 8.9. CV of A[1] before and after the addition of azide.
8.4 Kinetic Model Parameter Estimation from Experimental Data
Cyclic voltammetry is a unique tool to extract reagent kinetic parameters, i.e. rate constants
by detecting changes in peak currents or potentials versus time.
344-348
In situ cyclic voltammetry
data was used to determine changes in the free copper(I) concentration, which allowed for the
calculation of the kinetic parameters of the bismuth(III) acetylide coordination reactions during
1,2,3-triazolide P[X] formation. Since the reaction rate is a function of multiple parameters, a
direct (assumption-free) mechanistic analysis would allow accurate and reliable accounting of all
factors that influence the evolution of the rate-determining changes throughout the reaction.
349, 350

By direct analysis, one implies subsuming concentration variations of all involved catalytic species
and reactants. Thus, kinetic profiles of concentration versus time for the stepwise experiments can

177
be formulated in the form of ordinary differential rate equations (ODEs). For the direct and reverse
rate constant elucidation and their impact on the equilibrium processes, ODEs were formulated to
solve the mechanistic model of the copper(I)-catalyzed formation of 5-bismuth(III) triazolides
(Figure 8.10).

Figure 8.10. Reaction scheme of the first step, the π-intermediate complex formation of the
copper(I) catalyst and bismuth(III) acetylide.

Based on the Randles-Sevcik equation
351
, the peak current in a cyclic voltammogram is
directly proportional to a concentration of the species that participate in the electron transfer,
equation (8.1):
4
/
= (2.69·10)
0
√K
1
L√M√NO (8.1)
where 4
/
is the peak current (in our case of reduction current), n is the number of participating
electrons, A is the electrode area, D is a diffusion coefficient, N is the scan rate, C is the unknown
concentration of the [cat] species at a time corresponding to a certain peak current, 4
/
. Figures
8.11-8.16 show the current change over time after adding the acetylide (step 1) and after the
addition of the azide (step 2) derived from Figures 8.6 and 8.8, respectively.

178

Figure 8.11. Current change after the addition of (a) bismuth(III) acetylide[1] and (b) the azide.

Figure 8.12. Current change after the addition of (a) bismuth(III) acetylide[2] and (b) the azide.

Figure 8.13. Current change after the addition of (a) bismuth(III) acetylide[3] and (b) the azide.

179

Figure 8.14. Current change after the addition of (a) bismuth(III) acetylide[4] and (b) the azide.

Figure 8.15. Current change after the addition of (a) bismuth(III) acetylide[5] and (b) the azide.

Figure 8.16. Current change after the addition of (a) bismuth(III) acetylide[6] and (b) the azide.
For simplicity, we denote the copper triflate benzene complex catalyst [Cu
I
] hereafter as
[cat]. Considering the equilibrium process of the intermediate [A·cat] formation from [cat] and the
bismuth(III) acetylide in the theoretical model (Equation 8.2):

180
−
1[Q5R]
1R
= +T
2
[L] [Q5R] −T
*2
[L∙Q5R] (8.2)
Applying the conservation law for the catalyst and bismuth(III) acetylide components yields
equations 8.3 and 8.4:
[Q5R]
3
=[Q5R] + [L·Q5R] (8.3)
[L]
3
=[L] + [L·Q5R] (8.4)
where [Q5R]
3
and [L]
3
are the initial concentrations of the bismuth(III) acetylide and copper(I)
catalyst.
Considering the law of conservation of mass, equations (8.3) and (8.4), the concentration of the
catalyst can be expressed as equation 8.5:
352-354

1[Q5R]
1R
= −T
2
([L]
3
−[Q5R]
3
+ [Q5R] )[cat] + T
*2
([Q5R]
3
−[Q5R] ) (8.5)
The integration was done using a self-developed computer program relying on the fourth-
order classical Runge-Kutta method.
355, 356
The solution of the equation 8.5 was found via the grid
search using metrics of the LMS optimization:
min
4
!
,4
"!
]^[Q5R]
"6/,%
−[Q5R]
%
_
+
'
%73
(8.6)
where the concentration of the catalyst, [cat]exp, is derived from experimental kinetic data, and
values of [cat] are obtained from equation 8.5. n is the number of the experimentally derived
concentration values at n time intervals.
From the first set of kinetic data in Figures 8.11 to 8.16 for the bismuth(III) acetylide
addition, we derived a linear proportionality coefficient α for each kinetic experiment of the six
different acetylides while keeping the value of the initial copper(I) catalyst concentration constant
as given by equation (8.7):

181
4
/
= `[Q5R]
3
(8.7)

where 4
/
was denoted as the peak of the anode current and [Q5R]
3
is the initial concentration of the
catalyst.
The concentration of [cat] over time was obtained from 4
/
versus time data from the kinetic
CV experiments using equation 8.7. The complete details on the individual [cat] over time studies
for all six bismuth(III) acetylides are shown in Figure 8.17.

182

Figure 8.17. Concentration changes of the Cu
+
/Co
0
redox pair for step 1.
The equilibrium constant a
2
was determined from equation 8.5 while having results on
separate rate constants of the direct and reverse reaction of the [A·cat] formation. As the rate of

183
the [cat] coordinating to the bismuth(III)acetylide equals zero at equilibrium, we can solve the
quadratic equation (equation 8.5) for the equilibrium concentration [cat]eq.
The forward reaction rate parameter T
2
for the [A·cat] complex formation of the differently
substituted bismuth(III) acetylides follow a non-linear trend in activity when compared to the
corresponding equilibrium concentration [cat]eq. values. The reaction equilibrium parameter a
2
for
the [A·cat] complex formation was found to be inversely correlated in terms of values when
compared to the corresponding equilibrium concentrations [cat]eq. (Table 8.1).
Table 8.1. Rate Parameters Derived from the Cyclic Voltammogram Kinetic Studies
Entry

σpara
kA,
M
−1
sec
−1

σ(kA),
M
−1
sec
−1

kobs·10
−5
,
M·sec
−1

σ(kobs)·10
−7
,
M·sec
−1

KA·1
0
3
,
M
−1

σ(KA)·10
3
,
M
−1

[cat]eq·10
−5
, M

σ([cat]eq)·10
−5
, M
[1], R = OMe −0.27 2.39 0.0598 1.19 0.70 0.85 0.0215 24.32 0.033
[2], R =
t
Bu −0.20 1.95 0.0415 0.81 0.63 1.40 0.0344 16.32 0.020
[3], R = Me −0.17 2.80 0.0754 2.31 1.53 0.70 0.0189 27.97 0.016
[4], R = H 0.00 3.44 0.0749 2.24 1.47 0.92 0.0240 22.88 0.015
[5], R = Br 0.39 4.97 0.0923 0.34 0.17 1.36 0.0337 16.75 0.016
[6], R = CF3 0.43 4.36 0.0965 1.43 0.70 0.71 0.0190 27.83 0.013

Indeed, the more the equilibrium is shifted in the forward direction, the more [A·cat]
species are formed, thus the lower would be the equilibrium concentration of the free catalyst [cat].
This confirms our assumption that the activity of bismuth(III) acetylides of the copper(I)-catalyzed
cycloaddition with organic azides is not solely directed by para-phenyl substituents as has been
reported for proto and iodoalkynes.
357

After determining the equilibrium constant, a
2
, of the π-coordination of the bismuth(III)
acetylides to the copper(I) catalyst, we applied the same approach to the estimation of the azide
ligation/migratory insertion rate parameters. The addition of the azide to a solution of the [A·cat]
π-complex (stabilized CV response) resulted in an immediate increase in the reduction current.
Before the addition of the azide, the mixture contained dominantly the [A·cat] species. Upon

184
addition, the product, 5-bismuth(III) triazolide, was formed and a large amount of free copper
catalyst was released at once causing the sudden increase in peak current (Figure 8.8). Thereafter,
the peak current returned to the current levels observed for the [A·cat] complex formation. This
makes the copper(I) –bismuth(III) acetylide complex formation the faster process. Subsequently,
we can conclude that the azide ligation/migratory insertion is the rate-determining step (RDS), as
a continuous accumulation of the [A·cat] complex was observed. While the coordination of the ɑ-
nitrogen (N1) to copper(I) is influenced primarily by the electronic environment of the copper(I)
species, the formation of the first covalent bond between acetylene β-carbon and terminal nitrogen
of the azide moiety is also expected to influence the kinetics of the reaction as well (Figure 8.3).
While previous studies of (metal)acetylides in the copper(I)-catalyzed cycloaddition with azides
followed a linear (Hammett) reactivity trend,
358, 359
the observation of a non-Hammett-dependent
reactivity of the bismuth(III) acetylides indicates that, apart from para-phenyl substituents, other
effects such as transannular electron density donation of oxygen to the bismuth center play a role
in the control of the reaction kinetics. In the solid-state, the transannular Bi(1)-O(1) distance for
the least reactive acetylide, para-bromo substituent A[5], is one of the largest, whereas the more
reactive substrates, A[3], A[4], and A[6], have shorter Bi(1)-O(1) bond distances. We believe that
such perturbations not only affect the electron density of the C(β)-acetylene carbon but
subsequently also influence the electronic environment of the copper(I) species in the [A·cat]
complex. This causes different affinities of copper(I) species towards N→Cu(I) ligation (Figure
8.3). This supports our previous conclusion of the overall reaction to be azide-dependent.
Finally, increasing the potential window past the Cu(I)/Cu(0) redox pair allowed the
observation of electrochemically relevant processes at the bismuth center in the bismuth(III)-
acetylides as shown in Figure 8.18. The cyclic voltammograms were collected at 25°C in dry DMF

185
at 100 mV sec
−1
using the same electrochemical setup as before. As expected, the formation of the
copper(I) –bismuth(III) acetylide complex also impacts the electronic environment of the
bismuth(III) center

Figure 8.18. CV of [Cu] after the addition of (a) bismuth(III) acetylide and (b) azide. (c) A
comparison of the [Cu] CV before (blue) and after (orange) the addition of the bismuth(III)
acetylide complex and (purple) the addition of azide.

Our detailed mechanistic modeling performed in the context of CV independent reactivity
experiments revealed the electronic effects of para-phenyl functionalization together with
transannular interaction determine substrates rate parameters and reactivity. Since the reaction rate
is influenced by multiple parameters, a direct (assumption-free) mechanistic analysis allowed the
accurate and reliable examination of all factors that influence the complex catalytic formation of
5-bismuth(III)-triazolides.

186
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Abstract (if available)

Abstract Climate change is one of the most difficult challenges we have faced as humans. Addressing this issue requires a series of structural changes to our economy and way of life. Electrochemistry is one of the tools that will help us achieve zero emissions due to its ability to replace many greenhouse emitting processes with sustainable alternatives. For example, electric motors powered by lithium batteries are replacing many internal combustion engines that rely on fossil fuels. Hydrogen is increasingly being adopted as a clean fuel replacement as well.
One obstacle to lithium-ion battery (LIB) adoption is long charge times and low cycle life, an issue that many try to address by using conductive polymers instead of insulating ones as binders to improve electron transport. We have developed an in-situ technique that allows for the electrochemical characterization and electronic and ionic conductivity measurement of these polymers as a function of potential (Chapter 3). In Chapter 4, we demonstrate the use of bifunctional conducting polymers as binders for LIBs to greatly improve rate capability and cycle life. In Chapter 5, we utilize n-dopable conducting polymers as binders for the sulfur electrodes in lithium-sulfur batteries to limit the polysulfide shuttling and significantly enhance cycle life.
In Chapter 6, we probe a new class of transition metal perovskite chalcogenides as promising electrocatalysts for the splitting of water into hydrogen and oxygen. We demonstrate the advantages of using perovskite sulfides compared to the more commonly used perovskiteoxides. In Chapter 7, we develop robust and inexpensive iron-based catalysts for the hydrogen and oxygen evolution reactions. We also demonstrate an efficient all-iron alkaline electrolyzer that operates at 20 mA cm-2. Finally, we use electrochemistry to perform kinetic studies on bismuth complexes for drug delivery applications in Chapter 8.

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Creator Zayat, Billal (author)

Core Title Advancing lithium batteries and related electrochemical technologies for a sustainable future

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Degree Conferral Date 2022-12

Publication Date 08/15/2024

Defense Date 08/01/2022

Publisher University of Southern California (original), University of Southern California. Libraries (digital)

Tag conducting polymers,conductive binders,drug delivery,electrochemical impedance spectroscopy,electrochemistry,lithium-ion batteries,lithium-sulfur batteries,OAI-PMH Harvest,perovskite chalcogenides,sustainability,water electrolysis

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Language English

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Advisor Narayan, Sri (committee chair), Prakash, Surya (committee member), Sharada, Shaama (committee member)

Creator Email bmzayat@gmail.com,zayat@usc.edu

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