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Electrocatalytic thiolate- and selenolate-based coordination polymers for solar energy conversion
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Electrocatalytic thiolate- and selenolate-based coordination polymers for solar energy conversion
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Content
ELECTROCATALYTIC THIOLATE- AND SELENOLATE-BASED
COORDINATION POLYMERS FOR SOLAR ENERGY
CONVERSION
by
Courtney Ann Downes
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)
May 2018
ii
For Mom, Dad, and Kevin
iii
ACKNOWLEDGEMENTS
I will begin by thanking my advisor, Professor Smaranda C. Marinescu, for her support, guidance,
and mentorship over the last five years. She has supported my growth as a scientist and always
offered me the opportunity and flexibility to explore my interests both in and outside of the lab.
Without her support, I certainly would not be where I am today. I also want to thank the other
members of my committee, Professors Richard L. Brutchey, Mark E. Thompson, Sri R. Narayan,
and Kelly T. Sanders for providing important scientific insights and career advice throughout my
time at USC.
I want to acknowledge Alon Chapovetsky, Andrew Clough, Damir Popov, Eric Johnson, Nick
Orchanian, Ashley Hellman, Geo Rangel, and Keying Chen. My lab mates in the Marinescu group
have been constant sources of support and encouragement as we have attempted to navigate our
way through the ups and downs of graduate school. I always could count on them to help with any
problems I’ve encountered. When the inevitable lows of graduate school crept in, I knew that they
were there to lend their support and guidance as well as the occasional welcomed distraction. To
Alon, Andrew, and Damir, I am thankful to have gone through this entire graduate school
experience with them. The uncertainties and challenges associated with being the first students in
a new group made for a unique experience and I am happy to have had such supportive lab mates
during those times. To our former undergraduate students, specifically Joseph Yoo, Thomas Do,
and John Luna who started in the group about the same time I did, I want to thank them for all
their hard work and contributions toward my scientific achievements. It was a privilege to work
with them.
The USC Chemistry community has been a fantastic resource during my graduate career and I’ve
been fortunate enough to have benefited from the advice and help from a variety of faculty, staff,
and fellow graduate students. The Brutchey, Thompson, and Melot groups of the Inorganic
Division have been tremendous collaborators and supporters of my research and have been an
instrumental part of my success. I also want to thank the Narayan lab, particularly Lena Hoober-
Burkhardt and Buddhinie Jayathilake, who were so generous with their time and really helped me
change the trajectory of my research project in the last several years and expand my scientific skill
set in ways I was never expecting when I asked for their help.
iv
Finally, none of this would have been possible without the continued and relentless support from
my friends and family. Caitlin DeAngelo, my roommate for all five of my years here at USC, has
been a pillar of support for me both professionally and personally. She was always available to
lend a sympathetic ear and commiserate when our science was failing, but I could always trust her
to give me great advice and provide motivation for me to keep moving forward. I certainly lucked
out when getting paired with her as a roommate during our first year in Hillview and I could not
be more grateful. To Megan Bourne, she has been my best friend since our first year at Holy Cross
and has supported me through so many different challenges and hardships. She motivated me to
stay at USC during my first semester when I was questioning my decision to go to graduate school
particularly so far away from my family. I can’t thank her enough for all she has done for me over
the years. To my friends both new and old, I want to thank them for their continued friendship and
support.
To my family, their support has meant the world to me and is the reason I have made it this far.
While my decision to attend graduate school all the way in Los Angeles was not always
understood, they stood by me and offered any kind of assistance that I needed. My two living
grandparents, Nancy and Peter Caruso, are my inspiration and the sacrifices they have made for
our family is the reason I am where I am today. I only wish my Grandma and Grandpa Downes
could be here today to share in our family’s accomplishments and achievements.
Most importantly, I feel so lucky and fortunate to have been blessed with wonderful parents, Kevin
and Joan Downes, and brother, Kevin. My brother Kevin has been my oldest friend and confidante
and I am grateful for all he has done for me. My parents have always told me that any dream of
mine was never out of reach and they have offered so much support and love over the years to help
me achieve them all. They have made sacrifices to provide me with the education and tools needed
to succeed in this world and I hope they know I have never taken these sacrifices for granted. Their
belief in me and in the person I could become has been my motivation. As I graduate with my
Ph.D. from USC, I know they are proud and that means everything to me. I hope I can continue to
make them proud in my future endeavors as their pride in me has been my greatest motivation and
source of joy. I love them dearly and cannot adequately express my gratitude for all they’ve done
for me.
v
TABLE OF CONTENTS
Acknowledgements ......................................................................................................................... iii
Table of Contents .............................................................................................................................v
List of Figures .............................................................................................................................. viii
List of Schemes ............................................................................................................................. xiii
List of Tables ................................................................................................................................ xiv
Chapter 1. General Introduction ...................................................................................................1
1.1. Global Energy Outlook .................................................................................................2
1.2. Solar Fuels ....................................................................................................................3
1.3. Hydrogen Evolution Reaction (HER) ...........................................................................4
1.4. Molecular Electrocatalysts for Solar Fuel Production ..................................................7
1.5. Electrocatalytic Metal-Organic Frameworks for Solar Fuel Production ......................9
1.6. Electronic Structure of Metal Dithiolene Complexes .................................................10
1.7. Cobalt Dithiolenes as H2 Evolution Catalysts ............................................................14
1.8. Prospects for Dithiolene-based Coordination Polymers .............................................16
1.9. References ...................................................................................................................19
Chapter 2. Electrochemical and Photoelectrochemical H2 Production by a Cobalt Dithiolene-
based Coordination Polymer ..........................................................................................23
2.1. Introduction .................................................................................................................24
2.2. Results and Discussion ...............................................................................................25
2.3. Conclusions .................................................................................................................37
2.4. Experimental Details ...................................................................................................38
2.4.1. General Considerations ................................................................................38
2.4.2. Synthesis of CoBTT .....................................................................................39
2.4.3. Formation of CoBTT|GCE...........................................................................39
2.4.4. Formation of CoBTT|Si ...............................................................................39
2.4.5. Physical Methods .........................................................................................39
2.4.6. Electrode Fabrication ...................................................................................40
2.4.7. Electrochemical and Photoelectrochemical Methods ..................................40
2.5 References ....................................................................................................................42
Chapter 3. Metal-Dithiolene (M = Ni, Fe, Zn) Coordination Polymers for Electrocatalytic H2
Evolution ...........................................................................................................................46
3.1. Introduction .................................................................................................................47
3.2. Results .........................................................................................................................49
3.3. Discussion ...................................................................................................................61
3.4. Conclusions .................................................................................................................67
3.5. Experimental Details ...................................................................................................67
3.5.1. General Considerations ................................................................................67
3.5.2. Synthesis of NiBTT .....................................................................................68
3.5.3. Synthesis of FeBTT .....................................................................................68
3.5.4. Synthesis of ZnBTT .....................................................................................69
3.5.5. Immobilization of Coordination Polymers on GCE ....................................69
3.5.6. Electrochemical Methods.............................................................................69
3.5.7. Physical Methods .........................................................................................70
3.6. References ...................................................................................................................71
vi
Chapter 4. Bioinspired Benzenetetraselenolate-based Coordination Polymers for H2
Production ........................................................................................................................75
4.1. Introduction .................................................................................................................76
4.2. Results and Discussion ...............................................................................................78
4.3. Conclusion ..................................................................................................................92
4.4. Experimental Details ...................................................................................................92
4.4.1. General Considerations ................................................................................92
4.4.2. Synthesis of CoBTSe ...................................................................................93
4.4.3. Synthesis of NiBTSe ....................................................................................93
4.4.4. Deposition of CoBTSe and NiBTSe ............................................................93
4.4.5. Electrochemical Methods.............................................................................94
4.4.6. Physical Methods .........................................................................................95
4.5. References ...................................................................................................................95
Chapter 5. H2 Evolution by a Cobalt Selenolate Electrocatalyst and Related Mechanistic
Studies ...............................................................................................................................98
5.1. Introduction .................................................................................................................99
5.2. Results and Discussion ...............................................................................................99
5.3. Conclusions ...............................................................................................................115
5.4. Experimental Details .................................................................................................115
5.4.1. General Considerations ..............................................................................115
5.4.2. Synthesis of [Co(bds)2][PPh4] (1
TPP
) .........................................................116
5.4.3. Synthesis of [Co(bds)2][nBu4N]2 (2
TBA
) ....................................................116
5.4.4. Synthesis of the Black Precipitate ..............................................................116
5.4.5. Reduction of the Black Precipitate ............................................................116
5.4.6. Protonation of [Co(bds)2][nBu4N]2 (2
TBA
) with TFA ................................116
5.4.7. Physical Methods .......................................................................................117
5.4.8. Electrochemical Methods...........................................................................117
5.4.9. Computational Methods .............................................................................118
5.4.10. X-ray Structure Determination for 1
TPP
...................................................119
5.5. References .................................................................................................................119
Chapter 6. Understanding Variability in the H2 Evolving Activity of Dithiolene-based
Coordination Polymers ..................................................................................................124
6.1. Introduction ...............................................................................................................125
6.2. Results and Discussion .............................................................................................126
6.3. Conclusions ...............................................................................................................148
6.4. Experimental Details .................................................................................................148
6.4.1. General Considerations ..............................................................................148
6.4.2. Physical Methods .......................................................................................149
6.4.3. Synthesis of CoATT ..................................................................................149
6.4.4. Deposition of CoATT ................................................................................149
6.4.5. Electrochemical Methods...........................................................................150
6.4.6. Double-Layer Capacitance Measurements ................................................150
6.4.7. Electrochemical Impedance Spectroscopy ................................................150
6.4.8. Determination of Bulk Catalyst Loading ...................................................151
6.4.9. Controlled Potential Electrolysis and Chronoamperometry ......................151
6.5. References .................................................................................................................151
vii
Chapter 7. Electrocatalytic H2 Evolution from Benzenehexathiolate-based Coordination
Frameworks and the Effect of Film Thickness on H2 Production ............................154
7.1. Introduction ...............................................................................................................155
7.2. Results and Discussion .............................................................................................159
7.2.1. Electrochemical Analysis of FeBHT and NiBHT .....................................159
7.2.2. Comparison of the HER Activity of CoBHT, NiBHT, and FeBHT ..........166
7.2.3. Evaluation of the Thickness-Dependent HER Activity of CoBHT ...........169
7.3. Conclusions ...............................................................................................................175
7.4. Experimental Details .................................................................................................176
7.4.1. General Considerations ..............................................................................176
7.4.2. Synthesis of CoBHT ..................................................................................176
7.4.3. Synthesis of CoBHT Films with Varied Thickness ...................................177
7.4.4. Synthesis of FeBHT ...................................................................................177
7.4.5. Synthesis of NiBHT ...................................................................................177
7.4.6. Deposition of CoBHT, NiBHT, and FeBHT for Electrochemistry ...........178
7.4.7. Physical Methods .......................................................................................178
7.4.8. Electrochemical Methods...........................................................................178
7.4.9. Controlled Potential Electrolysis ...............................................................179
7.5. References .................................................................................................................180
Bibliography ................................................................................................................................183
viii
LIST OF FIGURES
Figure 1.1 Schematic of photoelectrochemical and electrochemical water splitting cell ...............3
Figure 1.2 Volcano plot of heterogenous HER catalysts ................................................................5
Figure 1.3 Mechanism of the HER on electrodes in acidic solutions .............................................7
Figure 1.4 Molecular orbital diagram of bis(dithiolate) dimer .....................................................12
Figure 1.5 Molecular orbital diagrams of [Fe(bdt)2]
2
⁻ and [Co(bdt)2]⁻ ........................................13
Figure 2.1 Experimental set-up for the synthesis of CoBTT ........................................................25
Figure 2.2 SEM of CoBTT ...........................................................................................................26
Figure 2.3 FTIR of [Co(bdt)2]⁻, BTT, and CoBTT .......................................................................26
Figure 2.4 XPS of CoBTT ............................................................................................................27
Figure 2.5 Polarization curve of CoBTT|GCE in pH 10.0 aqueous solution ................................28
Figure 2.6 Polarization curve of CoBTT|GCE in pH 7.0, 4.4, 2.6 and 1.3 aqueous solutions .....28
Figure 2.7 Dependence of the electrocatalytic current densities at -0.80 V versus SHE on the
CoBTT concentration.........................................................................................................29
Figure 2.8 Tafel plots of CoBTT|GCE ..........................................................................................30
Figure 2.9 CPE of CoBTT|GCE at -0.80 V versus SHE in pH 1.3 solution .................................30
Figure 2.10 CPE of CoBTT|GCE at -0.55 V versus SHE in pH 2.6 solution ...............................31
Figure 2.11 XPS of CoBTT after electrochemical studies ............................................................32
Figure 2.12 Polarization curves of CoBTT|Si at surface concentrations of 0.5 10
-6
and 0.7
10
-6
molCo/cm
2
in H2-saturated pH 1.3 ...............................................................................33
Figure 2.13 Polarization curves of CoBTT|Si at varying catalyst loadings ..................................33
Figure 2.14 Polarization curves of CoBTT|Si at surface concentrations of 4.0 10
-6
and 6.3
10
-6
molCo/cm
2
....................................................................................................................34
Figure 2.15 Linear sweep voltammograms of CoBTT|Si and Si in H2-saturated pH 1.3 .............34
Figure 2.16 Polarization curves of CoBTT|GCE and CoBTT|Si in H2-saturated pH 1.3 .............35
Figure 2.17 CPE of CoBTT|Si at -0.12 V versus RHE in pH 1.3 solution ...................................35
Figure 2.18 CPE of CoBTT|Si at -0.12 V versus RHE under chopped illumination at pH 1.3 ....36
Figure 2.19 Controlled potential electrolysis of CoBTT|Si at -0.12 V versus RHE in pH 1.3 for 2
hours ...................................................................................................................................36
Figure 2.20 XPS of CoBTT after photoelectrochemical studies ..................................................37
Figure 2.21 XPS spectra Si 2p region of Si, CoBTT|Si, and CoBTT|Si after photoelectrochemical
studies ................................................................................................................................37
Figure 3.1 SEM images of NiBTT, FeBTT, and ZnBTT .............................................................49
Figure 3.2 FTIR of NiBTT, FeBTT, ZnBTT, and benzene-1,2,4,5-tetrathiol ..............................50
Figure 3.3 Raman spectroscopy of NiBTT, FeBTT, and benzene-1,2,4,5-tetrathiol ....................50
Figure 3.4 X-ray photoelectron spectroscopy of NiBTT ..............................................................51
Figure 3.5 X-ray photoelectron spectroscopy of FeBTT ..............................................................52
Figure 3.6 X-ray photoelectron spectroscopy of [Fe(bdt)2]
2
⁻ .......................................................52
Figure 3.7 X-ray photoelectron spectroscopy of ZnBTT ..............................................................53
ix
Figure 3.8 Scan rate dependence studies for NiBTT at pH 10.0 ..................................................54
Figure 3.9 Polarization curves of FeBTT at pH 10.0 ....................................................................54
Figure 3.10 Polarization curves of NiBTT at pH 1.3 and generated current density as a function
of catalyst loading ..............................................................................................................56
Figure 3.11 Polarization curves of NiBTT at pH 1.3 at different catalyst loadings .....................56
Figure 3.12 Tafel plot of NiBTT at pH 2.6 ...................................................................................57
Figure 3.13 Polarization curves of FeBTT at pH 1.3 and generated current density as a function
of catalyst loading ..............................................................................................................57
Figure 3.14 Polarization curves of CoBTT, NiBTT, FeBTT, and ZnBTT at pH 1.3 ...................58
Figure 3.15 Polarization curves of ZnBTT in pH 10.0 and pH 1.3 solutions ...............................58
Figure 3.16 Controlled potential electrolysis of NiBTT in pH 1.3 solution .................................59
Figure 3.17 Controlled potential electrolysis of NiBTT in pH 2.6 solution .................................59
Figure 3.18 X-ray photoelectron spectroscopy of NiBTT after 1 h CPE at pH 1.3 ......................60
Figure 3.19 X-ray photoelectron spectroscopy of NiBTT after 6 h CPE at pH 1.3 ......................60
Figure 4.1 Structures of the active sites of [NiFe] and [NiFeSe] hydrogenases ...........................76
Figure 4.2 FTIR of CoBTSe and BTSeAc4...................................................................................79
Figure 4.3 X-ray photoelectron spectroscopy of CoBTSe ............................................................80
Figure 4.4 X-ray photoelectron spectroscopy of NiBTSe.............................................................80
Figure 4.5 Polarization curves of CoBTSe and CoBTT in pH 10.0 aqueous solutions ................81
Figure 4.6 Polarization curves of NiBTSe and NiBTT in pH 10.0 aqueous solutions .................81
Figure 4.7 Scan rate dependence studies of CoBTSe in pH 10.0 aqueous solutions ....................82
Figure 4.8 Polarization curves of CoBTSe in pH 10.0-pH 1.3 aqueous solutions........................82
Figure 4.9 Polarization curves of NiBTSe in pH 10.0-pH 1.3 aqueous solutions ........................83
Figure 4.10 Current densities of CoBTSe at pH 1.3 as a function of catalyst concentration .......83
Figure 4.11 Polarization curves of CoBTSe at different catalyst concentrations at pH 1.3 .........84
Figure 4.12 CPE of CoBTSe in pH 1.3 solution at different potentials ........................................85
Figure 4.13 Polarization curves of CoBTSe and CoBTT in pH 1.3 aqueous solutions ................86
Figure 4.14 Polarization curves of CoBTSe and NiBTSe in pH 1.3 aqueous solutions ...............87
Figure 4.15 Tafel analysis of CoBTSe in pH 1.3 aqueous solutions ............................................88
Figure 4.16 CPE of NiBTSe in pH 1.3 solution at -0.47 V versus RHE ......................................89
Figure 4.17 CPE of CoBTSe in pH 1.3 solution at -0.52 V versus RHE......................................90
Figure 4.18 Polarization curves of CoBTSe before and after CPE experiments at pH 1.3 ..........90
Figure 4.19 X-ray photoelectron spectroscopy of CoBTSe after CPE experiments .....................91
Figure 4.20 X-ray photoelectron spectroscopy of NiBTSe after CPE experiments .....................91
Figure 5.1 Crystal structure of 1
TPP
.............................................................................................100
Figure 5.2 UV-Vis spectrum of [Co(bds)2]⁻ ...............................................................................100
Figure 5.3 Cyclic voltammograms of [Co(bds)2]⁻ in 0.1 M TBAPF6 MeCN and 0.1 M KNO3 1:1
CH3CN/H2O solutions .....................................................................................................101
Figure 5.4 Scan rate dependence studies for [Co(bds)2]⁻ in 0.1 M TBAPF6 MeCN solution ....101
x
Figure 5.5 Scan rate dependence studies for [Co(bds)2]⁻ in
1:1 CH3CN/H2O solution with 0.1 M
KNO3................................................................................................................................102
Figure 5.6 Molecular orbital energy scheme for [Co(bds)2]⁻ .....................................................103
Figure 5.7 Relative molecular orbital energies for [Co(bds)2]
–
versus [Co(bdt)2]⁻.....................103
Figure 5.8 CVs of [Co(bds)2]⁻ in 0.1 M TBAPF6 MeCN solution at varying trifluoroacetic acid
concentrations ..................................................................................................................104
Figure 5.9 Plot of icat/ip versus [TFA] and plot of maximum current measured at –1.35 V vs. Fc
+/0
versus concentration of [Co(bds)2]⁻ in the presence of 2.2 mM TFA in 0.1 M TBAPF6
CH3CN solution ...............................................................................................................104
Figure 5.10 Cyclic voltammograms of [Co(bds)2]⁻ in 0.1 M KNO3 1:1 CH3CN/H2O solution at
varying trifluoroacetic acid concentrations ......................................................................105
Figure 5.11 Plot of icat/ip versus [TFA] in 0.1 M KNO3 1:1 CH3CN/H2O solution ....................105
Figure 5.12 Cyclic voltammograms of glassy carbon electrodes with 19.8 mM TFA in MeCN with
0.1 M TBAPF6 and 1:1 CH3CN/H2O with 0.1 M KNO3 .................................................106
Figure 5.13 CPE of [Co(bds)2]⁻
in 0.1 M KNO3 1:1 CH3CN/H2O solution and 8.8 mM TFA at
–1.02 V vs. Fc
+/0
..............................................................................................................106
Figure 5.14 UV-Vis spectra of [Co(bds)2]⁻
with varying TFA concentrations ...........................107
Figure 5.15 UV-Vis spectra of [Co(bds)2]⁻
before and after addition of [DMF(H)][OTf] .........107
Figure 5.16 XPS of the black precipitate ....................................................................................108
Figure 5.17 XPS of [Co(bds)2]⁻ ..................................................................................................109
Figure 5.18 Comparison of XPS spectra for [Co(bds)2]⁻
and the black precipitate ....................109
Figure 5.19
1
H NMR of filtered solution from reaction of [Co(bds)2]⁻
with [DMF(H)][OTf] ...110
Figure 5.20 FTIR of the [Co(bds)2]⁻
and the black precipitate ...................................................110
Figure 5.21 UV-Vis spectra of [Co(bds)2]⁻
and black precipitate after reduction with KC8 ......111
Figure 5.22 Cyclic voltammograms of adsorbed black precipitate on GCE in 0.1 M KNO 3 1:1
CH3CN/H2O solution .......................................................................................................111
Figure 5.23 Spin densities of [Co(bds)2]⁻
/2
⁻
and [Co(bdt)2]⁻
/2
⁻ ...................................................113
Figure 5.24 UV-Vis spectra of [Co(bds)2]⁻
and [Co(bds)2]
2
⁻ ......................................................114
Figure 6.1 FTIR of ATTAc4 and CoATT ...................................................................................127
Figure 6.2 SEM of CoATT .........................................................................................................127
Figure 6.3 XPS of CoATT ..........................................................................................................128
Figure 6.4 Polarization curves of CoATT-modified GCE in pH 1.3 solutions ..........................129
Figure 6.5 Cyclic voltammograms of CoATT-modified GCE to measure Cdl ...........................129
Figure 6.6 Current density difference at 0.15 V versus RHE plotted against the scan rate for
CoATT-modified GCE ....................................................................................................130
Figure 6.7 Cdl versus the overpotential to reach 10 mA/cm
2
for CoATT-modified GCE ..........130
Figure 6.8 Nyquist plots measured at -0.57 V versus RHE for CoATT-modified GCE ............131
Figure 6.9 Equivalent circuits used to fit EIS data .....................................................................131
Figure 6.10 Rct versus the overpotential to reach 10 mA/cm
2
for CoATT-modified GCE ........132
Figure 6.11 Polarization curves of CoATT-modified GR in pH 1.3 solutions ...........................133
xi
Figure 6.12 Cyclic voltammograms of CoATT-modified GR to measure Cdl ...........................134
Figure 6.13 Current density difference at 0.15 V versus RHE plotted against the scan rate for
CoATT-modified GR .......................................................................................................134
Figure 6.14 Cdl versus the overpotential to reach 10 mA/cm
2
for CoATT-modified GR ...........135
Figure 6.15 Relationship between [Co] and Cdl and [Co] and the overpotential ........................136
Figure 6.16 Current density at -0.40 and -0.50 V versus RHE versus [Co] ...............................137
Figure 6.17 Polarization curves of CoATT-modified GR plotted as a function of the generated
catalytic current per [Co] .................................................................................................137
Figure 6.18 Electrochemical impedance spectroscopy spectra of GR-2 ....................................138
Figure 6.19 Electrochemical impedance spectroscopy spectra of GR-4 ....................................138
Figure 6.20 Electrochemical impedance spectroscopy spectra of GR-5 ....................................138
Figure 6.21 Electrochemical impedance spectroscopy spectra of GR-6 ....................................139
Figure 6.22 EIS spectra measured at -0.57 V versus RHE for CoATT-modified GR ................140
Figure 6.23 Rs and R1 plotted with respect to overpotential for CoATT-modified GR..............141
Figure 6.24 Dependence of the Rct on the overpotential for CoATT-modified GR ...................142
Figure 6.25 Dependence of the overpotential to reach 10 mA/cm
2
on the Rct values measured at
-0.57 V versus RHE for CoATT-modified GR ................................................................143
Figure 6.26 Dependence of the overpotential to achieve 10 mA/cm
2
on the double layer
capacitance measured under HER conditions (C
*
dl) ........................................................144
Figure 6.27 Tafel plots of GR-2, GR-3, and GR-5 .....................................................................145
Figure 6.28 CPE of CoATT-modified GCE in pH 1.3 at -0.72 V versus RHE for 1 h ..............146
Figure 6.29 CPE of CoATT-modified GR in pH 1.3 at -0.72 V versus RHE for 7 h .................146
Figure 6.30 CPE of CoATT-modified GCE in pH 1.3 at -0.72 V versus RHE for 12 h ............146
Figure 6.31 XPS of CoATT after 1 h CPE .................................................................................147
Figure 6.32 XPS of CoATT after 12 h CPE ...............................................................................147
Figure 7.1 Equivalent circuit model (2TS) .................................................................................159
Figure 7.2 Electrochemical analysis of FeBHT-modified GCE .................................................160
Figure 7.3 Plots of Cdl versus overpotential and Rct versus Cdl for FeBHT-modified GCE .......161
Figure 7.4 Tafel analysis of FeBHT-modified GCE ...................................................................161
Figure 7.5 Electrochemical analysis of NiBHT-modified GCE .................................................162
Figure 7.6 Comparison of Bode plots for NiBHT-modified GCE ..............................................163
Figure 7.7 EIS spectra represented as Nyquist and Bode plots for NiBHT-modified GCE at
different overpotentials ....................................................................................................163
Figure 7.8 Tafel analysis of NiBHT-modified GCE ...................................................................164
Figure 7.9 Controlled potential electrolysis of NiBHT at -0.72 V versus RHE in pH 1.3 .........164
Figure 7.10 XPS of NiBHT.........................................................................................................165
Figure 7.11 XPS of NiBHT following CPE in pH 1.3 solution at -0.72 V versus RHE ............165
Figure 7.12 Comparison of polarizations curves and Cdl for CoBHT, NiBHT, and FeBHT ......166
Figure 7.13 EIS spectra for CoBHT, NiBHT, and FeBHT .........................................................167
Figure 7.14 Nyquist plots for CoBHT, NiBHT, and FeBHT at -0.27 V and -0.37 V ................167
xii
Figure 7.15 Bode plots for CoBHT, NiBHT, and FeBHT at -0.27 V and -0.37 V .....................167
Figure 7.16 Tafel analysis of CoBHT, NiBHT, and FeBHT ......................................................168
Figure 7.17 SEM of CoBHT-23 ..................................................................................................169
Figure 7.18 SEM of CoBHT-57 ..................................................................................................170
Figure 7.19 SEM of CoBHT-157 ................................................................................................170
Figure 7.20 SEM of CoBHT-244 ................................................................................................170
Figure 7.21 SEM of CoBHT-1000 ..............................................................................................170
Figure 7.22 Polarization curves and Cdl measurements for CoBHT-modified GCE with different
thicknesses .......................................................................................................................171
Figure 7.23 EIS spectra represented as Nyquist and Bode plots for CoBHT-modified GCE with
different thicknesses.........................................................................................................172
Figure 7.24 Comparison of the current density at various potentials as a function of film thickness
for CoBHT .......................................................................................................................173
Figure 7.25 Tafel analysis of CoBHT-modified GCE with different thicknesses ......................174
xiii
LIST OF SCHEMES
Scheme 1.1 Oxidation states of dithiolene ligand .........................................................................11
Scheme 1.2 Electronic structure of [Co(bdt)2]⁻ .............................................................................13
Scheme 1.3 Mechanism of H2 evolution from [Co(bdt)2]⁻............................................................15
Scheme 1.4 Structure of cobalt dithiolene-based MOFs ...............................................................16
Scheme 2.1 Synthesis of CoBTT...................................................................................................25
Scheme 3.1 Structures of cobalt dithiolene coordination polymers ..............................................48
Scheme 3.2 Synthesis of NiBTT, FeBTT, and ZnBTT .................................................................49
Scheme 3.3 Oxidation states of dithiolene ligand .........................................................................61
Scheme 3.4 Selection of surface-confined cobalt molecular catalysts ..........................................65
Scheme 3.5 Selection of surface-confined nickel molecular catalysts ..........................................66
Scheme 4.1 Synthesis of CoBTSe and NiBTSe ............................................................................78
Scheme 5.1 Mechanistic pathways for H2 evolution from [Co(bds)2]⁻ .......................................114
Scheme 6.1 Synthesis of CoATT ................................................................................................126
Scheme 7.1 Structure of CoBHT, NiBHT, and FeBHT ..............................................................157
xiv
LIST OF TABLES
Table 3.1 Electrocatalytic HER properties of metal dithiolene polymers.....................................62
Table 4.1 Overpotential to reach 10 mA/cm
2
at different CoBTSe loadings ................................84
Table 4.2 Electrocatalytic HER properties of the diselenolate- and dithiolate-based coordination
polymers .............................................................................................................................88
Table 5.1 Comparison of XPS binding energies for [Co(bds)2]
–
and the black precipitate ........109
Table 5.2 Crystal data and structure refinement for 1
TPP
............................................................120
Table 5.3 Bond lengths (Å) for 1
TPP
............................................................................................121
Table 5.4 Bond angles (°) for 1
TPP
..............................................................................................121
Tabe 6.1 Overview of CoATT-modified GCE............................................................................130
Table 6.2 Values from EIS fitting at -0.57 V versus RHE for CoATT-modified GCE ..............132
Table 6.3 Overview of CoATT-modified GR .............................................................................135
Table 6.4 Values from EIS fitting for GR-1................................................................................140
Table 6.5 Values from EIS fitting for GR-2................................................................................140
Table 6.6 Values from EIS fitting for GR-3................................................................................140
Table 6.7 Values from EIS fitting for GR-4................................................................................141
Table 6.8 Values from EIS fitting for GR-5................................................................................141
Table 6.9 Values from EIS fitting for GR-6................................................................................141
Table 7.1 Electrocatalytic HER properties for FeBHT-modified GCE ......................................161
Table 7.2 Electrocatalytic HER properties for NiBHT-modified GCE ......................................162
Table 7.3 Electrocatalytic HER properties for CoBHT, NiBHT, and FeBHT............................166
Table 7.4 Electrocatalytic HER properties of CoBHT at different film thicknesses ..................172
1
CHAPTER 1
General Introduction
A portion of this chapter has appeared in print:
Downes, C. A.; Marinescu, S. C. “Electrocatalytic Metal-Organic Frameworks for Energy
Applications.” ChemSusChem, 2017, 10, 4374-4392.
2
1.1. Global Energy Outlook
The transformation of our current fossil-fuel dominated energy system into a sustainable one where
energy is clean, affordable, accessible, and reliable is vital for preserving the environment and
combating the adverse effects of climate change.
1-3
The U.S. Energy Information Administration
(EIA) projects in the 2017 International Energy Outlook that total world energy consumption will
rise from 575 quadrillion British thermal units (Btu) in 2015 to 736 quadrillion Btu in 2040, an
increase of 28%.
4
Countries outside of the Organization for Economic Cooperation and
Development (OECD), specifically China and India, will account for more than half of the total
increase in global energy consumption from 2015-2040.
4
Rapidly rising population and economic
growth in Non-OECD nations necessitate greater resources and energy use. By 2040, almost two-
thirds of all energy will be consumed by Non-OECD nations.
4
The reliance on fossil fuels to meet this growing energy demand has raised many concerns because
of the exacerbation of climate change associated with continued fossil fuel use and greenhouse gas
emission. Fossil fuels are projected to account for 77% of energy use in 2040, with natural gas the
fastest growing fossil fuel source.
4
Energy efficiency and conservation, in addition to
decarbonizing our energy sources, are important strategies for sustainably meeting the world
energy demand.
2,3
This demand can be met, without compromising the environment, by
diversifying the global energy portfolio through the introduction of abundant renewable energy
sources such as solar, wind, and hydro power.
Renewable energy sources have been successfully integrated into the electricity sector and are
projected to be the fastest-growing source of energy for electricity generation from 2015 to 2040.
4
However, the intermittency of renewable energy may lead to a mismatch between energy supply
and demand.
1,3
As renewable resources begin to account for larger percentages of the electricity
supplied to the grid, the variability of renewable energy output needs to be addressed. While
renewable energy penetration in the electricity sector through the utilization of solar, wind, and
hydro power technologies has been significant, the transportation and industrial sectors require the
development of sustainable pathways moving toward decarbonization and electrification.
3,5
Liquid
fuels are expected to remain the dominant source of energy for transportation, although the share
of total transportation energy is projected to decline from 95% in 2015 to 88% in 2040.
4
In the
3
industrial sector, fossil fuels are the primary source of carbon feedstocks and provide the energy
needed to drive the chemical transformations.
1.2. Solar Fuels
One possible strategy to mitigate the intermittent nature of renewable energy sources and decrease
the reliance on traditional fossil fuels in the transportation and industrial sectors is through the
conversion and storage of renewable energy in chemical bonds.
1,5,6
Renewable energy can be
directly converted into high-value products such as hydrogen and hydrocarbons that can act as
energy carriers, fuels, and feedstocks (Figure 1.1a).
1,5,7,8
Alternatively, an indirect electrochemical
pathway where electricity generated from renewable energy drives the energy conversion is also
viable (Figure 1.1b).
5,7,9
The stored energy can be converted back to electricity to lessen the impact
of intermittent renewables on the grid, used as feedstocks to manufacture high value industrially
relevant chemicals, or directly as liquid fuels in the transportation sector.
1,5-7
Figure 1.1. Schematic of (a) photoelectrochemical (direct) and (b) electrochemical (indirect) water
splitting devices. Figure 1.1a reprinted from ref. 8 with permission from Elsevier. Figure 1.1b
reprinted from ref. 9.
The splitting of water into oxygen and hydrogen is an attractive strategy for the sustainable
production of hydrogen. Hydrogen, which has the highest energy density per mass of any fuel, has
received significant attention as an energy carrier that can be utilized in fuel cell technology to
convert chemical energy to electrical energy.
1,2
Carbon dioxide, a prominent greenhouse gas, can
also be converted to valuable products such as renewably produced carbon based fuels or
feedstocks for the synthesis of industrial chemicals.
6
Unfortunately, the conversions of H2O, H2,
O2, and CO2 into useful energy carriers and fuels using renewable energy are difficult and energy-
4
consuming. The high activation barriers associated with these transformations result in operation
at high overpotentials limiting their efficiencies and conversion rates. The design and
implementation of catalysts to facilitate these transformations is necessary for reducing the
activation barrier for the energy conversions of interest. For practical application and global
deployment of energy converting technologies, catalysts must be made from abundant and
inexpensive materials, operate in environmentally benign conditions, and perform at high
efficiencies with limited energy consumption to facilitate the conversion.
1.3. Hydrogen Evolution Reaction (HER)
One of the most important solar fuel producing reactions for transitioning from a fossil-fuel
dominated energy economy to a sustainable alternative is the hydrogen evolution reaction (HER),
the cathodic side of water splitting. Steam-methane reforming (eq. 1) is currently used industrially
to generate ~95% of the H2 needed in the United States.
10
Although this process offers the lowest
cost and highest energy efficiency route to hydrogen production, significant CO2 emissions
necessitate the development of a carbon-neutral pathway for H2 generation if a sustainable energy
future is to be realized.
CH4 + 2H2O → CO2 + 4H2 (1)
The splitting of water into oxygen and hydrogen (eq. 2) is one of the most attractive strategies for
renewably producing H2.
1,2
2H2O → O2 + 2H2 (2)
This sustainable source of hydrogen can be used as a mechanism for storing renewable energy in
chemical bonds to alleviate the problems associated with intermittent and variable energy sources.
The energy stored in the chemical bonds of H2 can be easily converted as demand rises to electrical
power in fuel cells with environmentally benign by-products (water) generating a carbon-neutral
fuel production and consumption cycle.
1
Additionally, renewably produced H2 can be utilized
directly as a fuel or in other areas of catalysis to synthesize valuable chemicals such as ammonia
and methanol.
6
Thermodynamically, water splitting into O2 and H2 requires 1.23 V necessitating
the development of electrocatalysts to affect this transformation. Focusing on the cathodic side of
water splitting, the hydrogen evolution reaction (eq. 3) is a two-electron transfer reaction.
5
2H
+
+ 2e
-
→ H2 (3)
The best-performing catalyst for the HER, platinum, is too scarce and expensive for global
deployment of water-splitting devices. Therefore, the development of homogeneous and
heterogeneous earth-abundant alternatives to platinum has been of great interest.
11-15
The Gibbs
free energy of hydrogen adsorption (ΔGH*) has been identified as a good description of the intrinsic
HER activity of heterogeneous and surface immobilized catalysts and is an important metric for
screening potential earth-abundant replacements for platinum.
16
According to the Sabatier
principle, an active catalyst must bind reaction intermediate(s) neither too strongly nor too weakly.
Plotting the experimentally measured exchange current density against the DFT-calculated ΔGH*
generates a volcano plot that reflects the importance of the Sabatier principle (Figure 1.2).
16
Platinum, the best performing HER catalyst, sits near the top of the volcano plot with a near
thermoneutral hydrogen adsorption energy. Catalysts to the right of platinum bind hydrogen too
weakly hindering the formation of any stable intermediates needed for H2 generation. Conversely,
catalysts to the left of platinum bind hydrogen too strongly inhibiting the H2 desorption/evolution
step. Therefore, the ΔGH* is an important metric to consider when designing earth-abundant HER
catalysts to replace platinum.
Figure 1.2. A volcano plot of the experimentally measured exchange current density as a function
of the DFT calculated Gibbs free energy of adsorption of hydrogen. Reproduced from ref. 16 with
permission of the Royal Society of Chemistry.
There are several other key parameters and metrics that are used to evaluate electrocatalysts for
the HER. The overpotential (η) is the difference between the standard potential and the actual
operating potential for the transformation of interest. Some electrochemical processes, such as the
6
HER, define the overpotential at a specific current density value, typically 10 mA/cm
2
.
17
Catalytic
potentials near the thermodynamic potential (small overpotentials) for the HER and efficiencies
near unity are necessary if replacement of platinum is to be realized. The efficiency of the
transformation is described by the Faradaic efficiency (FE), which compares the yield of the
electrochemical products with the charge supplied to the system to promote the conversion. For
practical systems, the total supplied charge must be used to form the products of interest with no
unwanted side-product formation. Extended electrolysis experiments are needed to assess the
stability and durability of the catalyst as many homogeneous and heterogeneous catalysts undergo
decomposition and/or deactivation following exposure to reductive acidic and/or alkaline
conditions. Relatedly, important metrics to evaluate the intrinsic activity of a catalyst are the
turnover number (TON), which is defined as the number of catalytic cycles achieved by a single
catalyst before deactivation, and the turnover frequency (TOF), which is the number of turnovers
per unit time. The density and reactivity of active sites must also be investigated and comparisons
between the bulk catalyst loading and the number of accessible and active catalytic sites should be
made. For heterogenous and surface-immobilized catalysts, the surface area of the materials is
important for understanding substrate and product diffusion and the accessibility of active sites.
As a two proton/two electron process, proton and H2 diffusion is not the only limiting factor for
high electrocatalytic activity. The ability to transport electrons from the electrode surface
throughout the heterogenous and surface immobilized catalyst is vital for generating large
electrocatalytic current densities and therefore, the charge transfer properties and conductivities of
the catalysts needs to be evaluated.
To interrogate different mechanistic pathways, the Tafel slope, which is extracted from Tafel
analysis experiments that examine the current response as a function of applied potential, is
dependent on the rate-limiting step.
18
The linear portions of the Tafel plots are fitted to the Tafel
equation, η = b log j + a (where η is the overpotential, a the Tafel constant, b the Tafel slope, and
j the current density). In acidic electrolytes, three reaction steps are possible for the HER (Figure
1.3): the Volmer (discharge) reaction (H3O
+
+ e
-
→ Hads + H2O; b = 120 mV/dec), the Heyrovsky
(ion + atom) reaction (H3O
+
+ e
-
+ cat-H → cat + H2 + H2O; b = 40 mV/dec), and the Tafel
(combination) reaction (cat-H + cat-H → 2cat + H2; b = 30 mV/dec).
18
The exchange current
density (j0) extracted from the Tafel equation is correlated to the rate of the electron transfer under
7
reversible conditions. The magnitude of j0 describes the electrochemical rate particularly the
efficacy of transferring electrons across catalytic interfaces.
16
Catalysts with low Tafel slopes and
high exchange current densities display high electrocatalytic activity. Determination of these
metrics is necessary for benchmarking new catalysts against the state of the art systems and
facilitates a better understanding of the properties that dictate the observed HER activity to aid in
designing and/or optimizing potential replacements for platinum.
Figure 1.3. Mechanism of the hydrogen evolution reaction on electrodes in acidic solutions.
Reproduced from ref. 16 with permission of the Royal Society of Chemistry.
1.4. Molecular Electrocatalysts for Solar Fuel Production
The design of earth abundant molecular electrocatalysts for solar fuel production to replace the
state of the art precious metals-based systems has been extensively explored. Advances in
molecular catalysts for the HER,
12,15,19
the oxygen evolution reaction (OER),
20-23
the oxygen
reduction reaction (ORR),
20,24,25
and CO2 reduction reaction (CO2 RR)
26-29
have been thoroughly
reviewed. Many molecular catalysts for energy converting applications have been designed as
models or biomimics of the active sites of enzymes that facilitate these reactions in nature. Most
of these active sites contain earth-abundant metals with intricate coordination environments poised
to facilitate operation at high efficiencies near the thermodynamic potential. The active sites of
[NiFe], [FeFe], and [Fe] hydrogenase enzymes, which catalyze hydrogen evolution and oxidation
8
at ambient conditions near the thermodynamic potential with turnover frequencies up to 9000 s⁻
1
for H2 production,
30,31
have inspired the design of earth-abundant molecular catalysts with
analogous coordination environments and structures to promote H
+
, H2, and e⁻ transfer in order to
generate high efficiencies at low overpotentials.
32-36
The use of these molecular catalysts in homogeneous solutions offers several advantages over
traditional heterogenous catalysts. Molecular catalysts have well-defined catalytic intermediates
that can be isolated, characterized, and studied to elucidate the mechanism of catalysis. The
structure of the catalyst is also extremely modular with catalyst selectivity, efficiency, and
mechanism tuned through simple synthetic modifications. These modifications include control
over the first, second, and outer coordination spheres. Molecular catalysts offer enhanced
understanding and control in comparison to heterogeneous systems. However, the use of
homogeneous catalysts has several disadvantages, such as poor recoverability and reusability. The
need for organic solvents because of solubility limitations hinders the integration of molecular
catalysts into energy conversion devices such as fuel cells and electrolyzers. Additionally,
molecular catalysts are only active when in the diffusion layer near the electrode surface resulting
in inactivity for most of the catalyst in solution.
37
This free diffusion of catalyst in solution is also
problematic for practical construction of energy conversion devices that require separation of the
reduction and oxidation catalysts to avoid degradation.
The incorporation of molecular catalysts into heterogeneous structures is a viable method for
maintaining the attractive properties of molecular systems while reducing the typical drawbacks
these systems suffer from. Heterogenized molecular catalysts can exhibit improved stability
because common bimolecular decomposition and deactivation pathways are eliminated due to site
isolation. Solubility is no longer a factor as the heterogenized molecular catalysts are immobilized
on the electrode surface allowing for operation in aqueous media. Because the catalyst is
immobilized directly on the electrode surface, the activity is not limited by catalyst diffusion;
therefore, the entirety of the immobilized catalyst could be active. Depending on the method of
immobilization, the heterogenized molecular catalyst can be removed from the electrode surface
and reused, allowing for recyclability of the system, which is necessary for energy converting
applications.
9
Methods such as covalent attachment,
35,38-40
surface polymerization,
41-45
and adsorption/non-
covalent attachment of metal complexes
46-50
have been used to heterogenize molecular catalysts
for energy conversion applications.
51-53
However, these methods generally suffer from low surface
concentrations, inability to facilitate charge transfer, poor control over the structure and orientation
of the catalysts (with respect to the electrode surface), and limited operational lifetimes.
Coordination polymers, particularly metal-organic frameworks (MOFs), have emerged as one of
the most promising platforms for incorporating well-defined molecular catalytic units thereby
enabling the realization of high active site densities in robust, crystalline, porous, and modular
architectures.
1.5. Electrocatalytic Metal-Organic Frameworks for Solar Fuel Production
Coordination polymers such as metal-organic frameworks (MOFs) offer the possibility of
combining the advantageous properties of homogeneous and heterogeneous electrocatalysts.
MOFs are composed of metal ions or clusters bridged by organic ligands to afford crystalline
structures with high surface areas and permanent porosity.
54-57
The hybrid nature of MOFs allows
for synthetic tunability leading to MOFs of varying pore sizes and chemical environments, which
can result in different chemical and physical properties.
54-57
The permanent porosity of MOFs
facilitates rapid substrate diffusion throughout the material and could result in unprecedented
active site densities for catalysis. Metal site isolation in MOFs is analogous to the well-defined
nature of molecular catalysts, while the MOF structure can introduce the stability (site isolation
limits bimolecular decomposition) and robustness typically associated with heterogeneous
catalysts. Although MOFs exhibit several attractive qualities for application as electrocatalysts,
most MOFs are insulating,
58,59
which limits their electrocatalytic activity.
Several strategies have been developed to overcome the poor conductivity of pristine MOFs and
improve their electrocatalytic performance. Highly conductive materials such as graphene or
graphene oxide can be integrated with MOFs to generate MOF composites that exhibit improved
conductivity and enhanced electrocatalytic activity.
60-62
Catalytic nanoparticles or guest species
can be loaded into the MOFs where the MOF acts as a porous template.
60-64
MOF derivatives are
formed following carbonization of MOFs to generate nanoporous carbon (NPC) structures with
high surface areas and narrow pore-size distributions.
61,65,66
In these strategies, the MOF typically
10
does not act as an electrocatalyst, but as a support to load functional moieties or precursors for
other heterogeneous catalytic materials.
Pristine MOFs, without any additives, have been successfully utilized as electrocatalysts through
control of the film thickness of deposited MOFs and/or incorporation of well-defined catalytic
units into the MOF structure.
67,68
Control over the film thickness of MOFs is important for
enhancing and facilitating electron transport through the catalytic active sites.
67,69,70
The
development of synthetic methods for growing thin film MOFs is vital because the utilization of
MOF powders requires polymeric binders, such as Nafion, to construct the electrodes. The use of
such binders can inhibit electrical integration between the catalyst and electrode resulting in poor
conductivity and charge transfer. Additionally, most MOFs use redox inactive ligands as linkers
limiting the ability for charge transport resulting in the insulating nature observed for many
MOFs.
58,71
The recent development of 2D conductive MOFs with redox active ligands connected by metal
ions that can facilitate charge transport has been a major breakthrough for the field of
electrocatalytic MOFs.
59
Nishihara and coworkers employed interfacial reaction conditions to
generate 2D MOFs based on the coordination of benzenhexathiolate and nickel.
72
The interfacially
grown nickel dithiolene MOF could be readily deposited on substrates of interest and displayed a
conductivity of 0.15 S/cm, thus confirming the use of redox-active ligands and metal ions is a
viable method for generating conductive MOFs. Inspired by this work, we sought to design
electrocatalytic 2D MOFs through the use of well-defined molecular catalytic units that utilize
charge transfer promoting redox-active ligands, such as dithiolenes, as the building block and
repeat unit for framework formation. By modulating the denticity of the redox-active ligand, 2D
MOFs and one-dimensional coordination polymer architectures can be generated whereby the
repeat unit is a well-known molecular catalyst for the H2 evolution reaction.
1.6. Electronic Structure of Metal Dithiolene Complexes
In the 1960s, the concept of non-innocence was developed to describe the electronic structure of
dithiolene ligands.
73-75
Metal complexes with these ligands were intensely colored and were
observed to undergo facile one-electron transfers. However, traditional oxidation state assignments
11
could not accurately describe their electronic structures because of the multiple oxidation states
available for the dithiolene ligand that facilitate ligand-based redox events (Scheme 1.1).
75
Scheme 1.1. The available oxidation states of a dithiolene ligand. Reprinted from ref. 75 with
permission from Elsevier.
Computational analysis and electronic structure calculations of metal dithiolene complexes
revealed extensive ligand-metal mixing in the frontier orbitals, which suggested the possibility of
one-electron redox events localized on the ligand (non-innocent) rather than the metal center
(innocent).
74,75
Rigorous spectroscopic investigations of metal dithiolene complexes utilizing
sulfur K-edge and metal K- and L-edge X-ray absorption spectroscopy (XAS) confirmed the non-
innocent nature of dithiolene ligands by allowing for the direct measurement of the oxidation state
of sulfur in the dithiolene ligand and of the metal center.
75
Evaluation of the electronic structure of a variety of transition metal bis(dithiolene) complexes has
led to the realization that the relative energies of the metal and dithiolene ligand orbitals dictate
whether the redox events are localized on the metal, ligand, or in some cases both. The importance
of the bis(dithiolene) and metal ligand orbitals necessitates an understanding of the molecular
orbital diagrams of the free ligand and of the metal bis(dithiolene) complexes with varying metal
centers. The hypothetical bis(dithiolate) dimer (L2)
4-
with a symmetry of D2h (Figure 1.4) has eight
orbitals corresponding to the symmetry adapted linear combinations of the S 3p orbitals: four -
type (1b3g, 1b1u, 1b2g, 1au) and four -type (1b2u, 1b3u, 1b1g, 1ag).
76,77
The orbitals with gerade
symmetry (1b3g, 1b1g, 1b2g, 1ag) can undergo symmetry-allowed interactions with the metal d-
orbitals with the appropriate symmetry. Most importantly, the HOMO
*
-b2g
can interact with the
dxz orbital of a transition metal. The innocent and non-innocent behavior of the dithiolene ligand
is dictated by the extent of the interaction of the ligand
*
-b2g with the metal dxz, which is
modulated by the relative energies of the orbitals.
76,77
Two bonding schemes, normal and inverted,
can be envisioned where the metal orbitals are destabilized with respect to the ligand orbitals in
the normal mode and stabilized in the inverted scenario.
12
Figure 1.4. Molecular orbital diagram of the hypothetical bis(dithiolate) dimer with D2h symmetry
obtained from BP86 DFT results.
77
Reproduced from ref. 76 with permission from the Royal
Society of Chemistry.
Of relevance for designing earth-abundant HER catalysts, the electronic structures of iron, cobalt,
and nickel bis(benzenedithiolate) complexes are discussed. As previously mentioned, the relative
energies of the
*
-b2g
of the benzene-1,2-dithiolate (bdt) dimer and the metal dxz orbital influence
the extent of metal-ligand mixing and the degree of non-innocent character in the complex. For
the metal d-orbitals, the effective nuclear charge (Zeff) of the metal, which increases from left to
right across the periodic table, dictates the energy of the orbitals. If the metal orbitals are
destabilized with respect to the ligand orbitals, a normal bonding scheme will occur. The 2b2g
SOMO of the metal complex will be metal based because the metal d orbitals are much higher in
energy than the ligand orbitals and cannot sufficiently engage in back-bonding interactions with
the ligand orbitals.
76
This is the case for Fe bis(dithiolene) complexes where the Fe 3d orbitals are
destabilized with respect to the benzene-1,2-dithiolate orbitals representing a normal bonding
scenario where the metal orbitals are higher in energy than the ligand orbitals. The higher in energy
metal orbitals contribute more to the 2b2g SOMO for the dianionic complex, [Fe(bdt)2]
2
⁻ (Figure
1.5). Therefore, benzene-1,2-dithiolate is considered innocent and the redox events are metal
centered because the 2b2g orbital is primarily Fe character (82%).
76,78
Because of the increase in
13
Zeff for cobalt in comparison to iron, the Co d-orbitals are stabilized and the dxz of Co and the b2g
of benzene-1,2-dithiolate are considered energetically equivalent. Because of the energy match,
the 2b2g orbital of [Co(bdt)2]⁻ is a mixture of metal and ligand based character (Figure 1.5).
78
Figure 1.5. Molecular orbital diagrams of [Co(L)2]⁻
and [Fe(L)2]
2
⁻ complexes (L = benzene-1,2,-
dithiolate). Reproduced from ref. 76 with permission from the Royal Society of Chemistry.
This results in a large amount of ligand-to-metal charge transfer in the electronic ground state of
[Co(bdt)2]⁻. The benzene-1,2-dithiolate ligands are no longer innocent as in the Fe case and the
electronic structure is assigned as having major contributions from three resonance forms (Scheme
1.2).
78
This leads to an inability to ambiguously assign the oxidation state of cobalt or the ligand
for the monoanionic ground state because of the large amount of ligand to metal charge transfer
occurring. Therefore, the electronic structure ground state is best described as a mixture of cobalt
in the +2 and +3 oxidation states with the corresponding benzene-1,2-dithiolate ligands as dianions
or containing radical character.
[Co
III
(bdt
2
⁻)2]⁻ ↔ [Co
II
(bdt
2
⁻)(bdt•⁻)]⁻
↔ [Co
II
(bdt•⁻)(bdt
2
⁻)]⁻
Scheme 1.2. Electronic structure of the ground state of [Co(bdt)2]⁻ is a combination of three
resonance structures that contain either oxidized cobalt or ligand (bdt = benzene-1,2-dithiolate).
Moving further across the periodic table, the effective nuclear charge increases even more and the
14
3d orbitals of Ni are stabilized and lower in energy than the Co 3d orbitals. Because the Ni 3d
orbitals are so low in energy, the 2b2g orbital of the resultant nickel bis(benzendithiolate) complex
is ligand based and the electronic structure of the ground state is best described as
[Ni
II
(bdt
2
⁻)(bdt•⁻)]⁻.
76,77
Sequential one electron oxidations occur on the dithiolene ligands. Nickel
retains the +2 oxidation state (d
8
) throughout the redox events. It was concluded that metal
bis(dithiolene) complexes with more than 50% S 3p character in the redox active 2b2g SOMO were
best described as mixed-valence ligand radicals with one electron redox events occurring on
orbitals of predominately ligand character.
1.7. Cobalt Dithiolenes as H2 Evolution Catalysts
Complexes of dithiolenes have received renewed interest because their unique electronic structures
featuring delocalized frontier orbitals are advantageous for use in multielectron catalytic
transformations. Redox active dithiolene ligands could facilitate storage of electrons on the ligand
rather than the metal center. This metal-ligand cooperativity whereby during catalysis, both the
metal and ligand participate in the activation and transformation of the substrate of interest, is
reminiscent of enzymatic catalysis and can allow for the emergence of new catalytic reactions.
Cobalt dithiolenes, which have significant metal-ligand mixing in their frontier orbitals have
shown great promise as electrocatalysts, photocatalysts, and photoelectrocatalysts for the 2H
+
/2e⁻
transfer reaction to produce H2.
79-82
Holland, Eisenberg, and coworkers first identified
[Co(bdt)2][TBA] (where bdt = benzene-1,2-dithiolate, TBA = tetrabutylammonium) as an active
electrocatalyst and photocatalyst in mixed aqueous/organic media. A TON of 2700 was measured
in 1:1 H2O/MeCN at pH 4.0 with Ru(bpy)3
2+
and ascorbic acid as the chromophore and sacrificial
donor, respectively. [Co(bdt)2][TBA] also displayed a reversible redox wave at -1.01 V versus
Fc
+/0
assigned to the [Co(bdt)2]⁻/[Co(bdt)2]
2
⁻ redox couple and the addition of trifluoracetic acid
triggered the appearance of a catalytic wave near this redox event. Controlled potential electrolysis
experiments at -1.0 V versus SCE in 1:1 H2O/MeCN with 65 mM of tosic acid revealed
electrocatalytic production of H2 at a Faradaic efficiency of >99%.
Computational analysis was employed by Hammes-Schiffer and coworkers to discern possible
mechanistic pathways for H2 evolution from [Co(bdt)2][TBA].
83
It was suggested that following
initial reduction to the dianionic complex, [Co(bdt)2]
2
⁻, two protonation events occur on different
15
sulfur moieties of the dithiolene ligand (Scheme 1.3). The protonation of the sulfur donors is
possible because of the extensive mixed metal-ligand character of the frontier orbitals of
[Co(bdt)2][TBA]. Additionally, more highly protonated complexes are easier to reduce thus
enabling a potential mechanism for reducing the overpotential necessary to drive H2 formation. A
second reduction of the ligand-protonated complex is followed by an intramolecular proton
transfer to generate a Co
III
-H with an adjacent protonated sulfur moiety. In the final step, H2 is
evolved through heterocoupling of H
+
and H⁻.
Scheme 1.3. Mechanism of H2 evolution from a cobalt bis(benzendithiolate) complex. Reprinted
with permission from ref. 83. Copyright 2014 American Chemical Society.
Following these initial reports, [Co(bdt)2][TBA] has been coupled with CdSe quantum dots and
ascorbic acid for photogeneration of H2 at pH 4.5 with a TON of 300,000 (60 h).
82
The system
displayed remarkable stability for over 30 h with a slight decline in activity observed as the
concentration of ascorbic acid decreased. Expanding on the use of CdSe nanocrystals as light
absorbers, CdSe quantum-dot sensitized NiO photocathodes were constructed and employed with
[Co(bdt)2][TBA] for photoelectrochemical H2 production. An onset of 0 V versus RHE was
observed under illumination in 0.1 M KCl with pH 6 buffer in 1:1 MeCN/H2O representing a 200
mV positive shift in the onset relative to the NiO-CdSe photocathodes without catalyst. Controlled
potential electrolysis under illumination at -0.28 V versus RHE generated 2 mA/cm
2
of catalytic
current density for 2 h with a FE of 100% and TON of 400 ± 30 with respect to [Co(bdt)2][TBA].
The versatility of [Co(bdt)2]⁻ as a H2 evolution catalyst in addition to the stability under reductive
16
and acidic aqueous conditions, the high efficiency for H2 production, and square planar
coordination geometry rendered [Co(bdt)2]⁻ an ideal candidate for incorporation into a 2D
extended framework architecture as a method to bypass many of the disadvantages associated with
traditional homogeneous catalysts that limit their practical feasibility.
1.8. Prospects for Dithiolene-based Coordination Polymers
To overcome the limitations of molecular catalysts, we explored the synthesis of extended
architectures with integrated cobalt dithiolene catalytic units. Trinucleating ligand scaffolds,
benzenehexathiolate (BHT) and triphenylene-2,3,6,7,10,11-hexathiolate (THT), were employed in
reactions with cobalt (II) to generate CoBHT and CoTHT MOFs (Scheme 1.4).
84
Upon evaporation
of the organic layer, a film was left at the newly formed gas-liquid interface and deposition of the
MOF films on the substrates of interest occurred via a bottom-up or top-down approach.
Scheme 1.4. Cobalt dithiolene-based MOFs (CoBHT and CoTHT) formed from the reaction of
Co
2+
and the trinucleating ligands, benzenehexathiolate or triphenylene-2,3,6,7,10,11-
hexathiolate.
17
Electrochemical measurements were carried out in aqueous solutions ranging from pH 10.0 to pH
1.3 using glassy carbon electrodes. In pH 1.3 aqueous solutions, to achieve the benchmarking
metric of 10 mA/cm
2
, CoBHT and CoTHT exhibited overpotentials of 340 and 530 mV,
respectively, at catalyst loadings of 0.7(1) × 10
-6
molCo/cm
2
and 1.1(1) × 10
-6
molCo/cm
2
. Tafel
slopes of 108 and 161 mV/dec and exchange current densities of 10
-5.3(1)
A/cm
2
were measured in
pH 2.6 aqueous solutions for CoBHT and CoTHT, respectively. The large Tafel slopes indicated
the Volmer discharge reaction, proton adsorption, was rate-limiting. The FE for H2 production was
97 ± 3%. Controlled potential electrolysis (CPE) experiments of CoBHT in pH 2.6 aqueous
solutions at -0.40 V and -0.50 V vs RHE displayed continuous charge build-up over 10 h with
decreases in current density associated with delamination of the catalyst from the electrode surface.
Decreases in the measured catalyst loading over the course of the experiment confirmed
delamination. X-ray photoelectron spectroscopy (XPS) of CoBHT and CoTHT after
electrochemical testing revealed Co, Na, and S peaks analogous to those observed before
electrolysis and inductively couple plasma mass spectrometry (ICP-MS) analysis of the resultant
electrolysis solutions indicated no solubilized cobalt. The integration of cobalt dithiolene catalytic
units into a network environment afforded stability under reductive and acidic conditions. Overall,
this study identified the first example of a MOF displaying high intrinsic electrocatalytic activity.
84
In this thesis, we look to expand this methodology to the use of dinucleating ligand scaffolds to
synthesize electrocatalytic thiolate- and selenolate-based coordination polymers for solar fuel
production. Chapter 2 describes the utilization of the dinucleating ligand scaffold, benzene-1,2,4,5-
tetrathiolate (BTT), to generate CoBTT, which was integrated with glassy carbon electrodes
(GCE) and silicon photocathodes for electrochemical and photoelectrocemical H2 evolution in pH
1.3 aqueous solutions. The conductive nature of metal dithiolene coordination polymers based on
BTT, which has been previously investigated, is important for achieving high electrocatalytic
activities.
85
Under simulated 1 sun illumination, the CoBTT-modified silicon photocathodes
generated photocurrents of up to 3.8 mA/cm
2
at 0 V vs RHE, demonstrating successful solar driven
hydrogen production. However, the CoBTT-modified GCE operated at large overpotentials (~560
mV) to achieve 10 mA/cm
2
with the deposition method offering limited durability during extended
electrolysis.
18
Chapter 3 and Chapter 4 explore strategies for reducing the overpotential of the coordination
polymers synthesized from dinucleating benzene-based ligands. Chapter 3 investigates the role of
the metal center in influencing the HER activity of BTT-based coordination polymers. NiBTT,
FeBTT, and ZnBTT were synthesized and electrochemically analyzed on GCE in pH 1.3 aqueous
solutions. The use of the dinucleating ligand BTT to synthesize NiBTT via an interfacial reaction
was first reported by Nishihara and coworkers.
86
NiBTT operated at a smaller overpotential, but
with a reduced efficiency in comparison to CoBTT. Another method for optimizing the HER
activity and reducing the overpotential is described in Chapter 4. The role of the chalcogen is
investigated by utilizing benzene-1,2,4,5-tetraselenolate (BTSe) as the ligand for the synthesis of
cobalt (CoBTSe) and nickel (NiBTSe) coordination polymers. In nature when one cysteine of
[NiFe] hydrogenase is replaced with a selenocysteine in [NiFeSe] hydrogenase, the H 2 evolution
activity is significantly improved upon incorporation of the selenium residue.
87,88
Relatedly, metal
selenides have also shown higher electrocatalytic H2 activity in comparison to metal sulfide
catalysts.
89
In Chapter 4, CoBTSe operates at ~200 mV smaller overpotential than the sulfur
analogue, CoBTT. Chapter 5 looks to understand the origin of the reduction in overpotential for
selenolate-based cobalt catalysts in comparison to their thiolate-based counterparts. The
mechanistic pathways for H2 evolution from a cobalt diselenolate molecular complex are
described.
Chapter 6 details the synthesis of a cobalt coordination polymer based on 9,10-dimethyl-2,3,6,7-
anthracenetetrathiolate. Utilizing this material, we look to understand the fundamental properties
that dictate the observed electrocatalytic HER activity. Previous investigations of thiolate- and
selnolate-based coordination polymers revealed variance in the HER activity for samples with the
same bulk catalyst loading and it was observed that magnitude of the change in bulk loading did
not always correlate to the magnitude of the changes in the measured overpotential. Metrics such
as the bulk catalyst loading, the electrochemically active surface area, and the charge transfer
resistance are interrogated in Chapter 6 to discern the factors that dictate and influence the variance
in the observed HER activity.
In Chapter 7, the effect of the metal center on the HER activity of BHT-based coordination
polymers is analyzed by electrochemically characterizing NiBHT and FeBHT. The
19
electrochemical behavior of CoBHT is also revisited with an emphasis on how the film thickness
influences the H2 production. Properties such as charge transfer, proton transfer, and the number
of accessible active sites are modulated as the film thickness increases and this relationship is
interrogated over a thickness range of 20 to 1000 nm. The coordination polymers synthesized in
this thesis all display different electrocatalytic HER behavior highlighting the retention of the
molecular nature and the ease at which the electrochemical properties can be tuned through simple
synthetic modifications. Coordination polymer architectures are therefore a viable method for
combining the advantageous properties of homogeneous and heterogenous catalysts, which
enables integration with practical solar-to-fuel converting devices.
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23
CHAPTER 2
Electrochemical and Photoelectrochemical H
2
Production
by a Cobalt Dithiolene-based Coordination Polymer
A portion of this chapter has appeared in print:
Downes, C. A.; Marinescu, S. C. “Efficient Electrochemical and Photoelectrochemical H2
Production from Water by a Cobalt Dithiolene One Dimensional Metal-Organic Surface.” J. Am.
Chem. Soc. 2015, 137, 13740-134743.
24
2.1. Introduction
Solar-driven hydrogen production through the reduction of water has emerged as an important
strategy for the development of clean-energy technologies.
1,2
The hydrogen evolution reaction
(HER) is efficiently catalyzed by noble metals, such as Pt, that generate large cathodic currents at
low overpotentials. However, the scarcity and high cost of platinum limits global deployment of
solar-to-fuel converting devices. Replacement of Pt with non-precious metal catalysts is desirable
for practical applications and global scalability of such potential clean-energy technologies. These
considerations have led to the development of homogeneous
3-9
and heterogeneous
10
catalysts that
employ non-precious metals. Efficient and practical clean-energy devices for the HER require the
immobilization of catalytically active species, based on inexpensive metals, onto electrodes.
11,12
Molecular catalysts are attractive because the ligand environment allows for tuning of their
reduction potentials and chemical properties. However, the reported methods for the
immobilization of molecular catalysts onto electrodes suffer from low surface coverage, poor
operational lifetimes, and limited substrate scope.
11,13-18
Of interest for practical artificial photosynthesis devices are integrated photoelectrochemical
systems that couple light harvesting and solar fuel generation to enable direct solar-to-hydrogen
production. Silicon is an excellent candidate as a light absorber in a photoelectrochemical system
due to its ideal band gap (Eg = 1.12 eV) for the absorption of low energy sunlight to drive the
HER.
19-21
Several photocathodes based on heterogeneous noble
21-26
and non-precious
27-38
metal
catalysts for solar-driven HER have been recently reported. In the realm of molecular catalysts, p-
type silicon has been used as a photocathode with free catalysts in solution
39-41
or with grafted
molecular catalysts
42
, albeit with low activity for the HER. Although new methodologies for the
modification of silicon surfaces through covalent attachment of molecular species show great
promise,
43-48
the development of integrated molecular catalysts|Si photocathode materials with
high activity for the solar-driven H2 production remains a major challenge.
We have recently reported that cobalt dithiolene units can be incorporated into 2D metal-organic
frameworks (MOFs) using the trinucleating ligands benzenehexathiolate (BHT) and triphenylene-
2,3,6,7,10,11-hexathiolate (THT). These cobalt dithiolene-based extended frameworks can be
readily immobilized onto carbon based electrodes and display high electrocatalytic H 2-evolving
25
activity and stability in acidic aqueous solutions.
49
We extend here the methodology to the
formation of a cobalt dithiolene coordination polymer (CoBTT) based on benzene-1,2,4,5-
tetrathiolate (BTT) and report its electrochemical and photoelectrochemical H2 evolving activity.
The utilization of the dinucleating ligand BTT following our previous investigations using the
trinucleating ligands BHT and THT demonstrates the versatility and modular nature of
coordination polymers as scaffolds for integrating cobalt dithiolene units.
Scheme 2.1. Synthesis of benzene-1,2,4,5-tetrathiolate-based cobalt coordination polymer
(CoBTT) through a liquid-liquid interfacial reaction.
2.2. Results and Discussion
Cobalt dithiolene species are among the most active molecular catalysts for the HER.
50-55
A cobalt
dithiolene polymer can be accessed through a liquid-liquid interfacial process whereby a cobalt(II)
precursor reacts with a dinucleating conjugated ligand, benzene-1,2,4,5-tetrathiol, which is
deprotonated in situ to generate BTT (Scheme 2.1). To that end, an ethyl acetate solution of
benzene-1,2,4,5-tetrathiol was gently layered on top of an aqueous solution of cobalt(II) acetate
and sodium acetate. The organic solvents evaporated over 1 h, leaving behind a black solid
(CoBTT) at the gas-liquid interface (Figure 2.1).
Figure 2.1. Typical set-up for the synthesis of the cobalt dithiolene coordination polymer based
on benzene-1,2,4,5-tetrathiolate, CoBTT.
26
Immersion of a carbon-based support (glassy carbon electrode – GCE, and highly oriented
pyrolytic graphite – HOPG) face down into the reaction mixture led to the deposition of CoBTT
generating CoBTT|GCE and CoBTT|HOPG, respectively, which were washed to remove the
excess starting materials and dried. Top-down scanning electron microscopy images of
CoBTT|HOPG, depict the rough and cracked morphology of deposited CoBTT (Figure 2.2).
Alternatively, CoBTT was collected, washed, suspended in water, and dropcast onto freshly
etched planar p-type Si(100) electrodes to generate CoBTT|Si.
Figure 2.2. A top-down scanning electron microscopy image of CoBTT|HOPG.
The measured elemental composition of CoBTT corresponds to a molecular formula of
[Co(C6H2S4)][Na].
49
The FTIR spectrum of CoBTT showed the disappearance of the strong S–H
stretching vibration present in benzene-1,2,4,5-tetrathiol at 2500 cm
–1
upon formation of the
coordination polymer (Figure 2.3).
Figure 2.3. FTIR spectra of [Co(bdt)2][nBu4N] (bdt = benzene-1,2-dithiolate; purple), benzene-
1,2,4,5-tetrathiol (teal), and CoBTT (red).
27
X-ray photoelectron spectroscopy (XPS) analysis of CoBTT revealed the presence of Co, S, and
Na (Figure 2.4), analogous with the cobalt dithiolene frameworks based on benzenehexathiolate
and triphenylene-2,3,6,7,10,11-hexathiolate.
49
Two sets of peaks were observed in the cobalt
region, with binding energies of ~780 eV and ~795 eV, which correspond to the 2p3/2 and 2p1/2
levels. Deconvolution of these signals generates four peaks; the peaks at 779.2 and 794.1 eV were
assigned to Co
III
, whereas the peaks at 780.7 and 796.6 eV were assigned to Co
II
.
56
This assignment
is in agreement with the electronic structure of cobalt bis(dithiolene) complexes, which have been
interrogated by a variety of spectroscopic and theoretical studies, and are best represented by the
resonances [Co
III
(bdt)2]⁻ [Co
II
(bdt)(bdt
•
)]⁻ (bdt = benzene-1,2-dithiolate).
57
Three additional
peaks were observed for CoBTT with binding energies of 1071.4 eV, ~228 eV, and ~163 eV,
which correspond to Na 1s, S 2s, and S 2p, respectively.
58-61
These data support assignment of the
coordination polymer as cobalt dithiolene moieties linked by tetrathiolate nodes.
Figure 2.4. X-ray photoelectron spectroscopy analysis of CoBTT showing the Co 2p, Na 1s, S 2s,
and S 2p core level XPS spectra.
The electrochemistry of CoBTT|GCE was investigated by cyclic voltammetry. A broad
irreversible reduction wave was observed in pH 10.0 aqueous solutions between 0.1 and –0.3 V
28
versus SHE (Figure 2.5). This reduction wave is similar with the one reported for the cobalt
dithiolene framework based on triphenylene-2,3,6,7,10,11-hexathiolate.
49
Figure 2.5. Polarization curve of CoBTT|GCE (catalyst loading: 4.4(4) × 10
–7
molCo/cm
2
) in N2-
saturated 0.1 M NaClO4 aqueous solutions at pH 10.0. Scan rate: 20 mV/s.
Maximum average surface catalyst concentrations of 5.5(6) × 10
–7
molCo/cm
2
were estimated from
the integration of the electrochemical wave observed in pH 10.0 solutions. Similar values were
obtained from inductively coupled plasma (ICP) measurements of the sonicated and digested
catalyst CoBTT, suggesting that most of the cobalt centers were electrochemically active.
Figure 2.6. Polarization curves of CoBTT (5.5(6) × 10
–7
molCo/cm
2
) in N2-saturated 0.1 M
NaClO4 aqueous solutions at pH 7.0 (purple), 4.4 (green), 2.6 (blue), 1.3 (red), and of blank GCE
at pH 1.3 (dashed black). Scan rate: 20 mV/s.
As the pH of the aqueous solution was lowered, an increase in current was observed (Figure 2.6),
indicating that a catalytic reaction is taking place. An overpotential of 560 mV was required to
29
reach a current density of 10 mA/cm
2
in pH 1.3 solutions. This overpotential is similar with the
one reported for the cobalt dithiolene framework based on triphenylene-2,3,6,7,10,11-hexathiolate
(530 mV) and ~200 mV more negative than the benzenehexathiolate system,
49
highlighting the
tunable nature of these cobalt dithiolene-based systems, an advantageous characteristic of
molecular catalysts. By comparison, unmodified GCE displayed insignificant increases in current.
The modified electrodes did not generate any soluble materials responsible for catalysis, as
indicated by negligible currents observed in the cyclic voltammetry studies of the solutions
resulted after the electrochemical studies of CoBTT|GCE, and by negligible soluble cobalt
concentrations determined from ICP measurements. Catalytic current densities of CoBTT|GCE
measured at potentials of –0.80 V versus SHE (–0.72 V versus RHE) in pH 1.3 solution increased
linearly with catalyst loading (Figure 2.7).
Figure 2.7. Current densities of CoBTT|GCE measured at –0.80 V versus SHE (–0.72 V versus
RHE) in N2-saturated pH 1.3 solutions as a function of the surface catalyst concentration. The
surface concentration was quantified by integrating the peak area of the redox wave observed in
pH 10.0 solutions or from ICP-MS studies of the sonicated and digested polymer CoBTT.
Tafel analyses gave slopes of 94(4) mV/dec, and exchange current densities of 10
–8.2(4)
A/cm
2
in
aqueous solutions at pH 4.4 and 2.6 (Figure 2.8), which are comparable to the values reported for
the cobalt dithiolene framework based on triphenylene-2,3,6,7,10,11-hexathiolate,
49
or cobalt
complexes onto amine-modified multiwalled carbon nanotubes.
13
30
Figure 2.8. Tafel plots of CoBTT|GCE (4.4(4) × 10
–7
molCo/cm
2
) in N2-saturated 0.1 M NaClO4
aqueous solutions at pH 4.4 (green; Tafel slope of 93 mV/dec; exhange current density of 10
–8.1
A/cm
2
), pH 2.6 (red; Tafel slope of 94 mV/dec; exchange current density of 10
–8.3
A/cm
2
), and pH
1.3 (blue; Tafel slope of 70 mV/dec, exchange current density of 10
–9.4
A/cm
2
). Scan rate: 0.5
mV/s.
Controlled potential electrolysis (CPE) of CoBTT|GCE in pH 1.3 H2SO4 solution, measured at –
0.80 V versus SHE (–0.72 V versus RHE), consumed 32 coulombs of charge after 2 h (Figure 2.9).
Analysis of the gas mixture in the headspace of the working compartment of the electrolysis cell
by gas chromatography confirmed production of H2 with a Faradaic yield of 97 ± 3%. By
comparison, unmodified GCE consumed only 4 coulombs of charge after a 2 h CPE in pH 1.3
solution at –0.80 V versus SHE (Figure 2.9).
Figure 2.9. Controlled potential electrolysis (CPE) at –0.80 V versus SHE (–0.72 V versus RHE)
of CoBTT|GCE (1.0(1) × 10
–7
molCo/cm
2
) in N2-saturated 0.1 M NaClO4 aqueous solutions at pH
1.3 (red) and of blank GCE (black dashed).
31
The durability of CoBTT|GCE in pH 2.6 aqueous solution was further assessed in a longer-
duration CPE experiment. CoBTT|GCE affords a continuous increase in charge build-up over a
10 h CPE at –0.55 V versus SHE (–0.40 V versus RHE) (Figure 2.10), however with moderate
stability. The observed decrease in activity over time was attributed to delamination of the catalyst
due to H2 bubble formation.
Figure 2.10. Controlled potential electrolysis at –0.55 V versus SHE (–0.40 V versus RHE) of
CoBTT|GCE (1.0(1) × 10
–7
molCo/cm
2
) in N2-saturated 0.1 M NaClO4 aqueous solutions at pH
2.6 (red) and of blank GCE (black dashed). It is important to note that, for longer controlled
potential electrolysis measurements, it is difficult to maintain equal pressure in both chambers of
the H-cell. This leads to variances in the surface area of the working electrode immersed in
solution, and therefore variances in the measured current. The spikes in current are due to the
formation of H2 bubbles, or due to short pause in the experiment to re-equilibrate the levels of the
liquids in the two compartments of the H-cell. The catalyst loading of CoBTT|GCE at the end of
the experiment is 0.51(5) × 10
–7
molCo/cm
2
.
XPS analyses of CoBTT|GCE after electrochemical studies displayed peaks similar with the ones
observed before electrolysis, suggesting that CoBTT is stable under reductive acidic conditions
(Figure 2.11). The growth of the peaks ~ 233 and ~169 eV in the S 2s and S 2p regions,
respectively, is indicative of oxidized sulfur most likely due to the sulfuric acid used in the pH 1.3
aqueous solutions.
32
Figure 2.11. X-ray photoelectron spectroscopy analysis of CoBTT after electrochemical studies
showing the Co 2p, Na 1s, S 2s, and S 2p core level XPS spectra.
The photoelectrochemistry of CoBTT|Si was investigated by linear sweep voltammetry (LSV) in
H2-saturated pH 1.3 H2SO4 aqueous solution under 100 mW/cm
2
of simulated air mass (AM)1.5G
(1 Sun) illumination. The surface concentration of CoBTT|Si was controlled by varying the
amount of CoBTT that was dropcast onto freshly etched Si(100) electrodes. The surface
concentrations of CoBTT|Si were obtained from ICP measurements of the sonicated and digested
catalyst. The maximum surface concentration of CoBTT|Si was 6.3(6) × 10
–6
molCo/cm
2
, which is
an order of magnitude higher than the surface concentration of CoBTT|GCE. Under 1 Sun
illumination, CoBTT|Si photocathodes with a surface concentration as little as 0.5(1) × 10
–6
molCo/cm
2
displayed a significant improvement in both the onset of photocurrent, as well as the
current density compared to that of bare p-Si(100) (Figure 2.12-2.13). Negligible current densities
were observed in the absence of light for CoBTT|Si photocathodes. At low surface concentrations,
slightly higher current densities were observed for photocathodes generated by dropcasting
sonicated rather than as synthesized suspensions of CoBTT in water (Figure 2.12), due to a higher
surface contact between the finely dispersed particles of CoBTT and Si. A similar behavior was
reported for photocathodes based on chemically exfoliated metallic MoS2.
34
33
Figure 2.12. Polarization curves of as synthesized (0.7(1) × 10
–6
molCo/cm
2
, pink) and sonicated
(0.5(1) × 10
–6
molCo/cm
2
, blue) CoBTT|Si and of Si (black) in H2-saturated pH 1.3 H2SO4 solutions
under simulated 1 Sun illumination. Polarization curves of CoBTT|Si (dashed black) and Si (cyan)
measured in the absence of light. Scan rate: 10 mV/s.
Gradually increasing the surface concentrations of CoBTT|Si to 4.0(4) × 10
–6
molCo/cm
2
shifted
the onset of photocurrent to more positive values and improved the current density at 0 V versus
RHE to 3.8 mA/cm
2
(Figure 2.13). Further increasing the surface concentration to 6.3(6) × 10
–6
molCo/cm
2
led to a decrease in catalytic activity (Figure 2.14), likely due to blockage of the light
absorption, as reported for photocathodes based on MoS2.
34
Figure 2.13. Polarization curves of CoBTT|Si with varying surface concentrations ranging from
0.5(1)–4.0(4) × 10
–6
molCo/cm
2
(blue–red) and of Si (black) in H2-saturated pH 1.3 H2SO4 solutions
under simulated 1 Sun illumination. Polarization curves of CoBTT|Si (dashed black) and bare Si
(cyan) measured in the absence of light. Scan rate: 10 mV/s.
34
Figure 2.14. Polarization curves of CoBTT|Si with surface concentrations of 4.0(4) × 10
–6
molCo/cm
2
(red) and 6.3(6) × 10
–6
molCo/cm
2
(purple) and of Si (black) in H2-saturated pH 1.3
H2SO4 solutions under simulated 1 Sun illumination. Polarization curves of CoBTT|Si (dashed
black) and Si (cyan) measured in the absence of light. Scan rate: 10 mV/s.
The light-limited current density of CoBTT|Si (Jph ≈ 21 mA/cm
2
)
was less than the value
anticipated from the band gap of Si due to losses associated with absorption of the incident photons
by the black CoBTT catalyst and the oxidation of the Si surface (Figure 2.15).
Figure 2.15. Linear sweep voltammograms of CoBTT|Si (4.0(4) × 10
–6
molCo/cm
2
) and bare Si
with and without simulated 1 Sun illumination in H2-saturated pH 1.3 H2SO4 solutions. Scan rate:
10 mV/s.
A 430 mV positive shift in the onset of activity to reach 1 mA/cm
2
was observed for CoBTT|Si
with the optimal surface concentration of 4.0(4) × 10
–6
molCo/cm
2
compared to the activity of the
Si photocathode (Figure 2.13). Moreover, to reach the same activity for H2 evolution from pH 1.3
35
H2SO4 solutions (1 mA/cm
2
), CoBTT|Si photocathodes operate at potentials 550 mV more
positive than CoBTT|GCE cathodes (Figure 2.16).
Figure 2.16. Polarization curves of illuminated (green) and dark (dashed black) CoBTT|Si (4.0(4)
× 10
–6
molCo/cm
2
) and CoBTT|GCE (7.6(7) × 10
–7
molCo/cm
2
, purple) in H2-saturated pH 1.3
H2SO4 solution. Scan rate: 10 mV/s.
CPE studies performed at –0.12 V versus RHE in pH 1.3 H2SO4 solution under simulated 1 Sun
illumination show that the current produced for CoBTT|Si was stable for 20 minutes (Figure 2.17-
2.18). Analysis of the gas mixture in the headspace of the working compartment of the
photoelectrolysis cell confirmed production of H2 with a Faradaic yield of 80 ± 3%. By
comparison, unmodified Si photocathodes displayed very little H2-evolving activity. Negligible
current densities were observed for CPE studies performed in the absence of light (Figure 2.18).
Figure 2.17 CPE studies of CoBTT|Si (0.9(1) × 10
–6
molCo/cm
2
, red) and bare Si (black) at –0.12
V versus RHE in N2-saturated pH 1.3 H2SO4 solutions under simulated 1 Sun illumination.
36
Figure 2.18 CPE studies of CoBTT|Si (0.6(1) × 10
–6
molCo/cm
2
, red) and bare Si (black) at –0.12
V versus RHE in H2-saturated pH 1.3 H2SO4 solutions under chopped (light on/light off) 1 Sun
illumination.
The durability of CoBTT|Si at –0.12 V versus RHE in pH 1.3 H2SO4 solution was further assessed
in a longer-duration CPE experiment. CoBTT|Si affords a continuous increase in charge build-up
over a 2 h CPE at –0.12 V versus RHE (Figure 2.19), however with limited stability, due to
delamination of the catalyst from the Si surface, as described above for CoBTT|GCE.
Figure 2.19. Controlled potential electrolysis studies of CoBTT|Si (0.91(1) × 10
–6
molCo/cm
2
, red)
and bare Si (black) at –0.12 V versus RHE in N2-saturated pH 1.3 H2SO4 solutions under simulated
1 Sun illumination.
XPS analyses of CoBTT|Si after photoelectrochemical studies displayed peaks similar with the
ones observed before, suggesting that CoBTT was stable during photoelectrochemical
experiments (Figure 2.20). Si 2p core level XPS spectra of Si and CoBTT|Si contained peaks at
~99 and ~102 eV, indicative of Si–H and Si–O, respectively, terminated surfaces (Figure 2.21).
Integration of CoBTT with freshly etched Si passivates the photocathode surface reducing Si–O
37
formation. However, the passivation was limited as oxidation was observed following
photoelectrochemical studies.
Figure 2.20. X-ray photoelectron spectroscopy (XPS) analysis of CoBTT after
photoelectrochemical studies showing the Co 2p, Na 1s, S 2s, and S 2p core level XPS spectra.
Figure 2.21. Si 2p core level X-ray photoelectron spectroscopy (XPS) spectra of Si (black) and
CoBTT|Si before (pink) and after (blue) photoelectrochemical studies.
2.3. Conclusions
In summary, we demonstrate here the successful integration of a cobalt dithiolene coordination
polymer based on benzene-1,2,4,5-tetrathiolate (CoBTT) with glassy carbon or Si electrodes to
38
generate catalyst modified cathodes. CoBTT|Si is an efficient photocathode material for solar-
driven hydrogen production from water, achieving photocurrents of 3.8 mA/cm
2
at 0 V versus
RHE under simulated 1 Sun illumination, which is among the highest photocurrents reported for
an immobilized molecular catalyst integrated with Si photocathode materials. The previously
reported covalent attachment of a nickel phosphine H2 evolving catalyst to a p-Si(111)
photoelectrode show promising activity at low overpotentials, although negligible current densities
are displayed at 0 V versus RHE under the conditions tested.
42
Moreover, the maximum surface
concentration of CoBTT|Si is 6.3(6) × 10
–6
molCo/cm
2
, which is four orders of magnitude higher
than the catalyst loadings reported for the covalently attached nickel phosphine H2 evolving
catalyst.
42
To reach a current density of 1 mA/cm
2
, CoBTT|Si photocathodes require potentials
550 mV less negative than CoBTT|GCE cathodes (photovoltage). Similar values were reported
for the photovoltage of several free catalysts in solution.
39-41
Other photocathode materials, such
as InP,
62
GaP,
17,18,63
or water-solubilized CdSe quantum dots,
53,64-67
have been investigated for
their ability to interface with molecular systems, and, despite their great promise, their chemical
stability and photoelectrochemical performance needs improvement. The generated CoBTT|GCE
and CoBTT|Si cathodes display high activities and moderate stabilities, suggesting that the
integration of a cobalt dithiolene molecular catalyst into a coordination polymer is a viable
immobilization strategy and thus paves the way towards development of practical devices.
2.4. Experimental Details
2.4.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity). Ethyl acetate and water were placed under vacuum and refilled with nitrogen (10 ×).
Benzene-1,2,4,5-tetrathiol
68
and [Co(bdt)2][nBu4N]
50
were prepared according to the reported
procedures. All other chemical reagents were purchased from commercial vendors and used
without further purification. Elemental analyses and ICP-MS studies (to confirm surface
concentrations of CoBTT immobilized on glassy carbon electrodes) were performed by Robertson
MicroLit Laboratories, Ledgewood, NJ, 07852. The pHs of the aqueous solutions were measured
with a benchtop Mettler Toledo pH meter.
39
Caution: Perchlorate salts are potential explosive chemicals and should only be used in very small
amounts, especially in the presence of the reductant H2. Please handle these mixtures using the
proper equipment and protection.
2.4.2. Synthesis of CoBTT
In a glove-box, Co(OAc)2(H2O)4 (13.3 mg, 0.053 mmol, 1.1 equiv) and anhydrous NaOAc (8.8
mg, 0.107 mmol, 2.2 equiv) were dissolved in 20 mL of degassed H2O. The aqueous solution was
transferred to a 70 mm x 50 mm crystallizing dish. An ethyl acetate solution (1 mL) of benzene-
1,2,4,5-tetrathiol (C6H6S4) (10.0 mg, 0.049 mmol, 1 equiv) was carefully layered via glass pipette
on top of the aqueous solution of cobalt(II) acetate and sodium acetate to cover ~80% of the
aqueous solution. The organic solvents were allowed to evaporate over one to two hours at room
temperature, leaving behind CoBTT as a black solid at the gas-liquid interface. The black solid
was collected, washed with water, methanol, and ethyl acetate. The resultant black powder was
dried under vacuum. Anal. Calcd for [Co(C6H2S4)][Na]•MeCO2H•H2O (C8H8CoNaO3S4): C,
26.52; H, 2.23; Co, 16.27; S, 35.39. Found: C, 25.68; H, 2.25; Co, 17.05; S, 32.29.
2.4.3. Formation of CoBTT|GCE
CoBTT was deposited onto glassy carbon electrodes by immersing a polished electrode face down
through the black film at the gas-liquid interface. The volatiles were allowed to evaporate at room
temperature. The modified electrode was washed with water, methanol, and ethyl acetate.
2.4.4. Formation of CoBTT|Si
CoBTT (15 mg) was suspended in water (2 mL) and dropcast onto the silicon electrodes. The
volatiles were allowed to evaporate at room temperature. The electrodes were washed with water
and ethyl acetate.
2.4.5. Physical Methods
FTIR spectra were acquired using a Bruker Vertex 80v spectrometer. Samples (2 mg) for analysis
were mixed into a KBr (100 mg) matrix and pressed into pellets.
40
Inductively coupled plasma optical emission spectroscopy (ICP-OES) measurements were
performed using a Thermo Scientific iCAP 7000 ICP-OES. A 1000 ppm cobalt standard in nitric
acid (Sigma Aldrich) was used to construct a calibration plot.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray
source was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
Scanning electron microscopy (SEM) images were attained using a JEOL JSM 7001F field-
emission microscope equipped for energy-dispersive X-ray spectroscopy (EDS) for elemental
analysis.
2.4.6. Electrode Fabrication
Prime (100) p-type silicon wafers with resistivity of 1–10 Ohm-cm were purchased from
University Wafer. A 50.8 mm diameter p-type silicon wafer was cut into 1 cm
2
pieces and were
cleaned by immersion in piranha acid (3:1 H2SO4/30% H2O2; caution: potentially explosive) for
30 minutes, rinsed with pure water (18.2 M ·cm resistivity), and dried under a stream of N2. A
gallium-indium eutectic (Aldrich) was applied to the back side of a freshly scratched Si chip. An
ohmic contact was established by embedding a coiled tinned copper wire and subsequently
covering the wire and eutectic with silver paint (SPI supplies). The Cu wire was passed through a
glass tube and epoxy (Loctite Hysol 1C) was used to insulate the entire back side of the Si from
the electrolyte and to define a 0.09 cm
2
surface area on the polished front side of the Si electrode.
The electrodes were etched for 1 minute in 6:1 buffered oxide etch (BOE), rinsed with pure water,
and dried under a stream of N2. The electrodes were immediately transferred to a N 2 filled glove-
box for preparation of CoBTT|Si photoelectrodes and assembly of the photoelectrochemical cell.
2.4.7. Electrochemical and Photoelectrochemical Methods
Electrochemistry experiments were carried out using a Pine potentiostat. Platinum wire used for
the electrochemical studies was purchased from Alfa Aesar. The electrochemical experiments
41
were carried out in a three-electrode configuration electrochemical cell under an inert atmosphere
(N2 or H2) using either glassy carbon electrodes (GCE) or silicon photoelectrodes as the working
electrode. A platinum wire, placed in a separate compartment, connected by a Vycor tip, and filled
with the electrolytic solution (0.1 M NaClO4) was used as the auxiliary (counter) electrode. The
reference electrode, placed in a separate compartment and connected by a Vycor tip, was based on
an aqueous Ag/AgCl/saturated KCl electrode. The reference electrode in aqueous media was
calibrated externally relative to ferrocenecarboxylic acid (Fc-COOH) at pH 7.0, with the Fe
3+/2+
couple at 0.28 V versus Ag/AgCl. All potentials reported in this paper were converted to the
standard hydrogen electrode (SHE) by adding a value of 0.205 V, or to the reversible hydrogen
electrode (RHE) by adding a value of (0.205 + 0.059 pH) V.
Simulated 1 Sun illumination was supplied by a 300 W Xe Arc lamp (Newport Corp. Model 66485;
AM 1.5G filter). The light intensity was calibrated to 100 mW/cm
2
using a Si reference cell
(Newport Corp. Model 91150V). A 25 mL Schlenk flask equipped with a magnetic stir bar and a
side arm was charged in a N2 glove-box with the Si photocathode, the aqueous Ag/AgCl/saturated
KCl electrode placed in a separate compartment and connected by a Vycor tip, and a platinum wire
also placed in a separate compartment and connected by a Vycor tip. 20 mL of pH 1.3 solution
were added to the three electrode photoelectrochemical (PEC) cell, which was then sealed in the
N2 glove-box before transferring to the Schlenk line where the N2 atmosphere was replaced by
research grade H2. Linear sweep voltammetry experiments were performed in pH 1.3 solutions
with constant purging of H2 and rapid stirring at a scan rate of 10 mV/s. Controlled potential
electrolysis (CPE) studies under chopped 1 Sun illumination were performed in pH 1.3 solutions
with constant purging of H2 and rapid stirring.
CPE measurements to determine Faradaic efficiency and study long-term stability were conducted
in a sealed two-chambered H-cell where the first chamber held the working and reference
electrodes in 50 mL of 0.1 M NaClO4 (aq) at the corresponding pH, and the second chamber held
the auxiliary electrode in 25 mL of 0.1 M NaClO4 (aq). The two chambers, which were both under
a N2 or H2 atmosphere, were separated by a fine porosity glass frit. Glassy carbon plate electrodes
(6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and auxiliary electrodes.
For photoelectrochemical CPE experiments, a silicon photocathode with a surface area of 0.25 cm
2
42
was used as the working electrode. The reference electrode was a Ag/AgCl/saturated KCl electrode
separated from the solution by a Vycor tip. Using a gas-tight syringe, 10 mL of gas were withdrawn
from the headspace of the H-cell and injected into a gas chromatography instrument (Shimadzu
GC-2010-Plus) equipped with a BID detector and a Restek ShinCarbon ST Micropacked column.
To determine the Faradaic efficiency of the electrochemical and photoelectrochemical cells, the
theoretical H2 amount based on total charge flowed is compared with the GC-detected H2 produced
from controlled-potential electrolysis.
To measure the catalyst loading and analyze the aqueous solutions after photoelectrochemical and
electrochemical experiments, inductively coupled plasma optical emission spectroscopy (ICP-
OES) was performed. To determine catalyst loading, CoBTT was washed from the silicon
electrode surface with acetone to remove all of the black powder. The acetone was allowed to
evaporate and the collected powder was digested in 2 mL of concentrated HNO3. Pure water was
added to the solution to reach a total volume of 25 mL. The samples were run in triplicate and the
presented catalyst loadings are the average of the three measurements. Following CPE
experiments, the pH 1.3 solutions were collected, filtered, and analyzed by ICP-OES to determine
if any cobalt material was solubilized during the experiment.
The aqueous solutions used in the electrochemical experiments have been prepared as follows. For
the pH 1.3 solution, 0.534 mL of 18.7 M H2SO4 were added to water (200 mL). For the pH 2.6
solution, citric acid (3.458 g) and Na2HPO4 (1.505 g) were dissolved in water (200 mL). For the
pH 4.4 solution, NaOAc (1.605 g) was dissolved in water (200 mL). Acetic acid (1.2 mL) was
added to reach the desired pH. For the pH 7.1 solution, NaH2PO4 (0.468 g) and Na2HPO4 (1.637
g) were dissolved in water (100 mL). For the pH 10.0 solution, NaHCO3 (0.339 g) and Na2CO3
(0.632 g) were dissolved in water (100 mL). The pHs of the solutions were measured with a
benchtop Mettler Toledo pH meter. All solutions were degassed and purged with N2.
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46
CHAPTER 3
Metal-Dithiolene (M = Ni, Fe, Zn) Coordination Polymers
for Electrocatalytic H
2
Evolution
A portion of this chapter has appeared in print:
Downes, C. A.; Marinescu, S. C. “One Dimensional Metal Dithiolene (M = Ni, Fe, Zn)
Coordination Polymers for Hydrogen Evolution Reaction.” Dalton Trans. 2016, 45, 19311-
19321.
47
3.1. Introduction
Energy harvested directly from sunlight offers a desirable approach to fulfilling, with minimal
environmental impact, the global need for clean energy.
1-3
Hydrogen produced through the
reduction of water is an attractive candidate as a clean and renewable energy carrier.
1-3
Precious
metals, such as Pt, efficiently catalyze hydrogen evolution from water with high activities at low
overpotentials. However, their paucity and high cost limit widespread use. Replacement of Pt with
non-precious metal catalysts would be desirable for practical applications and global scalability of
such potential clean-energy technologies. These considerations have led to the development of
catalytic systems for the HER that employ a variety of base metal homogeneous
4-10
and
heterogeneous
11-14
hydrogen evolving catalysts.
Immobilization of metal complexes to electrode surfaces
7,8,15
bridges the gap between
homogeneous and heterogeneous catalysis by maintaining the favorable properties of molecular
systems, such as tunability, while taking advantage of the robustness and efficiency present in
heterogeneous catalysts. The primary figure of merit for the HER is the overpotential required to
achieve a current density of 10 mA/cm
2
, which is the value expected for an integrated solar water-
splitting device under 1 sun illumination operating at 10% solar-to-fuel efficiency.
16-18
Recently,
several molecular catalysts have been immobilized onto carbon-based supports and display
moderate activity for the HER.
19-31
Catalysts based on nickel phosphine and cobalt dithiolene
complexes, immobilized through covalent and noncovalent methods, were the only catalytic
systems reported to achieve current densities higher than the benchmarking metric of 10 mA/cm
2
for the HER.
31-33
Current densities of 10 mA/cm
2
were achieved at an overpotential of only 150
mV in acidic aqueous media at room temperature for the immobilized nickel phosphine
complex.
32,33
Additionally, remarkable activities, similar to that of platinum, were observed at
elevated temperatures.
33
However, despite these promising results, the immobilization of
molecular electrocatalysts for the HER to carbon-based supports remains a great challenge due to
the limited successful expansion of the reported methods to other molecular systems and the low
surface concentration of catalyst achieved.
We have reported recently that metal dithiolenes incorporated into an extended framework act as
efficient electrocatalysts for the HER.
34
We have shown that ligand precursors such as benzene-
48
hexathiol or triphenylene-2,3,6,7,10,11-hexathiol, upon treatment with cobalt(II) and a weak base,
generate cobalt dithiolene units integrated into a 2D metal-organic framework (Scheme 3.1).
34
The
frameworks were deposited on glassy carbon electrodes (GCE) to generate electrocatalytic
materials with remarkable activity and stability for hydrogen evolution from water.
34
The nickel
analogue based on triphenylene-2,3,6,7,10,11-hexathiolate was also reported to be an efficient
electrocatalyst for the HER as a single-layer sheet, with current densities above 10 mA/cm
2
under
strongly acidic aqueous conditions (0.5 M H2SO4) at low overpotentials.
35
However, no prolonged
electrolysis studies were reported, making it hard to assess the stability of this material.
35
We have
extended this methodology to the formation of cobalt dithiolene systems based on benzene-1,2,4,5-
tetrathiolate.
36
This system was integrated with planar p-type Si to generate modified
photocathodes for efficient solar-driven hydrogen evolution from water.
36
Analogous nickel
coordination polymers based on BTT
37,38
and benzenehexathiolate
39
have been reported, but the
catalytic properties were not studied. We investigate here the catalytic properties of nickel, iron,
and zinc dithiolene coordination polymers based on benzene-1,2,4,5-tetrathiolate.
Scheme 3.1. Structures of the reported electrocatalytic cobalt dithiolene coordination polymers
(CoBHT, CoTHT, and CoBTT).
49
3.2. Results
The syntheses of nickel (NiBTT), iron (FeBTT), and zinc (ZnBTT) coordination polymers were
accomplished through liquid-liquid interfacial reactions similar to the reported procedure for the
generation of the cobalt benzene-1,2,4,5-tetrathiolate system (Scheme 3.2).
36
An ethyl acetate
solution of benzene-1,2,4,5-tetrathiol was gently layered on an aqueous solution of metal(II)
acetate and sodium acetate. Following evaporation of the organic layer, a solid film remains at the
gas-liquid interface which can either be collected as a powder or deposited on supports of interest,
such as GCE.
Scheme 3.2. Synthesis of benzene-1,2,4,5-tetrathiolate-based nickel (NiBTT), iron (FeBTT), and
zinc (ZnBTT) coordination polymers through a liquid-liquid interfacial reaction.
Top-down scanning electron microscopy images of NiBTT, FeBTT, and ZnBTT, are illustrated
in Figure 3.1.
40
FTIR of the generated polymers revealed the disappearance of the strong S–H
stretch at 2500 cm
–1
corresponding to the starting material benzene-1,2,4,5-tetrathiol (Figure 3.2).
Raman spectra of the electrochemically active polymers, NiBTT and FeBTT, prepared as films in
a similar manner as the electrodes used for electrochemical studies, showed a small feature
associated with the S-H stretch (Figure 3.3). Extensive washing of the polymers suggest the S-H
stretch is not from residual starting material, but possibly from S-H termination of the polymer
fragments.
Figure 3.1. Scanning electron microscopy images of (a) NiBTT, (b) FeBTT, and (c) ZnBTT.
50
Figure 3.2. FTIR spectra of NiBTT (pink), FeBTT (blue), ZnBTT (purple), and benzene-1,2,4,5-
tetrathiol (BTT) (cyan).
Figure 3.3. Raman spectroscopy of NiBTT (pink), FeBTT (blue), and benzene-1,2,4,5-tetrathiol
(BTT, cyan).
X-ray photoelectron spectroscopy (XPS) of NiBTT (Figure 3.4), FeBTT (Figure 3.5), and ZnBTT
(Figure 3.7) revealed the presence of Ni, Fe, and Zn, respectively. Additionally, all three polymers
displayed features indicative of Na and S as expected. Two sets of peaks were observed in the
nickel region of NiBTT attributed to one environment with binding energies of 853.5 and 870.7
eV corresponding to the 2p3/2 and 2p1/2 levels. Small satellites were detected at higher binding
energies relative to the major 2p3/2 and 2p1/2 peaks. Three additional features of NiBTT at 1071.4,
227.3, and 163.3 eV arise from the presence of Na 1s, S 2s, and S 2p. The S 2s region also displayed
a broad shake-up feature at ~232 eV which is often observed in metal bis(dithiolene) complexes.
51
These binding energies and the deconvoluted peaks are similar to what was observed in the XPS
spectra of the cobalt analogue.
36
Figure 3.4. XPS analysis of NiBTT displaying the Ni 2p, Na 1s, S 2s, and S 2p core level XPS
spectra.
In the iron region of FeBTT, two peaks with binding energies of 707.7 and 720.7 eV correspond
to the 2p3/2 and 2p1/2 levels. Deconvolution of the 2p3/2 region revealed two major peaks at 707.9
and 709.6 eV. The peak at 707.9 eV is attributed to Fe
II
, as similar binding energies of 707 eV
were observed for the molecular iron bis(benzenedithiolate) complex (Figure 3.6).
41,42
The major
peak at 709.6 eV may be a result of surface oxidation to Fe(III)-S species with a 2:1 ratio of Fe(II)
to Fe(III).
41,42
Three additional peaks at 1071.4, 226.7, and 162.5 eV were observed for FeBTT
corresponding to Na 1s, S 2s, and S 2p, respectively. A broad shake-up peak was observed in the
S 2s spectra around 228 eV, which is similar to the peak observed in the S 2s spectra of the cobalt
36
and nickel analogues.
52
Figure 3.5. XPS analysis of FeBTT displaying the Fe 2p, Na 1s, S 2s, and S 2p core level XPS
spectra.
Figure 3.6. XPS of the Fe 2p core level of [Fe(bdt)2]
2–
(where bdt = 1,2-benzenedithiolate).
Two peaks were observed in the zinc region of ZnBTT at binding energies of 1021.1 and 1044.1
eV corresponding to the 2p3/2 and 2p1/2 levels with a large spin orbit coupling of 23 eV. These
features are indicative of a Zn(II) oxidation state. Three additional broad features were seen at
1071.3, 226.5, and 162.5 eV, which represent Na 1s, S 2s, and S 2p.
53
Figure 3.7. XPS analysis of ZnBTT displaying Zn 2p, Na 1s, S 2s, and S 2p core level XPS
spectra.
The electrochemical properties of the nickel, iron, and zinc coordination polymers were studied
by cyclic voltammetry on GCE. In pH 10.0 aqueous solutions, a broad electrochemical wave was
observed between 0.2 and –0.1 V versus SHE for NiBTT and FeBTT, corresponding to the –1/–
2 redox couple of the metal bis(dithiolene) unit (Figure 3.8a, 3.9). The previously reported nickel
polymers supported by benzenehexathiolate
39
and benzene-1,2,4,5-tetrathiolate
38
ligands, display
broad redox waves at –0.05, and 0.20 V versus Fc
+/0
, respectively, in 0.1 M [nBu4N][ClO4]
dichloromethane solution. These positive redox waves (0.6 V versus SHE for nickel
benzenehexathiolate
39
and 0.84 V versus SHE for nickel benzene-1,2,4,5-tetrathiolate
38
) were
assigned as the 0/–1 couple. The wave derived from the –1/–2 couple was not observed in the
available potential window.
38,39
The molecular analogue, [Ni(bdt)2]
–
(where bdt = 1,2-
benzenedithiolate), displays a reversible wave with E1/2 at –0.94 V vs Fc
+/0
(–0.30 V versus SHE)
in 0.2 M [nBu4N][PF6] acetonitrile solution, and assigned to [Ni(bdt)2]
–1/–2
.
43
Therefore, these
previous studies inform our assignment of the broad redox waves between 0.2 and –0.1 V versus
SHE for NiBTT and FeBTT to the –1/–2 couple.
54
Figure 3.8. Scan rate dependence studies for NiBTT (4.2(4) × 10
–7
molNi/cm
2
) in 0.1 M NaClO4
aqueous solutions at pH 10.0. (a) Scan rates varied from 5 to 500 mV/s; (b) Cathodic peak current
density of NiBTT measured at 0.05 V versus SHE as a function of scan rate.
Figure 3.9. Polarization curves of FeBTT at differing catalyst loadings (8.3(8) × 10
–7
mol Fe/cm
2
(red) and 3.4(3) × 10
–7
mol Fe/cm
2
(blue)) in 0.1 M NaClO4 aqueous solutions at pH 10.0.
Relatedly, cobalt dithiolene coordination polymers based on benzenehexathiolate,
34
triphenylene-
2,3,6,7,10,11-hexathiolate,
34
and BTT,
36
and immobilized [Co(Cl2bdt)2]
–
(where Cl2bdt = 3,6-
dichloro-1,2-benzenedithiolate)
22,31
display broad redox waves at –0.30, ~ –0.3, ~ –0.1 V, and –
0.46 V versus SHE, respectively, in 0.1 M NaClO4 aqueous solutions at pH 10.0. These redox
waves were assigned to the –1/–2 couple. The values of the redox couples for the cobalt
coordination polymers were similar to the one reported for the molecular analogue, [Co(bdt)2]
–
,
which was observed at –0.4 V versus SHE in 0.1 M KNO3 1:1 H2O/CH3CN solutions, and assigned
to [Co(bdt)2]
–1/–2
.
44,45
55
The cathodic peak current density of NiBTT measured at 0.05 V versus SHE in pH 10.0 solutions
was directly proportional to the scan rate (Figure 3.8b), indicative of rapid electron transfer
between the glassy carbon electrode and NiBTT. The cathodic peak current density of NiBTT was
not directly proportional to the square root of the scan rate, as expected for a freely diffusing
species in solution obeying the Randles-Sevcik equation,
46
suggesting that the observed
electrochemical behaviour is consistent with a surface confined process. Integration of the
electrochemical wave at pH 10.0 allowed for the estimation of catalyst loadings for the systems
that are electrochemically active (NiBTT and FeBTT). These values were confirmed through
digestion of the deposited polymer in acid and subsequent ICP-OES measurements. The
correlation of the electrochemically (integrated) and ICP measured catalyst loadings indicate that
most of the metal centers were electrochemically active. The maximum catalyst loading achieved
for NiBTT was 6.8(7) × 10
–7
molNi/cm
2
, which is similar in magnitude to the surface concentration
observed for the cobalt analogue (5.5(6) × 10
–7
molCo/cm
2
).
36
The maximum catalyst loading
achieved for FeBTT was 8.3(8) × 10
–7
mol Fe/cm
2
(Figure 3.9). The change in intensity of the
electrochemical wave at pH 10.0 as a function of catalyst loading is clearly demonstrated for
FeBTT where the peak intensity was greater for the maximum achievable catalyst loading of
8.3(8) × 10
–7
mol Fe/cm
2
and decreased when the surface concentration was low (3.4(3) × 10
–7
mol Fe/cm
2
) (Figure 3.9).
As the pH of the aqueous solution was lowered, an increase in the catalytic current density was
observed for NiBTT (Figure 3.10a). This current increase corresponds to hydrogen evolution from
water, as verified by controlled potential electrolysis, further discussed below. Faster scan rates
were necessary in static pH 1.3 solutions to limit H2 bubble formation on the electrode surface due
to high electrocatalytic activity. Unmodified GCE displayed an insignificant increase in current.
In pH 1.3 solutions, an overpotential of 0.47 V was required to reach a current density of 10
mA/cm
2
at a surface concentration of 6.8(7) × 10
–7
molNi/cm
2
. In comparison, the cobalt analogue
required an overpotential of 0.56 V to reach a current density of 10 mA/cm
2
at a surface
concentration of 5.5(6) × 10
–7
molCo/cm
2
.
36
The observed current densities vary proportionally
with the surface concentration (Figure 3.10b-3.11). The current densities of NiBTT in pH 1.3
solutions measured at –0.80 V versus SHE increased linearly with catalyst loading (Figure 3.10b).
Tafel analyses of NiBTT in pH 2.6 aqueous solutions at 0.5 mV/s results in a Tafel slope of 76
56
mV/ dec (Figure 3.12). The measured exchange current density for NiBTT was 10
–8.1
A/cm
2
,
which is similar to the exchange current density of 10
–8.3
A/cm
2
reported for the cobalt analogue.
36
Figure 3.10. (a) Polarization curves of NiBTT (5.0(5) × 10
–7
molNi/cm
2
) in 0.1 M NaClO4 aqueous
solutions at pH 10.0 (orange), pH 7.0 (purple), 4.4 (green), 2.6 (blue), 1.3 (red), and of blank GCE
at pH 1.3 (dashed black); scan rate for pH 10.0–2.6: 20 mV/s; scan rate for pH 1.3: 100 mV/s. (b)
Current densities of NiBTT measured at –0.80 V versus SHE at pH 1.3 as a function of the surface
catalyst concentration. The surface concentration was quantified by integrating the peak area at
pH 10.0.
Figure 3.11. Polarization curves of NiBTT at catalyst loadings of 6.8(7) × 10
–7
mol Ni/cm
2
(purple),
5.2(5) × 10
–7
molNi/cm
2
(green), 5.0(5) × 10
–7
molNi/cm
2
(blue) and GCE (dashed black) in 0.1 M
NaClO4 aqueous solutions at pH 1.3.
57
Figure 3.12. Tafel plot of NiBTT (catalyst loading: 5.6(6) × 10
–7
molNi/cm
2
) in 0.1 M NaClO4
aqueous solutions at pH 2.6. Scan rate: 0.5 mV/s; Tafel slope of 76 mV/dec; exchange current
density of 10
-8.1
A/cm
2
.
FeBTT was investigated in pH 1.3 solution and displayed low activity toward the HER (Figure
3.13a, 3.14). The current densities of FeBTT measured at –0.80 V versus SHE increased linearly
with catalyst loading (Figure 3.13b). At the maximum catalyst loading achieved, 8.3(8) × 10
–7
mol Fe/cm
2
, a current density of only 4.8 mA/cm
2
was measured for FeBTT at –0.80 V versus SHE.
In comparison, the cobalt analogue achieved a current density of 23.3 mA/cm
2
at a catalyst loading
of 5.5(6) × 10
–7
molCo/cm
2
,
36
while NiBTT achieved 40.9 mA/cm
2
at a catalyst loading of 6.8(7)
× 10
–7
molNi/cm
2
at –0.80 V versus SHE. Therefore, the coordination polymer FeBTT exhibits
low activity toward the HER in comparison to the activity of the cobalt or nickel analogues (Figure
3.14). Because of this low activity, FeBTT was not subjected to further electrochemical studies.
Figure 3.13. (a) Polarization curves of FeBTT at differing catalyst loadings (8.3(8) 10
–7
mol Fe/cm
2
(red) and 3.4(3) 10
–7
mol Fe/cm
2
(blue)) in 0.1 M NaClO4 aqueous solutions at pH 1.3;
Scan rate: 20 mv/s. (b) Current densities of FeBTT measured at –0.80 V versus SHE at pH 1.3 as
a function of the surface catalyst concentration. The surface concentration was quantified by
integrating the peak area at pH 10.0.
´
´
58
Figure 3.14. Polarization curves of benzene-1,2,4,5-tetrathiolate-based coordination polymers in
0.1 M NaClO4 aqueous solutions at pH 1.3; Conditions: CoBTT (5.5(6) × 10
–7
molCo/cm
2
; teal);
NiBTT (6.8(7) × 10
–7
molNi/cm
2
; pink); FeBTT (8.3(8) × 10
–7
mol Fe/cm
2
; blue); ZnBTT (purple);
and blank GCE (dashed black).
ZnBTT was investigated by electrochemistry to discern any ligand-based features. The cyclic
voltammetry experiments in pH 10.0 aqueous solution showed no observable redox events
between potentials of 0.2 and –0.8 V versus SHE (Figure 3.15a). Additionally, ZnBTT displayed
negligible activity for the HER in pH 1.3 aqueous solutions (Figures 3.15b).
Figure 3.15. Polarization curves of ZnBTT (red) and GCE (black dashed) in 0.1 M NaClO4
aqueous solution at (a) pH 10.0 and (b) pH 1.3; Scan rate: 20 mV/s.
Controlled potential electrolysis (CPE) of NiBTT on GCE in pH 1.3 aqueous solutions at –0.80 V
versus SHE consumed 76 coulombs of charge after 1 h (Figure 3.16a). A constant current density
of approximately 3 mA/cm
2
was observed during these CPE studies (Figure 3.16a). Analysis of
the gas mixture in the headspace of the working compartment of the electrolysis cell confirmed
production of H2 with a Faradaic efficiency of 85 ± 5%. The durability of NiBTT was further
59
tested in longer-duration CPE studies (Figures 3.16b, 3.17). NiBTT affords in both pH 1.3 (Figure
3.16b) and pH 2.6 (Figure 3.17) a continuous increase in charge build-up over a 6 h CPE period.
The decrease in current was attributed to the delamination of NiBTT from the electrode surface,
due in part to both stirring and the evolution of H 2. The decrease in catalyst loading measured at
the end of the experiment was directly proportional to the decrease in current observed for NiBTT.
Figure 3.16. Controlled potential electrolysis of NiBTT (1.2(1) × 10
–7
molNi/cm
2
; red) and blank
GCE (black) at –0.80 V versus SHE in 0.1 M NaClO4 aqueous solutions at pH 1.3.
Figure 3.17. (a) CPE studies of NiBTT (1.4(1) × 10
–7
mol Ni/cm
2
) at –0.80 V versus SHE in 0.1 M
NaClO4 aqueous solutions at pH 2.6 (red) and of blank GCE (black). (b) Polarization curves of
NiBTT (1.4(1) × 10
–7
molNi/cm
2
) in 0.1 M NaClO4 aqueous solutions at pH 2.6 before CPE
experiment (red), after 1 h CPE (blue), and after 6 h CPE (green).
The amount of soluble nickel present in solution as determined by ICP-OES studies was negligible,
demonstrating that the modified electrode NiBTT does not generate any soluble materials during
electrocatalysis. Negligible current densities were detected in the cyclic voltammetry experiments
of the solutions resulting from the electrochemical studies of NiBTT. Therefore, the H2-evolving
activity investigated is not due to the formation of catalytically active soluble materials.
60
Additionally, the onset of catalysis for NiBTT was similar before and after 1 or 6 h CPE studies
(Figure 3.17b), indicating that the active form of the catalyst did not alter during electrocatalysis.
XPS analyses of NiBTT after 1 h (Figure 3.18) and 6 h (Figure 3.19) CPE studies displayed
features for Ni (835.5 and 870.6 eV), Na (1071.5), and S (227.4 and 163.4 eV) at the same binding
energies to the ones observed before electrocatalysis (Figure 3.4).
Figure 3.18. XPS analysis of NiBTT after 1 h CPE studies in pH 1.3 aqueous solutions.
Figure 3.19. XPS analysis of NiBTT after 6 h CPE studies in pH 1.3 aqueous solutions.
61
3.3. Discussion
Metal dithiolene molecular complexes have been extensively studied due to their highly unusual
physical properties and rich redox chemistry.
47,48
Both the metal and the ligands can exist in several
oxidation states and the redox non-innocent character of dithiolene ligands allows for reversible
acceptance or release of electrons (Scheme 3.3).
Scheme 3.3. Oxidation states of the dithiolene ligand.
Comprehensive study of metal dithiolenes found that cobalt, nickel and iron complexes display
most often square planar geometries,
47,48
whereas zinc complexes display tetrahedral
geometries.
49-51
Cobalt dithiolene species are among the most efficient molecular catalysts for the
hydrogen evolution reaction (HER).
44,45,52-54
A higher activity for the HER was observed for the
cobalt dithiolene complex, [Co(bdt)2]
–
(where bdt = benzene-1,2-dithiolate), in a 1:1 mixture of
water and acetonitrile, relative to the activity in a pure anhydrous organic solvent (100%
acetonitrile).
44
Related nickel dithiolene species display reduced activity for the HER
55-57
and/or
ill-defined chemistry in organic solvents.
43,44,58-61
Six-coordinate trigonal prismatic molybdenum
bis(benzenedithiolate) complexes were also shown to reduce protons to H2 photocatalytically and
electrocatalytically.
62
The proposed mechanism of hydrogen evolution for metal dithiolenes
involves protonation of the sulfur moiety on the dithiolene ligand, which is possible because of the
redox non-innocent character of the ligand. Protonation of the ligand framework may eventually
lead to ligand loss and decomposition, especially in organic media. The high stability of the
coordination polymers observed here is ascribed to the network environment and the absence of
an organic solvent.
The catalytic performance of several coordination polymers previously reported as well as the ones
developed here, are summarized in Table 3.1. We have reported that metal dithiolenes incorporated
into extended frameworks act as efficient electrocatalysts for the HER.
34,36
Studies in our
laboratory have shown that cobalt dithiolene coordination polymers based on benzenehexathiolate
ligands display remarkable H2-evolving activity and stability.
34
Overpotentials as low as 340 mV
were required to reach current densities of 10 mA/cm
2
at a pH of 1.3. The Tafel slope at pH 2.6
62
was 108 mV/decade, suggesting that the initial discharge step is rate-determining.
7,63
The exchange
current density (i0) was 10
–5.3
A/cm
2
,
34
which is superior to those reported for surface-confined
molecular catalysts discussed below. The reduction potentials of the cobalt dithiolene coordination
polymers based on triphenylene-2,3,6,7,10,11-hexathiolate and benzene-1,2,4,5-tetrathiolate are
similar, whereas the reduction potential of the benzenehexathiolate system is approximately 0.2 V
more positive, as expected for metal complexes with additional thiolate moieties.
64
These results
indicate that the developed coordination polymers are indeed tunable with distinct active sites, an
attractive feature of molecular systems. Because the catalytic moiety is easily tunable, these
materials have the same advantages as homogeneous, molecular systems with respect to the further
improvement of catalytic properties, while maintaining coordination in the robust extended
framework environment. Following our report, Feng and coworkers developed a nickel analogue
based on triphenylene-2,3,6,7,10,11-hexathiolate framework, which is also an efficient
electrocatalyst for the HER as a single-layer sheet.
35
However, no controlled-potential electrolysis
studies were reported,
35
making it hard to compare the stability of this material relative to the other
coordination polymers developed in our laboratory.
Table 3.1. Electrocatalytic HER properties of metal dithiolene coordination polymers.
M Ligand
a
Catalyst Loading
× 10
-7
mol M/cm
2
η
10 mA/cm
2 (mV)
b
Tafel slope
c
(mV/dec)
i 0
(A/cm
2
)
Ref.
Co BHT 7.0 340 108 10
–5.3
34
Co THT 11.0 530 161 10
–5.4
34
Ni THT — 413 80.5
d
6 ×10
–7
35
Co BTT 5.5 560 94
10
–9.4
36
Ni BTT 6.8 470 76 10
–8.1
this work
a
Where BTT = benzene-1,2,4,5-tetrathiolate, BHT = benzenhexathiolate, and THT = triphenylene-2,3,6,7,10,11-hexathiolate.
b
Overpotentials measured in pH 1.3
c
Tafel slopes and corresponding i0 recorded in pH 2.6.
d
Tafel slope and corresponding i0
recorded in 0.5 M H2SO4.
The nickel dithiolene coordination polymer based on the benzene-1,2,4,5-tetrathiolate ligand
(NiBTT) investigated here displayed an overpotential for the HER that is about 90 mV lower than
the overpotential of the cobalt analogue (CoBTT) as well as about 60 mV lower than the
overpotential of the cobalt triphenylene-2,3,6,7,10,11-hexathiolate system. The 76 mV/dec Tafel
slope of NiBTT is smaller than the 94 mV/dec reported for the cobalt analogue (CoBTT).
36
The
smaller Tafel slope suggests that the energy barrier of the discharge step (Volmer reaction) is
reduced for the nickel polymer. The larger Tafel slope of 94 mV/dec for the cobalt analogue
63
indicates the rate-determining step in the HER mechanism is the Volmer reaction, the discharge
step that converts protons into absorbed hydrogen on the catalyst surface.
65
This change in Tafel
slope for the cobalt and nickel analogues could be attributed to the differences in mechanisms
proposed for the corresponding molecular metal bis(dithiolene) complexes.
57
For nickel
bis(aryldithiolene) complexes, it is proposed that the protonation and redox events occur primarily
on the ligand because the frontier orbitals are mainly ligand based.
57
For the cobalt bis(dithiolene)
complexes, the cobalt orbitals contribute more significantly to the frontier orbitals so protonation
and redox events occur on both the ligand and the metal center.
54,57
Theoretical studies, performed
on the molecular cobalt dithiolene complex [Co(bdt)2]
–
, suggest that upon one electron-reduction
of [Co(bdt)2]
–
, two protonation events occur on different sulfur donors, followed by a second one
electron-reduction.
66
Prior to the release of hydrogen, an intramolecular proton shift takes place,
leading to the formation of a cobalt hydride with an adjacent protonated sulfur.
66
No stable species
protonated at both sulfur and nickel have been identified for nickel dithiolenes.
57
Additionally, minimal decrease in activity was observed for NiBTT during the first hour of the
CPE studies. The Faradaic efficiency of NiBTT was 85%, reduced from the 97% previously
reported for CoBTT.
36
Nickel bis(aryldithiolene) complexes display Faradaic efficiencies between
66% and 83%,
57
which are lower than those of analogous cobalt bis(aryldithiolene) complexes (90
± 10%).
54
Additionally, [Ni(bdt)2]
–
(where bdt = 1,2, benzenedithiolate) was reported to
decompose in organic acidic solutions under reducing conditions.
43
Using available spectroscopic
techniques, we do not detect catalyst decomposition, but the reduced efficiency could still be a
result of catalyst deactivation or incomplete turnover of the catalyst.
43
Additionally, two
mechanistic pathways for the HER have been proposed for nickel dithiolenes, ECEC and ECCE
(E = electron transfer and C = chemical reaction) with protonation occurring solely on the sulfur
moiety of the dithiolene ligand.
57
Calculations have suggested that geometric distortions occur to
accommodate the protonated ligand.
57
Because the catalysts must undergo distortion for H2 to
evolve, the H2 evolving activity may be impeded resulting in lower efficiencies. If this geometric
distortion is necessary, the incorporation of nickel dithiolene units in coordination polymers would
inhibit the degree of distortion thus reducing the H2 evolving efficiency.
64
A selection of surface-confined molecular catalysts and their reported catalytic activity for the
HER are illustrated in Scheme 3.4-3.5. Cobaloxime catalysts have been grafted onto multiwalled
carbon nanotubes through aryldiazonium reductive coupling followed by functionalization,
20,24
or
were incorporated into pyrene-bearing polymers.
28
However, low H2 evolving activities were
reported for these surface-confined cobaloxime catalysts. Nickel phosphine complexes with
pendant proton relays have also been grafted onto multiwalled carbon nanotubes through
aryldiazonium reductive and/or amide coupling, generating modified surfaces with remarkable
activity and stability under strongly acidic conditions (0.5 M H2SO4).
19,33
These systems displayed
current densities up to 10 mA/cm
2
at low potentials (–0.15 V versus SHE).
33
Additionally,
remarkable activities, similar to that of platinum, were observed at high temperatures for the
immobilized nickel phosphine complex.
33
A similar nickel phosphine complex with activated
esters was covalently attached to a glassy carbon surface terminated by organolithium anchor
points. This system exhibited low activity and decomposition of the surface-confined species was
observed in acidic acetonitrile.
21
Moreover, noncovalent modifications of the multiwalled carbon nanotubes with pyrene-
functionalized nickel phosphine complexes generated modified surfaces stable under strong acidic
conditions (0.5 M H2SO4), which displayed current densities up to 16 mA/cm
2
at –0.3 V versus
NHE.
32
Physisorption of metal complexes onto carbon based supports has not shown as much
success as covalent attachment because of low surface concentrations and activities. Cobalt
porphyrin complexes adsorbed onto carbon based supports
29
or incorporated into a nafion
membrane
27
displayed catalytic hydrogen evolution, but suffered from relatively low stability of
the attachment and/or catalytic performance. Cobalt phthalocyanine derivatives were incorporated
in poly(4-vinylpyridine-co-styrene) films reaching current densities of only 3.4 mA/cm
2
at –1.0 V
versus Ag/AgCl.
26,30
The molecular cobalt dithiolene complex was recently heterogenized via bulk
graphite adsorption to generate modified surfaces that achieve current densities of 10 mA/cm
2
at
an overpotential of 400 mV in pH 1.3 aqueous solutions showing remarkable improvements in the
viability of adsorption as a method of immobilizing molecular catalysts to carbon based
supports.
22,31
65
Scheme 3.4. A selection of surface-confined cobalt molecular catalysts and their reported catalytic
H2 evolving activity. Parameters are listed as follows: (a) support; (b) media; (c) current density
(mA/cm
2
) at a certain potential; (d) catalyst loading (mol/cm
2
); (e) reference. The following
abbreviations were used: MWCNT = multiwalled carbon nanotubes; BPG = basal-plane pyrolytic
graphite; GCE = glassy carbon electrode.
66
Scheme 3.5. A selection of surface-confined nickel molecular catalysts and their reported catalytic
H2 evolving activity. Parameters are listed as follows: (a) support; (b) media; (c) current density
(mA/cm
2
) at a certain potential; (d) catalyst loading (mol/cm
2
); (e) reference. The following
abbreviations were used: MWCNT = multiwalled carbon nanotubes; NHS = N-
hydroxysuccinimide ester; GCE = glassy carbon electrode.
Overall, only a couple of surface-confined molecular catalysts report the benchmarking metric of
10 mA/cm
2
of activity and/or prolonged catalysis in acidic aqueous media.
31-33
Significant success
in reaching high activity at low potentials in aqueous acidic media (up to 10 mA/cm
2
at –0.15 V
versus SHE) has been achieved utilizing nickel phosphine catalysts through both covalent and
noncovalent methods.
32,33
Studies in our laboratory have now shown that immobilization via
67
coordination polymers is another viable methodology to provide enhanced activity and stability in
comparison to solution-based molecular catalysts with systems reaching the desired 10 mA/cm
2
of
activity in acidic aqueous media at low overpotentials.
34,36
3.4. Conclusions
The nickel, iron, and zinc dithiolene coordination polymers reported here were synthesized
through liquid-liquid interfacial reactions between aqueous solutions of metal(II) acetate and
sodium acetate and organic solutions of benzene-1,2,4,5-tetrathiol. The generated metal dithiolene
coordination polymers were characterized by a variety of techniques. XPS analyses of these
materials revealed the presence of the expected elements: S, Na, and the corresponding metal (M
= Ni, Fe, or Zn). The electrochemical studies of the developed polymers were investigated under
fully aqueous conditions. Negligible activities were observed for the iron and zinc dithiolene
coordination polymers. However, the nickel dithiolene coordination polymer was an
electrocatalyst for the conversion of water into hydrogen at moderate overpotentials.
Overpotentials of 470 mV were required to reach current densities of 10 mA/cm
2
in pH 1.3 aqueous
solutions. In comparison, the cobalt analogue requires slightly higher overpotentials (560 mV) to
reach the same activity for the HER. In addition, the nickel system displayed no decrease in current
during the first hour of electrolysis. Replacement of the cobalt center with nickel results in a
reduction in the overpotential for H2 evolution, albeit with lower efficiency, displaying the ease in
which the catalytic unit can be tuned to modulate electrocatalytic performance because of the
retention of the molecular nature of the catalysts following integration into the coordination
polymer environment.
3.5. Experimental Details
3.5.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity). Ethyl acetate and water were placed under vacuum and refilled with nitrogen (10 ×).
Benzene-1,2,4,5-tetrathiol was prepared according to the reported procedure.
67
All other chemical
68
reagents were purchased from commercial vendors and used without further purification. The pHs
of the aqueous solutions were measured with a benchtop Mettler Toledo pH meter.
Caution: Perchlorate salts are potential explosive chemicals and should only be used in very small
amounts, especially in the presence of the reductant H2. Please handle these mixtures using the
proper equipment and protection.
3.5.2. Synthesis of NiBTT
A nickel polymer based on benzene-1,2,4,5-tetrathiolate has been previously reported,
37,38
however, the procedure detailed below has been modified and follows the synthetic method our
laboratory developed for the cobalt benzene-1,2,4,5-tetrathiolate system.
36
In a nitrogen filled
glove-box, Ni(OAc)2(H2O)4 (36.3 mg, 0.146 mmol) and anhydrous NaOAc (0.2 mL, 0.5 M
aqueous solution) were dissolved in 20 mL of degassed H2O. The aqueous solution was transferred
to a 70 mm 50 mm crystallizing dish. An ethyl acetate solution of benzene-1,2,4,5-tetrathiol
(C6H6S4) (1 mL, 7.3 mM) was carefully layered via glass pipette on top of the aqueous solution of
nickel(II) acetate and sodium acetate to cover ~80% of the aqueous solution. The organic solvents
were allowed to evaporate over one to two hours at room temperature, leaving behind NiBTT as
a black solid at the gas-liquid interface. The black solid was collected, washed with water,
methanol, and ethyl acetate. The resultant black powder was dried under vacuum. Anal. Calcd for
[Ni(C6H2S4)][H]•H2O (C6H5OS4Ni): C, 25.73; H, 1.80. Found: C, 24.33; H, 1.64.
3.5.3. Synthesis of FeBTT
An iron polymer based on benzene-1,2,4,5-tetrathiolate has been previously reported,
37
however,
the procedure detailed below has been modified and follows the synthetic method our laboratory
developed for the cobalt benzene-1,2,4,5-tetrathiolate system.
36
In a nitrogen filled glove-box,
Fe(OAc)2 (25.4 mg, 0.146 mmol) and anhydrous NaOAc (0.2 mL, 0.5 M aqueous solution) were
dissolved in 20 mL of degassed H2O. The aqueous solution was transferred to a 70 mm 50 mm
crystallizing dish. An ethyl acetate solution of benzene-1,2,4,5-tetrathiol (C6H6S4) (1 mL, 7.3 mM)
was carefully layered via glass pipette on top of the aqueous solution of iron(II) acetate and sodium
acetate to cover ~80% of the aqueous solution. The organic solvents were allowed to evaporate
over one to two hours at room temperature, leaving behind FeBTT as a black solid at the gas-
69
liquid interface. The black solid was collected, washed with water, methanol, and ethyl acetate.
The resultant black powder was dried under vacuum. Anal. Calcd for [Fe(C6H2S4)][H]•H2O
(C6H5OS4Fe): C, 26.00; H, 1.82. Found: C, 25.55; H, 1.43.
3.5.4. Synthesis of ZnBTT
In a nitrogen filled glove-box, Zn(OAc)2(H2O)2 (32.05 mg, 0.146 mmol) and anhydrous NaOAc
(0.2 mL, 0.5 M aqueous solution) were dissolved in 20 mL of degassed H2O. The aqueous solution
was transferred to a 70 mm 50 mm crystallizing dish. An ethyl acetate solution of benzene-
1,2,4,5-tetrathiol (C6H6S4) (1 mL, 7.3 mM) was carefully layered via glass pipette on top of the
aqueous solution of zinc(II) acetate and sodium acetate to cover ~80% of the aqueous solution.
The organic solvents were allowed to evaporate over one to two hours at room temperature, leaving
behind ZnBTT as a white solid at the gas-liquid interface. The white solid was collected, washed
with water, methanol, and ethyl acetate. The resultant white powder was dried under vacuum.
Anal. Calcd for [Zn(C6H2S4)][H]•H2O (C6H5OS4Zn): C, 25.13; H, 1.76. Found: C, 24.03; H, 1.78.
3.5.5. Immobilization of Coordination Polymers on GCE
The coordination polymers were deposited onto a GCE by immersing a polished electrode face
down into the reaction mixture. The volatiles were allowed to evaporate at room temperature. The
modified electrode was washed with water, methanol, and ethyl acetate.
3.5.6. Electrochemical Methods
Electrochemistry experiments were carried out using a Pine potentiostat. Platinum wire used for
the electrochemical studies was purchased from Alfa Aesar. The electrochemical experiments
were carried out in a three-electrode configuration electrochemical cell under an inert atmosphere
using GCE as the working electrode. A platinum wire, placed in a separate compartment,
connected by a Vycor tip, and filled with the electrolytic solution (0.1 M NaClO4) was used as the
auxiliary (counter) electrode. The reference electrode, placed in a separate compartment and
connected by a Vycor tip, was based on an aqueous Ag/AgCl/saturated KCl electrode. The
reference electrode in aqueous media was calibrated externally relative to ferrocenecarboxylic acid
(Fc-COOH) at pH 7.0, with the Fe
3+/2+
couple at 0.28 V vs Ag/AgCl. All potentials reported were
70
converted to the standard hydrogen electrode (SHE) by adding a value of 0.205 V, or to the
reversible hydrogen electrode (RHE) by adding a value of (0.205 + 0.059 pH) V.
Controlled potential electrolysis measurements to determine Faradaic efficiency and study long-
term stability were conducted in a sealed two-chambered H-cell where the first chamber held the
working and reference electrodes in 50 mL of 0.1 M NaClO4 (aq) at the corresponding pH, and
the second chamber held the auxiliary electrode in 25 mL of 0.1 M NaClO 4 (aq). The two
chambers, which were both under N2, were separated by a fine porosity glass frit. Plate-shaped
GCE (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and auxiliary
electrodes. The reference electrode was a Ag/AgCl/saturated KCl electrode separated from the
solution by a Vycor tip. Using a gas-tight syringe, 10 mL of gas were withdrawn from the
headspace of the H-cell and injected into a gas chromatography instrument (Shimadzu GC-2010-
Plus) equipped with a BID detector and a Restek ShinCarbon ST Micropacked column. To
determine the Faradaic efficiency of the electrochemical cell, the theoretical H2 amount based on
total charge flowed is compared with the GC-detected H2 produced from controlled-potential
electrolysis.
The aqueous solutions used in the electrochemical experiments have been prepared as follows. For
the pH 1.3 solution, 0.534 mL of 18.7 M H2SO4 were added to water (200 mL). For the pH 2.6
solution, citric acid (3.458 g) and Na2HPO4 (1.505 g) were dissolved in water (200 mL). For the
pH 4.4 solution, NaOAc (1.605 g) was dissolved in water (200 mL). Acetic acid (1.2 mL) was
added to reach the desired pH. For the pH 7.1 solution, NaH2PO4 (0.468 g) and Na2HPO4 (1.637
g) were dissolved in water (100 mL). For the pH 10.0 solution, NaHCO3 (0.339 g) and Na2CO3
(0.632 g) were dissolved in water (100 mL). The pHs of the solutions were measured with a
benchtop Mettler Toledo pH meter. All solutions were degassed and purged with nitrogen.
3.5.7. Physical Methods
FTIR spectra were acquired using a Bruker Vertex 80v spectrometer. Samples (2 mg) for analysis
were mixed into a KBr (100 mg) matrix and pressed into pellets.
71
Raman spectra were collected on a Horiba Raman Microscope using films of NiBTT, FeBTT, and
benzene-1,2,4,5-tetrathiol on glass microscope slides.
Inductively coupled plasma optical emission spectroscopy (ICP-OES) measurements were
performed using a Thermo Scientific iCAP 7000 ICP-OES.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray
source was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
Scanning electron microscopy (SEM) was performed on a JEOL JSM 7001F scanning electron
microscope.
3.6. References
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(4) Dempsey, J. L.; Brunschwig, B. S.; Winkler, J. R.; Gray, H. B. Acc. Chem. Res. 2009, 42,
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(5) Artero, V.; Chavarot-Kerlidou, M.; Fontecave, M. Angew. Chem. Int. Ed. 2011, 50, 7238-
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75
CHAPTER 4
Bioinspired Benzenetetraselenolate-based Coordination Polymers
for H
2
Evolution
A portion of this chapter has appeared in print:
Downes, C. A.; Marinescu, S. C. “Bioinspired Metal Selenolate Polymers with Tunable
Mechanistic Pathways for Efficient H2 Evolution.” ACS Catal. 2017, 7, 848-854.
76
4.1. Introduction
The ability to efficiently and cost-effectively convert solar energy into molecular hydrogen
through water splitting is necessary for meeting rising energy demands and to mitigate the adverse
effects of carbon-based fuels on the environment.
1-3
Numerous homogeneous
4-8
and
heterogeneous
9-12
systems have been developed to catalyze the hydrogen evolution reaction (HER)
in order to replace noble metals such as platinum. Hydrogenase enzymes catalyze reversible HER
from water with remarkably high turnover frequencies at low overpotentials and have inspired the
design of synthetic catalysts based on earth-abundant metals.
13-18
[NiFeSe] hydrogenases are a
subclass of [NiFe] hydrogenases with a selenocysteine (Sec) replacing a cysteine (Cys) residue
terminally bound to the Ni center (Figure 4.1).
19
These [NiFeSe] hydrogenases have emerged as
efficient catalysts for the HER, due to their increased activity (by a factor of 40) and oxygen-
tolerant H2 production in comparison with the conventional [NiFe] hydrogenases.
20,21
Additionally, integration of [NiFeSe] hydrogenases into photocatalytic and photoelectrochemical
systems has proven to be an effective and efficient method for direct light-to-hydrogen conversion
displaying the viability of photobiological fuel production.
22,23
Because of the successful
utilization of [NiFeSe] hydrogenases as catalysts in light-driven H2 production devices, these
enzymes have emerged as an attractive blueprint for designing synthetic H2 evolution catalysts.
Figure 4.1. Structures of the active sites of [NiFe] and [NiFeSe] hydrogenases.
Although the high catalytic activity of [NiFeSe] hydrogenases is well known, the role selenium
plays in catalysis is currently unclear, however, it has been proposed that the Sec residue acts as a
proton relay during H2 cycling similar to the Cys residue in [NiFe]-hydrogenases.
21
The enhanced
H2 production activity of [NiFeSe] hydrogenases has been attributed to the higher nucleophilicity
of selenocysteine than cysteine and the lower pKa of a selenocysteine selenol than a cysteine
thiol.
21
Detection of protonated selenocysteine in hydrogenase during H2 production is
challenging, therefore biomimetic molecular complexes bearing the Ni-Se moiety have been
77
developed to facilitate a better understanding of the function of selenium during the HER.
21,24-26
However, under electrocatalytic conditions, a solid deposit formed on the electrode surface and it
was established that the Ni-Se molecular mimics in solution were not responsible for the observed
HER activity.
21,24-26
Therefore, molecular systems active for electrocatalytic HER with a Ni-Se
motif have yet to be developed.
Heterogeneous systems based on pyrite-phase transition metal chalcogenides have been identified
as earth-abundant and low-cost alternatives to platinum that display high efficiencies and stability
for the HER in both acidic and basic media.
27-31
Interestingly, it has been shown that CoSe2
displays increased HER activity in comparison to CoS2 similar to the improvement in H2
production for [NiFe] hydrogenases when one of the cysteine residues is replaced by
selenocysteine.
21,24-26,29
It has been proposed that the strength of the E-H bond (where E = S or Se)
can influence the rate of hydrogen desorption from the catalytic site during the HER.
32
Therefore,
the enhanced activity of selenides versus sulfides could arise because the Se-H bond is weaker
than the S-H bond.
32
We have recently reported the integration of metal dithiolene units into extended frameworks as a
method of heterogenizing molecular catalysts.
33,34
These systems have been shown to retain their
molecular nature while displaying enhanced stability, activity, and durability in acidic aqueous
media in comparison to their molecular analogues. Cobalt and nickel coordination polymers (CPs)
based on benzene-1,2,4,5-tetrathiolate (BTT) exhibit high electrocatalytic activities and
efficiencies for the HER in pH 1.3 aqueous solutions.
33-35
The mechanism of H2 evolution for metal
dithiolenes is proposed to involve protonation of the sulfur moiety of the dithiolene ligand, which
overtime may result in ligand loss and catalyst decomposition.
36-38
We have shown that
incorporation of metal dithiolenes into extended frameworks can reduce the prevalence of
decomposition pathways and eliminates the use of organic solvents, which have been known to
contribute to catalyst decomposition.
33,34
Although the polymers were shown to operate at high efficiencies with no evidence of
decomposition, the overpotentials of 560 mV and 470 mV necessary to reach 10 mA/cm
2
of
activity in pH 1.3 aqueous solutions for the cobalt
34
and nickel
35
BTT polymers, respectively, were
78
undesirable. Because of the enhanced activity seen for selenium-based catalysts in comparison to
their sulfur counterparts, we set out to synthesize bioinspired selenolate analogues of our cobalt
and nickel coordination polymers as a strategy to improve the electrocatalytic activity and reduce
the overpotential for the HER. Herein, we report that a cobalt CP based on benzene-1,2,4,5-
tetraselenolate (BTSe) is an active HER electrocatalyst that displays enhanced activity and a
reduced overpotential in comparison to its sulfur only analogue. Additional electrochemical
studies on a nickel CP based on BTSe revealed high activity toward the HER in acidic aqueous
solutions offering possible insight into the role of selenium in the Ni-Se motif present in [NiFeSe]
hydrogenase.
4.2. Results and Discussion
Cobalt and nickel CPs were synthesized using a liquid-liquid interfacial reaction modified from
the reported BTT polymers previously synthesized in our laboratory (Scheme 4.1).
33,34
1,2,4,5-
tetrakis(acetylseleno)benzene (BTSeAc4) was synthesized following a reported procedure and the
isolation of the acetyl-protected selenol affords air-stability and reduces the difficulty associated
with handling air-sensitive selenols.
39
The acetyl groups in BTSeAc4 can be easily hydrolyzed in
the presence of base, such as NaO
t
Bu to form sodium BTSe (C6H2Se4Na4). An acetonitrile/ethyl
acetate solution of [M(MeCN)6][BF4]2 (where M = Co, Ni) was gently layered on top of the
aqueous solution of sodium BTSe. The organic solvents were allowed to evaporate over several
hours leaving a black film (CoBTSe and NiBTSe) at the gas-liquid interface, which was
subsequently deposited on the support of interest.
Scheme 4.1. Synthesis of benzene-1,2,4,5-tetraselenolate-based cobalt (CoBTSe) and nickel
(NiBTSe) coordination polymers through a liquid-liquid interfacial reaction.
79
Deposition was carried out by two methods: (a) immersing the support face down in the reaction
mixture and (b) synthesizing CoBTSe and NiBTSe with the support at the bottom of the reaction
vessel which was subsequently lifted up and through the film formed at the gas-liquid interface.
Following deposition, the modified electrodes were washed with water and methanol to remove
the residual starting materials. The poor film quality and low yield of NiBTSe due to the interfacial
reaction conditions limited its extensive characterization; therefore, electrochemical studies were
focused on CoBTSe whose synthetic preparation resulted in a high quality film and yield.
Figure 4.2. FTIR spectra of CoBTSe (purple) and 1,2,4,5-tetrakis(acetylseleno)benzene
(BTSeAc4, cyan).
The FTIR spectrum of CoBTSe showed the disappearance of the strong C=O stretch of the acetyl
protecting group of BTSeAc4 around ~1720 cm
–1
demonstrating the successful hydrolysis of the
acetyl group (Figure 4.2). XPS studies of CoBTSe revealed the presence of Co, Se, and Na (Figure
4.3). In the cobalt region, two peaks at binding energies of 778.5 and 793.6 eV were observed for
the 2p3/2 and 2p1/2 levels with the corresponding satellite features similar to the previously reported
cobalt CPs based on BTT, benzenehexathiolate (BHT), and triphenylene-2,3,6,7,10,11-
hexathiolate (THT).
33,34
Two additional peaks were observed at 1071.3 and 55.1 eV which arise
from the presence of Na 1s and Se 3d. Deconvolution of the Se 3d region revealed two peaks at
54.7 and 55.6 eV corresponding to 3d5/2 and 3d3/2 levels in agreement with the expected metal-
selenolate binding motif.
40
XPS studies of NiBTSe revealed the presence of Ni with two peaks at
binding energies of 853.5 and 870.7 eV along with the corresponding satellite features (Figure
4.4), which is similar with the XPS spectra of reported Ni
2+
complexes.
41
Deconvolution of the
selenium signal resulted in two peaks at 55.1 and 56 eV for the 3d5/2 and 3d3/2 levels (Figure 4.4).
A peak at 1071.1 eV indicated the presence of Na (Figure 4.4).
80
Figure 4.3. X-ray photoelectron spectroscopy analysis of CoBTSe.
Figure 4.4. X-ray photoelectron spectroscopy analysis of NiBTSe.
81
The electrochemistry of CoBTSe and NiBTSe were investigated by cyclic voltammetry (CV) on
glassy carbon electrodes (GCE). No differences in electrochemistry were observed for electrodes
prepared from different synthetic batches or using deposition methods (a) or (b). A broad
electrochemical wave was observed for CoBTSe in pH 10.0 aqueous solutions between 0 and –
0.3 V versus SHE (Figure 4.5). This wave was negatively shifted compared to the analogous cobalt
BTT CP
34
(Figure 4.5), which is consistent with previous reports.
24-26
Because selenium increases
the electron density on the cobalt center, the reduction process is thermodynamically less
favorable. Electrochemical investigation of the related molecular complexes, [Co(bdt)2]
–
and
[Co(bds)2]
–
(where bdt = 1,2-benzenedithiolate and bds = 1,2-benzenediselenolate), revealed the
reduction of [Co(bds)2]
–
occurred 40 mV more negative than [Co(bdt)2]
–
.
42
An analogous negative
shift in the broad electrochemical wave in pH 10.0 aqueous solutions was observed for NiBTSe in
comparison to the related nickel BTT CP (Figure 4.6).
Figure 4.5. Polarization curves of CoBTSe (purple, 9.2 × 10
–7
molCo/cm
2
) and CoBTT (black,
4.4(4) × 10
–7
molCo/cm
2
) in 0.1 M NaClO4 aqueous solution at pH 10.0. Scan rate: 20 mV/s.
Figure 4.6. Polarization curves of NiBTSe (blue, 4.5 × 10
–7
molNi/cm
2
) and NiBTT (black, 4.2 ×
10
–7
molNi/cm
2
) in 0.1 M NaClO4 aqueous solutions at pH 10.0. Scan rate: 20 mV/s.
82
Figure 4.7. (a) Scan rate dependence studies of CoBTSe (catalyst loading: 7.7 × 10
–7
molCo/cm
2
)
in 0.1 M NaClO4 aqueous solutions at pH 10.0. Scan rates varied from 5 to 100 mV/s. (b) Cathodic
current density at -0.22 V versus SHE of CoBTSE plotted as a function of scan rate (mV/s).
The cathodic wave of CoBTSe in pH 10.0 solutions varied linearly with the scan rate as expected
for a surface confined process (Figures 4.7). Large cathodic currents were generated as the pH of
the solution was decreased (Figures 4.8), which is attributed to the HER, as described below. This
catalytic onset began in pH 4.4 solutions (Figure 4.7) and reached significant activity in the most
acidic media tested, pH 1.3.
Figure 4.8. Polarization curves of CoBTSe (7.3 × 10
–7
molCo/cm
2
) in 0.1 M NaClO4 aqueous
solutions at pH 10.0 (orange), pH 7.0 (purple), 4.4 (green), 2.6 (blue), 1.3 (red), and of blank GCE
at pH 1.3 (dashed black); Scan rate: 20 mV/s (pH 10.0–2.6) and 100 mV/s (pH 1.3).
Interestingly, the onset of catalysis emerged at the potential of the cathodic wave seen in pH 10.0
solutions (~ –0.25 V versus SHE) in comparison to the sulfur analog, which occurs ~0.5 V more
negative than its corresponding cathodic wave. The first onset was followed by a second at
approximately –0.6 V versus SHE possibly indicating a change in the mechanism of catalysis, as
83
described below. Similar enhancements in catalytic current upon lowering the pH were seen for
NiBTSe, which displayed H2 evolving activity beginning in pH 4.4 aqueous solutions (Figure 4.9).
However, only one onset was observed for NiBTSe indicating that the change in mechanism as a
function of applied potential is unique to CoBTSe.
Figure 4.9. Polarization curves of NiBTSe (4.5 × 10
–7
molNi/cm
2
) in 0.1 M NaClO4 aqueous
solutions at pH 10.0 (orange), pH 7.0 (purple), 4.4 (green), 2.6 (blue), 1.3 (red), and of blank GCE
at pH 1.3 (dashed black); Scan rate: 20 mV/s (pH 10.0–7.0) and 100 mV/s (pH 4.4–1.3).
Figure 4.10. Current densities of CoBTT measured at –0.80 V versus SHE (–0.72 V versus RHE)
at pH 1.3 as a function of the surface catalyst concentration. The surface concentration was
quantified by integrating the peak area of the redox wave observed in pH 10.0.
The catalytic activity at –0.80 V versus SHE (–0.72 V versus RHE) was directly proportional to
the catalyst loading for CoBTSe (Figure 4.10), which was estimated through integration of the
electrochemical wave in pH 10.0 aqueous solutions. As the catalyst concentration increased, the
intensity of the first onset at approximately –0.25 V versus SHE (–0.17 V versus RHE) was greatly
84
enhanced (Figure 4.11) suggesting the prominence of the mechanism of H2 evolution at more
positive potentials may be related to the catalyst loading. Analyzing the generated catalytic current
as a function of the catalyst concentration revealed the intensity of the first onset was not solely
due to the catalyst loading (Figure 4.11b). The highest catalyst loading achieved, 9.2 × 10
–7
molCo/cm
2
, generated the benchmarking metric of 10 mA/cm
2
of activity
11
at an overpotential
(𝜂 10 𝑚𝐴 /𝑐𝑚
2) of only 343 mV in pH 1.3 aqueous solutions. The 𝜂 10 𝑚𝐴 /𝑐𝑚
2 for CoBTSe ranged
from ~600–350 mV as a function of its corresponding catalyst loading (Table 4.1).
Figure 4.11. Polarization curves of CoBTSe in pH 1.3 at different catalyst loadings; 3.7 × 10
–7
molCo/cm
2
(green), 7.3 × 10
–7
molCo/cm
2
(blue), 9.2 × 10
–7
molCo/cm
2
(red), and blank GCE (black
dashed) with the y-axis represented as (a) current density (mA/cm
2
) and (b) current per
concentration of cobalt (mA/molCo × 10
5
) measured via integration of the redox wave in pH 10.0.
Scan rate: 100 mV/s.
Table 4.1. Overpotential necessary to reach 10 mA/cm
2
of activity in pH 1.3 aqueous solutions at
various catalyst loadings of CoBTSe.
Catalyst Loading
× 10
-7
molCo/cm
2
η
10 mA/cm
2 (mV)
3.7 602
5.0 572
7.3 517
9.2 343
To ensure production of H2 was responsible for the increase in current density in pH 1.3 solutions,
controlled potential electrolysis experiments (CPE) were performed. CPE of CoBTSe performed
at –0.72 V versus RHE in pH 1.3 aqueous solutions showed the electrocatalytic response was
stable for 1 h (Figure 4.12). Over 1 h, 81 coulombs of charge were consumed and analysis of the
85
gaseous headspace of the electrolysis cell by gas chromatography confirmed the production of H2
with a Faradaic efficiency (FE) of 97±2%. Additionally, because two onsets of catalysis were
observed in CV experiments, it was important to determine if the earlier onset was also due to H2
production by CoBTSe. CPE experiments were conducted at –0.32 V versus RHE in pH 1.3
aqueous solutions. After 1 h, 53 coulombs of charge were passed with a 98±2% FE for H2
production, suggesting that CoBTSe is an efficient catalyst for the HER at low ɳ.
Figure 4.12. Controlled potential electrolysis of CoBTSe in pH 1.3 at –0.72 V versus RHE (1.2 ×
10
–7
molCo/cm
2
; red), –0.32 V versus RHE (4.0 × 10
–7
molCo/cm
2
; blue), and blank GCE (black
dashed).
Since electrocatalytic proton reduction was responsible for the two onsets seen for CoBTSe and
not a secondary process, we propose the two onsets are a result of a change in mechanism as the
potential is swept more negative. Following a one-electron reduction and a single protonation at
one of the selenolate moieties of the ligand, catalysis can proceed through two pathways. At more
positive potentials, a second protonation followed by the final reduction to produce H2 is favored
(ECCE where E = electron transfer and C = chemical reaction). As the potential is scanned more
negative, reduction followed by protonation and release of H2 is preferred (ECEC). Similar
changes in mechanism as a function of applied potential have been previously observed for nickel
bis(aryldithiolene) complexes for HER.
43
Because the intensity of the first onset increases as a
function of catalyst loading (Table 4.1, Figure 4.11), we propose that the first mechanism can be
favored as the catalyst loading increases. Additionally, it should be noted that the two onsets could
also be due a change in the substrate used for H2 generation with proton reduction favored at
smaller overpotentials and water reduction at larger overpotentials.
86
Figure 4.13. Polarization curves of CoBTSe (9.2 × 10
–7
molCo/cm
2
, purple), benzene-1,2,4,5-
tetrathiolate-based cobalt CP (CoBTT, 5.5 × 10
–7
molCo/cm
2
, green), and blank GCE (black
dashed) in 0.1 M NaOCl4 aqueous solutions at pH 1.3 with the y-axis represented as (a) current
density (mA/cm
2
) and (b) current per concentration of cobalt (mA/molCo × 10
5
) measured via
integration of the redox wave in pH 10.0. Scan rate: 100 mV/s.
Theoretical studies of the cobalt dithiolene complex, [Co(bdt)2]
–
, suggest that upon one-electron
reduction of [Co(bdt)2]
–
, two protonations occur on different sulfur moieties of the
benzenedithiolate ligand followed by a second one-electron reduction.
37
A subsequent
intramolecular proton shift to form a cobalt hydride adjacent to a protonated sulfur leads to release
of H2.
37
Two mechanistic pathways for the HER have been proposed for nickel dithiolenes, ECEC
and ECCE (E = electron transfer and C = chemical reaction) with protonation occurring solely on
the sulfur moiety of the dithiolene ligand.
43
Calculations have suggested that geometric distortions
occur to accommodate the protonated ligand and to allow for H2 evolution.
43
No stable species
protonated at both nickel and sulfur have been observed for nickel dithiolenes. The mechanistic
pathways for the cobalt and nickel dithiolenes are determined by the frontier orbitals of the
complexes. For nickel bis(aryldithiolene) complexes, it is proposed that the protonation and redox
events occur primarily on the ligand because the frontier orbitals are mainly ligand based.
43
For
the cobalt bis(dithiolene) complexes, however, the cobalt orbitals contribute more significantly to
the frontier orbitals so protonation and redox events occur on both the ligand and the metal center.
37
We propose that an ECCE mechanism, favored at higher catalyst loadings, would involve
protonation of the selenium moieties, since selenium is known to be more readily protonated than
its corresponding sulfur analogue.
21,25,32
Similar with the mechanism of H2 evolution by nickel
bis(aryldithiolene) complexes, we propose that H2 evolution by cobalt selenolate polymers occurs
87
from recombination of two protonated selenium moieties, which are in closer proximity at high
catalyst loadings. This type of mechanism does not necessitate an intramolecular proton transfer
to form a metal hydride. For nickel bis(aryldithiolene) complexes, following a geometric
distortion, the two protonated sulfurs recombine from the same catalytic unit to evolve H2.
43
Because of the rigidity of the coordination polymer network, we suggest that two protonated
selenium moieties from two different catalytic units can combine following a second one-electron
reduction and evolve H2. As the catalyst loading increases, the availability of protonated selenium
moieties in close enough proximity to recombine upon one-electron reduction increases.
Figure 4.14. Polarization curves of CoBTSe (pink; 9.2 × 10
–7
molCo/cm
2
) and NiBTSe (blue; 9.0
× 10
–7
mol Ni/cm
2
) in 0.1 M NaOCl4 aqueous solutions at pH 1.3.
The appearance of two onsets and the possibility of two available mechanisms was not observed
for the related cobalt BTT CP indicating that this dual mechanism is unique to CoBTSe. The
𝜂 10 𝑚𝐴 /𝑐𝑚
2 of 343 mV for CoBTSe is a significant improvement from the 𝜂 10 𝑚𝐴 /𝑐𝑚
2 of 560 mV
measured for the cobalt BTT derivative (Figure 4.13).
34
We attribute this enhanced activity to the
larger exchange current density of CoBTSe, further discussed below, and the availability of an
alternative mechanism. Catalysis occurs at the potential of the cathodic wave seen in pH 10.0
solutions (~ –0.25 V versus SHE) in comparison to the HER for the sulfur analog, which only
occurs at potentials ~0.5 V more negative than its corresponding cathodic wave. The activity of
CoBTSe is similar to that of the cobalt BHT framework reported by us (𝜂 10 𝑚𝐴 /𝑐𝑚
2 = 340 mV).
33
At a catalyst loading of 9.0 × 10
–7
molNi/cm
2
, NiBTSe displayed an 𝜂 10 𝑚𝐴 /𝑐𝑚
2 of 353 mV in pH
1.3 solution (Figure 4.14). This is an improvement from the overpotential of 470 mV necessary
for the related nickel BTT CPs at a catalyst loading of 6.8 × 10
–7
molNi/cm
2
.
35
An improvement in
88
catalytic activity has previously been reported for [NiFeSe] hydrogenase and CoSe2 in comparison
to their sulfur only analogues.
21,24-26,29
Table 4.2. Electrocatalytic HER properties of diselenolate and dithiolate coordination polymers.
M Ligand
a
Catalyst Loading
× 10
-7
mol M/cm
2
η
10 mA/cm
2 (mV)
Tafel slope
(mV/dec)
i 0
(A/cm
2
)
Ref.
Co BTSe 9.2 343 97 (pH 1.3) 10
–4.4
this work
Ni BTSe 9.0 353 ̶ ̶ this work
Co BTT 5.5 560 70 (pH 1.3)
b
10
–9.4
34
Ni BTT 6.8 470 76 (pH 2.6) 10
–8.1
35
Co BHT 7.0 340 108 (pH 2.6) 10
–5.3
33
Co THT 11.0 530 161 (pH 2.6) 10
–5.4
33
a
Where BTSe = benzene-1,2,4,5-tetraselenolate, BTT = benzene-1,2,4,5-tetrathiolate, BHT = benzenhexathiolate, and THT =
triphenylene-2,3,6,7,10,11-hexathiolate.
b
Tafel slope of 94 mV/dec recorded in pH 2.6.
Tafel analysis of CoBTSe gave a Tafel slope of 97 mV/dec and an exchange current density of
10
–4.4
A/cm
2
(Figure 4.15). Tafel slopes and exchange current densities of the previously reported
cobalt and nickel coordination polymers have been included in Table 4.2.
33-35
The cobalt BTT CP
exhibits the lowest Tafel slope which indicates the highest intrinsic activity for the HER, however,
it is the least active of the four cobalt based polymers studied in our laboratory. The higher Tafel
slopes of CoBTSe, cobalt BHT, and cobalt THT signifies the rate-determining step in the HER
mechanism is the Volmer reaction, the discharge step that converts protons into absorbed hydrogen
on the catalyst surface. The reduction in the Tafel slope for the cobalt BTT polymers suggests a
reduction in the energy barrier associated with the Volmer reaction. Further investigation is
necessary to determine the origin of the large variance (70–160 mV/dec) in the experimentally
determined Tafel slopes.
Figure 4.15. Tafel plot of CoBTSe (3.8 × 10
–7
molCo/cm
2
) in pH 1.3; Scan rate: 0.5 mV/s.
89
Although the Tafel slope predicts high intrinsic activity, the exchange current density for the cobalt
BTT polymer is the lowest of the four materials. High exchange current densities are characteristic
of high electrocatalytic activity. The cobalt BTSe, BHT, and THT polymers possess similar
exchange current densities. This can explain the comparative activity of the cobalt BTSe and BHT
frameworks that exhibit similar Tafel slopes and exchange current densities, and can reach 10
mA/cm
2
of activity at overpotentials ~350 mV. The cobalt THT system, on the other hand, requires
an overpotential of 530 mV. Although the cobalt THT system has a high exchange current density
in comparison to the cobalt BTT system, the overpotential to reach 10 mA/cm
2
is very similar
(530–560 mV) which may be attributed to its large Tafel slope of 161 mV/dec. The Tafel slope of
161 mV/dec is the largest of the coordination polymers synthesized in our laboratory suggesting
the energy barrier of desorption of H2 from the catalyst surface (Volmer reaction) is greatest for
the cobalt THT system.
Further CPE studies of NiBTSe also revealed continuous charge build up over 1 h in pH 1.3 at –
0.47 V versus RHE, however, the system operated at a reduced efficiency of ~70% (Figure 4.16).
This reduction in efficiency for the nickel analogue relative to cobalt was also observed for the
BTT derivatives, with the nickel BTT CP operating at an 85% FE compared to 97% for cobalt.
34,35
Figure 4.16. Controlled potential electrolysis of NiBTSe (2.1 × 10
–7
molNi/cm
2
; red) in pH 1.3
H2SO4 aqueous solution at –0.47 V versus RHE and blank GCE (black dashed).
CoBTSe was also subjected to 4 h of CPE in pH 1.3 solution at –0.52 V versus RHE and displayed
moderate stability (Figure 4.17). CV experiments before and after exposure to strong reductive
acidic conditions revealed no significant change (Figures 4.18). XPS analyses following 5 h of
CPE for CoBTSe and 1 h of CPE for NiBTSe in pH 1.3 revealed no change in the features observed
90
prior to electrochemical testing (Figures 4.19-4.20). Negligible amounts of cobalt and selenium
were observed in ICP-OES measurements of the pH 1.3 solution used for longer duration CPE
experiments of CoBTSe indicating no cobalt or selenium species were solubilized over the course
of electrochemical analysis.
Figure 4.17. Controlled potential electrolysis of CoBTSe (1.2 × 10
–7
molCo/cm
2
; red) in pH 1.3
H2SO4 aqueous solution at –0.52 V versus RHE.
Figure 4.18. Polarization curves of CoBTSe (4.0 × 10
–7
molCo/cm
2
) (a) before controlled potential
electrolysis (red), after 2 h (blue), and 4 (green) h of CPE in pH 1.3 at –0.52 V versus RHE and
(b) before controlled potential electrolysis (red) and after 1 h of CPE (blue) in pH 1.3 at –0.36 V
versus RHE.
91
Figure 4.19. X-ray photoelectron spectroscopy (XPS) analysis of CoBTSe after 5 h of controlled
potential electrolysis in pH 1.3 aqueous solutions at -0.52 V versus RHE.
Figure 4.20. X-ray photoelectron spectroscopy (XPS) analysis of NiBTSe after 1 h of controlled
potential electrolysis in pH 1.3 aqueous solutions at -0.47 V versus RHE.
92
4.3. Conclusions
In conclusion, we have demonstrated the preparation of bioinspired cobalt and nickel selenolate
polymers, which act as efficient electrocatalysts for the HER from water. We report here the first
molecular system utilizing the Co-Se motif to be investigated as a highly active and stable catalyst
for the HER in acidic aqueous media representing an important step in understanding the role of
selenium in facilitating and improving the H2 evolving activity in aqueous solutions. CoBTSe
exhibits an 𝜂 10 𝑚𝐴 /𝑐𝑚
2 of only 343 mV in pH 1.3 aqueous solutions at a catalyst loading of 9.2 ×
10
–7
molCo/cm
2
.
The replacement of sulfur with selenium results in a 217 mV improvement in the
ɳ to reach the same H2-evolving activity. This improvement arises from the presence of two
catalytic regimes for CoBTSe. We propose two potential dependent mechanisms, ECCE and
ECEC, with ECCE greatly reducing the overpotential for the HER at higher catalyst loadings.
Improvements in the H2-evolving activity were also observed for NiBTSe, which exhibits a 117
mV enhancement in 𝜂 10 𝑚𝐴 /𝑐𝑚
2 compared to the analogous nickel BTT CP. The significant
improvement in the HER activity originating from the simple modification of the ligand
framework displays the advantages of maintaining the molecular nature of catalysts, while the low
ɳ and durability under prolonged reductive acidic conditions documents the advantages of
heterogenization of molecular catalysts. The incorporation and immobilization of molecular
catalysts via CPs has proven a viable method to systematically tune important catalytic metrics
such as mechanistic pathways, overpotential, stability, and activity, aiding in the design of earth-
abundant and efficient H2 evolving catalysts.
4.4. Experimental Details
4.4.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity). Ethyl acetate and water were placed under vacuum and refilled with nitrogen (10 ×).
Acetonitrile was degassed with nitrogen and passed through activated alumina columns and stored
over 4Å Linde-type molecular sieves. 1,2,4,5-tetrakis(acetylseleno)benzene was prepared
according to the reported procedure.
39
All other chemical reagents were purchased from
93
commercial vendors and used without further purification. The pHs of the aqueous solutions were
measured with a benchtop Mettler Toledo pH meter.
Caution: Perchlorate salts are potential explosive chemicals and should only be used in very small
amounts, especially in the presence of the reductant H2. Please handle these mixtures using the
proper equipment and protection.
4.4.2. Synthesis of CoBTSe
In a glove-box, Na
t
OBu (0.18 mL of 0.5 M solution in methanol, 4 equivalents) was added to
1,2,4,5-tetrakis(acetylseleno)benzene (12.88 mg, 0.023 mmol). Once dissolved, 20 mL of
degassed H2O was added to the solution. The aqueous solution was transferred to a 70 mm × 50
mm crystallizing dish. [Co(MeCN)6][BF4]2 (36.3 mg, 0.146 mmol) was dissolved in a 1:4 mixture
of acetonitrile and ethyl acetate and 1 mL of this solution was carefully layered via glass pipette
on top of the aqueous solution to cover ~80% of the surface area. The organic solvents were
allowed to evaporate over one to two hours at room temperature, leaving behind CoBTSe as a
black solid at the gas-liquid interface. The black solid was collected, washed with water, methanol.
The resultant black powder was dried under vacuum.
4.4.3. Synthesis of NiBTSe
In a glove-box, Na
t
OBu (0.18 mL of 0.5 M solution in methanol, 4 equivalents) was added to
1,2,4,5-tetrakis(acetylseleno)benzene (12.88 mg, 0.023 mmol). Once dissolved, 20 mL of
degassed H2O was added to the solution. The aqueous solution was transferred to a 70 mm × 50
mm crystallizing dish. [Ni(MeCN)6][BF4]2 (36.3 mg, 0.146 mmol) was dissolved in a 1:4 mixture
of acetonitrile and ethyl acetate and 1 mL of this solution was carefully layered via glass pipette
on top of the aqueous solution to cover ~80% of the surface area. The organic solvents were
allowed to evaporate over one to two hours at room temperature, leaving behind NiBTSe as a
black solid at the gas-liquid interface.
4.4.4. Deposition of CoBTSe and NiBTSe
Deposition was carried out by two methods: (a) immersing the support face down in the reaction
mixture and (b) synthesizing CoBTSe or NiBTSe with the support at the bottom of the reaction
94
vessel which was subsequently lifted up and through the film formed at the gas-liquid interface.
Following deposition, the substrate was washed with water and methanol.
4.4.5. Electrochemical Methods
Electrochemistry experiments were carried out using a Pine potentiostat. Platinum wire used for
the electrochemical studies was purchased from Alfa Aesar. The electrochemical experiments
were carried out in a three-electrode configuration electrochemical cell under an inert atmosphere
using glassy carbon electrodes (GCE) as the working electrode. A platinum wire, placed in a
separate compartment, connected by a Vycor tip, and filled with the electrolytic solution (0.1 M
NaClO4) was used as the auxiliary (counter) electrode. The reference electrode, placed in a separate
compartment and connected by a Vycor tip, was based on an aqueous Ag/AgCl/saturated KCl
electrode. The reference electrode in aqueous media was calibrated externally relative to
ferrocenecarboxylic acid (Fc-COOH) at pH 7.0, with the Fe
3+/2+
couple at 0.28 V vs Ag/AgCl. All
potentials reported in this paper were converted to the standard hydrogen electrode (SHE) by
adding a value of 0.205 V, or to the reversible hydrogen electrode (RHE) by adding a value of
(0.205 + 0.059 × pH) V.
Controlled potential electrolysis measurements to determine Faradaic efficiency and study long-
term stability were conducted in a sealed two-chambered H-cell where the first chamber held the
working and reference electrodes in 50 mL of 0.1 M NaClO4 (aq) at the corresponding pH, and
the second chamber held the auxiliary electrode in 25 mL of 0.1 M NaClO 4 (aq). The two
chambers, which were both under N2, were separated by a fine porosity glass frit. Glassy carbon
plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and auxiliary
electrodes. The reference electrode was a Ag/AgCl/saturated KCl(aq) electrode separated from the
solution by a Vycor tip. Using a gas-tight syringe, 2 mL of gas were withdrawn from the headspace
of the H-cell and injected into a gas chromatography instrument (Shimadzu GC-2010-Plus)
equipped with a BID detector and a Restek ShinCarbon ST Micropacked column. To determine
the Faradaic efficiency, the theoretical H2 amount based on total charge flowed is compared with
the GC-detected H2 produced from controlled-potential electrolysis.
95
4.4.6. Physical Methods
FT-IR spectra were acquired using a Bruker Vertex 80v spectrometer. Samples (2 mg) for analysis
were mixed into a KBr (100 mg) matrix and pressed into pellets.
Inductively coupled plasma optical emission spectroscopy (ICP-OES) measurements were
performed using a Thermo Scientific iCAP 7000 ICP-OES.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray
source was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
The aqueous solutions used in the electrochemical experiments have been prepared as follows. For
the pH 1.3 solution, 0.534 mL of 18.7 M H2SO4 were added to water (200 mL). For the pH 2.6
solution, citric acid (3.458 g) and Na2HPO4 (1.505 g) were dissolved in water (200 mL). For the
pH 4.4 solution, NaOAc (1.605 g) was dissolved in water (200 mL). Acetic acid (1.2 mL) was
added to reach the desired pH. For the pH 7.1 solution, NaH2PO4 (0.468 g) and Na2HPO4 (1.637
g) were dissolved in water (100 mL). For the pH 10.0 solution, NaHCO3 (0.339 g) and Na2CO3
(0.632 g) were dissolved in water (100 mL). The pHs of the solutions were measured with a
benchtop Mettler Toledo pH meter. All solutions were degassed and purged with nitrogen.
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98
CHAPTER 5
H
2
Evolution by a Cobalt Selenolate Electrocatalyst
and Related Mechanistic Studies
A portion of this chapter has appeared in print:
Downes, C. A.; Yoo, J. W.; Orchanian, N. M.; Marinescu S. C. “H2 Evolution by a Cobalt
Selenolate Electrocatalyst and Related Mechanistic Studies.” Chem. Commun. 2017, 53, 7306-
7309.
99
5.1. Introduction
The development of efficient and inexpensive catalysts capable of solar-to-fuel energy conversion
is vital for mitigating the adverse effects of carbon-based fuels on the environment and meeting
the rising global energy demand.
1-3
Hydrogenase enzymes reversibly catalyze the hydrogen
evolution reaction (HER) at high efficiencies near the thermodynamic potential.
4
The coordination
environment and prominence of earth-abundant metals in the active sites of hydrogenase enzymes
has led to the development of H2 evolving catalysts based on iron, nickel, and cobalt.
5,6
Aryldithiolate metal complexes have been extensively explored as catalysts for the HER with
mechanistic studies indicating sulfur as the site of protonation analogous to the thiolate donors
acting as proton relays in the hydrogenase enzymes.
7-16
[NiFeSe] hydrogenases, a subclass of [NiFe] hydrogenases with a selenocysteine (Sec) replacing
a cysteine (Cys) residue, exhibit increased HER activity in comparison to conventional [NiFe]
hydrogenases.
17,18
Several molecular models have been developed to understand the role selenium
plays in enhancing the HER activity, however, these systems were shown to decompose during
catalytic testing.
17,19,20
Recent work in our laboratory has investigated one and two dimensional
coordination polymers based on hexathiolate,
21
tetrathiolate,
22,23
and tetraselenolate
24
ligand
frameworks as catalysts for the HER in pH 1.3 aqueous solutions. The cobalt benzene-1,2,4,5-
tetraselenolate polymer exhibited a ~200 mV reduction in the overpotential to reach 10 mA/cm
2
of activity in comparison to the cobalt benzene-1,2,4,5-tetrathiolate-based polymer.
24
This
improvement in overpotential is attributed to the promotion of an alternative mechanism at more
positive potentials, dependent on the catalyst concentration, which was not observed for the sulfur-
only analogue. To probe the possible catalytic intermediates and mechanistic pathways, the
synthesis, reactivity, and H2 evolving activity of the analogous molecular analogue,
[Co(bds)2][nBu4N] (1
TBA
; where bds = benzene-1,2-diselenolate and TBA = tetrabutylammonium)
is explored here.
5.2. Results and Discussion
Complex 1
TBA
was synthesized according to a reported literature procedure.
25
Difficulties growing
X-ray quality crystals of 1
TBA
led to the isolation of the cation-exchanged [Co(bds)2][PPh4] (1
TPP
;
100
where TPP = tetraphenylphosphonium), which adopts a square planar geometry with Co-Se bond
lengths of 2.29(1) Å (Figure 5.1). In comparison, [Co(bdt)2]
–
(where bdt = benzene-1,2-dithiolate)
features Co-S bond lengths of 2.16(1) Å.
7
The UV-Vis spectrum of 1
TBA
(Figure 5.2) in acetonitrile
displays absorption bands at 340, 398, 682, and 729 nm similar to previously investigated metal
benzenedithiolates.
9,26
Figure 5.1. Solid state structure of [Co(bds)2][PPh4] (1
TPP
). Hydrogen atoms and the non-
coordinating cation are omitted for clarity.
Figure 5.2. UV-Vis absorption spectrum of [Co(bds)2]
–
in acetonitrile.
Cyclic voltammetry (CV) experiments of 1
TBA
were performed using a glassy carbon electrode
(GCE) in an acetonitrile solution of 0.1 M tetrabutylammonium hexafluorophosphate (TBAPF6)
or a 1:1 CH3CN/H2O solution of 0.1 M potassium nitrate (KNO3), due to the previously reported
enhanced activity of [Co(bdt)2]
–
in mixed organic/aqueous media.
7
All the potentials listed here
are referenced to Fc
+/0
. Complex 1
TBA
featured a reversible one electron wave assigned to the
[Co(bds)2]
–/2–
redox couple, appearing at –1.33 V in acetonitrile (Figure 5.3a) and –1.04 V in 1:1
CH3CN/H2O (Figure 5.3b). This feature occurs 60 and 30 mV more negative than the S-only
analogue, [Co(bdt)2]
–
, due to an increase in the electron density of the Se analogue.
26
Scan rate
dependence studies of the one electron redox couple associated with [Co(bds)2]
–/2–
in both CH3CN
101
and 1:1 CH3CN/H2O revealed that the cathodic and anodic peak currents increased linearly as a
function of the square root of the scan rate (Figure 5.4-5.5), as expected for a species diffusing
freely in solution.
Figure 5.3. Cyclic voltammograms of [Co(bds)2]
–
(0.5 mM) in (a) MeCN with 0.1 M TBAPF6 and
(b) 1:1 CH3CN/H2O with 0.1 M KNO3 at a scan rate of 100 mV/s. Ferrocene is used as an internal
standard. An irreversible feature was seen upon oxidation of [Co(bds)2]
–
at –0.32 V vs. Fc
+/0
in
MeCN indicating the low chemical stability of oxidized [Co(bds)2]
–
. A similar irreversible feature
upon oxidation is observed for [Co(bdt)2]
–
.
7,26
Figure 5.4. (a) Cyclic voltammograms of [Co(bds)2]
–
(0.5 mM) in MeCN with 0.1 M TBAPF6 at
scan rates ranging from 100 mV/s (red) to 2000 mV/s (purple). (b) Cathodic peak current (red) and
anodic peak current (blue) as a function of the square root of the scan rate.
102
Figure 5.5. (a) Cyclic voltammograms of [Co(bds)2]
–
(0.5 mM) in 1:1 CH3CN/H2O with 0.1 M
KNO3 at scan rates ranging from 100 mV/s (red) to 1000 mV/s (blue). (b) Cathodic peak current
(red) and anodic peak current (blue) as a function of the square root of the scan rate.
While extensive theoretical studies on [Co(bdt)2]
–
have been previously reported by others,
13
similar studies on the benzenediselenolate system have not yet been explored. We report here a
series of unrestricted DFT calculations on [Co(bds)2]
–
at the 6-31++G(d,p)/wB97XD level of
theory. The calculated reduction potential for [Co(bds)2]
–/2–
, –1.04 V, is in agreement with the
experimental value observed in 1:1 CH3CN/H2O solution (–1.04 V). Previous reports for
[Co(bdt)2]
–
suggest that the frontier orbitals (LUMO, SOMO, and SOMO-1) are substantially
influenced by the orbital overlap between the cobalt d-orbitals and the ligand p-orbitals.
27
Our
computational studies indicate similar metal-ligand orbital overlap for [Co(bds)2]
–
(Figure 5.6). As
the benzenediselenolate ligand orbitals are expected to be higher in energy than those of
benzenedithiolate, the non-bonding ligand orbitals, 2au and 1b1u, are higher in energy and introduce
increased ligand-character to the resulting frontier molecular orbitals (Figure 5.7). As such, the
[Co(bds)2]
–/2–
redox couple likely involves population of an orbital with significant ligand
character. As the SOMO-1 orbital is the lowest-lying orbital available for reduction of [Co(bds)2]
–
to [Co(bds)2]
2–
, the higher energy of this orbital is consistent with the increased reduction potential
observed for [Co(bds)2]
–
relative to [Co(bdt)2]
–
.
103
Figure 5.6. Molecular orbitals energy scheme for [Co(bds)2]
–
. The yellow box indicates Kohn-
Sham orbitals with significant ligand character, the blue box indicates metal-dominant orbitals,
and orbitals outside these boxes exhibit mixed metal-ligand character.
Figure 5.7. Relative molecular orbital energies of [Co(bds)2]
–
(blue) versus [Co(bdt)2]
–
(red)
(where bdt = benzene-1,2-dithiolate).
Hydrogen evolution studies of 1
TBA
were conducted in 0.1 M TBAPF6 acetonitrile solutions using
trifluoroacetic acid (TFA) as the proton source (pKa = 12.7 in MeCN). The addition of 2.2 mM
TFA triggered the appearance of a catalytic wave that grows from the [Co(bds)2]
–/2–
reversible
104
redox couple (Figure 5.8). As with [Co(bdt)2]
–
,
the increase in current upon addition of TFA was
small in anhydrous MeCN solutions (Figure 5.8, 5.9a).
7
Saturation at high [TFA] was not achieved
due to the formation of a black precipitate, which is further discussed below. Formation of black
particles has also been observed for related metal dithiolate complexes.
9,11,19,28
The onset of a
catalytic wave began at concentrations of 1
TBA
as low as 0.05 mM and catalytic currents varied
linearly with catalyst concentration at low acid concentrations consistent with a first order reaction
with respect to catalyst (Figure 5.9b).
Figure 5.8. Cyclic voltammograms of [Co(bds)2]
–
(0.5 mM) in MeCN with 0.1 M TBAPF6 at
varying concentrations of trifluoroacetic acid. Scan rate: 100 mV/s.
Figure 5.9. (a) Plot of icat/ip versus [TFA] taken at 100 mV/s in MeCN with 0.1 M TBAPF6 and
0.5 mM of [Co(bds)2]
–
. (b) Plot of maximum current measured at –1.35 V versus concentration of
[Co(bds)2]
–
in 0.1 M TBAPF6 in CH3CN in the presence of 2.2 mM TFA.
Because of the small current enhancements achieved upon the addition of TFA in pure organic
media, the hydrogen evolution ability of 1
TBA
was investigated in 0.1 M KNO3 solutions of 1:1
CH3CN/H2O. As seen previously in MeCN, addition of TFA triggered the appearance of a catalytic
105
wave at the reversible [Co(bds)2]
–/2–
redox couple, which is attributed to hydrogen evolution
(Figure 5.10). The catalytic peak was shifted slightly more negative upon successive TFA titrations
(Figure 5.10). Larger current enhancements were obtained in the mixed organic/aqueous media in
comparison to anhydrous MeCN. A current density of 10 mA/cm
2
was achieved at –1.14 V in the
presence of 1
TBA
(0.5 mM) and TFA (19.8 mM) (Figure 5.10-5.11). With bare glassy carbon
electrodes (GCE), negligible current densities were generated under analogous conditions (Figure
5.12). However, saturation in acid was not observed in either solvent mixture due to the formation
of a black precipitate.
Figure 5.10. Cyclic voltammograms of 1
TBA
(0.5 mM) in 1:1 CH3CN/H2O with 0.1 M KNO3 at
varying concentrations of trifluoroacetic acid. Scan rate: 100 mV/s.
Figure 5.11. Plot of icat/ip versus the concentration of trifluoroacetic acid taken at 100 mV/s in 1:1
CH3CN/H2O with 0.1 M KNO3 and 0.5 mM of [Co(bds)2]
–
.
106
Figure 5.12. Cyclic voltammograms of [Co(bds)2]
–
(0.5 mM, red) and a bare glassy carbon
electrode (GCE, black dashed) with 19.8 mM TFA in (a) MeCN with 0.1 M TBAPF6 and (b) 1:1
CH3CN/H2O with 0.1 M KNO3. Scan rate: 100 mV/s.
Controlled potential electrolysis (CPE) performed in 0.1 M KNO3 solution of 1:1 CH3CN/H2O at
–1.02 V
in the presence of 8.8 mM TFA confirmed production of H2 with a Faradaic efficiency
(FE) of 80% (Figure 5.13). Over the course of the 1 h CPE experiment, a black precipitate formed.
The black precipitate was collected after electrolysis, however, its amorphous nature and lack of
solubility in common solvents inhibited rigorous structural characterization.
Figure 5.13. Controlled potential electrolysis (CPE) of 0.5 mM [Co(bds)2]
–
(red) and GCE (black)
in 0.1 M KNO3 1:1 CH3CN/H2O solution and 8.8 mM TFA at –1.02 V. The decrease in current
over time is due to the formation of the black precipitate during the CPE experiment.
A reducing potential was not required for the formation of this precipitate as it can be generated
from only 1
TBA
and acid (TFA or [DMF(H)][OTf]) and its rate of formation was related to the acid
107
strength. The prominent absorption features of 1
TBA
observed by UV-Vis spectroscopy decreased
upon addition of TFA and a precipitate was observed in solution (Figure 5.14). The absence of
soluble [Co(bds)2]
–
in the resultant solution suggests complete consumption of [Co(bds)2]
–
(Figure
5.15).
Figure 5.14. UV-Vis absorption spectra of [Co(bds)2]
–
in DMF after the addition of TFA (15.2 μL
– 220.4 μL) displaying the disappearance of the diagnostic UV-Vis features of [Co(bds)2]
–
.
Figure 5.15. UV-Vis absorption spectra of [Co(bds)2]
–
(purple) and the remaining acetonitrile
solution (green) after treating [Co(bds)2]
–
with [DMF(H)][OTf] to form the black precipitate. The
solution was filtered before analysis to remove the black particles.
X-ray photoelectron spectroscopy (XPS) of the black precipitate revealed the presence of cobalt,
selenium, fluorine, and sulfur (Figure 5.16). Binding energies (BE) of 778.9 and 794.0 eV were
observed for the Co 2p3/2 and 2p1/2 levels (Figure 5.16). The Se region displayed two features at
108
55.2 and 56.1 eV corresponding to the 3d5/2 and 3d3/2 peaks (Figure 5.16). The F 1s peak at 688
eV and the S 2s peak at 230.5 eV are indicative of triflate anion(s) (Figure 5.16). In comparison,
XPS of 1
TBA
revealed cobalt, selenium, and nitrogen features (Figure 5.17). Co 2p3/2 and 2p1/2
peaks at 779.4 and 794.6 eV and Se 3d5/2 and 3d3/2 features at 54.6 and 55.5 eV were observed
(Figure 5.17). These shifts in BE in the Co and Se regions of 1
TBA
with respect to the black
precipitate indicate changes in the electronic states of the cobalt and selenium atoms (Figure 5.18,
Table 5.1). In the black precipitate generated from [Co(bds)2]
–
, the Co 2p features were shifted to
lower BE indicative of an increased valence electron charge for cobalt while the Se 3d features
were shifted to higher BE suggesting a decrease in the valence electron charge for selenium.
9
Additionally, the N 1s feature at 401.9 eV corresponding to the TBA cation of 1
TBA
were not
observed for the black precipitate suggesting the overall charge of the complex changed.
1
H NMR
spectroscopy studies of the reaction mixture following isolation of the black precipitate confirmed
the loss of the TBA cation, but did not suggest any ligand loss (Figure 5.19).
Figure 5.16. XPS analysis of the black precipitate formed from reaction of [Co(bds)2]
–
with
[DMF(H)][OTf]. Co 2p, Se 3d, F 1s, and S 2s core level XPS spectra are displayed.
109
Figure 5.17. X-ray photoelectron spectroscopy analysis of [Co(bds)2]
–
. Co 2p, Se 3d, and N 1s
core level XPS spectra are displayed.
Figure 5.18. X-ray photoelectron spectroscopy of the Co 2p and Se 3d regions of 1
TBA
(purple)
and the black precipitate (black).
Table 5.1. Comparison of the XPS binding energies of the Co 2p and Se 3d regions for
[Co(bds)2][nBu4N] and the black precipitate.
XPS Region
Binding Energy (eV) for
[Co(bds)2][nBu4N]
Binding Energy (eV)
for Black Precipitate
Δ Binding Energy (eV)
Co 2p3/2 779.4 778.9 0.5
Co 2p1/2 794.6 794.0 0.6
Se 3d5/2 54.6 55.2 –0.6
Se 3d3/2 55.5 56.1 –0.6
110
Figure 5.19.
1
H NMR of the resultant acetonitrile-d3 solution after treatment of [Co(bds)2]
–
with
excess [DMF(H)][OTf]. The generated black precipitate was collected by filtration, and the
resultant acetonitrile-d3 solution was analyzed by
1
H NMR spectroscopy. Tetrabutylammonium
(TBA) triflate is observed in the resultant solution, confirming the loss of TBA during the
formation of the black precipitate. Additional peaks are related to excess [DMF(H)][OTf] and
acetonitrile-d3.
Figure 5.20. FTIR spectra of [Co(bds)2]
–
(purple) and the black precipitate formed in the presence
of acid (black).
111
The reactivity of the isolated black precipitate with reducing equivalents was investigated.
Addition of excess KC8 regenerated [Co(bds)2]
–
, which was confirmed by UV-Vis spectroscopy
(Figure 5.21), indicating the cobalt diselenolate molecular unit is maintained in the black
precipitate.
Figure 5.21. UV-Vis absorption spectra of [Co(bds)2]
–
in DMF (purple) and a filtered solution
from the reaction of KC8 and black particles in DMF displaying the regeneration [Co(bds)2]
–
upon
chemical reduction of the black particles (teal).
Electrochemical reduction of the black precipitate following adsorption onto glassy carbon
electrodes was also accomplished. The electrode adsorbed black precipitate was electrochemically
reduced generating [Co(bds)2]
–
(Figure 5.22).
Figure 5.22. Cyclic voltammograms of a glassy carbon electrode soaked for 30 minutes in a
solution of [Co(bds)2]
–
and excess [DMF(H)][OTf] and a blank glassy carbon electrode (black
dashed, GCE) in 1:1 CH3CN/H2O solution with 0.1 M KNO3 (scan rate: 100 mV/s).
The large current density generated on the first scan indicated removal of the electrode-adsorbed
black precipitate upon reduction. A broad reduction feature was observed during successive CV
112
scans, which was assigned to the [Co(bds)2]
–
complex. The drop in current density for sequential
scans was attributed to diffusion of [Co(bds)2]
–
away from the electrode. Analogous
electrochemical behavior of an adsorbed black precipitate has been reported for the cobalt
benzenedithiolate complex, [Co(bdt)2]
–
.
9
Performing the chemical reduction of the black precipitate on a larger scale in a sealed vessel
allowed for sampling of the headspace of the reaction mixture by gas chromatography and revealed
formation of H2. Due to this revelation, as well as the insoluble nature of this material, we propose
that the precipitate contains oligomeric, protonated selenolate moieties, {[Co(bds)2(H)x]
x–1
}m.
Related metal selenolate and thiolate complexes have also been shown to yield black precipitates
in the presence of acids.
9,11,19,28
Several studies have identified the black precipitates as oligomeric
structures that result from the formation of metal-chalcogenide bonds between molecular
units.
9,19,28
An analogous nickel selenolate complex
19
was reported to undergo protonation of a
monodentate selenolate ligand resulting in cleavage of free selenol, which induced formation of a
black precipitate identified as oligomeric in nature. As we have used a bidentate ligand (benzene-
1,2-diselenolate) and confirmed the lack of free ligand, we propose that the monomeric cobalt
diselenolate moiety undergoes one or two protonation events (x = 1 or 2), which induces
oligomerization.
Theoretical studies of [Co(bdt)2]
–
have suggested that the first step of the HER is a one electron
reduction to form [Co(bdt)2]
2–
, followed by two protonations occurring on different sulfur moieties
of the benzenedithiolate ligands. A second one-electron reduction and a subsequent intramolecular
proton shift to form a cobalt hydride adjacent to a protonated sulfur leads to release of H2.
13
To
gain insight into the site of protonation during catalysis for [Co(bds)2]
–
, we investigated the
localization of the spin densities
29
for [Co(bds)2]
–
and [Co(bds)2]
2–
. Unrestricted DFT calculations
for [Co(bds)2]
–
indicate that the α spin-orbitals for SOMO and SOMO-1 are both occupied, while
their corresponding β spin-orbitals are vacant, as expected for the ground-state triplet species. The
spin density of [Co(bds)2]
–
is representative of the spin-localization for both the α-spin electrons
in the SOMO and SOMO-1, and a visualization of the spin density for [Co(bds)2]
–
indicates spin-
localization on both the cobalt and selenium atoms. Upon reduction of [Co(bds)2]
–
to [Co(bds)2]
2–
, the SOMO-1 β spin-orbital
becomes occupied, leading to a net spin contribution of zero from the
113
SOMO-1 (Figure 5.23). A visualization of the spin density for [Co(bds)2]
2–
exhibits a lack of spin-
localization on the selenium atoms, suggesting that the newly-introduced β-spin electron sits on a
ligand-localized orbital and that the selenium atoms carry an excess of electron density. The
analogous study on [Co(bdt)2]
–
and [Co(bdt)2]
2–
revealed identical behavior. Together, the
substantial ligand character of the frontier orbitals as well as the introduction of a β-spin to the
selenium atoms upon reduction of [Co(bds)2]
–
strongly suggest that protonation after reduction is
likely to occur at a selenium site in an analogous fashion to protonation at the sulfur site in
[Co(bdt)2]
–
.
Figure 5.23. Spin densities
29
of [Co(bds)2]
–
and [Co(bdt)2]
–
along with their one electron reduced
forms, [Co(bds)2]
2–
and [Co(bdt)2]
2–
.
However, the formation of the black precipitate from 1
TBA
and acid in the absence of reducing
equivalents indicates that reduction is not necessary for protonation of [Co(bds)2]
–
to occur.
Additionally, subsequent chemical reduction of the black precipitate releases H2 and regenerates
[Co(bds)2]
–
,
confirming the black precipitate is a viable intermediate of the catalytic cycle. To
determine if the reduced species is a viable catalytic intermediate, 1
TBA
was reduced in the presence
of KC8 forming [Co(bds)2][nBu4N]2 (2
TBA
) (Figure 5.24). Addition of TFA to a solution of 2
TBA
,
resulted in the regeneration of [Co(bds)2]
–
and the production of H2, which was confirmed by gas
chromatographic analysis.
114
Figure 5.24. UV-Vis absorption spectra of [Co(bds)2]
–
(purple) and [Co(bds)2]
2–
(pink) in DMF.
Because the black precipitate and [Co(bds)2]
2–
are competent H2 evolution catalysts, we propose
the following two mechanistic pathways (Scheme 5.1) for H2 evolution – EC and CE (E =
electrochemical, C = chemical step). Complex [Co(bds)2]
–
can be reduced by one electron to form
[Co(bds)2]
2–
followed by protonation, which most likely occurs on the selenolate moiety of the
ligand. This protonated intermediate can eliminate H2 through a bimolecular pathway, via a second
protonation, or by undergoing a second reduction and subsequent protonation.
Scheme 5.1. Two proposed mechanistic pathways for H2 evolution from [Co(bds)2]
–
in the
presence of protons and electrons (E = electrochemical, C = chemical step).
Electrochemical studies at low concentrations of weak acids indicate reduction as the first step
followed by protonation (EC) since the catalytic wave grows from the [Co(bds)2]
–/2–
redox couple.
At high concentrations of weak acids or low concentrations of strong acids, a black precipitate
forms during electrochemical testing. This precipitate can also be generated in the presence of acid
without a reducing potential. Thus, we propose that the second mechanism of H2 evolution
115
involves initial protonation of [Co(bds)2]
–
at the selenolate moiety of benzenediselenolate, which
induces the formation of an insoluble black precipitate. Although it has been challenging to
identify the exact structure of this species, we suggest that this precipitate is oligomeric
{[Co(bds)2(H)x]
x–1
}m. This isolated species is reduced to eliminate H2 and regenerate [Co(bds)2]
–
suggesting a CE mechanism.
5.3. Conclusions
In conclusion, we have synthesized [Co(bds)2]
–
and investigated its electrocatalytic hydrogen
evolving activity in organic and mixed aqueous/organic media. [Co(bds)2][nBu4N] is an active
HER catalyst in 1:1 CH3CN/H2O and 19.8 mM TFA, with current densities up to 10 mA/cm
2
at –
1.14 V. A black precipitate was formed in the presence of acid, independent of a reducing potential.
The proposed oligomeric {[Co(bds)2(H)x]
x–1
}m was found to reenter the catalytic cycle in the
presence of reducing equivalents, release H2, and regenerate [Co(bds)2]
–
. Thus, our related
mechanistic studies indicate that the black precipitate, oligomeric {[Co(bds)2(H)x]
x–1
}m, was a
viable catalytic intermediate. Furthermore, mechanistic studies show [Co(bds)2]
2–
was also a
competent catalytic intermediate for H2 evolution. Two mechanisms have been proposed for the
HER, EC and CE, which are dependent on the strength and concentration of the proton source.
5.4. Experimental Details
5.4.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity) and placed under vacuum and refilled with nitrogen (10 ×). Excluding water, all other
solvents used were degassed with nitrogen and passed through activated alumina columns and
stored over 4Å Linde-type molecular sieves. [Co(bds)2][nBu4N] (1
TBA
) (where bds = 1,2-
benzenediselenolate) was synthesized according to literature procedure.
25
Proton NMR spectra
were acquired at room temperature using a Varian 400-MR 2-Channel spectrometer and referenced
to the residual
1
H resonances of the deuterated solvent (
1
H: CD3CN, δ 1.94 ppm). All other
chemical reagents were purchased from commercial vendors and used without further purification.
116
5.4.2. Synthesis of [Co(bds)2][PPh4] (1
TPP
)
[Co(bds)2][PPh4] was synthesized following the same literature procedure as 1
TBA
.
25
Tetraphenylphosphonium chloride was used for the cation exchange rather than
tetrabutylammonium bromide. Single crystals were grown by slow diffusion of diethyl ether into
a methylene chloride solution of [Co(bds)2][PPh4].
5.4.3. Synthesis of [Co(bds)2][nBu4N]2 (2
TBA
)
KC8 (9 mg, 0.234 mmol, 6 eq.) was added to a stirred solution of 1
TBA
(30 mg, 0.039 mmol) in
THF at –30°C. The solution was allowed to warm to room temperature and stirred for 3-4 h,
resulting in a color change of the solution from blue to dark yellow. After 4 h, the solution was
filtered and [NBu4]Br (25 mg, 0.078 mmol) was added to the dark-yellow filtrate. The mixture
was stirred for ~20 h. The solution was again filtered through a microfiber filter with Celite. The
filtrate was evaporated to dryness and recrystallized from a mixture of tetrahydrofuran (THF) and
diethyl ether.
5.4.4. Synthesis of the Black Precipitate
Excess [DMF(H)][OTf] or trifluoroacetic acid (TFA) was added to a solution of 1
TBA
in
acetonitrile and was stirred overnight. The solution was filtered and the collected black precipitate
was washed with acetonitrile and acetone, and dried under vacuum.
5.4.5. Reduction of the Black Precipitate
The isolated black precipitate was suspended in DMF and the reaction flask was sealed with a
septum. A suspension of excess KC8 in DMF was added to the flask via syringe and the mixture
was allowed to stir for 30 minutes. Over the course of 30 minutes, the solution turned blue and 2
mL of the headspace of the reaction mixture was injected into the gas chromatograph confirming
H2 production.
5.4.6. Protonation of [Co(bds)2][nBu4N]2 (2
TBA
) with TFA
A solution of 2
TBA
in DMF was added to a round bottom flask and sealed with a septum.
Trifluoroacetic acid (1-2 equivalents) was added to the flask via syringe and the reaction mixture
117
was allowed to stir for 30 minutes. Over the course of 30 minutes, the solution turned blue and 2
mL of the headspace of the reaction mixture was injected into the gas chromatograph confirming
H2 production.
5.4.7. Physical Methods
UV-Vis spectra were taken using a UV-1800 Shimadzu UV spectrophotometer and quartz
cuvettes.
FT-IR spectra were acquired using a Bruker Vertex 80v spectrometer. Samples (2 mg) for analysis
were mixed into a KBr (100 mg) matrix and pressed into pellets.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray
source was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
5.4.8. Electrochemical Methods
Electrochemistry experiments were carried out using a Pine potentiostat. Cyclic voltammetry
experiments were carried out in a single compartment cell under N2 using a 3.0 mm diameter
glassy carbon electrode as the working electrode, platinum wire purchased from Alfa Aesar as the
auxiliary electrode, and silver wire as the reference electrode. Ferrocene was used as an internal
standard in all electrochemical experiments. Electrochemical experiments were carried out in
either 0.1 M TBAPF6 CH3CN solutions or 0.1 M KNO3 1:1 CH3CN/H2O solutions.
Controlled potential electrolysis experiments were carried out in a two-chambered H-cell. The first
chamber held the working and reference electrodes in 50 mL of 0.5 mM [Co(bds)2][nBu4N] in 0.1
M KNO3 1:1 CH3CN/H2O solution. The second chamber contained the counter electrode in 25 mL
of 0.1 M KNO3 1:1 CH3CN/H2O solution. The two chambers were separated from each by a fine
porosity frit. The reference electrode was placed in a separate compartment and connected by a
Vycor frit. Glassy carbon plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used
as the working and auxiliary electrodes. Ferrocene was used as an internal standard for all
118
controlled potential electrolysis experiments. Using a gas-tight syringe, 2 mL of gas were
withdrawn from the headspace of the H cell and injected into a gas chromatography instrument
(Shimadzu GC-2010-Plus) equipped with a BID detector and a Restek ShinCarbon ST
Micropacked column. To determine the Faradaic efficiency, the theoretical H2 amount based on
total charge flowed is compared with the GC-detected H2 produced from controlled-potential
electrolysis.
5.4.9. Computational Methods
All calculations were run using the Q-CHEM program package.
30
Geometry optimizations were
run with unrestricted DFT calculations at the ω-B97x-D level of theory using a relatively small 6-
31+G* basis for a low-cost analysis of the system, and were verified as stable geometries with
frequency calculations at the same level of theory. All single-point energy calculations were
carried out in the 6-31++G** basis, with additional polarization functions and further
augmentation. The ω-B97x-D functional was used throughout this study, as it provides reduced
self-interaction errors through long-range Hartree-Fock corrections and some empirical fitting for
accuracy, which is beneficial for determining the ionization potentials of transition metal-
containing systems.
Solvation effects were considered for redox potential analysis with the conductor-like screening
model (COSMO).
31,32
Solvation effects were considered for a 1:1 mixture of CH3CN/H2O (using
a dielectric constant of 50.15), as this is a convenient solvent for the dissolution and
electrochemical characterization of [Co(bds)2]
–
. The redox potential of the species was determined
through calculation of the Gibbs free energy for the gas-phase ionization process.
33
This was
achieved through a simple Hess cycle calculation using the adiabatic ionization energy of the
dianionic species (calculated by the ΔSCF procedure) as well as the solvation energies of the
oxidized and reduced species. All calculated redox potentials are presented with respect to the
ferrocene/ferrocenium couple for straightforward comparison between calculated potentials and
electrochemical measurements. The standard accepted values of 4.281 V for SHE and 0.400 V
versus SHE for the Fc/Fc
+
couple were applied for referencing of absolute potentials. Spin density
plots were generated as implemented in the QChem software package for a visualization of spin
distribution across the metal site and redox non-innocent ligands.
29
119
5.4.10. X-ray Structure Determination for 1
TPP
The X-ray intensity data were measured on a Bruker APEX DUO system equipped with a multi-
layer optics monochromator and a CuKα IuS microsource (λ = 1.54178 Å). The structure was
solved and refined using the Bruker SHELXTL Software Package, using the space group C 1 2/c
1, with Z = 4 for the formula unit, C36H28CoPSe4.
A total of 5904 frames were collected. The total exposure time was 23.02 hours. The frames were
integrated with the Bruker SAINT software package using a SAINT V8.37A (Bruker AXS, 2013)
algorithm. The integration of the data using a monoclinic unit cell yielded a total of 36109
reflections to a maximum θ angle of 68.28° (0.83 Å resolution), of which 2760 were independent
(average redundancy 13.083, completeness = 94.7%, Rint = 6.54%, Rsig = 2.90%) and 2538
(91.96%) were greater than 2σ(F
2
). The final cell constants of a = 16.1283(8) Å, b = 12.1688(6)
Å, c = 16.3580(8) Å, β = 96.567(3)°, volume = 3189.4(3) Å
3
, are based upon the refinement of the
XYZ-centroids of 9947 reflections above 20 σ(I) with 10.30° < 2θ < 136.4°. Data were corrected
for absorption effects using the multi-scan method (SADABS). The ratio of minimum to maximum
apparent transmission was 0.726. The calculated minimum and maximum transmission
coefficients (based on crystal size) are 0.2970 and 0.7000.
The structure was solved and refined using the Bruker SHELXTL Software Package, using the
space group C 1 2/c 1, with Z = 4 for the formula unit, C 36H28CoPSe4. The final anisotropic full-
matrix least-squares refinement on F
2
with 192 variables converged at R1 = 2.32%, for the observed
data and wR2 = 5.63% for all data. The goodness-of-fit was 1.037. The largest peak in the final
difference electron density synthesis was 0.369 e
–
/Å
3
and the largest hole was –0.376 e
–
/Å
3
with
an RMS deviation of 0.076 e
–
/Å
3
. On the basis of the final model, the calculated density was 1.804
g/cm
3
and F(000), 1688 e
–
.
120
Table 5.2. Crystal data and structure refinement for 1
TPP
.
Chemical formula C 36H 28CoPSe 4
Formula weight 866.32 g/mol
Temperature 100(2) K
Wavelength 1.54178 Å
Crystal size 0.038 × 0.070 × 0.159 mm
Crystal habit dark brown blade
Crystal system monoclinic
Space group C 1 2/c 1
Unit cell dimensions a = 16.1283(8) Å α = 90°
b = 12.1688(6) Å β = 96.567(3)°
c = 16.3580(8) Å γ = 90°
Volume 3189.4(3) Å
3
Z 4
Density (calculated) 1.804 g/cm
3
Absorption coefficient 10.100 mm
–1
F(000) 1688
Diffractometer Bruker APEX DUO
Radiation source IuS microsource, CuKα
Theta range for data collection 4.56 to 70.32°
Index ranges –19 ≤ h ≤ 19, –14 ≤ k ≤ 14, –19 ≤ l ≤ 19
Reflections collected 41906
Independent reflections 3036 [R(int) = 0.0653]
Coverage of independent reflections 99.7%
Absorption correction multi-scan
Max. and min. transmission 0.7000 and 0.2970
Structure solution technique direct methods
Structure solution program SHELXTL XT 2014/5 (Bruker AXS, 2014)
Refinement method Full-matrix least-squares on F
2
Refinement program SHELXTL XL 2014/7 (Bruker AXS, 2014)
Function minimized Σ w(F o
2
– F c
2
)
2
Data / restraints / parameters 3039 / 0 / 192
Goodness-of-fit on F
2
1.038
Δ/σmax 0.002
Final R indices 2799 data; I > 2σ(I) R 1 = 0.0234, wR 2 = 0.0551
all data R 1 = 0.0268, wR 2 = 0.0569
Weighting scheme
w=1/[σ
2
(F o
2
)+(0.0295P)
2
+4.0287P]
where P=(F o
2
+2F c
2
)/3
Largest diff. peak and hole 0.344 and –0.402 eÅ
–3
R.M.S. deviation from mean 0.077 eÅ
–3
121
Table 5.3. Bond lengths (Å) for 1
TPP
.
C1-C2 1.388(3) C1-C6 1.402(3)
C1-Se1 1.908(2) C2-C3 1.394(3)
C2-Se2 1.912(2) C3-C4 1.383(4)
C3-H3 0.95 C4-C5 1.387(4)
C4-H4 0.95 C5-C6 1.386(4)
C5-H5 0.95 C6-H6 0.95
C7-C12 1.388(3) C7-C8 1.398(3)
C7-P1 1.796(2) C8-C9 1.386(3)
C8-H8 0.95 C9-C10 1.385(4)
C9-H9 0.95 C10-C11 1.382(4)
C10-H10 0.95 C11-C12 1.392(4)
C11-H11 0.95 C12-H12 0.95
C13-C18 1.395(3) C13-C14 1.397(3)
C13-P1 1.794(2) C14-C15 1.388(3)
C14-H14 0.95 C15-C16 1.383(4)
C15-H15 0.95 C16-C17 1.387(4)
C16-H16 0.95 C17-C18 1.390(3)
C17-H17 0.95 C18-H18 0.95
Co1-Se2 2.2868(2) Co1-Se2 2.2869(2)
Co1-Se1 2.2921(2) Co1-Se1 2.2922(2)
P1-C13 1.794(2) P1-C7 1.796(2)
Table 5.4. Bond angles (°) for 1
TPP
.
C2-C1-C6 119.6(2) C2-C1-Se1 120.54(17)
C6-C1-Se1 119.81(19) C1-C2-C3 120.1(2)
C1-C2-Se2 119.31(17) C3-C2-Se2 120.64(17)
C4-C3-C2 120.0(2) C4-C3-H3 120.0
C2-C3-H3 120.0 C3-C4-C5 120.3(2)
C3-C4-H4 119.9 C5-C4-H4 119.9
C4-C5-C6 120.6(2) C4-C5-H5 119.9
C6-C5-H5 119.9 C5-C6-C1 119.9(2)
C5-C6-H6 120.1 C1-C6-H6 120.1
C12-C7-C8 120.6(2) C12-C7-P1 120.90(18)
C8-C7-P1 117.89(17) C9-C8-C7 119.3(2)
C9-C8-H8 120.4 C7-C8-H8 120.4
C10-C9-C8 120.1(2) C10-C9-H9 120.0
C8-C9-H9 120.0 C11-C10-C9 120.6(2)
C11-C10-H10 119.7 C9-C10-H10 119.7
C10-C11-C12 119.9(2) C10-C11-H11 120.0
122
C12-C11-H11 120.0 C7-C12-C11 119.4(2)
C7-C12-H12 120.3 C11-C12-H12 120.3
C18-C13-C14 120.4(2) C18-C13-P1 119.47(17)
C14-C13-P1 119.07(18) C15-C14-C13 119.5(2)
C15-C14-H14 120.2 C13-C14-H14 120.2
C16-C15-C14 119.9(2) C16-C15-H15 120.0
C14-C15-H15 120.0 C15-C16-C17 120.7(2)
C15-C16-H16 119.7 C17-C16-H16 119.7
C16-C17-C18 120.1(2) C16-C17-H17 120.0
C18-C17-H17 120.0 C17-C18-C13 119.4(2)
C17-C18-H18 120.3 C13-C18-H18 120.3
Se2-Co1-Se2 180.0 Se2-Co1-Se1 92.008(8)
Se2-Co1-Se1 87.991(9) Se2-Co1-Se1 87.991(9)
Se2-Co1-Se1 92.010(9) Se1-Co1-Se1 179.999(12)
C13-P1-C13 111.38(15) C13-P1-C7 112.37(10)
C13-P1-C7 105.95(10) C13-P1-C7 105.94(10)
C13-P1-C7 112.37(10) C7-P1-C7 108.92(15)
C1-Se1-Co1 103.78(7) C2-Se2-Co1 104.30(7)
5.5. References
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(2) Seh, Z. W.; Kibsgaard, J.; Dickens, C. F.; Chorkendorff, I.; Nørskov, J. K.; Jaramillo, T.
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R. J. Am. Chem. Soc. 2011, 133, 15368-15371.
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2011, 2, 2593-2597.
124
CHAPTER 6
Understanding Variability in the H
2
Evolving Activity of
Dithiolene-based Coordination Polymers
A portion of this chapter has appeared in print:
Downes, C. A.; Marinescu, S. C. “Understanding Variability in the Hydrogen Evolution Activity
of a Cobalt Anthracenetetrathiolate Coordination Polymer.” ACS Catal. 2017, 7, 8605-8612.
125
6.1. Introduction
Sustainable hydrogen production from water splitting (2H2O → 2H2 + O2) has emerged as a
promising pathway for the storage and conversion of renewable energy resources. Solar energy, in
particular, is an attractive alternative to fossil fuels owning to its unparalleled abundance. The
ability to store solar energy in H2, a clean and carbon-neutral energy carrier, as a method to mitigate
its intermittent nature, is vital for the implementation of a sustainable alternative to the fossil-fuel
dominated economy.
1
The hydrogen evolution reaction (HER, i.e., 2H
+
+ 2e
-
→ H2) requires the
development of inexpensive catalysts capable of achieving high current densities at low
overpotentials.
2,3
The replacement of Pt-based catalysts with earth-abundant materials is necessary
for large-scale application of H2-based technologies.
2,3
The immobilization of molecular catalytic units through adsorption
4-6
, covalent attachment
7-9
, or
non-covalent interactions
10-16
have emerged as useful strategies for combining the attractive
attributes of both homogeneous and heterogeneous electrocatalysis in an effort to replace noble
metals in solar-to-fuel converting devices.
17,18
Heterogenizing molecular catalysts improves the
stability and durability of these systems while maintaining well-defined catalytic units. Recent
work in our laboratory has utilized dinucleating
11-13
and trinucleating
14
ligand scaffolds to
synthesize coordination polymers (CPs) through an interfacial reaction allowing for the
immobilization of catalytic units via CPs. Initial investigations focused on CPs with a cobalt
dithiolene catalytic unit
11,14
because it had been well-established that cobalt bis(benzene-1,2-
dithiolene) was an efficient HER catalyst.
19,20
These investigations were expanded to include
replacement of the cobalt center with other earth-abundant metals such as iron and nickel
12
and
modifications to the ligand backbone by replacing sulfur with selenium generating diselenolate-
based CPs.
13
Electrochemical studies of dithiolate- and diselenolate-based CPs in acidic aqueous media revealed
these materials were active HER catalysts that could operate over a wide range of overpotentials
depending on the catalyst loading. The cobalt dithiolene CPs based on benzenehexathiolate,
14
triphenylene-2,3,6,7,10,11-hexathiolate,
14
and benzene-1,2,4,5-tetrathiolate
11
exhibited
overpotentials of 340, 530, and 560 mV to reach 10 mA/cm
2
of HER activity at catalyst loadings
of 7.0 × 10
-7
, 11 × 10
-7
, and 5.5 × 10
-7
molCo/cm
2
, respectively. For the cobalt bezenetetraselenolate
126
CP, a range of overpotentials were measured (602-343 mV) as the catalyst loading was increased
from 3.7 × 10
-7
to 9.2 × 10
-7
molCo/cm
2
.
13
However, it was observed that very small changes in
catalyst loading could lead to drastic reductions in the overpotential and even samples with the
same catalyst loading exhibited different overpotentials. The aim of our present investigation is to
identify the origins of the large variance in achievable overpotentials for dithiolene-based CPs as
bulk catalyst loading, which we have previously used to explain differences in the HER activity
of these systems, may not be the dominating factor in the resultant overpotential.
Here, a novel cobalt dithiolene coordination polymer (CoATT) based on 9,10-dimethyl-2,3,6,7-
anthracenetetrathiolate was synthesized and electrochemically characterized on both glassy carbon
and graphite electrodes to identify the origins of the large variability in the overpotential.
Numerous glassy carbon and graphite electrodes modified with CoATT were characterized by
capacitance measurements, electrochemical impedance spectroscopy studies, and Tafel analyses.
The role of electrochemically accessible active sites, electron and charge transfer, and electrical
integration between the catalyst and the support were evaluated and used to develop strategies for
optimizing the HER activity of dithiolate- and diselenolate-based coordination polymers.
Scheme 6.1. Synthesis of 9,10-dimethyl-2,3,6,7-anthracenetetrathiolate-based cobalt coordination
polymer (CoATT) through a liquid-liquid interfacial reaction.
6.2. Results and Discussion
CoATT was synthesized using a liquid-liquid interfacial reaction modified from the reported
synthetic procedure for the cobalt benzene-1,2,4,5-tetraselenolate CP.
13
9,10-Dimethyl-2,3,6,7-
anthracenetetra(thioacetate) (ATTAc4) was prepared according to the reported procedure
21
and the
acetyl protecting group was removed in the presence of base, such as NaOH, at 65°C to form
127
sodium ATT. An acetonitrile/ethyl acetate solution of [Co(MeCN)6][BF4]2 was gently layered on
top of the aqueous solution of sodium ATT. The organic solvents were allowed to evaporate over
several hours leaving a film at the gas-liquid interface, which was deposited on the support of
interest. The FT-IR spectrum (Figure 6.1) of CoATT showed the disappearance of the strong C=O
stretch of the acetyl protecting group around ~1700 cm
-1
of ATTAc4 indicating successful
deprotection prior to the formation of CoATT.
Figure 6.1. FT-IR spectra of 9,10-dimethyl-2,3,6,7-anthracenetetra(thioacetate) (teal) and CoATT
(pink).
Top down scanning electron microscopy images of CoATT (Figure 6.2) show the film-like nature
resultant from the deposition method presented here. Morphological differences on the edges of
the cracked film may influence the electrochemical behavior, which is further discussed below.
Figure 6.2. Top down scanning electron microscopy images of CoATT.
X-ray photoelectron spectroscopy (XPS) studies of CoATT revealed the presence of cobalt, sulfur,
and sodium (Figure 6.3). Deconvolution of the cobalt 2p region generated four peaks at binding
energies (BE) of 778.7, 781.7, 793.9, and 796.6 eV with the lower BE peaks corresponding to the
128
2p3/2 levels and the higher BE peaks corresponding to the 2p1/2 levels (Figure 6.3). Analogous
cobalt features have been observed for the previously reported cobalt dithiolene coordination
polymers
11,14
and surface adsorbed cobalt dithiolene molecular catalysts.
4
Additional features at
BEs of 1071.6, 227.2, and 162.8 eV correspond to the Na 1s, S 2s, and S 2p regions (Figure 6.3).
Figure 6.3. X-ray photoelectron spectroscopy analysis of CoATT showing the Co 2p, Na 1s, S
2s, and S 2p core level XPS spectra.
To investigate electrocatalytic hydrogen evolution from CoATT, we performed detailed
electrochemical measurements in a three-electrode configuration with N2-saturated pH 1.3 H2SO4
electrolyte, where all of the potentials here are referenced to the reversible hydrogen electrode
(RHE). Initial electrochemical testing was performed using glassy carbon electrodes (GCE)
modified with CoATT, denoted GCE-1, GCE-2, GCE-3, and GCE-4. Polarization curves of GCE-
1, GCE-2, GCE-3, and GCE-4 are shown in Figure 6.4. All polarization curves were compensated
for Ohmic drop (see Experimental for details). To compare the HER activities exhibited by
different catalysts, the overpotential (η) needed to reach a current density of 10 mA/cm
2
was
evaluated.
22
The measured overpotentials ranged from 445-571 mV (Table 6.1).
129
Figure 6.4. Polarization curves for GCE-1 (red), GCE-2 (green), GCE-3 (blue), and GCE-4
(purple). All measurements were carried out in N2-saturated pH 1.3. Scan rate: 100 mV/s.
To explore the origin of the large range of overpotentials for CoATT, the double layer capacitance
(Cdl), which can be used to estimate the electrochemically active surface area (ECSA), was
measured via cyclic voltammetry (Figures 6.5). Current response in the potential window 0.1-0.2
V versus RHE at different scan rates (20-150 mV/s) should be due only to double-layer charging
and discharging. The capacitance can be calculated from the scan rate dependence of the charging
current density at 0.15 V versus RHE, where the slope of Δ j versus scan rate plot is twice the Cdl.
The measured Cdl values (Figure 6.6, Table 6.1) revealed that GCE-1, which operates at the lowest
overpotential, has the highest accessible surface area. Increased Cdl corresponds to a reduction in
the overpotential to reach 10 mA/cm
2
(Figure 6.7). However, such small changes in Cdl from 5.77
to 5.81 mF/cm
2
for GCE-4 and GCE-3 did not alone explain the 67 mV reduction in the
overpotential.
Figure 6.5. Cyclic voltammograms in pH 1.3 at scan rates ranging from 20, 40, 60, 100, and 150
mV/s in the region of 0.1-0.2 V versus RHE for (a) GCE-1 (b) GCE-3 and (c) GCE-4.
130
Figure 6.6. Current density difference (Δj = ja – jc) at 0.15 V versus RHE plotted against the scan
rate (squares) for GCE-1 (red), GCE-3 (blue), and GCE-4 (purple); the Cdl values are estimated
through linear fitting of the plots (dashed lines). All measurements were carried out in N2-saturated
pH 1.3 H2SO4 solutions at room temperature.
Figure 6.7 Double layer capacitance (Cdl) measured from cyclic voltammetry versus the
overpotential to reach 10 mA/cm
2
of HER activity for GCE-1 (red), GCE-3 (blue), and GCE-4
(purple).
Table 6.1. Overview of CoATT-modified glassy carbon electrodes.
Sample Cdl (mF/cm
2
) Rct (Ω) η
10 mA/cm
2 (mV)
GCE-1 8.60 714.6 445
GCE-2 – 853.3 461
GCE-3 5.81 1061 504
GCE-4 5.77 1365 571
To further probe the variance in overpotentials, electrochemical impedance spectroscopy (EIS)
was performed to investigate the HER electrode kinetics and interfacial properties. Nyquist plots
of GCE-1, GCE-2, GCE-3, and GCE-4 recorded at -0.57 V versus RHE are shown in Figure 6.8.
131
Figure 6.8. Nyquist plots (markers) with respective fits (solid line) measured at -0.57 V versus
RHE for GCE-1 (red), GCE-2 (green), GCE-3 (blue), and GCE-4 (purple). All measurements were
carried out in N2-saturated pH 1.3 H2SO4 solutions at room temperature.
Only one semicircle was observed in each Nyquist plot suggesting the catalytic reaction was
characterized by one time constant. The absence of Warburg impedance indicated that mass
transport was rapid enough so the reaction was kinetically controlled. Therefore, the equivalent
circuit based on one time constant, 1T, was used to fit the EIS data (Figure 6.9). Rs is attributed to
the uncompensated solution resistance, Rct is the charge transfer resistance related to the kinetics
of electrocatalysis, and CPE is a constant phase element which represents the double-layer
capacitance under HER conditions. To obtain a satisfactory simulation of the experimental EIS
data, it was necessary to use CPE rather than a capacitor in the equivalent circuit. The use of a
CPE is required to explain depressed semicircles because of uneven charging of the double layer
due to microscopic surface roughness and inhomogeneity.
23-26
Figure 6.9. Equivalent circuit models used to explain the EIS response associated with the HER:
(a) one-time constant model (1T), (b) two-time constant parallel model (2TP), and (c) two-time
constant serial model (2TS).
132
The extracted parameters for the fits to this equivalent circuit for the four CoATT-modified GCE
are summarized in Table 6.2. The Rct values at -0.57 V versus RHE for GCE-1, GCE-2, GCE-3,
and GCE-4 were 714.6, 853.3, 1061.0, and 1365.0 Ω (Table 6.1). The lower value of R ct
corresponds to a faster reaction rate and higher catalytic activity for the HER (Figure 6.10) as
observed for the most active sample, GCE-1, which exhibits the lowest Rct, 714.6 Ω, and
overpotential, 445 mV.
Table 6.2. Values from fitting EIS data to equivalent circuit at -0.57 V vs RHE for CoATT-
modified GCE.
Sample η (V) R s (Ω) R ct (Ω) n 1 CPE 1 (F)
GCE-1 0.57 57.33 714.6 0.8527 1.109 × 10
-6
GCE-2 0.57 57.65 853.3 0.9227 1.221 × 10
-6
GCE-3 0.57 69.42 1061.0 0.881 7.917 × 10
-7
GCE-4 0.57 57.26 1365.0 0.7983 1.752 × 10
-6
Figure 6.10. Dependence of the overpotential to reach 10 mA/cm
2
of HER activity on the
calculated values of Rct (Ω) from fitted EIS data recorded at -0.57 V versus RHE for GCE-1 (red),
GCE-2 (green), GCE-3 (blue), and GCE-4 (purple).
The observed Rct and Rs (ranging from 57-70 Ω) values for CoATT on GCE are much larger than
desired for promotion of fast electron transfer and high HER activity. The large R s values may
indicate poor electrical integration of the catalyst and the conductive glassy carbon support. The
Rs values were also observed to fluctuate significantly as a function of applied overpotential, which
is not expected. These discrepancies may result from poor physical contact between CoATT and
GCE and the use of an electrode with an ill-defined active surface area. The inadequate physical
contact is observed as CoATT is delaminated from the surface during electrochemical testing. The
poor contact results in large and variable Rs and Rct values. Additionally, the GCE used in these
133
experiments have an active area of 0.07065 cm
2
, however, the total surface area is 0.32 cm
2
with
the active glassy carbon area surrounded by a non-conductive support. The deposition method
presented here, immersing the electrode face-down through the film at the gas-liquid interface,
offers little control over how the catalyst is dispersed over the conductive and non-conductive
regions of the GCE. This lack of control over the area of deposition and the conductive nature of
metal dithiolene coordination polymers
27,28
could result in variance in the measured resistance
values. With these considerations, further electrochemical testing of CoATT was performed on
bulk graphite electrodes.
Polarization curves of six graphite electrodes (GR) modified with CoATT, denoted GR-1, GR-2,
GR-3, GR-4, GR-5, and GR-6, are presented in Figure 6.11. As with the glassy carbon electrodes,
a wide range of overpotentials to reach 10 mA/cm
2
(388-527 mV) were achieved for CoATT-
modified GR. Further exploration into the origins of this wide range of overpotentials revealed
differences in the double layer capacitance values (Figures 6.12-6.13, Table 6.3). The overpotential
necessary to reach 10 mA/cm
2
of HER activity decreased with increasing Cdl (Figure 6.14).
Figure 6.11. Polarization curves for GR-1 (red), GR-2 (orange), GR-3 (yellow), GR-4 (green),
GR-5 (blue), and GR-6 (purple). All measurements were carried out in N2-saturated pH 1.3 H2SO4
solutions at room temperature. Scan rate: 100 mV/s.
134
Figure 6.12. Cyclic voltammograms in pH 1.3 solutions at scan rates ranging from 20, 40, 60, 100,
and 150 mV/s in the region of 0.1-0.2 V versus RHE for (a) GR-2 (b) GR-3 (c) GR-4 (d) GR-5
and (e) GR-6.
Figure 6.13. The current density difference (Δj = ja – jc) at 0.15 V versus RHE plotted against the
scan rate (squares); the Cdl values are estimated through linear fitting of the plots (dashed lines).
All measurements were carried out in N2-saturated pH 1.3 H2SO4 solutions at room temperature.
135
Figure 6.14. Double layer capacitance (Cdl) measured from cyclic voltammetry versus the
overpotential to reach 10 mA/cm
2
of HER activity for GR-2 (orange), GR-3 (yellow), GR-4
(green), GR-5 (blue), and GR-6 (purple).
The double layer capacitance measurements are related to the electrochemically accessible surface
area, however, the Cdl values do not always trend with the bulk catalyst loading.
29
To examine the
relationship between Cdl and bulk catalyst loading, CoATT was removed from the surface of the
graphite electrodes through washing following electrochemical testing and the cobalt
concentration was measured via ICP-MS. Measuring the catalyst loading following
electrochemical testing and through collection of the catalyst via washing may lead to lower than
expected cobalt concentrations. These coordination polymers have been shown to delaminate from
the surface during electrochemical characterization
11,14
and the method of collection (washing)
does not guarantee complete removal of CoATT from the surface.
Table 6.3. Overview of CoATT-modified graphite electrodes.
Sample Cdl (mF/cm
2
)
Bulk Loading
(× 10
-7
molCo/cm
2
)
Rct (Ω) η
10 mA/cm
2 (mV)
GR-1 – – 13.05 388
GR-2 6.34 5.78 26.32 404
GR-3 3.63 4.81 86.27 432
GR-4 1.25 2.49 190 486
GR-5 1.21 1.17 235.7 507
GR-6 0.94 2.49 270.6 527
In Table 6.3, the ICP-MS measured cobalt concentrations are displayed. Overall, Cdl generally
increased with bulk catalyst loading (Figure 6.15a) resulting in lower overpotentials to achieve 10
mA/cm
2
at higher concentrations of CoATT (Figure 6.15b). The current density at -0.4 V and -0.5
136
V versus RHE also increased with bulk catalyst loading and electrochemically accessible surface
area (Figure 6.16). Measured catalyst loadings for GR-4 and GR-6 were both 2.49 × 10
-7
molCo/cm
2
, but the Cdl values were 1.25 and 0.94 mF/cm
2
, respectively, indicating
electrochemically accessible surface areas can differ even for samples with the same bulk loading.
These differences may result from the varying morphologies observed for CoATT at the edges of
the deposited film (Figure 6.2).
Additionally, small changes in bulk catalyst loading did not always result in small changes in Cdl
values. For GR-3 and GR-2, increases in catalyst loading from 4.81 × 10
-7
to 5.78 × 10
-7
molCo/cm
2
led to Cdl values of 3.63 and 6.34 mF/cm
2
, respectively. However, for GR-5 and GR-4, the increase
in catalyst loading from 1.17 × 10
-7
to 2.49 × 10
-7
molCo/cm
2
only resulted in a Cdl change of 1.21
to 1.25 mF/cm
2
. Therefore, while Cdl increases with catalyst loading, the magnitude of the increase
in Cdl is not correlated to the magnitude of the increase in catalyst loading as other factors such as
morphology influence the number of accessible active sites. Conversion of the current density in
the polarization curves (Figure 6.11) to current per moles of cobalt (Figure 6.17) revealed the
catalytic current trends similarly when analyzed as a function of geometric surface area or cobalt
concentration. Double layer capacitance and bulk catalyst loading measurements were not
performed for GR-1.
Figure 6.15. (a) Cobalt concentration (molCo/cm
2
) measured via ICP versus the Cdl (mF/cm
2
)
measured via cyclic voltammetry and (b) cobalt concentration (molCo/cm
2
) measured via ICP
versus the overpotential to reach 10 mA/cm
2
of HER activity for GR-2 (orange), GR-3 (yellow),
GR-4 (green), GR-5 (blue), and GR-6 (purple).
137
Figure 6.16. Current density at -0.40 V versus RHE versus (a) cobalt concentration (molCo/cm
2
)
and (b) the Cdl (mF/cm
2
) and current density at -0.50 V versus RHE versus (c) cobalt concentration
(molCo/cm
2
) and (d) the Cdl (mF/cm
2
) for GR-2 (orange), GR-3 (yellow), GR-4 (green), GR-5
(blue), and GR-6 (purple).
Figure 6.17. Polarization curves for GR-2 (orange), GR-3 (yellow), GR-4 (green), GR-5 (blue),
and GR-6 (purple) where the y-axis represents current per concentration of cobalt measured via
ICP (mA/molCo × 10
5
).
EIS measurements were performed over a range of overpotentials (170-570 mV) and the
corresponding Nyquist and Bode plots for GR-2, GR-4, GR-5, and GR-6 are shown in Figures
6.18-6.21. The modified graphite electrodes exhibit comparable EIS responses indicating similar
mechanisms for the HER (Figure 6.22). The appearance of a depressed semicircle at high
frequencies (Figure 6.22b), in addition to the low frequency semicircle, required the use of an
138
equivalent circuit with two-time constants as opposed to the 1T model used to fit the GCE data.
To fit this data, the two-time constant parallel (2TP) and the two-time constant serial (2TS) models
were employed (Figure 6.9).
23,24,30
Figure 6.18. (a) Nyquist plot of GR-2 showing experimental EIS response (markers) and fits (solid
lines) and (b,c) Bode plots showing EIS response of GR-2 at various overpotentials at pH 1.3.
Figure 6.19. (a) Nyquist plot of GR-4 showing experimental EIS response (markers) and fits (solid
lines) and (b,c) Bode plots showing EIS response of GR-4 at various overpotentials at pH 1.3.
Figure 6.20. (a) Nyquist plot of GR-5 showing experimental EIS response (markers) and fits (solid
lines) and (b,c) Bode plots showing EIS response of GR-5 at various overpotentials at pH 1.3.
139
Figure 6.21. (a) Nyquist plot of GR-6 showing experimental EIS response (markers) and fits (solid
lines) and (b,c) Bode plots showing EIS response of GR-6 at various overpotentials at pH 1.3.
For the 2TP model, both time constants change with overpotential and are related to the kinetics
of the HER. The two-time constants (R1-CPE1 and R2-CPE2) are attributed to the adsorption of
hydrogen on the electrode surface
23,24,30,31
and the charge transfer kinetics, respectively. The 2TS
model has been used to describe the HER response on porous electrodes, with only one-time
constant at low frequencies, (Rct-CPE2), related to charge transfer kinetics, changing as a function
of overpotential. The second time constant (R1-CPE1) at high frequencies is overpotential
independent and has been assigned to several phenomena such as the porosity of the
electrode
23,30,32-34
or the interfacial resistance resulting from electron transfer between the catalyst
and the electrode.
35-37
Both the 2TP and the 2TS models were used to fit the experimental EIS data,
however, the Bode plots indicated the depressed semicircle observed at high frequencies was
overpotential independent proving the 2TS model was suitable (Figures 6.18-6.21).
The values from this fit are summarized in Tables 6.4-6.9. Rs and R1 values were unaffected by
changes in the overpotential (Figure 6.23) and the Rs values were significantly reduced (10-12 Ω)
in comparison to the recorded Rs values on GCE (57-70 Ω) indicating enhanced electrical
integration between the catalyst and electrode. The Rct values decreased with increased
overpotential (Figure 6.24) corresponding to a faster reaction rate. Additionally, as the
overpotential was increased, the Bode plots for the CoATT-modified GR displayed a reduction in
the lower frequency semicircle, resulting from an increase in the reaction rate, and a shift of the
lower frequency peak to higher frequencies, indicating a shorter reaction time constant (Figures
6.18-6.22).
140
Figure 6.22. EIS spectra measured at -0.57 V versus RHE in pH 1.3 solutions presented as (a)
Nyquist plots (markers) with respective fits (solid line), (b) expansion of the high frequency range
of the Nyquist plots, and (c,d) Bode plots for GR-1 (red), GR-2 (orange), GR-3 (yellow), GR-4
(green), GR-5 (blue), and GR-6 (purple).
Table 6.4. Values from fitting EIS data to equivalent circuit shown in Figure 6.22 for GR-1.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 11.04 3.479 0.7599 2.606 × 10
-4
13.05 0.8475 1.817 × 10
-4
Table 6.5. Values from fitting EIS data to equivalent circuit shown in Figure 6.18 for GR-2.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 11.04 2.125 0.7667 2.77 × 10
-4
26.32 0.891 7.854 × 10
-5
0.47 10.42 2.638 0.726 3.894 × 10
-4
27.15 0.9172 7.079 × 10
-5
0.37 10.03 2.421 0.753 3.687 × 10
-4
74.26 0.9139 7.908 × 10
-5
0.17 9.947 0.5087 0.9408 1.149 × 10
-4
607.2 0.7583 2.536 × 10
-4
Table 6.6. Values from fitting EIS data to equivalent circuit shown in Figure 6.22 for GR-3.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 11.81 3.611 0.767 2.14 × 10
-4
86.27 0.8822 5.23 × 10
-5
141
Table 6.7. Values from fitting EIS data to equivalent circuit shown in Figure 6.19 for GR-4.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 11.53 3.725 0.7838 1.478 × 10
-4
191.2 0.9385 5.882 × 10
-5
0.47 11.46 3.242 0.8204 1.28 × 10
-4
192.1 0.9411 5.777 × 10
-5
0.37 11.37 3.348 0.7939 1.635 × 10
-4
1245 0.9437 5.718 × 10
-5
0.17 11.44 2.519 0.8945 6.339 × 10
-5
12050 0.9132 7.24 × 10
-5
Table 6.8. Values from fitting EIS data to equivalent circuit shown in Figure 6.20 for GR-5.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 10.94 2.221 0.8628 1.301 × 10
-4
235.7 0.934 5.318 × 10
-5
0.47 10.75 1.943 0.9012 1.082 × 10
-4
266.5 0.9351 5.194 × 10
-5
0.37 10.69 1.873 0.9073 1.012 × 10
-4
1738 0.9342 5.147 × 10
-5
0.27 10.67 1.417 0.9697 5.462 × 10
-5
7019 0.915 5.872 × 10
-5
Table 6.9. Values from fitting EIS data to equivalent circuit shown in Figure 6.21 for GR-6.
η (V) R s (Ω) R 1 (Ω) n 1 CPE 1 (F) R ct (Ω) n 2 CPE 2 (F)
0.57 12.01 2.71 0.8426 1.114 × 10
-4
270.6 0.9352 4.424 × 10
-5
0.47 11.91 2.271 0.8994 6.723 × 10
-5
294.7 0.9288 4.494 × 10
-5
0.37 11.85 2.573 0.8528 1.05 × 10
-4
1847 0.9362 4.254 × 10
-5
0.27 11.81 2.251 0.8849 7.359 × 10
-5
8227 0.9238 4.563 × 10
-5
0.17 11.93 1.441 1.0 2.568 × 10
-5
13740 0.8969 5.381 × 10
-5
Figure 6.23. Values for Rs (blue) and R1 (red) from the EIS data in Tables 6.5, 6.7, 6.8, and 6.9
for (a) GR-2 (b) GR-4 (c) GR-5 (d) GR-6.
142
Figure 6.24. Values for Rct from the EIS data in Tables 6.5, 6.7, 6.8, and 6.9 for (a) GR-2 (b) GR-
4 (c) GR-5 (d) GR-6.
Comparison of the EIS response for all six graphite modified electrodes at -0.57 V versus RHE in
Figure 6.22 clearly demonstrates that a lower Rct is related to improved HER performance. Rct
values of 13.05, 26.32, 86.27, 191.2, 235.7 and 270.6 Ω were found for GR-1, GR-2, GR-3, GR-
4, GR-5, and GR-6, respectively. These Rct values are much lower than those observed on GCE
indicating much faster electron transfer and HER kinetics. The electrode with the lowest Rct, GR-
1, exhibits the smallest η
10 mA/cm
2
of 388 mV. A linear relationship between Rct and η
10 mA/cm
2 was
observed (Figure 6.25) confirming a lower Rct is necessary for high HER activity. The Bode plots
(Figure 6.22b,c) also indicated an increased reaction rate as the lower frequency semicircle was
reduced in intensity and shifted to higher frequencies as the HER performance was improved from
GR-6 (η
10 mA/cm
2 = 527 mV) to GR-1 (η
10 mA/cm
2 = 388 mV).
143
Figure 6.25. Calculated values of Rct (Ω) from fitted EIS data recorded at -0.57 V versus RHE
versus the overpotential to reach 10 mA/cm
2
of HER activity for GR-1 (red), GR-2 (orange), GR-
3 (yellow), GR-4 (green), GR-5 (blue), and GR-6 (purple).
As we have previously discussed, increasing the Cdl values of the CoATT-modified GR led to
decreased overpotentials, which is expected for a higher number of electrochemically accessible
active sites. However, determining the Cdl using cyclic voltammetry experiments measured in the
non-faradaic region can result in an underestimation or overestimation of the catalytically active
area. Such phenomena as surface coordination and intercalation of ions and potential dependent
conductivities can result in the measurement of capacitance that is not related to the
electrochemically accessible surface area of the catalyst.
23,26,38-40
Therefore, the double-layer
capacitance under HER conditions (C
dl
*
) was extrapolated from the EIS data to provide a second
measure of the accessible surface area to confirm the results obtained from the CV experiments.
C
dl
*
values from EIS experiments performed at lower overpotentials revealed higher surface areas
for the more active electrodes (Figure 6.26) analogous to the trend observed for the Cdl values
measured from the CV experiments. For the less active samples, the C
dl
*
values were constant with
overpotential. However, for the more active samples, C
dl
*
decreased with increasing overpotential
because of surface adsorption of hydrogen occluding the catalyst surface (Figure 6.26). This is
clearly seen as the C
dl
*
for GR-2 decreased at more negative overpotentials.
144
Figure 6.26. Double layer capacitance measured under HER conditions (C
dl
*
) and calculated from
the EIS data collected at the following overpotentials (a) 570 mV (b) 470 mV (c) 370 mV and (d)
170 mV versus the overpotential to reach 10 mA/cm
2
of HER activity for GR-1 (red), GR-2
(orange), GR-3 (yellow), GR-4 (green), GR-5 (blue), and GR-6 (purple).
The Tafel plots for GR-2, GR-3, and GR-5 are displayed in Figure 6.27. Tafel analysis can be used
to determine different mechanistic pathways with the Tafel slope dependent on the rate-limiting
step. In acidic electrolytes, three reaction steps are possible for the HER: the Volmer (discharge)
reaction (H3O
+
+ e
-
→ Hads + H2O, b = 120 mV/dec), the Heyrovsky (ion + atom) reaction (H3O
+
+ e
-
+ cat-H → cat + H2 + H2O, b = 40 mV/dec), and the Tafel (combination) reaction (cat-H + cat-
H → 2cat + H2, b = 30 mV/dec).
41
The linear portions of the Tafel plots were fitted to the Tafel
equation, η = b log j + a (where η is the overpotential, a the Tafel constant, b the Tafel slope, and
j the current density), yielding Tafel slopes of 130, 134, and 144 mV/dec for GR-2, GR-3, and GR-
5. The decrease in Tafel slope was correlated with a reduction in the Rct measured by EIS and the
overpotential to reach 10 mA/cm
2
resulting in improved HER kinetics and electron transport
through the material.
42
The large Tafel slopes for the CoATT-modified GR suggest a rate-limiting
Volmer-like step, which has been previously observed for related surface immobilized metal
dithiolene catalysts.
6,11,12,14
Theoretical studies
43
on cobalt dithiolene complexes have proposed
145
that reduction and thiolate protonation are the first steps for the HER, which corresponds to the
Volmer discharge reaction. The decrease in Tafel slope was also accompanied by an increase in
Cdl with GR-2 exhibiting the highest electrochemically active surface area of the electrodes
evaluated by Tafel analysis. The reduction in Tafel slope suggests a smaller barrier for the
discharge reaction (catalyst reduction and thiolate protonation), which may be directly attributed
to the increase in the electrochemically accessible catalytic units necessary for reduction and
protonation.
Figure 6.27. Tafel analysis of GR-2 (orange), GR-3 (yellow), and GR-5 (blue). Exchange current
densities of 10
-6.35
, 10
-6.1
, and 10
-7.01
A/cm
2
were obtained for GR-2, GR-3, and GR-5. All
measurements were carried out in N2-saturated pH 1.3 H2SO4 solutions at room temperature.
To assess the stability and efficiency of CoATT, controlled potential electrolysis (CPE) and
chronoamperometry experiments were performed. Shorter duration CPE experiments performed
at -0.72 V versus RHE revealed continuous charge build up over 1 h with little loss of activity and
Faradaic efficiencies of 97% were achieved for CoATT (Figure 6.28). Longer electrolysis
experiments performed at -0.72 V versus RHE on graphite (Figure 6.29) and glassy carbon (Figure
6.30a) electrodes also demonstrated stable H2 production with oscillations in the generated current
arising from removal of adsorbed H2 from the electrode surface. Cyclic voltammetry experiments
following 7 and 12 h of CPE demonstrated no significant change in the electrochemical behavior
(Figure 6.30b). X-ray photoelectron spectroscopy analysis after 1 h (Figure 6.31) and 12 h (Figure
6.32) of CPE at -0.72 V versus RHE in pH 1.3 solutions revealed analogous Co, Na, and S features
to the ones observed before, indicating the stability of CoATT during electrochemical testing.
146
Figure 6.28. Controlled potential electrolysis of CoATT|GCE (red) and GCE (black dashed) in
pH 1.3 solution at -0.72 V versus RHE.
Figure 6.29. Chronoamperometry of CoATT|GR measured in pH 1.3 solutions at -0.72 V versus
RHE.
Figure 6.30. (a) Controlled potential electrolysis of CoATT|GCE (red) in pH 1.3 solution at -0.72
V versus RHE (b) Cyclic voltammetry experiments before CPE, after 7 h of CPE (green), and after
12 h of CPE (blue). Fluctuations in the CPE data are a result of removal of large H2 bubbles from
the electrode surface exposing more catalyst and jostling of the H-cell. Fresh pH 1.3 solution was
added after 7 h of CPE.
147
Figure 6.31. X-ray photoelectron spectroscopy analysis of CoATT after 1 h of controlled potential
electrolysis in pH 1.3 solution at -0.72 V versus RHE.
Figure 6.32. X-ray photoelectron spectroscopy analysis of CoATT after 12 h of controlled
potential electrolysis in pH 1.3 solution at -0.72 V versus RHE.
148
6.3. Conclusions
In summary, we report the synthesis of a cobalt dithiolene coordination polymer based on 9,10-
dimethyl-2,3,6,7-anthracenetetrathiolate via an interfacial reaction. The electrochemical hydrogen
evolution ability of CoATT was investigated on glassy carbon and graphite electrodes. Previous
reports on related coordination polymers revealed large variances in overpotential with small
changes in catalyst loading. Through analysis of the double layer capacitance and electrochemical
impedance spectroscopy for CoATT on glassy carbon and graphite, it was revealed that
electrochemically accessible surface area, rather than bulk catalyst loading, and charge transfer
resistance were better indicators for changes in overpotential. EIS measurements also
demonstrated large and variable Rs values (56-70 Ω) for CoATT-modified GCE suggesting poor
electrical integration between the substrate and catalyst. Graphite proved a more suitable electrode
material resulting in stable and low (~10 Ω) Rs values indicating better electrical integration. For
CoATT-modified GR systems, overpotentials ranging from 388-527 mV were achieved and this
variance in overpotential correlated well with changes in the Cdl and Rct values. Investigating the
fundamental electrochemical characteristics of the cobalt anthracenetetrathiolate CP has provided
insight into the critical factors that influence the HER activity. This understanding enables the
design of coordination polymers with improved catalytic performance and furthers the viability of
CPs as a method for immobilization of molecular catalysts for solar-to-fuel converting devices.
6.4. Experimental Details
6.4.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity) and placed under vacuum and refilled with nitrogen (10 ×). Excluding water, all other
solvents used were degassed with nitrogen and passed through activated alumina columns and
stored over 4Å Linde-type molecular sieves. The pHs of the aqueous solutions were measured with
a benchtop Mettler Toledo pH meter. 9,10-dimethyl-2,3,6,7-anthracenetetra(thioacetate) was
synthesized according to literature procedure.
21
All other chemical reagents were purchased from
149
commercial vendors and used without further purification. ICP-MS studies were performed by
Robertson MicroLit Laboratories, Ledgewood, NJ, 07852.
6.4.2. Physical Methods
FTIR spectra were acquired using a Bruker Vertex 80v spectrometer. Samples (2 mg) for
analysis were mixed into a KBr (100 mg) matrix and pressed into pellets.
Scanning electron microscopy (SEM) was performed on a JEOL JSM 7001F scanning electron
microscope.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray
source was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
6.4.3. Synthesis of CoATT
In a glove-box, NaOH (1 mL of 0.25 M aqueous solution) was added to a 2 mL methanol solution
of 9,10-dimethyl-2,3,6,7-anthracenetetra(thioacetate) (10.3 mg, 0.021 mmol). This solution was
heated at 65°C overnight. Following the overnight deprotection, 18 mL of degassed H 2O was
added to the solution. The aqueous solution was transferred to a 70 mm × 50 mm crystallizing
dish. [Co(MeCN)6][BF4]2 (36.3 mg, 0.146 mmol) was dissolved in a 1:4 mixture of acetonitrile
and ethyl acetate and 1 mL of this solution was carefully layered via glass pipette on top of the
aqueous solution to cover ~80% of the surface area. The organic solvents were allowed to
evaporate over several hours at room temperature, leaving behind CoATT as a black solid at the
gas-liquid interface. The black solid was collected and washed with water and methanol. The
resultant black powder was dried under vacuum.
6.4.4. Deposition of CoATT
Deposition was carried out by immersing the support face down through the film formed at the
gas-liquid interface. Following deposition, the substrate was washed with water and methanol.
150
6.4.5. Electrochemical Methods
Electrochemistry experiments were carried out using a VersaSTAT 3 potentiostat. Platinum wire
used for the electrochemical studies was purchased from Alfa Aesar. The electrochemical
experiments were carried out in a three-electrode configuration electrochemical cell under an inert
atmosphere using glassy carbon electrodes (GCE, 0.07065 cm
2
surface area) or graphite rods (GR,
0.32 cm
2
surface area) as the working electrode. Graphite rods were purchased from Graphite
Machining, Inc. (Grade NAC-500 Purified, < 10 ppm ash level). The graphite rods were cut to the
appropriate length and polished with sandpaper and alumina. The rod length was wrapped in
Teflon or coated in epoxy (Loctite Hysol 1C) to define a 0.32 cm
2
surface area. XPS analysis of
the polished graphite rods indicate no measurable amounts of transition metals, which could
potentially participate in proton reduction. A platinum wire, placed in a separate compartment,
connected by a Vycor tip, and filled with the electrolytic solution (0.1 M NaClO4) was used as the
auxiliary (counter) electrode. The reference electrode, placed in a separate compartment and
connected by a Vycor tip, was based on an aqueous Ag/AgCl/saturated KCl electrode. The
reference electrode in aqueous media was calibrated externally relative to ferrocenecarboxylic acid
(Fc-COOH) at pH 7.0, with the Fe
3+/2+
couple at 0.28 V versus Ag/AgCl. All potentials reported
in this paper were converted to the standard hydrogen electrode (SHE) by adding a value of 0.205
V or to the reversible hydrogen electrode (RHE) by adding a value of (0.205 + 0.059 × pH) V.
6.4.6. Double-Layer Capacitance Measurements
Cyclic voltammograms for double layer capacitance measurements were taken in a potential
window between 0.1 and 0.2 V versus RHE in pH 1.3 at scan rates of 20, 40, 60, 80, 100, and 150
mV/s. The total current density obtained from the current density difference (Δj = ja – jc) at 0.15 V
versus RHE was plotted against the scan rate. The slope is twice the value of the double layer
capacitance. For Tafel analysis, polarization curves were measured at pH 1.3 with a scan rate of 2
mV/s.
6.4.7. Electrochemical Impedance Spectroscopy
Electrochemical impedance spectroscopy (EIS) measurements were carried out at different
overpotentials ( η = 570-70 mV) in the frequency range of 500 kHz – 0.1 Hz with 10 mV sinusoidal
perturbations in pH 1.3 solutions. Experimental EIS data were analyzed and fitted with the
151
ZSimpWin software. The obtained polarization curves were corrected by the iR loss according to
the following equation:
Ecorr = Emea – iRs
Where Ecorr is the iR-corrected potential, Emea is the experimentally measured potential, and Rs is
the solution resistance extracted from the fitted EIS data.
6.4.8. Determination of Bulk Catalyst Loading
To obtain the bulk catalyst loading, CoATT was removed from the electrode surface through
washing with acetone. Collected samples were dried and sent to Robertson MicroLit Laboratories
(Ledgewood, NJ, 07852) for ICP-MS analysis for determination of the cobalt concentration.
6.4.9. Controlled Potential Electrolysis and Chronoamperometry
Controlled potential electrolysis (CPE) and chronoamperometry measurements to determine
Faradaic efficiency and study long-term stability were conducted in a sealed two-chambered H-
cell where the first chamber held the working and reference electrodes in 50 mL of 0.1 M NaClO4
(aq) pH 1.3 solution, and the second chamber held the auxiliary electrode in 25 mL of 0.1 M
NaClO4 (aq). The two chambers, which were both under N2, were separated by a fine porosity
glass frit. CPE experiments were performed with the following set-ups: (1) glassy carbon plate
electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) as the working and auxiliary electrodes
and (2) graphite rods (0.32 cm
2
surface area) as the working and auxiliary electrodes. The reference
electrode was a Ag/AgCl/saturated KCl (aq) electrode separated from the solution by a Vycor tip.
Using a gas-tight syringe, 2 mL of gas were withdrawn from the headspace of the H-cell and
injected into a gas chromatography instrument (Shimadzu GC-2010-Plus) equipped with a BID
detector and a Restek ShinCarbon ST Micropacked column. To determine the Faradaic efficiency,
the theoretical H2 amount based on total charge flowed is compared with the GC-detected H2
produced from controlled-potential electrolysis.
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154
CHAPTER 7
Electrocatalytic H
2
Evolution from Benzenehexathiolate-based Coordination
Frameworks and the Effect of Film Thickness on H
2
Production
A portion of this chapter has appeared in print:
Downes, C. A.; Clough, A. J.; Chen, K.; Yoo, J. W.; Marinescu, S. C. “Evaluation of the H2
Evolving Activity of Benzenehexathiolate Coordination Frameworks and the Effect of Film
Thickness on H2 Production.” ACS Appl. Mater. Interfaces 2017, 10, 1719-1727.
155
7.1. Introduction
To sustainably meet the rising energy demand associated with rapid global population and
economic growth, continued development and deployment of renewable energy resources is
paramount.
1
The unparalleled abundance of solar energy, while attractive to meet the massive
global energy demand, is hindered by its intermittent nature, resulting in a mismatch between
energy supply and demand.
2
Sustainable hydrogen production from water splitting (2H2O → 2H2
+ O2) has emerged as a promising pathway for the storage and conversion of solar energy.
2,3
Hydrogen is a valuable energy carrier that can be transformed into electricity using fuel cell
technology or used in the production of industrially relevant chemicals such as ammonia and
methanol.
4,5
Hydrogen is currently generated from steam methane reforming whereby four
molecules of H2 are produced per CH4; however, one molecule of CO2 is also produced.
6
A
sustainable energy future will require the carbon-neutral production of H2 and water splitting is a
viable pathway to do so. Electrocatalysts that are cost-effective and scalable must be developed
for the hydrogen evolution (HER) and oxygen evolution reactions (OER) to facilitate water
splitting at high efficiencies with minimal energy input.
Earth-abundant homogeneous
7,8
and heterogeneous
9-11
catalysts have been extensively explored as
alternatives to platinum, the benchmark electrocatalyst for the HER. Inherent problems with
homogeneous catalysts such as the lack of solubility and stability in aqueous media and the
diffusion (to the electrode surface) dependent activation of the solubilized catalyst limit their
practical feasibility. Similarly, heterogeneous catalysts have ill-defined active sites making
mechanistic understanding, rational design, and optimization difficult. These short-comings can
be overcome by combining the advantageous properties of homogeneous and heterogeneous
catalysts. The heterogenization of molecular catalysts has been accomplished through a variety of
methods
12,13
such as surface adsorption,
14-18
electropolymerization,
19-21
covalent attachment,
22-24
and incorporation into extended frameworks.
25-31
These systems can display the improved stability
and robustness associated with heterogeneous catalysts while the well-defined catalytic active sites
reminiscent of molecular systems can be extensively studied and modified to give insight into the
mechanism of catalysis.
156
Molecular catalysts incorporated into extended frameworks such as coordination polymers and
metal-organic frameworks (MOFs) have been investigated as electrocatalysts for energy
converting applications.
32-34
The use of MOFs as intrinsic electrocatalysts, however, has been
limited because of their traditionally insulating nature. The facilitation of efficient charge and
electron transfer is an important property for the promotion of high electrocatalytic activity. The
ability to deposit thin film coordination frameworks directly onto electrode surfaces has improved
their electrocatalytic performance because thin film MOFs exhibit higher conductivities than the
bulk materials.
35,36
However, most of these systems still exhibit insufficient charge transfer
properties inhibiting high electrocatalytic activity. Additionally, the weak coordination bonds in
the frameworks lead to structural instability in aqueous conditions.
The development of conductive coordination frameworks has been a breakthrough in the
application of MOFs for electrocatalytic energy conversion.
37
The use of redox active ligands has
facilitated charge transport through the frameworks and stronger metal-ligand bonds have
improved stability in aqueous acidic and/or alkaline media. The few successful electrocatalytic
MOFs displaying high activity and stability for the HER,
28-30
oxygen reduction reaction (ORR),
31
OER,
38-41
and CO2 reduction (CO2RR)
42,43
all exhibit intrinsic electrical conductivity. The
conductive nature of dithiolene-based coordination frameworks makes these systems attractive
candidates for electrocatalytic applications.
36,44-46
We have extensively explored the use of
trinucleating and dinucleating dithiolene ligand scaffolds to synthesize 2D and 1D metal dithiolene
coordination frameworks
28
and polymers,
25,26
which perform as robust electrocatalysts for the
HER.
To further investigate the variables that dictate the HER activity of dithiolene-based coordination
frameworks, we examine here the effect of the coordinated metal by performing detailed
electrochemical analysis of CoBHT, NiBHT, and FeBHT (BHT = benzenhexathiolate, Scheme
7.1). FeBHT was synthesized through an interfacial reaction previously reported in the literature
for the isolation of BHT-based coordination frameworks.
28,44
The synthesis and conductivity of
NiBHT was reported by Nishihara and coworkers,
44,45
however, the electrocatalytic HER activity
was not studied. We have previously explored the electrocatalytic HER performance of CoBHT
157
with the system displaying an overpotential of 340 mV and Faradaic efficiency of 97% in pH 1.3
aqueous solutions and a Tafel slope of 108 mV/dec (pH 2.6).
28
Scheme 7.1. Structure of the benzenehexathiolate (BHT) coordination frameworks (CoBHT,
NiBHT, and FeBHT) studied here.
A modified synthetic method was utilized here for the isolation of CoBHT and electrochemical
analysis was performed using glassy carbon electrodes (GCE). Electrochemical techniques such
as electrochemical impedance spectroscopy (EIS) and double-layer capacitance (Cdl)
measurements, which provide insight into the charge transfer/HER kinetics and an estimation of
the electrochemically active surface area, respectively, were not employed in the previous report
on CoBHT. This analysis is used here to improve the fundamental understanding of the parameters
that dictate the observed HER activity of CoBHT.
Additionally, we explore the effect of film thickness on the HER activity of CoBHT. For layered
materials, thicker films correspond to higher bulk catalyst loadings (higher number of available
active sites) leading to improved HER performance.
47,48
However, it has been observed that a
plateau can be reached whereby increases in catalyst loading do not generate higher
electrocatalytic current densities.
18,42,47,49
Poor diffusion of electrons and protons through thick
158
films limits the number of electrochemically accessible active sites even as the bulk catalyst
loading increases, thus inhibiting the HER activity.
47,49,50
The long-term durability of thicker films
can also be problematic. High mass loadings can lead to catalyst modified electrodes exhibiting
limited mechanical robustness with cracking and delamination from the electrode surface
occurring during electrochemical testing.
Previous investigation into the HER performance of monolayer and bulk powders of cobalt
dithiolene and mixed dithiolene-diamine coordination frameworks revealed increases in the
overpotential to achieve 10 mA/cm
2
and the Tafel slope as the method of catalyst deposition was
changed from a monolayer (0.8 ± 0.1 nm thickness) to bulk powder (0.12 mg catalyst prepared as
an ink).
30
These differences are expected as thin film dithiolene- and diamine-based coordination
frameworks are more conductive than the bulk materials because of a reduction in the number of
grain boundaries, which inhibit charge transfer.
35,36,44,45
Further investigation into the effect of film
thickness on the HER activity by increasing the amount of deposited catalyst from a monolayer to
multilayer films was not performed.
The influence of film thickness on the activity of intrinsic MOF electrocatalysts, although an
important parameter for optimizing catalytic performance, has not been extensively explored. To
the best of our knowledge, the most comprehensive study on the effect of MOF film thickness on
the electrocatalytic activity was performed by Yaghi, Yang, and coworkers on a cobalt porphyrin
based MOF, Al2(OH)2TCPP-Co (TCPP-H2 = 4,4’,4”,4”(porphyrin-5,10,15,20-
tetrayl)tetrabenzoate).
42
The CO2 reduction performance increased with the number of atomic
layer deposition (ALD) cycles before reaching maximum activity at 50 cycles (MOF film thickness
of ~30-70 nm). At higher loadings, the performance decreased because of poor charge and mass
transport properties.
Extensive analysis of the influence of film thickness on the HER performance of coordination
frameworks is necessary to facilitate the continued development of these systems as practical
catalysts for energy converting devices. By varying the film thickness from ~20-1000 nm, we look
to understand how efficiently charges and protons move through CoBHT. We can then determine
159
the optimal conditions to maximize the number of accessible active sites and HER activity without
increasing the resistance of the system and limiting electron and proton transfer.
7.2. Results and Discussion
7.2.1. Electrochemical Analysis of FeBHT and NiBHT
To investigate the electrocatalytic hydrogen evolving activity of CoBHT, NiBHT, and FeBHT, we
performed detailed electrochemical measurements in a three-electrode configuration with N2-
saturated pH 1.3 H2SO4 electrolyte, where all of the potentials are referenced to the reversible
hydrogen electrode (RHE). Multiple electrodes modified with the BHT-based frameworks were
prepared and electrochemically tested. We have recently shown that the deposition method used
here, immersion of the electrode through the interfacially grown film, leads to small variances in
the bulk catalyst loading even for films of the same thickness.
36
These small variances in bulk
loading and the morphological differences the deposition method presented here induces for each
sample (how the sample adheres to the substrate) can influence the measured Cdl and Rct values.
Thus, it is important to establish the trends observed for the electrocatalytic HER activity are valid
for multiple modified electrodes.
The EIS response for all three systems studied was best described by the two-time constant serial
model (2TS), where the high frequency semi-circle is related to the surface porosity of the
electrode or the contact between the electrode and the catalyst layer and the low frequency semi-
circle is related to the charge transfer kinetics of HER (Figure 7.1).
48,51-53
NiBHT and FeBHT were
readily deposited on GCE following the interfacial synthesis and electrochemically explored as
hydrogen evolving catalysts.
Figure 7.1. Equivalent circuit model used to fit the EIS data. The two-time constant serial model
(2TS) was employed where Rs is the solution resistance and Rct-CPE2 is related to the charge
transfer reaction. R1-CPE1 is unrelated to HER kinetics and remains relatively unchanged with
overpotential. R1-CPE1 has been associated with the porosity of the electrode or the contact
between the electrode and the catalyst layer.
160
FeBHT exhibits overpotentials ranging from 473-541 mV (Figure 7.2a, Table 7.1). Increases in
Cdl resulted in reduced HER activity for FeBHT, which was not expected (Figure 7.2, Table 7.1).
The rise in Cdl was accompanied by an increase in the Rct and Tafel slope (Figures 7.2-7.4).
Enhancements in the Cdl (1.4 to 2.6 mF/cm
2
) corresponded to increases in the Rct from 102.8 to
221.3 Ω (η = 470 mV) and Tafel slope from 112 to 143 mV/dec (Figure 7.3-7.4, Table 7.1). These
results indicate poor charge and proton transfer for the FeBHT samples with higher catalyst
loading. The reproducibility of the polarization curves was also limited for FeBHT. Successive
scans generated smaller overpotentials and improved charge transfer properties. This could be
correlated to the observed removal of FeBHT from the electrode surface during electrochemical
testing as we have shown that FeBHT is more active for the HER at lower catalyst loadings.
Figure 7.2. (a) Polarization curves for Fe-GCE-1 (red), Fe-GCE-2 (green), and Fe-GCE-3 (blue);
scan rate: 100 mV/s, (b) Plot showing the extraction of the double layer capacitance (Cdl). The
current density difference (Δj = ja – jc) at 0.15 V versus RHE plotted against the scan rate (markers);
the Cdl values are estimated through linear fitting of the plots (dashed lines). EIS spectra measured
at (c) -0.37 V and (d) -0.47 V versus RHE presented as Nyquist plots (markers) with respective
fits (solid line).
161
Table 7.1. Overpotential, double-layer capacitance, and charge transfer resistance values for Fe-
GCE-1, Fe-GCE-2, and Fe-GCE-3.
Sample η @ 10
mA/cm
2
(mV)
Cdl
(mF/cm
2
)
Rct (Ω) at η =
270 mV
Rct (Ω) at η =
370 mV
Rct (Ω) at η =
470 mV
Fe-GCE-1 473 1.4 1.843 × 10
4
1.309 × 10
3
102.8
Fe-GCE-2 496 2.0 2.158 × 10
4
2.484 × 10
3
134.7
Fe-GCE-3 541 2.6 5.278 × 10
4
3.603 × 10
3
221.3
Figure 7.3. Plots of Fe-GCE-1 (red), Fe-GCE-2 (green), and Fe-GCE-3 (blue) displaying (a) Cdl
values (mF/cm
2
) calculated from cyclic voltammetry measurements versus the overpotential to
reach 10 mA/cm
2
of HER activity. (b) Calculated values of Rct (Ω) from fitted EIS data recorded
at -0.47 V versus RHE versus Cdl values (mF/cm
2
).
Figure 7.4. Tafel analysis of Fe-GCE-1 (red), Fe-GCE-2 (green), and Fe-GCE-3 (blue). Exchange
current densities of 10
-6.8
, 10
-6.8
, and 10
-6.5
A/cm
2
were obtained for Fe-GCE-1, Fe-GCE-2, and Fe-
GCE-3.
162
For NiBHT, overpotentials as low as 331 mV were achieved at a Cdl value of 6.29 mF/cm
2
(Figure
7.5, Table 7.2). A broad reduction feature was observed at ~0 V for NiBHT and was assigned to
reduction of the framework (Figure 7.5a). This feature has been previously reported for NiBHT
and the related nickel benzene-1,2,4,5-tetrathiolate coordination polymer.
26,44
Higher
electrocatalytic activity was associated with larger Cdl values and reduced Rct (Table 7.2).
Figure 7.5. (a) Polarization curves for Ni-GCE-1 (red) and Ni-GCE-2 (blue); scan rate: 100 mV/s.
(b) Plot showing the extraction of the double layer capacitance (Cdl). The current density difference
(Δj = ja – jc) at 0.15 V versus RHE plotted against the scan rate (markers); the Cdl values are
estimated through linear fitting of the plots (dashed lines). EIS spectra measured at (c) -0.27 V and
(d) -0.37 V versus RHE presented as Nyquist plots (markers) with respective fits (solid line).
Table 7.2. Overpotential, double-layer capacitance, and charge transfer resistance values for Ni-
GCE-1 and Ni-GCE-2.
Sample η @ 10
mA/cm
2
(mV)
Cdl
(mF/cm
2
)
Rct (Ω) at η =
270 mV
Rct (Ω) at η =
370 mV
Ni-GCE-1 331
6.29 5.316 × 10
3
218.5
Ni-GCE-2 405 1.65 1.582 × 10
4
752.2
The value of Rct was dependent on the applied overpotential as at more negative potentials the
charge transfer resistance was reduced indicating faster HER kinetics (Figures 7.6-7.7). The
intensity of the low frequency feature in the Bode plot (Figure 7.7) was diminished and shifted to
163
higher frequencies at larger overpotentials. The Tafel slope did not change with Cdl, but the
exchange current density (j0) increased with more electrochemically accessible active sites (Figure
7.8).
49
Figure 7.6. Bode plots of Ni-GCE-1 (red) and Ni-GCE-2 (blue) at (a) -0.27 V and (b) -0.37 V
versus RHE.
Figure 7.7. EIS spectra presented as Nyquist plots (markers) with respective fits (solid line) for
(a) Ni-GCE-1 and (b) Ni-GCE-2 and Bode plots for (c) Ni-GCE-1 and (d) Ni-GCE-2 at -0.27 V
(green) and -0.37 V versus RHE (purple).
164
Figure 7.8. Tafel analysis of Ni-GCE-1 (red) and Ni-GCE-2 (blue). Exchange current densities of
10
-7.8
and 10
-8.2
A/cm
2
were obtained for Ni-GCE-1 and Ni-GCE-2.
Controlled potential electrolysis (CPE) of NiBHT at -0.72 V revealed a steady current response
for the first hour followed by a gradual decline in the generated current (Figure 7.9). The current
drop-off was a response to the delamination of NiBHT from the electrode surface as the black
insoluble catalyst was seen suspended in solution and minimal NiBHT was observed on the GCE
surface after 7 h of electrolysis. ICP-OES studies revealed no solubilized nickel species confirming
the decline in activity was due to catalyst delamination not dissolution. Analysis of the gaseous
products by gas chromatography after 1 h of CPE confirmed H2 production with the Faradaic
efficiency (FE) ranging from 65-75%. Most nickel dithiolene based catalysts operate at low to
moderate efficiencies (FE ~ 60-80%)
26,54,55
in comparison to cobalt dithiolene catalysts, which
perform at high efficiencies (FE > 90%).
14,25,28,56
X-ray photoelectron spectroscopy (XPS)
following CPE revealed nickel and sulfur features analogous to those of the as-prepared NiBHT
(Figure 7.10-7.11).
44
Figure 7.9. CPE of NiBHT performed at -0.72 V versus RHE in pH 1.3 aqueous solution.
165
Figure 7.10. X-ray photoelectron spectroscopy of NiBHT. Binding energies of 853.2/870.6 eV (Ni
2p), 226.6 eV (S 2s), and 162.6/163.6 eV (S 2p) were observed for NiBHT.
Figure 7.11. X-ray photoelectron spectroscopy of NiBHT following controlled potential
electrolysis in pH 1.3 solution at -0.72 V versus RHE. Significant delamination of NiBHT from
the electrode surface was observed during the CPE experiment, which results in the observed
current drop off in Figure 7.9 and the low intensity of the XPS peaks. Binding energies of
853.1/870.2 eV (Ni 2p), 226.5 eV (S 2s), and 162.7 eV (S 2p) were observed for NiBHT after CPE.
166
7.2.2. Comparison of the HER Activity of CoBHT, NiBHT, and FeBHT
Polarization curves of CoBHT, NiBHT, and FeBHT are presented in Figure 7.12a highlighting the
influence of the metal center on the HER activity of benzenehexathiolate-based coordination
frameworks. CoBHT displayed the lowest overpotential, 185 mV, in comparison to NiBHT (331
mV) and FeBHT (405 mV). The increased activity of CoBHT is correlated to a higher
electrochemically accessible surface area (Figure 7.12b) and lower charge transfer resistance at
smaller overpotentials (Table 7.3, Figure 7.13-7.14).
Figure 7.12. (a) Polarization curves for CoBHT (pink), NiBHT (cyan), and FeBHT (blue); scan
rate: 100 mV/s. (b) Plot showing the extraction of the double layer capacitance (Cdl). The current
density difference (Δj = ja – jc) at 0.15 V versus RHE plotted against the scan rate (markers); the
Cdl values are estimated through linear fitting of the plots (dashed lines).
NiBHT and FeBHT exhibit very large Rct values at -0.27 V versus RHE inhibiting high
electrocatalytic HER at low overpotentials (Figure 7.14-7.15, Table 7.3), however, at larger
overpotentials, the Rct values decreased significantly (Figure 7.13-7.14, Table 7.3). For NiBHT,
Rct decreased from 5,316 (η = 270 mV)
to 218.5 Ω (η = 370 mV), while for FeBHT, a comparable
Rct to CoBHT and NiBHT was not achieved until η = 470 mV. The Rct for FeBHT decreased from
1.8 × 10
4
(η = 270 mV) to 102.8 Ω (η = 470 mV).
Table 7.3. Electrocatalytic HER properties for BHT-based coordination frameworks.
Sample
η
10 mA/cm
2
(mV)
Cdl
(mF/cm
2
)
Rct (Ω) at η
= 270 mV
Rct (Ω) at η
= 370 mV
Rct (Ω) at η
= 470 mV
Tafel slope
(mV/dec)
j0
(A/cm
2
)
CoBHT 185 27.1 64.4 − − 88 10
-4.9
NiBHT 331 6.29 5.316 × 10
3
218.5 − 67 10
-7.8
FeBHT 473 1.43 1.843 x 10
4
1.309 × 10
3
102.8 119 10
-6.7
167
Figure 7.13. (a) EIS spectra and (b) Bode plots measured at -0.27 V , -0.37 V , and -0.47 V versus
RHE presented as Nyquist plots (markers) with respective fits (solid line) for CoBHT (pink),
NiBHT (cyan), and FeBHT (blue), respectively.
Figure 7.14. EIS spectra presented as Nyquist plots (markers) with respective fits (solid line) for
CoBHT (pink), NiBHT (cyan), and FeBHT (blue) at (a) -0.27 V and (b) -0.37 V versus RHE.
Figure 7.15. Bode plots at (a) -0.27 V and (b) -0.37 V versus RHE for CoBHT (pink), NiBHT
(cyan), and FeBHT (blue).
168
CoBHT is the superior HER catalyst stemming from the significantly enhanced charge transfer
through the material at modest overpotentials and greater number of electrochemically accessible
active sites. Higher electrocatalytic HER activity was achieved for CoBHT at all film thicknesses
investigated (~20-1000 nm) in comparison to NiBHT (131 nm) even as the number of
electrochemically accessible active sites decreased for CoBHT (Table 7.4). The larger j0 of CoBHT
at all thicknesses compared to NiBHT highlights the higher intrinsic electrocatalytic HER activity
of CoBHT (Table 7.3-7.4). To further improve our understanding of how the metal center
influences the electrocatalytic HER activity, the conductivity of the BHT-based coordination
frameworks is currently under investigation in our laboratory to gain a complete understanding of
the origin of the distinct electrochemical responses observed for CoBHT, NiBHT, and FeBHT.
Tafel analysis revealed Tafel slopes of 88, 67, and 119 mV/dec for CoBHT, NiBHT, and FeBHT,
respectively (Figure 7.16). As with previously reported dithiolene-based coordination frameworks
and polymers, the Tafel slopes indicate the Volmer discharge reaction is the rate-limiting
step.
25,26,28-30
The Tafel slope of FeBHT, 119 mV/dec, is much larger than that of CoBHT and
NiBHT. The electronic structure of the analogous molecular complex, [Fe(bdt)2]
2
⁻, could provide
insight into the origins of the larger Tafel slope and greater barrier for Hads. The Fe 3d orbitals are
much higher in energy than the ligand (benzene-1,2-dithiolate) orbitals resulting in a
predominately metal centered singly occupied molecular orbital (SOMO). The electronic structure
is best described as a spin triplet Fe
II
coordinated to two innocent benzene-1,2-dithiolate ligands,
[Fe
II
(bdt
2
⁻)2]
2
⁻.
57,58
Figure 7.16. Tafel analysis of CoBHT (pink), NiBHT (cyan), and FeBHT (blue).
169
Computational studies propose the molecular complexes, [Co(bdt)2]⁻ and [Ni(bdt)2]⁻ (bdt =
benzene-1,2-dithiolate), undergo relatively similar HER mechanisms with the thiolate moieties of
the ligand scaffold the most likely site for initial protonation.
54,59
The non-innocent nature of these
complexes arises because the SOMO of [Co(bdt)2]⁻ contains significant mixed-metal ligand
character, while the SOMO for [Ni(bdt)2]⁻ is mostly ligand-based.
57,58
The lack of non-innocence
in the benzene-1,2-dithiolate ligands and metal dominant SOMO of [Fe(bdt)2]
2
⁻
could give rise to
a different mechanism for the HER than the computational studies propose for [Co(bdt)2]⁻ and
[Ni(bdt)2]⁻. The increase in effective nuclear charge as the metal is changed from Fe to Co to Ni
for the benzene-1,2-dithiolate complexes modulates the metal and ligand character in the
molecular orbitals, potentially resulting in different HER mechanisms and rate-limiting steps for
the frameworks. More comprehensive theoretical studies are necessary to further understand the
differences in the measured Tafel slopes.
7.2.3. Evaluation of the Thickness-Dependent HER Activity of CoBHT
Looking to understand proton and charge transfer in CoBHT to optimize the HER performance,
we varied the amount of BHT used in the synthesis of CoBHT, which resulted in films of different
thicknesses. The average film thickness ranged from 23 nm to 1 μm. It has been shown that
hydrogen production can be dependent on the film thickness because of changes to the electron,
charge, and proton transfer properties and the number of available and accessible active sites.
47-
50,60,61
The effect of film thickness on the electrochemical hydrogen evolving activity of CoBHT
was investigated on glassy carbon electrodes. To distinguish the electrochemical results of films
of different thicknesses, the electrodes studied are denoted as CoBHT-X where X corresponds to
the film thickness in nanometers.
Figure 7.17. Top-down scanning electron microscopy image of CoBHT-23.
170
Figure 7.18. Top-down scanning electron microscopy image of CoBHT-57.
Figure 7.19. Top-down scanning electron microscopy image of CoBHT-157.
Figure 7.20. Top-down scanning electron microscopy image of CoBHT-244.
Figure 7.21. Top-down scanning electron microscopy image of CoBHT-1000.
Scanning electron microscopy (SEM) images revealed morphological differences for the CoBHT
films of different thicknesses (Figures 7.18-7.21). Films within the thickness range of 20-250 nm
171
exhibited a sheet like morphology expected for thin films. Puckering and wrinkling of the film was
observed, which can help to expose more active sites for the HER. The 1000 nm films displayed
an extensively cracked and rough morphology with puckered and overlapping flakes of film. This
type of morphology has the potential to expose more active sites and the interior of the thick film,
which can induce higher electrocatalytic activity. However, the cracked morphology gives rise to
extensive grain boundaries that can potentially inhibit electron transfer from the electrode surface
to the exposed active sites, thus reducing the number of electroactive catalytic sites.
Polarization curves of CoBHT-modified GCE with 5 different film thicknesses, CoBHT-23,
CoBHT-57, CoBHT-157, CoBHT-244, and CoBHT-1000 are presented in Figure 22a. As the film
thickness was increased from 23 to 244 nm, the overpotential to achieve 10 mA/cm
2
decreased
from 246 mV to 185 mV (Figure 7.22a, Table 7.4). The reduced overpotential was accompanied
by an increase in the Cdl from 2.1 to 27.1 mF/cm
2
. Within the thickness range 23 to 244 nm, an
increase in the amount of deposited CoBHT resulted in the expected increase in the number of
accessible active sites (Figure 7.22b, Table 7.4).
Figure 7.22. (a) Polarization curves for CoBHT-23 (purple), CoBHT-57 (blue), CoBHT-157
(orange), CoBHT-244 (red), and CoBHT-1000 (green); scan rate: 100 mV/s, (b) Plot showing the
extraction of the double layer capacitance (Cdl). The current density difference (Δj = ja – jc) at 0.15
V versus RHE plotted against the scan rate (markers); the Cdl values are estimated through linear
fitting of the plots (dashed lines).
As the thickness of the films increased to 1000 nm, reduced HER activity was observed and an
overpotential of 213 mV was measured. Concomitantly, the number of electrochemically
accessible active sites did not improve as the film thickness was increased from 244 to 1000 nm
172
indicating not all the cobalt sites were electrocatalytically active. The substantial increase in mass
loading for CoBHT-1000 did not result in a greater number of electrochemically accessible active
sites as compared to films in the thickness range of 20-250 nm, suggesting insufficient electron
migration from the electrode surface through the 1000 nm thick film. The inability to
electrochemically access the greater number of cobalt active sites associated with the large increase
in film thickness from 244 to 1000 nm inhibits the electrocatalytic HER activity of CoBHT-1000.
Table 7.4. Electrocatalytic HER properties for CoBHT-modified electrodes with different film
thicknesses.
Sample
Average Film
Thickness (nm)
η
10 mA/cm
2
(mV)
Cdl
(mF/cm
2
)
Rct (Ω) @ η
= 270 mV
Tafel slope
(mV/dec)
j0
(A/cm
2
)
CoBHT-1000 1000 213 15.9 144.6 62 10
-6.3
CoBHT-244 244 185 27.1 64.4 88 10
-4.9
CoBHT-157 157 192 22.8 78.3 92 10
-4.8
CoBHT-57 57 225 3.10 100.8 76 10
-5.6
CoBHT-23 23 246 2.11 116.2 77 10
-5.6
The charge transfer resistance, extrapolated from EIS experiments, can reveal differences in the
ability to move charge through CoBHT as the film thickness is modulated. EIS analysis at -0.27 V
versus RHE for CoBHT-23, CoBHT-57, CoBHT-157, CoBHT-244, and CoBHT-1000 is
presented in Figure 7.23. The Rct decreased from 116 to 64 Ω as the film thickness increased from
23 to 244 nm. The improved charge transfer kinetics for CoBHT-244 facilitates electrocatalytic
hydrogen evolution at smaller overpotentials.
Figure 7.23. EIS spectra measured at -0.27 V versus RHE presented as (a) Nyquist plots (markers)
with respective fits (solid line) and (b) Bode plot for CoBHT-23 (purple), CoBHT-57 (blue),
CoBHT-157 (orange), CoBHT-244 (red), and CoBHT-1000 (green).
173
The intensity of the major feature of the Bode plot was reduced and shifted to lower frequencies
for the most active electrode, CoBHT-244, confirming the superior HER kinetics (Figure 7.23b).
Further increases in film thickness to 1000 nm (CoBHT-1000) led to diminished charge transfer
properties and HER kinetics as evidenced by the larger Rct (144.6 Ω) and the greater intensity of
the major feature of the Bode plot (Figure 7.23). The larger charge transfer resistance as the film
thickness reached 1000 nm indicates the charge transfer from the electrode to the CoBHT surface
and finally to the protons in solution is faster for thinner films.
50
CoBHT-1000 exhibited an overpotential superior to CoBHT-23 and CoBHT-57, however, the Rct
indicates sluggish charge transfer properties as the HER reaction rate is enhanced at larger
overpotentials. Analyzing the electrocatalytic current density versus the film thickness at four
potentials (-0.20, -0.22, -0.25 and -0.27 V vs RHE) revealed the HER activity increased until a
thickness of 244 nm (Figure 7.24). At 1000 nm, smaller current densities were generated at all
potentials compared to CoBHT-157 and CoBHT-244. For CoBHT-1000, the higher catalyst
loadings associated with increased film thickness did not result in larger current densities and
lower overpotentials as compared to films in the thickness range of ~20-250 nm because of poor
charge transfer properties and fewer electrochemically accessible active sites (Table 7.4). At small
overpotentials, CoBHT-1000 was catalytically competent and generated higher current densities
than CoBHT-57. However, at larger overpotentials (η = 270 mV), the HER activity of CoBHT-57
(25.1 mA/cm
2
) surpassed the activity of CoBHT-1000 (23.1 mA/cm
2
) confirming the results of
the EIS experiments performed at this potential (Table 7.4, Figure 7.23-7.24).
Figure 7.24. Comparison of the current density (mA/cm
2
) versus thickness (nm) at -0.20 V (blue),
-0.22 V (green), -0.25 V (pink), and -0.27 V versus RHE (purple).
174
Sluggish electron and charge transfer and poor proton permeation through the thick film limited
the catalytic performance at larger overpotentials (Figures 7.23-7.24).
50
CoBHT-1000 modified
electrodes reproducibly exhibited insufficient charge transfer and proton diffusion properties at
large overpotentials limiting the HER activity (Figures 7.23-7.24). The larger electrochemically
accessible surface areas achieved for CoBHT-1000 modified electrodes in comparison to CoBHT-
23 and CoBHT-57 (Table 7.4) initially resulted in higher HER performance (smaller overpotentials
to achieve 10 mA/cm
2
), however, this high activity could not be sustained at larger overpotentials
as poor charge and proton diffusion rendered CoBHT-1000 catalytically incompetent. Therefore,
an optimal film thickness of ~250 nm has been identified for CoBHT whereby proton diffusion,
charge transport, and the number of electrochemically accessible active sites are maximized.
The Tafel slopes were extrapolated from similar potential windows and ranged from 62-92 mV/dec
for CoBHT modified GCE (Figure 7.25). The Tafel slopes correspond to a rate-liming Volmer
discharge step. The cobalt dithiolene-based coordination frameworks and polymers previously
reported as HER catalysts all displayed large Tafel slopes suggesting the reaction was limited by
insufficient adsorption of H
+
.
25,28,30
Similar Tafel slopes and exchange current densities (Table 7.4)
were measured for samples with relatively analogous film thicknesses (CoBHT-23/CoBHT-57 and
CoBHT-157/CoBHT-244). Identifying if different barriers for Hads arise as the thickness is
modulated from ~20-1000 nm is currently under investigation.
Figure 7.25. Tafel analysis of CoBHT-23 (purple), CoBHT-57 (blue), CoBHT-157 (orange),
CoBHT-244 (red), and CoBHT-1000 (green) in pH 1.3 aqueous solutions within the same
potential window.
175
The examination into the role of film thickness on the HER activity for CoBHT revealed an optimal
thickness range (~250 nm) where charge transport, proton transfer, and the number of
electrochemically accessible active sites were maximized. We have shown that the diminished
charge and proton transfer properties and reduced number of electrochemically accessible active
sites (even at higher bulk catalyst loadings) led to the poorer than expected HER performance of
CoBHT-1000. In addition to the fundamental properties such as charge and proton transfer that
dictate the electrocatalytic activity, practical device construction is an important factor to consider
when determining the optimal catalyst loading and deposition method. The mechanical robustness
of the deposited catalyst was greatly decreased as the film thickness was increased to 1000 nm.
Cracking of the film and delamination from the electrode surface during electrochemical testing
was common for CoBHT-1000, thereby limiting the practical feasibility of electrodes constructed
with such high mass loadings. Therefore, it is important to consider the mechanical robustness of
the catalyst-modified electrodes, in addition to electrocatalytic performance, when selecting an
appropriate deposition method and catalyst thickness.
7.3. Conclusions
In summary, cobalt, nickel, and iron coordination frameworks based on benzenehexathiolate were
synthesized via an interfacial reaction and electrochemically investigated as hydrogen evolving
catalysts in pH 1.3 aqueous solutions. CoBHT exhibited the lowest overpotential of 185 mV. The
ability to operate at lower overpotentials for CoBHT was related to the improved charge transfer
properties in comparison to NiBHT and FeBHT. Understanding the effect of film thickness on
electrocatalytic activity, an insufficiently explored area in coordination framework
electrocatalysis, is vital for determining the scalability and practicality of these systems as catalysts
for energy conversion applications. Our exploration into the movement of charge and protons
through the CoBHT films revealed thickness dependent HER activity. As the film thickness was
increased, the overpotential initially decreased as higher accessible active sites were realized. An
optimal thickness was achieved at 244 nm where charge transport, proton diffusion and transfer,
and the number of electrochemically accessible active sites were optimized. As the film thickness
was increased further to ~1000 nm, the charge transfer and proton permeation through the thick
films was insufficient to promote high electrocatalytic activity especially when the HER kinetics
were enhanced at larger overpotentials.
176
Although the cracked and rough morphology of CoBHT-1000 allowed for the exposure of more
active sites, the introduction of significant grain boundaries diminished the charge transfer
properties and hindered the electrochemical accessibility of many of the available active sites.
Diffusion of protons and electrons through the thick film limited the overall HER performance
canceling out any potential gains associated with increased mass loading. The identification of an
optimized thickness (~250 nm) is valuable for informing device construction as the higher mass
loading films (~1000 nm) display minimal mechanical robustness limiting their practical viability.
The optimized CoBHT modified GCE, presented here, operate at overpotentials ranging from 185-
246 mV representing, to the best of our knowledge, the highest performing HER activity of metal
dithiolene based coordination frameworks. Understanding the ability to move charge and protons
through these systems has allowed for the design and construction of heterogenized molecular
catalysts displaying HER activity more reminiscent of traditional heterogeneous catalysts.
7.4. Experimental Details
7.4.1. General Considerations
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres glovebox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. Water was deionized with the Millipore Synergy system (18.2 M ·cm
resistivity) and placed under vacuum and refilled with nitrogen (10 ×). Excluding water, all other
solvents used were degassed with nitrogen and passed through activated alumina columns and
stored over 4Å Linde-type molecular sieves. A benchtop Mettler Toledo pH meter was used to
measure the pH of the aqueous solutions used for electrochemistry. Benzenehexathiol (BHT) was
synthesized according to literature procedure.
62
All other chemical reagents were purchased from
commercial vendors and used without further purification.
7.4.2. Synthesis of CoBHT
The synthesis of CoBHT was previously reported by our laboratory,
28
however, a modified
synthetic procedure is presented here. In a nitrogen filled glovebox, 2 mL of dichloromethane was
added to 2 mg of BHT (0.0074 mmol) in a glass vial. The vial was sealed and sonicated for 5
minutes to form a cloudy white suspension. The suspension was diluted with 40 mL of
177
dichloromethane in a 120 mL jar and water (40 mL) was carefully layered on top of the organic
phase. An aqueous solution of CoCl2·6H2O (2 mL, 0.084 M, 0.17 mmol) was then gently added
via a glass pipette to the aqueous phase. The reaction was left undisturbed for 1-2 days to obtain a
black film at the liquid/liquid interface.
7.4.3. Synthesis of CoBHT Films with Varied Thicknesses
In a nitrogen filled glovebox, 20 mL of degassed dichloromethane was added to 2 mg of BHT in
a glass vial to prepare a saturated stock solution. The film thickness was modified by varying the
volume of the BHT stock solution used in the synthesis (3, 2, 1, or 0.5 mL). The aliquot of the
BHT stock solution was added to 40 mL of dichloromethane in a 120 mL jar. Water (40 mL) was
carefully layered on top of the organic phase followed by the addition of an aqueous solution of
CoCl2·6H2O (2 mL, 0.084 M, 0.17 mmol). The reaction was then left undisturbed for 1-2 days to
obtain a black film at the liquid/liquid interface.
7.4.4. Synthesis of FeBHT
In a nitrogen filled glovebox, 2 mL of dichloromethane was added to 2 mg of BHT (0.0074 mmol)
in a glass vial. The vial was sealed and sonicated for 5 minutes to form a cloudy white suspension.
The suspension was diluted with 40 mL of dichloromethane in a 120 mL jar and water (40 mL)
was carefully layered on top of the organic phase. An aqueous solution of FeCl 2·4H2O (2 mL,
0.084 M, 0.17 mmol) was then gently added via a glass pipette to the aqueous phase. The reaction
was left undisturbed for 1-2 days to obtain a black film at the liquid/liquid interface.
7.4.5. Synthesis of NiBHT
The synthesis of NiBHT was adapted from a reported literature procedure.
44
In a nitrogen filled glovebox, 20 mL of degassed dichloromethane was added to 2 mg of BHT in
a glass vial to prepare a saturated stock solution. 1 mL of the BHT stock solution was added to 40
mL of dichloromethane in a 120 mL jar. Water (40 mL) was carefully layered on top of the organic
phase followed by the addition of an aqueous solution of NiCl2·6H2O (2 mL, 0.084 M, 0.17 mmol).
The reaction was then left undisturbed for 1-2 days to obtain a black film at the liquid/liquid
interface.
178
7.4.6. Deposition of CoBHT, NiBHT, and FeBHT for Electrochemistry
Deposition was carried out by immersing the electrode face down through the film formed at the
liquid-liquid interface. Following deposition, the substrate was washed with water and methanol.
*The following notations are used to describe the modified electrodes:
M-GCE-Y where M = metal, GCE = glassy carbon electrode, and Y = sample number.
CoBHT-X where X = average film thickness.
7.4.7. Physical Methods
Atomic Force Microscopy (AFM) topography images were collected in tapping mode using an
Agilent 5420 SPM instrument S3 operating in tapping mode. The probe tips were Tap300-G
Silicon AFM probes (resonant frequency 300 kHz, force constant 40 N/m) purchased from
Budgetsensors.com and aligned prior to use. Images were collected with a scan rate of 0.1 lines
per second and over an area of at least 40 µm. All samples were imaged under one atmosphere of
air at room temperature.
Scanning electron microscopy (SEM) was performed on a JEOL JSM 7001F scanning electron
microscope using an accelerating voltage of 15 kV.
Inductively coupled plasma optical emission spectroscopy (ICP-OES) measurements were
performed using a Thermo Scientific iCAP 7000 ICP-OES.
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray source
was the Al K α line at 1486.6 eV. Low-resolution survey spectra were acquired between binding
energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV, were
collected on individual XPS lines of interest. The sample chamber was maintained at < 2 × 10
–9
Torr. The XPS data were analyzed using the CasaXPS software.
7.4.8. Electrochemical Methods
Electrochemistry experiments were carried out using a VersaSTAT 3 potentiostat. Platinum wire
used for the electrochemical studies was purchased from Alfa Aesar. The electrochemical
experiments were carried out in a three-electrode configuration electrochemical cell under an inert
atmosphere using glassy carbon electrodes (GCE, 0.07065 cm
2
surface area) as the working
179
electrode. A platinum wire, placed in a separate compartment, connected by a Vycor tip, and filled
with the electrolytic solution (0.1 M NaClO4) was used as the auxiliary (counter) electrode. The
reference electrode, placed in a separate compartment and connected by a Vycor tip, was based on
an aqueous Ag/AgCl/saturated KCl electrode. The reference electrode in aqueous media was
calibrated externally relative to ferrocenecarboxylic acid (Fc-COOH) at pH 7.0, with the Fe
3+/2+
couple at 0.28 V vs Ag/AgCl. All potentials reported in this paper were converted to the reversible
hydrogen electrode (RHE) by adding a value of (0.205 + 0.059 × pH) V.
Cyclic voltammograms for double layer capacitance measurements were taken in a potential
window between 0.1 and 0.2 V vs RHE in pH 1.3 at scan rates of 20, 40, 60, 80, 100, and 120
mV/s. The total current density obtained from the current density difference (Δj = ja – jc) at 0.15 V
vs RHE was plotted against the scan rate. The slope is twice the value of the double layer
capacitance. For Tafel analysis, polarization curves were measured at pH 1.3 with a scan rate of 2
mV/s.
Electrochemical impedance spectroscopy (EIS) measurements were carried out at different
overpotentials (η = 170-470 mV) in the frequency range of 200 kHz – 0.1 Hz with 10 mV
sinusoidal perturbations in pH 1.3 solutions. Experimental EIS data were analyzed and fitted with
the ZSimpWin software.
The obtained polarization curves were corrected for iR loss according to the following equation:
Ecorr = Emea – iRs
Where Ecorr is the iR-corrected potential, Emea is the experimentally measured potential, and Rs is
the solution resistance extracted from the fitted EIS data.
7.4.9. Controlled Potential Electrolysis
Controlled potential electrolysis (CPE) measurements to determine Faradaic efficiency and study
long-term stability were conducted in a sealed two-chambered H-cell where the first chamber held
the working and reference electrodes in 50 mL of 0.1 M NaClO4 (aq) pH 1.3 solution and the
second chamber held the auxiliary electrode in 25 mL of 0.1 M NaClO4 (aq). The two chambers
were both under N2 and separated by a fine porosity glass frit. CPE experiments were performed
180
with glassy carbon plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) as the working
and auxiliary electrodes. The reference electrode was a Ag/AgCl/saturated KCl (aq) electrode
separated from the solution by a Vycor tip. Using a gas-tight syringe, 2 mL of gas were withdrawn
from the headspace of the H-cell and injected into a gas chromatography instrument (Shimadzu
GC-2010-Plus) equipped with a BID detector and a Restek ShinCarbon ST Micropacked column.
To determine the Faradaic efficiency, the theoretical H2 amount based on total charge flowed is
compared with the GC-detected H2 produced from controlled-potential electrolysis.
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Abstract (if available)
Abstract
Sustainable hydrogen production from water splitting (2H₂O → 2H₂ + O₂) has emerged as a promising pathway for the storage and conversion of solar energy. Hydrogen is a valuable energy carrier that can be transformed into electricity using fuel cell technology or used in the production of industrially relevant chemicals such as ammonia and methanol. The conversion of water into H₂ requires the design of earth‐abundant catalysts to replace platinum, whose scarcity and high cost limit global deployment of solar‐to‐fuel converting devices. Practical electrochemical and photoelectrochemical energy converting devices also require the immobilization of stable and active hydrogen evolving catalysts onto electrodes and photocathode materials, which remains a significant challenge. Heterogenization of molecular catalysts combines the advantageous properties of homogeneous and heterogeneous catalysis and facilitates catalyst immobilization onto electrodes of interest. Within this field, the incorporation of molecular catalysts into extended frameworks has emerged as a viable method for improving the catalytic performance and durability of molecular catalysts while maintaining a high degree of control over the structure and properties of the catalytic active site. By using dinucleating and trinucleating ligand scaffolds, we have generated a library of thiolate‐ and selenolate‐based coordination polymers and frameworks and have investigated their use as electrocatalysts for the hydrogen evolution reaction in acidic aqueous media. Because the molecular nature of these systems is retained, we have systematically investigated the ability to tune important catalytic properties of the coordination frameworks such as the overpotential, charge and proton transfer, and efficiency through simple modifications of the ligand scaffold and/or metal center.
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Downes, Courtney Ann
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Electrocatalytic thiolate- and selenolate-based coordination polymers for solar energy conversion
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Doctor of Philosophy
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Chemistry
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06/18/2018
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