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Design of catalysts for the transformation of abundant small molecules into solar fuels

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DESIGN OF CATALYSTS FOR THE TRANSFORMATION OF
ABUNDANT SMALL MOLECULES INTO SOLAR FUELS

by

Eric Michael Johnson

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 2020

ii

“They don’t think it be like it is, but it do”
Oscar Gamble

iii

ACKNOWLEDGEMENTS
I would first like to thank my advisor, Professor Smaranda C. Marinescu, for allowing me
to further my knowledge of chemistry and truly grow both as a chemist and as a person. She
provided me with the freedom to pursue my own interests but also was a much needed guide in
times of confusion. I also would like to thank the other members of my thesis committee,
Professors Travis J. Williams, Richard L. Brutchey, Sri R. Narayan, Sarah Feakins, Mark
Thompson, and Jongseung Yoon for being a helpful hand and aiding in my development as a
doctoral candidate in chemistry at USC.
I would also like to thank the MarinesCrew, Dr. Courtney Downes, Dr. Alon Chapovetsky,
Dr. Andrew Clough, Dr. Damir Popov, Geo Rangel, Nick Orchanian, Ashley Hellman, Keying
Chen, Jeremy Intrator, Dr. Jeff Liu, Adam Samuel, and David Velazquez. You all have provided
some truly wonderful times and memories that I will cherish forever. To Courtney, you were both
a great friend and great mentor in the group. To this day I still believe you are one of the best
chemists I know, and being in a group with you made me want to be a better chemist. I’m still
sorry we both thought we hated each other when I first joined. I really am that awkward. But I am
glad we got over it and became great friends. Alon, you were always a great… friend. You were
my best bro in the group and coming to work every day was made better knowing you would be
there to talk about chemistry and stupid stuff with me. The group just wasn’t the same without
you.
The USC chemistry department has become like an extended family for me, and I wouldn’t
be where I am today without my friends in the Melot, Brutchey, M. Thompson, Williams, and
other groups. Above all else, I must say that Erica Howard is a real one. She has been a great friend
since our first year and has helped me tremendously throughout these long years. To Ariel Wein,
you were always a great friend to me and I am sad that we weren’t able to hang out more during
our time here. You are a beautiful soul and made the transition to LA less painful. Abegail Tadle
has been a homie since day one. Somehow every time I come to borrow a chemical from your lab
it feels like we spend 5 minutes sorting that out and then another hour talking. I appreciate you
immensely and still owe you a lunch. To Nikki Pedowitz, we seem to have a trend of talking a lot
for brief periods of time before not speaking for months or years. Still, you’ve been a great friend
and helped me get through things I had been ignoring for a long time, and I can’t thank you enough
for that.
iv

With as great as my USC friends have been, my friends from Virginia have been even
greater. This first acknowledgement goes out to my oldest friend, Max Deleon. We’ve known each
other for over half of our lives now, and while I think there are times where we don’t really
understand each other, you are family to me and I love you. I am so proud of the person you have
become and I hope I can make you proud as well. To Kathryn Milburn, my favorite Russian, I
don’t know where I would be without you. While it takes a long time for us to finally get to hang
out, I am happy that even if life takes us in different directions you are still willing to put up with
me and let me meet your hot friends. To Tasha Horton, the best girlfriend I never had, I will always
miss the times in undergrad when we could hang out pretty much any time we wanted and weren’t
limited by the fact that we have to be real adults and live in very different areas. Hopefully one
day we can end up within at least 4 states of each other and hang out in the future. And last but
certainly not least, the G.O.A.T. of this friend stuff, July Laszakovits, I have to give a massive
shout out to you. In addition to being an absolutely stellar person that put up with a lot of my
complaining throughout the years, you also understand science things and can support me with a
sincerity of only someone who has gone through the same thing. I don’t know if I could have made
it through this process without you. I love you all.
I must, of course, give a brief thanks to my therapist, Dr. Terence Ford, for helping me deal
with the strain of graduate school and life in general. While our time together was short, you helped
open my eyes to a lot of problems both in my environment and in my behavior and I think I am a
happier person because of you. Therefore I must sincerely thank you. I would also like to thank
my unofficial therapist, Faheem Rasheed Najm a.k.a. the one and only T-Pain, without whom I
would not have had the mental strength to come and do this every day for this many years. You
are a shining beacon in this dark world.
Finally, I have to thank my family. I wouldn’t be here without them. To my parents Ken
and Michelle, I thank you for everything you ever did for me. You always pushed me to do my
best and more importantly that it’s okay when I fail. You believed in me when I didn’t believe in
myself and helped motivate me to make it this far. To my siblings, Tori and Nick, I am so proud
of the young adults you have become. I know we have fought a lot throughout the years but I am
glad we have reached a point where we still have each other in our lives, even though Nick never
answers his phone. Thank you both for looking up to me and making me strive to be someone
worthy of it.
v

TABLE OF CONTENTS

Epigraph ......................................................................................................................................... ii
Acknowledgements ......................................................................................................................... iii
Table of Contents .............................................................................................................................v
List of Tables .................................................................................................................................. ix
List of Figures ................................................................................................................................ xi
List of Schemes ........................................................................................................................... xviii
Abstract ........................................................................................................................................ xix
Chapter 1. General Introduction ...................................................................................................1
1.1 Outlook on Global Energy .............................................................................................1
1.2 Solar Fuel Production ....................................................................................................1
1.3 Transformation of Abundant Molecules to Solar Fuels .................................................3
1.3.1 Carbon Dioxide Reduction Reaction ..............................................................3
1.3.2 Oxygen Reduction Reaction (ORR) ..............................................................4
1.4 Homogeneous Electrocatalysts for CO2 Reduction .......................................................5
1.4.1 Manganese Catalysts .......................................................................................5
1.4.2 Iron Catalysts ..................................................................................................7
1.4.2.1 Porphyrins ........................................................................................7
1.4.2.2 Iron Carbonyl Clusters ...................................................................10
1.4.2.3 Hydroxybenzene Bipyridines.........................................................11
1.4.3 Cobalt Catalysts ............................................................................................12
1.4.3.1 Tetra-Aza Macrocycles ..................................................................12
1.4.3.2 Cyclopentadienyl Complexes ........................................................13
1.4.4 Nickel Catalysts ............................................................................................14
1.4.4.1 Cyclam ...........................................................................................14
1.4.4.2 Polypyridyl .....................................................................................16
1.4.5 Copper Catalysts ...........................................................................................16
1.5 Heterogenization of Molecular Catalysts.....................................................................17
1.5.1 CO2 Reduction ..............................................................................................18
1.5.1.1 Cobalt Porphyrins ..........................................................................18
1.5.1.2 Iron Porphyrins ..............................................................................19
1.5.1.3 Rhenium Bipyridines .....................................................................20
1.5.2 O2 Reduction .................................................................................................21
1.5.2.1 Triphenylene-based MOFs.............................................................21
1.5.2.2 Iron Porphyrins ..............................................................................23
1.6 Outline of Thesis ..........................................................................................................24
1.7 References ....................................................................................................................26
Chapter 2. Switching Catalyst Selectivity via the Introduction of a Pendant Nitro Group...34
2.1 Abstract ........................................................................................................................34
2.2 Introduction ..................................................................................................................34
2.3 Results and Discussion ................................................................................................36
2.3.1 Synthesis and Characterization .....................................................................36
2.3.2 Cyclic Voltammetry in Acetonitrile ..............................................................38
2.3.3 Controlled Potential Electrolysis ..................................................................40
2.3.4 Cyclic Voltammetry Dependence Studies ....................................................42
vi

2.4 Conclusion ...................................................................................................................43
2.5 Acknowledgements ......................................................................................................43
2.6 Experimental Methods .................................................................................................44
2.6.1 Physical Methods ..........................................................................................44
2.6.2 Electrochemical Methods..............................................................................45
2.6.3 Cyclic Voltammetry of L
NO2
and CoL
NO2
....................................................45
2.6.4 Controlled Potential Electrolysis ..................................................................45
2.7 Synthetic Methods .......................................................................................................46
2.7.1 Synthesis of L
NO2
..........................................................................................46
2.7.2 Synthesis of Co complex ..............................................................................46
2.8 Additional Figures .......................................................................................................48
2.9 References ....................................................................................................................73
Chapter 3. Electrochemistry of a Cobalt Nitrophenyl Substituted Aminopyridine Complex
in Dimethylformamide Solution .................................................................................................78
3.1 Abstract ........................................................................................................................78
3.2 Introduction ..................................................................................................................78
3.3 Results and Discussion ................................................................................................80
3.3.1 Cyclic Voltammetry ......................................................................................80
3.3.2 Controlled Potential Electrolysis ..................................................................81
3.3.3 Determination of CO Source ........................................................................83
3.4 Conclusion ...................................................................................................................84
3.5 Acknowledgements ......................................................................................................84
3.6 Experimental Methods .................................................................................................85
3.6.1 Physical Methods ..........................................................................................85
3.6.2 Electrochemical Methods..............................................................................85
3.6.3 Cyclic Voltammetry of L
NO2
and CoL
NO2
....................................................85

3.6.4 Controlled Potential Electrolysis ..................................................................86
3.7 Additional Figures .......................................................................................................87
3.8 References ....................................................................................................................93
Chapter 4. Covalent Organic Frameworks Composed of Rhenium Bipyridine and Metal
Porphyrins: Designing Heterobimetallic Frameworks with Two Distinct Metal Sites .........98
4.1 Abstract ........................................................................................................................98
4.2 Introduction ..................................................................................................................98
4.3 Results and Discussion ..............................................................................................101
4.3.1 Synthesis and Characterization of Rhenium Complex ...............................101
4.3.2 Electrochemical Studies of Molecular Species ...........................................103
4.3.3 Immobilization of Rhenium Complex via COFs ........................................106
4.4 Conclusions ................................................................................................................115
4.5 Acknowledgements ....................................................................................................116
4.6 Experimental Methods ...............................................................................................116
4.6.1 Physical Methods ........................................................................................116
4.6.2 Electrochemical Methods............................................................................119
vii

4.6.3 Preparation of Carbon Fabric COF-modified Working Electrodes ...........119
4.6.4 Cyclic Voltammetry of 1 and 2 ...................................................................119
4.6.5 Cyclic Voltammetry of COFs .....................................................................119
4.6.6 Controlled Potential Electrolysis ................................................................120
4.6.7 Controlled Potential Electrolysis of 1 and 2 ...............................................120
4.6.8 Controlled Potential Electrolysis of COFs ..................................................120
4.7 Synthetic Methods .....................................................................................................121
4.7.1 2,2'-Bipyridine-5,5'-dicarbaldehyde (1). .....................................................121
4.7.2 2,2'-Bipyridine-5,5'-dicarbaldehydetricarbonylchlororhenium(I) (2). ........121
4.7.3 COF-Bpy Synthesis....................................................................................121
4.7.4 COF-Re Synthesis. .....................................................................................122
4.7.5 COF-Mix Synthesis....................................................................................122
4.7.6 Synthesis of COF-Re_Co...........................................................................122
4.7.7 Synthesis of COF-Re_Fe. ..........................................................................123
4.8 References ..................................................................................................................147
Chapter 5. A Cobalt Triphenylene MOF for Oxygen Reduction ...........................................152
5.1 Abstract ......................................................................................................................152
5.2 Introduction ................................................................................................................152
5.3 Results and Discussion ..............................................................................................154
5.3.1 Cyclic Voltammetry of CoTHT ..................................................................154
5.3.2 Cyclic Voltammetry of CoTHT-Nafion-Carbon Black Ink Mixture ..........157
5.3.3 X-Ray Photoelectron Spectroscopy of CoTHT-Nafion-Carbon Black
Ink Mixture .................................................................................................159
5.3.4 Cyclic Voltammetry of the Oxidized CoTHT-Nafion-Carbon Black Ink
Mixture .......................................................................................................161
5.3.5 Theoretical ORR Activity of CoTHT .........................................................162
5.4 Conclusions ................................................................................................................163
5.5 Acknowledgements ....................................................................................................164
5.6 Experimental Methods ...............................................................................................165
5.6.1 Electrochemical Methods............................................................................165
5.6.2 Physical Characterization Methods ............................................................165
5.6.3 Koutecký-Levich Equation .........................................................................166
5.7 Synthetic Methods .....................................................................................................167
5.7.1 Synthesis of CoTHT ...................................................................................167
5.7.2 Deposition of CoTHT for Electrochemical Study ......................................167
5.8 Additional Figures .....................................................................................................168
5.9 References ..................................................................................................................171
Chapter 6. Future Directions and Outlook ...............................................................................174
6.1 Lariat Macrocycle for Solar Energy Storage .............................................................174
6.2 Decarbonylation Reactions using Lariat ....................................................................179
6.3 Bimetallic COF ..........................................................................................................179
viii

6.4 Cobalt Triphenylenehexathiol MOF ..........................................................................182
Bibliography ................................................................................................................................183

ix

LIST OF TABLES

Table 2.1. Comparison of atomic distances in the metallated lariat and parent complexes as
determined from the crystal structures. ..............................................................................38
Table 2.2. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile
under N2 with 1 M TFE and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. ......................................56
Table 2.3. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile
under CO2 with 1 M TFE and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. ....................................56
Table 2.4. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile
under CO2 with 1 M H2O and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. ....................................56
Table 2.5. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile
under CO2 with 1 M Phenol and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. ................................57
Table 2.6. Results of 2 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile
under CO2 with 1 M Phenol and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
(catalyst solution) and
the solvent-washed electrode with no added CoL
NO2
and otherwise identical conditions
(washed electrode). ............................................................................................................58
Table 2.7. The slope of the lines generated from plotting the shift in the -1.65 V vs Fc
+/0
couple vs
added acid concentration and their corresponding gas and acid environments. ................62
Table 2.8. Sample and crystal data for CoL
NO2
. ..........................................................................67
Table 2.9. Bond lengths (Å) for CoL
NO2
. ....................................................................................68
Table 2.10. Bond angles (°) for CoL
NO2
. .....................................................................................70
Table 3.1. Results of 2 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF with 0.1
M TBAPF6 and with varied atmospheres and TFE concentrations. .................................83
Table 3.2. Results of 2 hour controlled potential electrolysis of 0.5 mM L
NO2
, Co(ClO4)∙6H2O,
and CoL
NO2
in DMF under CO2 with 1 M TFE and 0.1 M TBAPF6. ..............................89
Table 3.3. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
N2 with 1 M TFE and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. ................................................89
Table 3.4. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M TFE and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. .............................................90
Table 3.5. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M TFE and 0.1 M TBAPF6 at -2.25 V vs Fc
+/0
. .............................................90
Table 3.6. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M TFE and 0.1 M TBAPF6 at -2.45 V vs Fc
+/0
. .............................................91
Table 3.7. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M TFE and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. .............................................91
Table 3.8. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M H2O and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. .............................................92
Table 3.9. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M phenol and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. .........................................92
x

Table 3.10. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under
CO2 with 1 M methanol and 0.1 M TBAPF6 at -2.70 V vs Fc
+/0
. .....................................92
Table 4.1. Gaseous products measured using gas chromatography after 2 hours of bulk electrolysis
at -2.0 V vs SCE for a blank solution, 1, and 2 under N2 and CO2 atmospheres with 4 M
MeOH. .............................................................................................................................104
Table 4.2. Gaseous products measured by GC after 120 minutes of bulk electrolysis with bare
carbon fabric, COF-Re, COF-Re_Co, and COF-Re_Fe under CO2 atmosphere. .........114
Table 4.3. BET surface areas of the synthesized COFs. .............................................................140
Table 4.4. Crystal data and structure refinement for 2. ...............................................................143
Table 4.5. Bond lengths (Å) for 2. ..............................................................................................144
Table 4.6. Bond angles (°) for 2. .................................................................................................145

xi

LIST OF FIGURES

Figure 1.1. Design of a standard H2 fuel cell, where H2 gas is oxidized at the anode (left side) and
O2 is reduced at the cathode (right side). ............................................................................2
Figure 1.2. CO2 reduction processes and their standard reduction potentials. ................................3

Figure 1.3. Half reactions and potentials for the water splitting reaction. ......................................4
Figure 1.4. Catalytic mechanism of Mn bipyridine with (red) and without (green) the added
methoxy substituent. Reprinted with permission from reference 6. ....................................7
Figure 1.5. Structures of the hydroxy (left) and methoxy (right) iron porphyrin complexes.
Reprinted from reference 53 with permission from AAAS. ................................................8
Figure 1.6. Family of iron porphyrins containing different electron withdrawing groups. Adapted
with permission from references 55 and 57. ........................................................................9
Figure 1.7. Structures of the iron carbonyl clusters with interstitial N atoms where L = CO, PPh3,
or PPh2(CH2)2OH. Reprinted with permission from reference 61. Published by the Royal
Society of Chemistry..........................................................................................................10
Figure 1.8. Structure of the cobalt tetra-aza macrocycle. Reprinted with permission from reference
65........................................................................................................................................13
Figure 1.9. The trans I (left) and trans III (right) forms of [Ni(cyclam)]
2+
. Reprinted with
permission from reference 73. ...........................................................................................14
Figure 1.10. Methylated variants of cyclam. Reprinted with permission from reference 80. ......15
Figure 1.11. CO2 reduction catalysts (a) COF-366-Co and COF-367-Co, (b) COF-366-Co
modified through the organic linker, and (c) the MOF [Al2(OH)2TCPP-Co]. Adapted with
permission from reference 91. Published in Nano Research by Springer Nature, 2019. ..19
Figure 1.12. MOF-525 on an FTO electrode where the electrons can hop through the available
iron sites to reach the CO2 substrate. Reprinted with permission from reference 96. .......20
Figure 1.13. Formation of a rhenium bipyridine-based MOF on an FTO surface through liquid-
phase epitaxy. Reprinted with permission of Journal of Materials Chemistry A from
reference 101; permission conveyed through Copyright Clearance Center, Inc. ..............21
Figure 1.14. General structures showing the packing of triphenylene MOFs in hexagonal (a) and
trigonal (b) phases. Reprinted with permission from reference 103. Published by the Royal
Society of Chemistry..........................................................................................................22
Figure 1.15. The structure of Ni-CAT where the two axial water species are omitted for clarity.
Reprinted with permission from reference 104. Copyright 2017 American Chemical
Society................................................................................................................................23
Figure 1.16. The (a) (001) and (b) (100) planes of PCN-223-Fe showing its similarity to MOF-
525 and how charge hopping may occur through the porphyrin sites. Reprinted with
permission from reference 105. .........................................................................................23
Figure 2.1. Parent complex CoL
5
(left) and nitrophenyl lariat modified complex CoL
NO2
(right).
Counter anions removed for clarity. ..................................................................................37
Figure 2.2. Crystal structure of CoL
NO2
with side (A and B) and top (C) views. Noncoordinating
counter ions, unbound solvent species, and hydrogen atoms omitted for clarity. .............38
xii

Figure 2.3. (A) CVs of 0.5 mM CoL
NO2
under N2 (blue) and CO2 (red) atmospheres in 0.1 M
TBAPF6 MeCN solution. (B) CVs of 0.5 mM CoL
NO2
under an N2 atmosphere with 0 M
TFE (green) and 0.87 M TFE (purple) added in 0.1 M TBAPF6 MeCN solution. (C) CVs
of 0.5 mM CoL
NO2
under a CO2 atmosphere with 0 M TFE (green) and 0.87 M TFE
(purple) added in 0.1 M TBAPF6 MeCN solution. (D) Controlled potential electrolysis of
0.5 mM CoL
NO2
under N2 (blue) and CO2 (red) with 0.87 M TFE in 0.1 M TBAPF6 MeCN
solution. ..............................................................................................................................39
Figure 2.4. CPE of 0.5 mM CoL
NO2
under CO2 in 0.1 M TBAPF6 MeCN solution at -2.70 V vs
Fc
+/0
with 1 M TFE (red), water (blue), and phenol (green) as an added proton source. ..41
Figure 2.5.
1
H NMR spectrum (400 MHz) of CoL
NO2
in CD3CN................................................47
Figure 2.6.
1
H NMR spectra (400 MHz) overlay of CoL
NO2
(top) and CoL
5
(bottom) in deuterated
pyridine (Py-d5). .................................................................................................................48
Figure 2.7. UV-Vis spectrum of CoL
NO2
in DMF solution. .........................................................48
Figure 2.8. FTIR spectrum of L
NO2
...............................................................................................49
Figure 2.9. FTIR spectrum of CoL
NO2
. .........................................................................................49
Figure 2.10. FTIR spectrum overlay of the NO2 regions of L
NO2
and CoL
NO2
. ...........................50
Figure 2.11. 0.5 mM L
NO2
in 0.1 M TBAPF6 MeCN under N2 (blue) and CO2 (red) atmospheres
Scan rate = 100 mV/s. ........................................................................................................50
Figure 2.12. 0.5 mM L
NO2
in 0.1 M TBAPF6 MeCN under a CO2 atmosphere titrated with TFE.
Scan rate = 100 mV/s. ........................................................................................................51
Figure 2.13. CV overlay of L
NO2
(blue), CoL
NO2
(red), and ZnL
NO2
(green) in 0.1 M TBAPF6
acetonitrile solution under N2. Scan rate = 100 mV/s. .......................................................51
Figure 2.14. (A) Cyclic voltammogram of CoL
NO2
(0.5 mM) in MeCN with 0.1 M TBAPF6 under
N2 with varied scan rates with (B) a corresponding plot of log(current density) vs log(scan
rate). ...................................................................................................................................52
Figure 2.15. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under an N2
atmosphere with increasing amounts of TFE. Scan rate = 100 mV/s. ...............................52
Figure 2.16. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2
atmosphere with increasing amounts of TFE. Scan rate = 100 mV/s. ...............................53
Figure 2.17. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under an N2
atmosphere with increasing amounts of H2O. Scan rate = 100 mV/s. ...............................53
Figure 2.18. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2
atmosphere with increasing amounts of H2O. Scan rate = 100 mV/s. ...............................54
Figure 2.19. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under an N2
atmosphere with increasing amounts of Phenol. Scan rate = 100 mV/s. ...........................54
Figure 2.20. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2
atmosphere with increasing amounts of Phenol. Scan rate = 100 mV/s. ...........................55
Figure 2.21. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2
atmosphere with increasing amounts of Phenol. Low concentrations of Phenol only. Scan
rate = 100 mV/s. .................................................................................................................55
xiii

Figure 2.22. XPS spectra of the Co 2p region for the acetonitrile (A) and acetonitrile and DMF
(B) washed sections of the glassy carbon working electrode following a two hour controlled
potential electrolysis of CoL
NO2
with 1 M phenol.............................................................57
Figure 2.23. Controlled potential electrolysis of 1 M Phenol acetonitrile solution containing 0.5
mM CoL
NO2
and a clean electrode (red) and a washed electrode in the absence of CoL
NO2

(black). 0.1 M TBAPF6 acetonitrile solution at -2.70 V vs Fc
+/0
. ......................................58
Figure 2.24. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under an N2
atmosphere with increasing amounts of D2O. Scan rate = 100 mV/s. ...............................59
Figure 2.25. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2
atmosphere with increasing amounts of D2O. Scan rate = 100 mV/s. ...............................59
Figure 2.26. Titration of CoL
NO2
in 0.1 M TBAPF6 MeCN solution under a CO2 atmosphere with
increasing amounts of TFE focused on the reduction event at -1.65 V vs Fc
+/0
shifting
anodically. ..........................................................................................................................60
Figure 2.27. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the
concentration of the added proton source in solution under different conditions. Red circles
= under CO2 with titrated TFE. Blue circles = under N2 with titrated TFE. ......................60
Figure 2.28. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the
concentration of the added proton source in solution under different conditions. Red circles
= under CO2 with titrated H2O. Blue circles = under N2 with titrated H2O. ......................61
Figure 2.29. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the
concentration of the added proton source in solution under different conditions. Red circles
= under CO2 with titrated phenol. Blue circles = under N2 with titrated phenol. ............. 61
Figure 2.30. Plot of potential vs pKa for CoL
NO2
under N2 based on the shift of the couple at -1.65
V vs Fc
+/0
when titrated with water, TFE, and phenol. ......................................................62
Figure 2.31. Plot of current density vs TFE concentration for CoL
NO2
in MeCN solution under N2
titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all
concentrations. ...................................................................................................................63
Figure 2.32. Plot of current density vs H2O concentration for CoL
NO2
in MeCN solution under N2
titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all
concentrations. ...................................................................................................................63
Figure 2.33. Plot of current density vs Phenol concentration for CoL
NO2
in MeCN solution under
N2 titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for
all concentrations. ..............................................................................................................64
Figure 2.34. Plot of current density vs TFE concentration for CoL
NO2
in MeCN solution under
CO2 titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for
all concentrations. ..............................................................................................................64
Figure 2.35. Plot of current density vs H2O concentration for CoL
NO2
in MeCN solution under
CO2 titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for
all concentrations. ..............................................................................................................65
xiv

Figure 2.36. Plot of current density vs Phenol concentration for CoL
NO2
in MeCN solution under
CO2 titrated with corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for
all concentrations. ..............................................................................................................65
Figure 2.37. Titrations of CoL
NO2
into 0.1 M TBAPF6 acetonitrile solution under N2 with 0.5 M
TFE. ...................................................................................................................................66
Figure 2.38. Current density vs catalyst concentration for CoL
NO2
. ............................................66
Figure 3.1. Cyclic voltammograms of L
NO2
(0.5 mM) in DMF with 0.1 M TBAPF6. (A) Overlay
of scans under N2 and CO2. (B) Titrations with TFE under a CO2 atmosphere. Scan rate =
100 mV/s. ...........................................................................................................................80
Figure 3.2. Electrochemical analysis of CoL
NO2
(0.5 mM) in DMF with 0.1 M TBAPF6. (A)
Overlay of cyclic voltammetry scans under N2 (blue) and CO2 (red), (B) under an N2
atmosphere with 0 M (green) and 1.02 M (purple) TFE, and (C) under a CO2 atmosphere
with 0 M (green) and 0.93 M (purple) TFE. (D) Bulk electrolysis of CoL
NO2
in DMF with
1 M TFE under N2 (blue), 1 M TFE under CO2 (red), and 0 M TFE under CO2 (green).
[CoL
NO2
] = 0.5 mM. Scan rates = 100 mV/s. ....................................................................82
Figure 3.3. Cyclic voltammogram of CoL
NO2
(0.5 mM) in DMF with 0.1 M TBAPF6 under N2
with varying scan rates.......................................................................................................87
Figure 3.4. Plot of log(current) vs log(scan rate) for CoL
NO2
. .....................................................87
Figure 3.5. Cyclic voltammograms of CoL
NO2
(0.5 mM) under CO2 in saturated DMF with 0.1 M
TBAPF6and 1 M TFE after electrolysis (red) and of rinsed electrode (3 × 10 mL DMF) in
a fresh DMF solution containing 0.1 M TBAPF6and 1 M TFE under CO2 (blue). Scan rate
= 100 mV/s. Measured in an H cell. Working electrode surface area approximated as
8 cm
2
. .................................................................................................................................88
Figure 3.6. Controlled potential electrolysis of L
NO2
(red), Co(ClO4)∙6H2O (blue), and CoL
NO2

(green) under CO2 in DMF with 0.1 M TBAPF6 and 1 M TFE. All concentrations were 0.5
mM. All electrolyses were performed at -2.70 V vs Fc
+/0
. Experiments were performed for
2 hours. ...............................................................................................................................88
Figure 3.7. Controlled potential electrolysis of CoL
NO2
under N2 (blue) and CO2 (red) in DMF
with 0.1 M TBAPF6 and 1 M TFE. All concentrations were 0.5 mM. All electrolyses were
performed at -2.70 V vs Fc
+/0
. Experiments were performed for 24 hours. ......................89
Figure 3.8. Controlled potential electrolysis of CoL
NO2
(0.5 mM) under CO2 in DMF with 0.1 M
TBAPF6 and 1 M TFE for 6 hours at -2.25 V (red), -2.45 V (blue), and -2.70 V (green) vs
Fc
+/0
. ...................................................................................................................................90
Figure 3.9. Controlled potential electrolysis of CoL
NO2
(0.5 mM) under CO2 in DMF with 0.1 M
TBAPF6 and 1 M of H2O (red), TFE (blue), phenol (green), or methanol (purple) as the
proton source. All electrolyses were performed at -2.70 V vs Fc
+/0
. Experiments were
performed for 6 hours. .......................................................................................................91
Figure 4.1. X-ray crystal structure of 2. Selected hydrogen atoms have been omitted from the
bipyridine ligand for clarity. ............................................................................................101
xv

Figure 4.2. FT-IR spectrum of 2. The black box highlights the carbonyl stretching frequency
region. ..............................................................................................................................102
Figure 4.3. Electrochemical studies of 1 and 2 in DMF. (A) Overlay of the cyclic voltammograms
of 1 (0.5 mM, blue) and 2 (0.5 mM, red) under an N2 atmosphere and (B) cyclic
voltammograms of 2 (0.5 mM) under N2 (purple) and CO2 (green) atmospheres. (C) Cyclic
voltammogram of 2 (0.5 mM) under 1 atm of CO2 with increasing concentrations of MeOH
(from 0 to 4 M). Cyclic voltammetry experiments were performed in DMF with 0.1 M
TBAPF6 at a scan rate of 20 mV/s. (D) Controlled potential electrolysis of 1 (blue), 2 (red),
and a blank solution (dashed black) at -2.0 V versus SCE in DMF with 4 M MeOH under
a CO2 atmosphere. ............................................................................................................104
Figure 4.4. PXRD patterns of COF-Bpy (green), COF-Re (yellow), COF-Re_Co (light red), and
COF-Re_Fe (black).........................................................................................................108
Figure 4.5. P21212 space group used to model COF-Re-Co. .......................................................109
Figure 4.6. Overlay of the P21212 space group predicted (blue) and experimental (magenta) PXRD
patterns. ............................................................................................................................109
Figure 4.7. PCC2 space group used to model COF-Re-Co. .........................................................110
Figure 4.8. Overlay of the PCC2 space group predicted (blue) and experimental (magenta) PXRD
patterns. ............................................................................................................................110
Figure 4.9. XPS analyses of COFs. (A) Re 4f XPS spectrum of COF-Re. (B) Co 2p XPS spectrum
of COF-Re_Co. (C) Fe 2p XPS spectrum of COF-Re_Fe. ............................................111
Figure 4.10. FT-IR spectra of COF-Bpy (green), COF-Re (yellow), COF-Re_Co (light red), and
COF-Re_Fe (dark red). ...................................................................................................112
Figure 4.11. Cyclic voltammograms of (A) COF-Re, (B) COF-Re_Co, and (C) COF-Re_Fe in
pH 7.2 aqueous phosphate buffer solutions with 0.5 M KHCO3 under N2 (blue) and CO2
(red) atmosphere. Scan rate = 100 mV/s. .........................................................................113
Figure 4.12. CPE of bare carbon fabric (dashed black), COF-Re (red), COF-Re_Co (blue), and
COF-Re_Fe (green) at -1.1 V vs SHE under CO2 in pH 7.2 aqueous phosphate buffer with
0.5 M KHCO3. .................................................................................................................114
Figure 4.13. 400 MHz
1
H NMR spectrum of 1 in DMSO-d6. ....................................................124
Figure 4.14. 100 MHz
13
C{
1
H} NMR spectrum of 1 in DMSO-d6. ...........................................125
Figure 4.15. 400 MHz
1
H NMR spectrum of 2 in DMSO-d6. ....................................................126
Figure 4.16. 100 MHz
13
C{
1
H} NMR spectrum of 2 in DMSO-d6. ...........................................127
Figure 4.17. FTIR spectrum of 1.................................................................................................128
Figure 4.18. Variable scan rate of the first redox couple of 2. Scan rates are all mV/s. .............128
Figure 4.19. Plot of current density vs square root of the scan rate for the first redox couple of 2.
Current densities were measured at -0.585 V and -0.542 V vs SCE. ..............................129
Figure 4.20. Variable scan rate of both redox couple of 2. Scan rates are all mV/s. ..................129
Figure 4.21. Plot of current density vs square root of the scan rate for the second redox couple of
2. Current densities were measured at -0.931 V and -0.869 V vs SCE. ..........................130

xvi

Figure 4.22. Cyclic voltammograms of 0.5 mM 1 under N2 and CO2 environment in DMF with
0.1 M TBAPF6. ................................................................................................................130
Figure 4.23. Cyclic voltammograms of 1 and 2 (0.5 mM) under CO2 in DMF with 0.1 M TBAPF6.
..........................................................................................................................................131
Figure 4.24. Cyclic voltammograms of 1 (blue) and 2 (red) (0.5 mM) under N2 in DMF with 4 M
MeOH and 0.1 M TBAPF6. .............................................................................................131
Figure 4.25. Cyclic voltammograms of 2 (0.5 mM) under N2 in DMF and 0.1 M TBAPF6 with 0
M (purple) and 4 M (green) MeOH. ................................................................................132
Figure 4.26. Cyclic voltammogram of bare glassy carbon (black dashed), 1 (blue), and 2 (red) (0.5
mM) in CO2 saturated DMF with 0.1 M TBAPF6 and 4 M MeOH. Scan rate = 100 mV/s.
Measured in an H cell. Working electrode surface area = 8 cm
2
. ....................................132
Figure 4.27. Cyclic voltammograms of 2 (0.5 mM) under CO2 in saturated DMF with 0.1 M
TBAPF6 and 4 M MeOH before CPE (green) and of rinsed electrode (3  10 mL DMF) in
a fresh DMF solution containing 0.1 M TBAPF6 and 4 M MeOH under CO2 (purple). Scan
rate = 100 mV/s. Measured in an H cell. Working electrode surface area = 8
cm
2
. ..................................................................................................................................133
Figure 4.28. Controlled potential electrolyses of 1 (blue), 2 (red), and a blank solution (dashed
black) at -2.0 V versus SCE in DMF with 4 M MeOH under a N2 atmosphere. .............133
Figure 4.29. PXRD pattern of COF-Bpy. ..................................................................................134
Figure 4.30. PXRD pattern of COF-Re. .....................................................................................134
Figure 4.31. PXRD pattern of COF-Mix. ..................................................................................135
Figure 4.32. PXRD pattern of COF-Re_Co. ..............................................................................135
Figure 4.33. PXRD pattern of COF-Re_Fe. ..............................................................................136
Figure 4.34. N 1s XPS spectrum of COF-Bpy. ..........................................................................136
Figure 4.35. N 1s XPS spectrum of COF-Re. ............................................................................137
Figure 4.36. XPS analyses of COF-Re_Co. (A) N 1s XPS spectrum. (B) Re 4f XPS
spectrum. ..........................................................................................................................137
Figure 4.37. XPS analyses of COF-Re_Fe. (A) N 1s XPS spectrum. (B) Re 4f XPS
spectrum. ..........................................................................................................................137
Figure 4.38. IR spectrum of COF-Bpy.......................................................................................138
Figure 4.39. IR spectrum of COF-Re. ........................................................................................138
Figure 4.40. IR spectrum of COF-Re_Co. .................................................................................139
Figure 4.41. IR spectrum of COF-Re_Fe...................................................................................139
Figure 4.42. N2 sorption isotherm of COF-Re at 77K. Red = adsorption, Blue = desorption. ..140
Figure 4.43. N2 sorption isotherm of COF-Re_Co at 77K. Red = adsorption,
Blue = desorption. ...........................................................................................................140
Figure 4.44. N2 sorption isotherm of COF-Re_Fe at 77K. Red = adsorption,
Blue = desorption. ...........................................................................................................141
Figure 4.45. TGA trace of the COF-Re activated sample under N2. .........................................141
Figure 4.46. TGA trace of the COF-Re_Co activated sample under N2. ...................................142

xvii

Figure 4.47. TGA trace of the COF-Re_Fe activated sample under N2. ...................................142
Figure 5.1. CV of CoTHT in pH 13 aqueous solution under N2 (blue) and O2 (red). Scan rate = 10
mV/s. Rotation rate = static. ............................................................................................156
Figure 5.2. CVs of the bare electrode (black dashed), CoCl2 metal precursor (blue), THT ligand
precursor (green), and CoTHT MOF (red) in pH 13 solution under O2. Scan rate = 10 mV/s.
Rotation rate = static. .......................................................................................................156
Figure 5.3. CVs of CoTHT at various rotation rates ranging from 0 (red) to 2500 (purple) rpm in
pH 13 solution under O2. Scan rate = 10 mV/s. ...............................................................157
Figure 5.4. CVs of CoTHT films prepared with Nafion (A) and Nafion and carbon black (B) at
pH 13 under O2. Scan rate = 10 mV/s. Rotation rate = static. .........................................158
Figure 5.5. 25 consecutive CVs of CoTHT ink mixture under O2 in pH 13 KOH solution, where
the current decreases and the overpotential increases after each scan. Rotation rate = 1600
rpm. Scan rate = 10 mV/s. ...............................................................................................159
Figure 5.6. XPS spectra of the S 2s and S 2p regions before (A and C) and after (B and D)
CVs. ................................................................................................................................160
Figure 5.7. XPS spectra of the Co 2p region before (A) and after (B) CVs. ..............................160
Figure 5.8. Co 2p (a), S 2s (b), and S 2p (c) XPS spectra of CoTHT following 10 days of O 2
exposure. ..........................................................................................................................161
Figure 5.9. 25 consecutive CVs of CoTHT ink prepared air free (red) and pre-oxidized CoTHT
ink (black) in pH 13 solution under O2. Scan rate = 10 mV/s. Rotation rate =
1600 rpm. .........................................................................................................................161
Figure 5.10. CV of CoTHT ink (red), bare glassy carbon electrode (black dashed), and carbon
black ink mixture (purple dotted) in pH 10 aqueous solution under O2 (red). Scan rate = 10
mV/s. Rotation rate = static. ............................................................................................168
Figure 5.11. Subsequent CVs of CoTHT ink in pH 10 aqueous solution under O2. Scan rate = 10
mV/s. Rotation rate = 1600 rpm. .....................................................................................168
Figure 5.12. CV of glassy carbon in pH 0.28 aqueous solution under N2 (blue) and O2 (red). Scan
rate = 10 mV/s. Rotation rate = 1600 rpm. ......................................................................169
Figure 5.13. CV of CoTHT in pH 0.28 aqueous solution under N2 (blue) and O2 (red). Scan rate
= 10 mV/s. Rotation rate = 1600 rpm. .............................................................................169
Figure 5.14. CV of glassy carbon in pH 1.16 aqueous solution under N2 (blue) and O2 (red). Scan
rate = 10 mV/s. Rotation rate = 1600 rpm. ......................................................................170
Figure 5.15. CV of CoTHT in pH 1.16 aqueous solution under N2 (blue) and O2 (red). Scan rate
= 10 mV/s. Rotation rate = 1600 rpm. .............................................................................170
Figure 6.1. Different functional groups that can be explored as direct proton relays to bound CO2.
..........................................................................................................................................177
Figure 6.2. General photocatalysis setup for COFs using a sacrificial donor species and lamp
illumination. .....................................................................................................................181

xviii

LIST OF SCHEMES

Scheme 1.1. Proposed pathway for the formation of a hydride bridged between two iron atoms
within the iron carbonyl cluster, followed by reduction of CO2 to formate. Reprinted with
permission from reference 59. ...........................................................................................11
Scheme 1.2. Proposed pathway for the formation of an iron-hydride species followed by reduction
of CO2 to formate. Reprinted with permission from reference 62. ....................................12
Scheme 1.3. Catalytic cycle of the Co cyclopentadienyl complex to reduce CO 2 to formic acid.
Reprinted with permission from reference 68. ..................................................................14
Scheme 1.4. Proposed pathway for the reduction of CO2 by NiTPEN and decomposition pathway.
Reprinted with permission from reference 83; permission conveyed through Copyright
Clearance Center, Inc. ........................................................................................................16
Scheme 1.5. Catalytic cycle to form oxalate from air using copper catalyst dimers. Reprinted from
reference 90. Published in Science by The American Association for the Advancement of
Science and Copyright Clearance Center, 2010. ...............................................................17
Scheme 2.1. Synthesis of L
NO2
from L
5
and synthesis of CoL
NO2
from L
NO2
. .............................36
Scheme 4.1. Syntheses of COF-Re, COF-Re_Co, and COF-Re_Fe. .......................................107
Scheme 4.2. Synthesis of 1 and 2. ...............................................................................................123
Scheme 5.1. Synthesis of CoTHT from cobalt and triphenylene hexathiol. ...............................155
Scheme 5.2. Deposition of CoTHT film onto glassy carbon through physical adsorption. ........155
Scheme 5.3. Proposed mechanism for the reaction between CoTHT (truncated for clarity) and O2
based on DFT calculations. ..............................................................................................163
Scheme 6.1. Synthesis of longer lariat arm macrocycles which are more air stable. ..................176
Scheme 6.2. Formation of diazonium group via the lariat arm and subsequent attachment to an
electrode surface. .............................................................................................................178

xix

ABSTRACT

The use of renewable energy in place of current fossil fuel technology has grown rapidly in its
practicality, but is still limited by insufficient storage methods. Due to the inability to store energy
long term, energy generated from solar panels must be used in the area and time frame within
which it is generated. This hinders the availability of energy at night, which is when a large amount
of energy is used. By developing catalysts which can use this energy to generate chemical fuels or
valuable industrial products, the reliance on fossil fuels can still be minimized. Success has been
seen with both solubilized homogeneous catalysts and surface-bound heterogeneous catalysts. In
this work, the development of novel homogeneous catalysts and the heterogenization of well-
known homogeneous catalysts into covalent organic and metal-organic frameworks will be
presented. These species will be applied for CO2 reduction and O2 reduction.

1

CHAPTER 1. GENERAL INTRODUCTION

1.1 Outlook on Global Energy
The current global dependence on fossil fuels for energy purposes has resulted in multiple
energy and environmental problems that will continue to grow in severity if left unchecked,
making the creation of sustainable energy systems one of the most pressing issues to date.
1-6
The
International Energy Agency has reported that the global energy demand had reached 18 TW in
2013, and that this number is predicted to grow by 25% by the year 2040 due to rising human
populations and growing industrialization, particularly in areas such as India, China, and African
countries.
7
In 2013, roughly 80% of global energy was produced from the consumption of fossil
fuels (coal, oil, and gas), resulting in a release of 32 Gt/year of carbon dioxide (CO2).
7
Rising
levels of atmospheric CO2 have already started impacting the environment through global warming
and ocean acidification, and the continued and increasing use of CO2-producing fuel sources will
only serve to exacerbate these issues.
8, 9
In order to most effectively combat CO2 emissions while
meeting global energy requirements, it is necessary to diversify our energy sources through the
use of carbon-free, renewable resources such as solar, wind, hydro, and geothermal power.
2, 3, 10-12

While integration of renewable energy sources into the global energy market has already
begun, there are growing concerns due to the intermittency of these energy producers.
1, 2
For
example, solar energy is available during the daylight hours, making night time energy
requirements difficult to meet. The electricity, transportation, and industrial sectors contributed to
39% of the global energy demand in 2010, and are expected to continue to grow, making the
development of reliable methods of energy storage paramount for meeting the needs of these
fields.
4
Transportation, in particular, relies heavily on liquid chemical fuels, so producing
sustainable methods of liquid fuel generation from renewable energy resources can be seen as a
crucial step in decarbonization of our energy infrastructure.
4

1.2 Solar Fuel Production
To combat intermittency and make renewable energy sources more viable for
transportation and industrial sectors, chemists have taken the approach of storing that energy in
the form of chemical bonds.
1, 4, 5, 13
An attractive strategy is to use renewable energy in the
formation of hydrogen, hydrocarbons, and other valuable fuel and feedstock chemicals.
1, 4-6, 13, 14

2

These chemical fuels and feedstocks would then serve as substitutes for areas typically dominated
by fossil fuels such as transportation (liquid fuels), chemical production industries (plastics and
fertilizers), and electricity.
4
For this goal, attention is drawn to chemicals which are already
abundant in the atmosphere such as carbon dioxide and water. Carbon dioxide, which is harmful
as well as relatively abundant, can be converted into hydrocarbons to serve as liquid fuel
replacements and into feedstock chemicals needed in the chemical production industry.
6, 8, 15

Splitting water (H2O) into oxygen and hydrogen is a potential pathway for generating hydrogen
for use as an energy dense, carbon-free fuel through the use of fuel cell technology.
16-21

Figure 1.1. Design of a standard H 2 fuel cell, where H 2 gas is oxidized at the anode (left side) and O 2 is reduced at
the cathode (right side).

However, converting CO2 and H2O into value-added products is complicated by the high energetic
cost of the conversion.
19-23
These chemicals have high thermodynamic barriers, often requiring the
formation or breakage of bonds with bond orders of two or higher. Selectivity may also be poor as
multiple products may be formed via various routes.
6, 21, 23
The design and implementation of
catalysts is needed to lower the activation barrier of these processes while also increasing the
selectivity and efficiency to obtain the desired products. For global deployment purposes, catalysts
materials must be abundant, inexpensive, operate under benign or environmentally considerate
conditions, and form the desired products with high yields and high energetic efficiency.

3

1.3 Transformation of Abundant Molecules to Solar Fuels
1.3.1 Carbon Dioxide Reduction Reaction
To understand how to appropriately design catalyst to facilitate carbon dioxide reduction
and water splitting, the barriers of these processes must first be understood. Carbon dioxide is a
linear molecule consisting of three atoms: a central carbon and two oxygen atoms that are both
double bonded to the carbon. All atoms that compromise CO2 have filled valence shells, making
the molecule resistant to reduction and oxidation.
6
Single electron reduction of CO2 requires the
charge neutral, linear molecule to rearrange into a bent radical anion. This change in bond order
and orbital rehybridization to accommodate an additional electron results in a high thermodynamic
barrier. CO2 has access to numerous reduction products such as carbon monoxide, formate,
methanol, methane, and other multi-carbon species such as hydrocarbons, requiring varying
amounts of protons and electrons.
24-27
As the number of protons and electrons used for the
reduction increase, the corresponding thermodynamic barrier decreases, but concerted multi-
electron, multi-proton processes are instead kinetically limited.
22

CO2 + e
-
 CO2·
-
E˚' = -1.90 V (Eq. 1)
CO2 + 2H
+
+ 2e
-
 CO + H2O E˚' = -0.53 V (Eq. 2)
CO2 + 2H
+
+ 2e
-
 HCOOH E˚' = -0.61 V (Eq. 3)
CO2 + 4H
+
+ 4e
-
 HCOH + H2O E˚' = -0.48 V (Eq. 4)
CO2 + 6H
+
+ 6e
-
 CH3OH + H2O E˚' = -0.38 V (Eq. 5)
CO2 + 8H
+
+ 8e
-
 CH4 + H2O E˚' = -0.24 V (Eq. 6)
Figure 1.2. CO 2 reduction processes and their standard reduction potentials.

Additionally, CO2 reduction often competes with the hydrogen evolution reaction, further
increasing the list of possible products and making selectivity for CO2 reduction challenging.
24-27

In literature, the most common CO2 reduction products are carbon monoxide (CO) and
formate (HCO2
-
).
6
Both have usefulness as solar fuel products, with carbon monoxide being used
for industrial processes via the Fischer-Tropsch process to access a variety of hydrocarbon
chemicals and with formate and its conjugate acid formic acid being used as reagents and fuels.
6,
28-30
Reduction to CO and HCO2
-
are both two electron, two proton processes, where the product
formed depends largely on whether a CO2 molecule or proton are bound to the metal center first.
4

CO formation occurs when a metal carboxylate adduct undergoes protonation at one of the oxygen
atoms, breaking the bonds between that oxygen and the carbon species and releasing CO and H2O
as products. Alternatively, insertion of CO2 into a metal hydride results in the formation of of a
formate adduct (HCO2
-
). However, formation of metal hydrides may also lead to insertion of a
proton to generate hydrogen through the hydrogen evolution reaction, again hindering selectivity.
For a reduction that has a high energetic barrier and multiple competing reactions, careful
consideration must be put into design choices for catalysts to effectively produce desired solar
fuels.

1.3.2 Oxygen Reduction Reaction (ORR)
The use of hydrogen in fuel cell applications relies on two reactions: the oxidation of
hydrogen and the reduction of oxygen to overall produce water.
31, 32
Hydrogen is the supplier of
energy, and its formation from the reduction of protons has been optimized in part due to its
simplicity. The oxygen reduction half reaction, on the other hand, is sluggish and complex and
therefore seen as the bottleneck of this process.
33
Compared to the hydrogen side of the reaction,
which takes place at 0 V vs RHE, the oxygen half reaction has a large overpotential at 1.23 V vs
RHE.

Figure 1.3. Half reactions and potentials for the water splitting reaction.

The reduction of oxygen to water is formally a four electron, four proton process, but like CO 2,
there are multiple products and pathways that have been observed.
19-21
A two electron reduction
to form peroxides may also take place, producing peroxide intermediates.
19-21
While hydrogen
peroxide is an industrially valuable chemical,
21
the addition of a competing reaction complicates
the oxygen reduction process.
In addition to the barriers of the reaction, the most successful catalysts for oxygen reduction
are rare and expensive noble metals. The use of platinum and platinum group metals (PGMs) is
cost prohibitive for global deployment, but other metals are too inefficient to perform ORR as
5

needed.
32, 33
Even as catalysts with the highest efficiency for ORR, PGMs are not able to perform
ORR at rates sufficient to compete with hydrogen oxidation, thus keeping ORR as the bottleneck
of fuel cells.
33
In order for hydrogen fuel cell technology to be applicable on a global scale, greater
progress must be made in employing cheaper and more abundant materials.
In the realm of catalysis, there are typically considered to be two types of catalysts:
homogeneous and heterogeneous. Homogeneous catalysts are easily tuned and optimized through
ligand synthesis and modification and are isolable in various intermediate stages, providing
mechanistic insight. However, they suffer from limited stability and poor solubility in aqueous
media. Heterogeneous catalysts are seen to have the opposite problem; they perform in various
media and with long lifetimes but display poor selectivity and cannot be easily modified.
Understanding both types of catalysis is crucial for developing scalable catalysts for solar fuels.

1.4 Homogeneous Electrocatalysts for CO2 Reduction
A wide variety of homogeneous catalysts exist in the literature for solar fuel production.
6,
8, 18, 34-39
For the purposes of limiting the scope of this section to only applicable catalysts, only
species containing late first row transition metals which perform CO2 reduction will be discussed.

1.4.1 Manganese Catalysts
Manganese bipyridine catalysts have recently gained popularity as alternatives to rhenium
bipyridine complexes as manganese is more abundant by multiple orders of magnitude. Initial
studies investigated the fac-[Mn(bpy-R)(CO)3Br] species (R = H, Me), where bpy-R is 4,4’-
disubstituted 2,2’-bipyridine.
40
The [Mn(bpy)(CO)3Br] complex operates at overpotentials roughly
400 mV more favorable than its Re counterpart and still selectively produces CO from CO 2, and
substituting the 4 and 4’ positions with methyl groups leads to an increase in selectivity and
stability lasting several hours.
40
However, while Re complexes reduce CO2 without any added
proton source, Mn complexes are incapable of this. Building on the improvement noted for the
methyl substituted species, bipyridine complexes with t-butyl groups in the 4,4’ positions which
resulted in turnover frequencies (TOFmax) of up to 340 s
-1
in the presence of 1.4 M 2,2,2-
trifluoroethanol (TFE).
41
For these complexes, mechanistic insight was gained through study of
two irreversible, one electron reduction event seen by cyclic voltammetry (CV) combined with
spectroscopic insight, DFT calculations, and comparison to the well-studied rhenium complexes.
6

The proposed mechanism suggests a release of the axial halide ligand after the first reduction
event, followed by dimerization. The dimer is then reduced again, producing the five-coordinate
monomer species.
The bipyridine ligand was next modified in the 6,6’ positions with bulky mesityl groups
(Mes-bpy), resulting in a change in the CV spectrum from two irreversible, one electron events to
a single quasireversible two electron event.
42
Due to the steric bulk of the mesityl groups, the
reduced species progresses from the dehalogenated five coordinate species directly to the doubly
reduced species without forming and breaking the dimer species. As with the previous manganese
complexes, the Mes-bpy complex does not reduce CO2 without an added proton source, but it is
capable of reducing CO2 to CO with 98% faradaic efficiency while remaining stable for over 7
hours with it. With added methanol (3.2 M) and TFE (1.4 M), the complex reduces CO 2 with
TOFmax of 2000 s
-1
and 5000 s
-1
, respectively, making this catalyst a full order of magnitude faster
than the t-Bu-bpy species. In addition to enhanced performance with Brønsted acids, added Lewis
acids such as Mg
2+
also have an effect on catalysis.
43
CO2 can be reduced to CO and CO3 through
this method, but in aprotic media the carbonate salts precipitate from solution and may hinder
catalysis by blocking the electrode surface.
43

By replacing the mesityl groups with 2,6-dimethoxyphenyl groups, dimerization is still
hindered but the methoxy groups facilitate the quickening of the rate-determining dehydration
step.
44
Through the methoxy groups, hydrogen bonding interactions occur which stabilize the
metal-carboxylic acid intermediate. Coupling this with the electronic substituent effects makes the
cleavage of the C-OH bond becomes much more favorable, allowing for protonation to occur
before reduction.
44

7

Figure 1.4. Catalytic mechanism of Mn bipyridine with (red) and without (green) the added methoxy substituent.
Reprinted with permission from reference 6.

Further improvements can be made by changing the methoxy groups to hydroxyl
groups.
45,46
These complexes similarly reduce the energy barrier for the C-OH bond cleavage step
and show improved rates and overpotentials from the parent [Mn(bpy)(CO)3Br] complex, but also
exhibit formate generation which was previously unobserved for manganese bipyridines.
45

1.4.2 Iron Catalysts
1.4.2.1 Porphyrins
Iron porphyrins are one of the most well studied classes of CO2 reduction catalysts to
date.
47-57
The first studied species was iron tetraphenyl porphyrin, [(tpp)Fe
III
]Cl, which was found
to reduce CO2 to CO in protic solvents but rapidly deactivate in aprotic ones.
47, 48
The addition of
Lewis acids, both mono- and divalent, increased the activity of the porphyrin complex similarly to
8

the manganese bipyridine complexes.
48, 51
Of the Lewis acids studies, a trend was observed where
Mg
2+
≈ Ca
2+
> Ba
2+
> Li
+
> Na
+
, with the identity of the Lewis acid playing a role in the amount
of byproduct formed (10-30%) by bulk electrolysis.
51
The proposed catalytic mechanism begins
with reduction to the Fe
0
species, which nucleophilically attacks the carbon of CO2 to form an
adduct.
51
From there, the metal carboxylate can either be protonated twice to form CO, or, in the
presence of Mg
2+
as a Lewis acid, reacts with a second CO2 to form a carbonate and Fe-carbonyl
complex.
51
The addition of weak Brønsted acids also enhanced catalysis, with selectivity
increasing with acid strength. Using 1-propanol as a proton source gave a 60/35% mixture of CO
and formate, respectively, while TFE produced CO with 96% faradaic efficiency.
47, 50
Production
of formate with weaker acids is hypothesized to be due to the weak hydrogen bonding interactions
promoting a more basic carbon adduct, making protonation of the carbon more favorable than the
oxygen.
52

Similar to Mn bpy compexes, the substitution of pendant phenolic groups in place of the
arene groups results in a substantial improvement in catalytic activity for iron porphyrins. The
phenol-substituted porphyrin outperformed a methoxy-substituted porphyrin through all potentials
according to Tafel analysis, reaching TOF values which were 9 orders of magnitude higher.
53
The
presence of the phenolic groups influences the electron transfer steps; the stabilization of the bound
CO2 results in the second reduction being more difficult than the first, which is the opposite of
Fe(tpp).
54

Figure 1.5. Structures of the hydroxy (left) and methoxy (right) iron porphyrin complexes. Reprinted from reference
53 with permission from AAAS.

9

Detailed study of the effect of the phenolic groups showed that the benefits of these substituents
are two-fold: the Fe-CO2 adduct is stabilized and the local proton concentration around the active
site is increased.
54

The effects of electron withdrawing groups on the activity of Fe(tpp) was investigated
through multiple methods. The installment of perfluorophenyl groups in various amounts (1, 2,
and 4) in place of the arenes caused two competing properties to be imparted onto the complex.
55

By withdrawing electron density from the metal center, reduction of the metal center was
accomplished at more positive potentials which results in lower overpotentials. The loss in electron
density is also detrimental, however, due to the lowered nucleophilicity of the metal center
resulting in slower rate constants. Porphyrins with trimethylammonium groups in the para position
perform CO2 reduction in DMF similar to other Fe(tpp) complexes, but have the added benefit of
operating in aqueous conditions.
56, 57
Selective CO2 reduction in water is valuable but rare due to
the competing hydrogen evolution pathway. Moving the trimethylammonium groups from the para
position to the ortho position results in a marked improvement in overpotential (220 mV) and
TOFmax (10
6
s
-1
).
57
This is contrary to the results of the perfluorophenyl groups, and the
trmethylammonium substituted porphyrin complex remains as one of the most efficient
homogeneous CO2 reduction catalysts to date.

Figure 1.6. Family of iron porphyrins containing different electron withdrawing groups. Adapted from references.
Adapted with permission from references 55 and 57.

10

1.4.2.2 Iron Carbonyl Clusters
Tetrairon clusters held together through an interstitial carbon or nitrogen have been
successfully used as CO2 to formate catalysts in both organic and aqueous media.
58-61
The species
with an interstitial nitrogen atom had the best performance of the catalysts studied, displaying high
selectivity (95%) for CO2 to formate even in aqueous media with a pH range of 5 to 13.
59
In the
absence of CO2, hydrogen evolution occurs which suggests that formate production is kinetically
favorable with this system.
59
The appearance of a new oxidation feature by CV at 0.5 V vs SCE
suggests the formation of a hydride, which would be expected for a catalyst which produces
formate and hydrogen. This species is further confirmed to be the active species by use of
nonelectrochemical means to produce formate from the hydride intermediate [HFe4N(CO)12]
2-
.
59

Catalyst selectivity displays a dependence on H2O, but the selectivity uniquely increases with
added water. The hydricity of [HFe4N(CO)12]
2-
is dependent on solvent, with the implication that
there are more favorable thermodynamics for catalysis in aqueous media than organic. Selectivity
is further impacted by hydricity in the case of [Fe4C(CO)12]
2-
, where the interstitial carbon atom
species has a higher hydride-donating ability and therefore only forms hydrogen even under a CO2
atmosphere.
59

Figure 1.7. Structures of the iron carbonyl clusters with interstitial N atoms where L = CO, PPh 3, or PPh 2(CH 2) 2OH.
Reprinted with permission from reference 61. Published by the Royal Society of Chemistry.
11

Scheme 1.1. Proposed pathway for the formation of a hydride bridged between two iron atoms within the iron
carbonyl cluster, followed by reduction of CO 2 to formate. Reprinted with permission from reference 59.

Attempts to introduce proton donors into the nitrogen bridged cluster met with mixed
success. [Fe4N(CO)11(PPh3)]
-
and [Fe4N(CO)11(PPh2(CH2)2OH)]
-
saw negative shifts in reduction
potential due to the donating nature of the added ligands, driving the reduction potential into a
realm where competing HER from the glassy carbon working electrode is now present.
61
The
addition of the PPh3 ligand still generates formate with H2 formation attributed entirely to the
glassy carbon electrode, but the PPh2(CH2)2OH containing species forms H2 only.
61
This drastic
change in selectivity is attributed to the proximity of the proton donor to the active site as well as
proton delivery kinetics.

1.4.2.3 Hydroxybenzene Bipyridines
Iron complexes generated from 6,6’-di(3,5-di-tert-butyl-2-hydroxybenzene)-2,2’-
bipyridine displayed competency as CO2 reduction catalysts, but are notable in that the products
formed are largely dependent on the proton environment in solution.
62
In the absence of any added
protons, CO and CO3 are generated via the disproportionation of two CO2 molecules. However,
when phenol is added to the solution, a mixture of products is detected with formate (68 ± 4% FE)
being the dominant product. H2 (30 ± 10% FE) and CO (1.1 ± 0.3% FE) were also detected. Results
from UV-Vis spectroscopy, SEC-IR, cyclic voltammetry, and controlled potential electrolysis
helped inform a mechanism where a protonated hydroxybenzene version of the complex serves as
12

a precatalyst. Following further reduction, a chloride is lost from the iron which is then capable of
forming an iron-hydride. This hydride is what promotes the formation of both H2 and formate.
Kinetic isotope effects using deuterated phenol are also consistent with the insertion of CO2 into a
metal hydride.

Scheme 1.2. Proposed pathway for the formation of an iron-hydride species followed by reduction of CO 2 to
formate. Reprinted with permission from reference 62.

1.4.3 Cobalt Catalysts
1.4.3.1 Tetra-Aza Macrocycles
Cobalt complexes are rarely seen to be highly selective CO 2 reduction catalysts, often
having competitive HER be more favorable.
63-65
The cobalt containing macrocycle
[Co
III
N4H(Br)2]Br forms CO with 45% faradaic efficiency in wet acetonitrile, but also produces
H2 with 30% FE.
65

13

Figure 1.8. Structure of the cobalt tetra-aza macrocycle. Reprinted with permission from reference 65.

The catalytic activity increased upon the addition of added H2O. Further spectroscopic-theoretical
studies aimed to elucidate a possible mechanism for this catalyst, suggesting that CO 2 binding
through the carbon atom results in the adduct being only partially charged and slightly bent.
65, 66

This catalyst is also capable of turning over using photochemical conditions when photosensitizer
[Ir(ppy)3] and sacrificial donor NEt3 were added to the catalyst in solution with CO2.
65, 67

1.4.3.2 Cyclopentadienyl Complexes
A cobalt complex containing a cyclopentadiene ligand and diphosphine ligand P
R
2N
R
’2,
which contains two pendant amine moieties, has been seen to selectively produce formic acid from
CO2 with TOFmax of 1000 s
-1
and overpotentials in the range of 500-700 mV, making this complex
one of the best for homogeneous formic acid catalysts.
68
The properties of the phosphines and
amines were modulated, revealing that the catalyst performance increases with amine basicity and
phosphine electron-donating ability.
68
The proposed mechanism involves an initial one electron
reduction coupled with the loss of the iodide ligand to enter the catalytic cycle.
68
A second
reduction is coupled with proton addition to the complex at the pendant amine, which is then
transferred to the metal to form the metal hydride following the addition of another electron.
68
CO2
can then coordinate to the hydride, being stabilized by H2O as a proton source which is being
directed by the pendant amine groups.
68
The formation of the C-H bond from CO2 in this system
is proposed to occur via hydride transfer from Co to C instead of the typical metal-hydride
insertion, resembling traditional hydride transfer catalysts.
69, 70

14

Scheme 1.3. Catalytic cycle of the Co cyclopentadienyl complex to reduce CO 2 to formic acid. Reprinted with
permission from reference 68.

1.4.4 Nickel Catalysts
1.4.4.1 Cyclam
Nickel cyclam systems are some of the first studied catalysts for CO 2 reduction, and also
one of the most well studied.
71-81
[Ni(cyclam)]
2+
can exist in various conformations, most notably
trans I and trans III conformations (15% and 85%, respectively) in water.
82

Figure 1.9. The trans I (left) and trans III (right) forms of [Ni(cyclam)]
2+
. Reprinted with permission from reference
73.
15

Trans I places all amino protons in the same plane and appears geometrically more bowl-shaped
than the trans III conformation, which has two protons in one plane and two in the other and
appears more flat. This catalyst is unique in that its selectivity can be dictated by the electrolyte
and media in which catalysis is performed; formate is favored in aprotic systems (75% FE) while
CO is favored in aqueous conditions with pH values between 4 and 5.
76
Further studies of catalyst
behavior with a mercury working electrode showed little dependence on the catalyst concentration,
implying that the catalyst adsorbed onto the Hg electrode was the main active species.
77
When the
working electrode is changed to a glassy carbon to inhibit adsorption to the electrode, the catalyst
could still reduce CO2 to CO but with slower rates and more negative potentials.
78

While Ni(cyclam) has fascinating properties, it is limited in its deployment due to issues
of deactivation. The reduction of CO2 to CO results in catalyst poisoning as the active Ni species
binds CO, lowering the amount of available Ni(cyclam) available for catalysis as it is trapped in a
non-catalytic pathway.
79
Further reduction to the Ni(0) species represents an irreversible
deactivation, further limiting the available amount of catalyst.
79
These deactivation pathways can
be prevented, however, through the use of CO scavenger species. Addition of a methylated variant
of Ni(cyclam) has been used to prevent deactivation as it has a stronger affinity for CO and lower
affinity for CO2, meaning it does not directly compete with Ni(cyclam) as a catalyst but instead
prevents poisoning.
72

The addition of methyl groups to the carbon backbone of cyclam results in even better
performance in water, achieving higher currents and lower overpotentials while maintaining
selectivity for CO.
80, 81

Figure 1.10. Methylated variants of cyclam. Reprinted with permission from reference 80.

The selectivity can also be tuned through the manipulation of the pH. At acidic pH values (<2), H2
becomes the dominant product, allowing for syngas formation when combined with the CO2 to
16

CO reduction process.
81
The catalytic activity is highly dependent on the stereochemistry of the
added methyl groups. While RRSS-[Ni(HTIM)]
2+
outperforms the parent complex [Ni(cyclam)]
2+
,
its stereoisomer RSSR-[Ni(HTIM)]
2+
performs worse.
81

1.4.4.2 Polypyridyl
The ligand N,N,N′,N′-tetrakis(2-pyridylmethyl)ethylenediamine (TPEN) coordinated to nickel
is capable of reducing CO2 to CO.
83
In both the absence and presence of phenol as a proton source, CO
is formed. Similarly to the cyclam examples, generation of CO leads to catalyst poisoning, and Ni(CO)4
is rapidly produced. The use of a CO sponge can help to mitigate some of the poisoning, but
decomposition cannot be eliminated entirely without further modifications of the system.

Scheme 1.4. Proposed pathway for the reduction of CO 2 by NiTPEN and decomposition pathway. Reprinted with
permission from reference 83; permission conveyed through Copyright Clearance Center, Inc.

1.4.5 Copper Catalysts
Homogeneous copper catalysts for CO2 reduction are extremely rare,
84-89
but one such case
exists where a bimetallic copper catalyst is capable of reducing CO2 to oxalate.
90
This complex
can form oxalate from the CO2 in air, a rare feat for homogeneous catalysts.
90
However, the
performance of this complex is poor, reaching only 6 turnovers after 7 hours of electrolysis.
90
The
poor number of turnovers is likely in part due to the thermodynamic stability of the oxalate bridged
17

dimeric complex.
90
Treatment of the oxalate containing complex with LiClO4 is necessary to
remove oxalate as the lithium salt in order to continue catalysis.
90

Scheme 1.5. Catalytic cycle to form oxalate from air using copper catalyst dimers. Reprinted from reference 90.
Published in Science by The American Association for the Advancement of Science and Copyright Clearance
Center, 2010.

1.5 Heterogenization of Molecular Catalysts
Study of homogeneous catalysts is crucial for understanding how to best design selective
catalysts and understand their catalytic pathways, but global deployment of catalysts for solar fuel
production relies more heavily on long lifetimes and use of benign solvents. To make use of the
valuable aspects of both homogeneous and heterogeneous catalysts, homogeneous catalysts can
be heterogenized by immobilizing them in covalent organic frameworks (COFs) and metal organic
frameworks (MOFs). Through this approach, well-studied homogeneous catalysts can gain the
longer lifetimes and solvent operability needed to be industrially viable. While many such COFs
and MOFs have already been investigated for solar fuel applications, the scope of this work will
be limited to CO2 and O2 reduction reactions.

18

1.5.1 CO2 Reduction
1.5.1.1 Cobalt Porphyrins
Numerous studies incorporating cobalt porphyrins into extended frameworks have been
reported.
91
The first example of an immobilized homogeneous catalyst capable of CO2 reduction
was a COF comprised of a cobalt porphyrin with a phenyl or biphenyl linker (COF-366-Co and
COF-367-Co, respectively).
92
While cobalt porphyrins are known to be less selective than their
iron porphyrin counterparts, often forming H2 in high amounts, this COF forms CO from CO2 with
high selectivity (90+%) even under aqueous conditions.
92
The activity of the catalyst was proven
to increase as a function of organic linker length, with COF-367-Co reaching turnover numbers
(TON) of 3901 versus TON of 1352 for COF-366-Co.
92
Determination of the number of
electrocatalytically active sites revealed that the number of Co sites actually contributing to
catalysis was <10%.
92
The synthesis was modified such that the amount of Co was diluted by
introducing controlled amounts of copper porphyrin as an inactive porphyrin species. Diluting
COF-367-Co to 10% and 1% Co resulted in large increases in TOF to 4400 and 9400, respectively,
though the selectivity dropped to 70% and 40% FE for CO.
92
As these frameworks are effectively
heterogeneous, performing long-term electrolysis for 136 hours resulted in an overall TON of
24,000 which shows a marked improvement over homogeneous catalysts.
92

COF-366-Co was further investigated by modifying the electron withdrawing and donating
effects of the organic linker. By introducing electron withdrawing fluorinated organic linkers into
COF-366-Co, the current density increases from 45 mA/mg to 65 mA/mg while maintaining high
selectivity for CO2 reduction to CO in water.
93
Physical characterization studies proved that the
electronic properties of the linker influenced the electronics of the active site.
93
While a
tetrafluorinated linker as the strongest electron withdrawing would be expected to have the best
influence on the catalyst, it was actually one of the worst tested, which is believed to be due to
increased hydrophobicity of the COF.
93

A cobalt porphyrin MOF, [Al2(OH)2TCPP-Co], was grown through atomic layer
deposition onto an electrode surface and tested electrocatalytically under identical conditions.
94

The selectivity of this MOF was noticeably less than that of the COFs, reaching only 76% FE for
CO.
94
Controlled potential electrolysis studies showed stability for up to 7 hours, resulting in a
TON of 1400.
94
Physical characterization after 7 hours of controlled potential electrolysis showed
retention of features observed prior to electrolysis, again supporting the stability of the MOF.
94
By
19

using atomic layer deposition to grow the MOF, the electrochemical properties affected by film
thickness could be optimized, reaching the best charge transport, active site density, and mass
transport properties after 50 cycles.
94

Figure 1.11. CO 2 reduction catalysts (a) COF-366-Co and COF-367-Co, (b) COF-366-Co modified through the
organic linker, and (c) the MOF [Al 2(OH) 2TCPP-Co]. Adapted with permission from reference 91. Published in
Nano Research by Springer Nature, 2019.

1.5.1.2 Iron Porphyrins
Iron porphyrins have also been immobilized through incorporation into COFs and MOFs.
An iron porphyrin COF was used in a solvent-free synthesis with 2,6-dihydroxyterephthalaldehyde
to grow a thin film onto a gold substrate.
95
The porphyrin and linker are dissolved in volatile
solvents, and the substrate is dipped in one, allowed to dry, and dipped in the other before being
heated to produce the desired COF film.
95
The COF displayed a respectable 80% faradaic
efficiency for CO2 to CO with a TOF of >600 h
-1
mol
-1
based only on electroactive iron sites in
MeCN electrolyte solution.
95
Changing the solvent to DMF resulted in a loss in selectivity for CO
which was explained through the higher solubility of free iron porphyrin in DMF, resulting in the
active material effectively leaching off the electrode.
95
The COF was also tested under aqueous
conditions, but saw a drastic shift in selectivity to 80% H2 and only 20% CO.
95
As iron porphyrins
are effective catalysts on their own, incorporating the molecular species into a COF worsened its
catalytic activity but instead kept the material on the electrode which produced a greater TON of
CO.
An iron porphyrin MOF was synthesized by post synthetically modifying a porphyrin MOF
linked through hexazirconium(IV) nodes with FeCl3, forming the desired iron porphyrins with
93% incorporation of iron.
96
The MOF, MOF-525, was then deposited onto FTO substrates
20

through electrophoretic deposition.
96
Controlled potential electrolysis of the MOF showed
formation of CO and H2 in nearly equal ratios (54% and 45%, respectively), which has value as a
syngas mixture.
96
Upon the addition of TFE as a proton source, the selectivity for CO decreased
to 41% by FE with H2 formation increasing by the same amount.
96
The catalyst lifetime was also
shortened upon the addition of TFE, as proven by postcatalysis studies.
96
With many COFs and
MOFs, there is a concern over effective electron transport as most frameworks are considered
insulating.
96
MOF-525 utilizes the positioning of iron sites in the MOF to facilitate redox charge
hopping, allowing electrons to make it from the electrode to the diffusion layer.
96
Even with charge
hopping, though, the electron transport in MOF-525 is still slow and thus the molecular complex
has a TOF 16 times higher than the MOF.
96

Figure 1.12. MOF-525 on an FTO electrode where the electrons can hop through the available iron sites to reach the
CO 2 substrate. Reprinted with permission from reference 96.

1.5.1.3 Rhenium Bipyridines
While many photocatalytic frameworks based on rhenium bipyridine are present in the
literature,
97-100
this work focuses on electrocatalysis and thus they are outside the scope of this
work. Rhenium bipyridine was modified with carboxylic acid moieties in the 5 and 5’ positions,
and then integrated into a surface-grafted Zn-based MOF through the use of liquid-phase
epitaxy.
101

21

Figure 1.13. Formation of a rhenium bipyridine-based MOF on an FTO surface through liquid-phase epitaxy.
Reprinted with permission of Journal of Materials Chemistry A from reference 101; permission conveyed through
Copyright Clearance Center, Inc.

Electrolysis studies were conducted in acetonitrile solution with 5% TFE as an added proton
source, and CO2 reduction to CO was the dominant process with only a small amount of detectable
H2.
101
The catalyst was only stable for about 30 minutes before rapidly decomposing and
deactivating, as shown by the loss of diffraction peaks in the X-ray diffraction spectrum.
101
Due to
decomposition, TON could only be measured for up to 2 hours, reaching a value of 580.
101

1.5.2 O2 Reduction
1.5.2.1 Triphenylene-based MOFs
Ni3(HITP)2 (HITP = 2,3,6,7,10,11-hexaiminotriphenylene) was studied for ORR in 0.1 M
KOH solution.
33
This material exhibited relatively high conductivity for MOFs, reaching 40 S cm
-
1
as a film and 2 S cm
-1
as a bulk pellet.
33
Films were grown through a solvothermal synthesis
directly onto glassy carbon electrodes, and analysis by CV showed an onset of catalysis at 0.82 V
vs RHE which is a 0.18 V overpotential relative to platinum (1.00 V vs RHE).
33
The MOF showed
high stability, maintaining 88% of its current density after 8 hours of controlled potential
electrolysis.
33
The catalyst is largely a peroxide forming catalyst as the calculated number of
electrons passed was 2.25, or an 87.5% faradaic efficiency for H2O2 with the rest being attributed
to H2O formation.
33
Conducting electrolysis at more positive potentials shifts the selectivity of the
catalyst, with water formation reaching its peak at 0.75 V vs RHE.
33
DFT analysis of the catalyst
coupled with experimental analysis of the Ni oxidation state during catalysis revealed that the
active site of catalysis is most likely the β carbon of the triphenylene ring instead of the Ni metal
center or binding N atoms as is typical for this class of MOFs.
102

22

From this initial study, a family of catalysts including Cu 3(HITP)2, Cu3(HHTP)2,
Ni3(HHTP)2, and Co3(HHTP)2 (HHTP = 2,3,6,7,10,11-hexahydroxytriphenylene) were
synthesized and evaluated for ORR capabilities.
103
While the Cu3(HITP)2 species has the highest
current by CV of this family, it rapidly deactivates under catalytic conditions.
103
This study
demonstrated the importance of stacking on catalysis, as Ni3(HITP)2, Cu3(HITP)2, and
Cu3(HHTP)2 (hexagonal crystal system) where the layers are slip stacked outperformed
Ni3(HHTP)2, and Co3(HHTP)2 (trigonal crystal system) where each layer is separated by trinuclear
clusters rotated 60˚ from the 2D honeycomb lattice layers.
103

Figure 1.14. General structures showing the packing of triphenylene MOFs in hexagonal (a) and trigonal (b) phases.
Reprinted with permission from reference 103. Published by the Royal Society of Chemistry.

The HHTP systems all deactivated at faster rates than those comprised of HITP, which is
believed to be due to better energetic and/or spatial overlap between HITP and the metal.
103
The
HHTP bonds are weaker and could therefore make decomposition more favorable.
103
In all cases,
the MOFs produced H2O at more negative potentials (~300 mV overpotential or more) and H2O2
at more positive ones.
103
This class of materials is unique in that attempts to replicate the reactivity
with a molecular analogue were unsuccessful, making it one of few examples where the MOF is
not based on an already successful homogeneous catalyst.
Additional studies on the nickel and cobalt HHTP MOFs also confirmed ORR activity. The
Co-CAT and Ni-CAT species showed current enhancements via CV under an atmosphere of O2 at
-0.576 V and -0.666 V vs Ag/AgCl, respectively.
104
When carbon black was added as a conductive
support, the onsets of these current increase improved by roughly 300 mV for both complexes,
23

achieving a roughly 200 mV overpotential relative to Pt/C.
104
Both catalysts produced H2O as the
dominant product, with Ni-CAT being the more selective of the two.
104

Figure 1.15. The structure of Ni-CAT where the two axial water species are omitted for clarity. Adapted with
permission from reference 104. Copyright 2017 American Chemical Society.

1.5.2.2 Iron Porphyrins
Porphyrins have also been shown to be effective oxygen reduction catalysts, as an iron
porphyrin MOF constructed with Zr6 oxo clusters as nodes (PCN-223-Fe) was grown
solvothermally on FTO for electrochemical experimentation.
105
As with the iron porphyrins used
for CO2 reduction, this MOF relies on charge hopping through available metal sites to facilitate
catalysis.
105

Figure 1.16. The (a) (001) and (b) (100) planes of PCN-223-Fe showing its similarity to MOF-525 and how charge
hopping may occur through the porphyrin sites. Reprinted with permission from reference 105.
24

While most ORR catalysts are tested under aqueous conditions, this catalyst was tested in DMF.
A current increase was observed at potentials below -0.4 V vs NHE when subjected to an O2
atmosphere, and the current increased even further when acetic acid or trichloroacetic acid was
added as a proton source.
105
Electrolysis at various potentials revealed that H2O production was
more favorable at more negative potentials (-0.6 V vs NHE and beyond).
105
Product distribution
is also dependent on proton source as acetic acid at negative potentials only produces 6% H2O2
while trichloroacetic acid produces 34% at similar potentials.
105
Use of trichloroacetic acid also
results in rapid current decreases compared to acetic acid, which is attributed to the formation of
an amorphous coating on the electrode surface which may be inhibiting catalysis.
105

1.6 Outline of Thesis
While much progress has been made in the field of catalysts for solar fuel production, there
is still new information being discovered every day. Each new catalyst developed brings with it
unique activity or qualities that further our understanding of how to design the best catalyst for the
desired application. This work will present insights into both homogeneous and heterogeneous
catalyst design for both CO2 and O2 reduction reactions. Modifications to known catalysts will be
discussed, and their influence on the selectivity of said catalysts will be explored. The synthesis of
heterogenized materials for CO2 and O2 reduction will then be examined, and their applicability to
these reactions will be discussed.
In chapter two, the synthesis of a cobalt tetra-aza macrocycle containing a pendant nitro
group will be presented. This work builds upon previously published cobalt tetra-aza macrocycle
work in our group, where the activity was directly related to the number of secondary amine groups
in the catalyst. The effect of the nitro group on CO2 reduction catalysis is examined and explained
in detail in order to better understand the influence that charged and potentially chelating
functionalities can have on known catalysts. The dependence of the catalyst on solvent will be
demonstrated, as will the shift in selectivity from the parent complex with an analysis on similar
responses in the literature.
Chapter three will discuss further studies with the cobalt lariat macrocycle, where CO 2
catalysis was tested in dimethylformamide to more closely resemble the conditions used for the
macrocycles without the lariat arm. Under these conditions, HER was still the dominant reaction
25

pathway but unique chemistry was observed under these conditions. The controlled potential
electrolysis under an inert atmosphere (N2) is shown to produce CO. The methods of determining
where the CO originates from are discussed.
Chapter four will examine the design of a covalent organic framework with two distinct
metal sites comprised of two well studied CO2 reduction catalysts and its application towards CO2
reduction. In this work, a novel homogeneous catalyst was synthesized as was necessary to
synthesize the desired COF, and the characterization and electrochemical analysis of this complex
is presented here as well. Finally, the synthesis of three COFs containing rhenium and either cobalt
or iron metal sites is provided along with structural characterization through powder X-ray
diffraction, infrared spectroscopy, and X-ray photoelectron spectroscopy. The COFs were tested
for CO2 reduction activity in aqueous conditions similar to those previously reported for porphyrin
COFs.
Chapter five looks at a cobalt 2,3,6,7,10,11-hexathiotriphenylene metal organic framework
for oxygen reduction. The catalyst is compared to the known hexaimino- and
hexahydroxytriphenylene systems as structural analogues, and a comparison between
experimental results and theoretical predictions will be provided. The catalyst was tested under
standard oxygen reduction conditions (0.1 M KOH solution) and the results of that
experimentation will be explained and future directions discussed.
Chapter six will encompass all projects as future directions are suggested. The successes
and failures of each project will be addressed in regards to where the project can go and what
experiments may have the most value.

26

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95. Cheung, P. L.; Lee, S. K.; Kubiak, C. P., Facile Solvent-Free Synthesis of Thin Iron
Porphyrin COFs on Carbon Cloth Electrodes for CO2 Reduction. Chemistry of Materials
2019, 31 (6), 1908-1919.
96. Hod, I.; Sampson, M. D.; Deria, P.; Kubiak, C. P.; Farha, O. K.; Hupp, J. T., Fe-
Porphyrin-Based Metal–Organic Framework Films as High-Surface Concentration,
Heterogeneous Catalysts for Electrochemical Reduction of CO2. ACS Catalysis 2015, 5
(11), 6302-6309.
97. Choi, K. M.; Kim, D.; Rungtaweevoranit, B.; Trickett, C. A.; Barmanbek, J. T. D.;
Alshammari, A. S.; Yang, P.; Yaghi, O. M., Plasmon-Enhanced Photocatalytic CO2
Conversion within Metal–Organic Frameworks under Visible Light. Journal of the
American Chemical Society 2017, 139 (1), 356-362.
98. Deng, X.; Albero, J.; Xu, L.; García, H.; Li, Z., Construction of a Stable Ru–Re Hybrid
System Based on Multifunctional MOF-253 for Efficient Photocatalytic CO2 Reduction.
Inorganic Chemistry 2018, 57 (14), 8276-8286.
99. Easun, T. L.; Jia, J.; Calladine, J. A.; Blackmore, D. L.; Stapleton, C. S.; Vuong, K.
Q.; Champness, N. R.; George, M. W., Photochemistry in a 3D Metal–Organic
Framework (MOF): Monitoring Intermediates and Reactivity of the fac-to-mer
Photoisomerization of Re(diimine)(CO)3Cl Incorporated in a MOF. Inorganic Chemistry
2014, 53 (5), 2606-2612.
100. Wang, C.; Xie, Z.; deKrafft, K. E.; Lin, W., Doping Metal–Organic Frameworks for
Water Oxidation, Carbon Dioxide Reduction, and Organic Photocatalysis. Journal of the
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101. Ye, L.; Liu, J.; Gao, Y.; Gong, C.; Addicoat, M.; Heine, T.; Woll, C.; Sun, L., Highly
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15320-15326.
33

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(39), 33539-33543.

34

CHAPTER 2. SWITCHING CATALYST SELECTIVITY VIA THE INTRODUCTION
OF A PENDANT NITRO GROUP

2.1 ABSTRACT
The conversion of abundant small molecules to value-added products serves as an
attractive method to store renewable energy in chemical bonds. A family of cobalt aminopyridine
complexes was previously reported to reduce CO2 to CO with 98% faradaic efficiency by
generating a high local acid concentration through hydrogen bonding networks. The addition of a
nitrophenyl moiety to the ligand backbone results in a drastic shift in selectivity. Large current
increases in the presence of added protons are seen under both N2 and CO2, and a reduction couple
initially at -1.65 V vs Fc
+/0
shifts anodically with the extent depending on the strength of the acid.
H2 is determined to be the dominant reduction product by gas chromatography (FE > 60%) in
addition to small amounts of CO2 reduction products.

2.2 INTRODUCTION
The divestment in fossil fuels in favor of renewable sources such as solar power is limited
by the limited ability to store this energy for use when it is most needed.
1, 2
To combat these spatio-
temporal issues, the storage of energy in the form of chemical bonds is seen as an attractive
method.
1, 3-7
Solar energy can be used to drive chemical reactions such as the hydrogen evolution
reaction (HER) and carbon dioxide reduction to produce value-added chemicals.
5, 8-14
Of these two
processes, CO2 reduction is more complicated due to a higher energetic barrier, multiple possible
reduction products, and competition with hydrogen evolution.
5, 13-16
In order to develop ideal
catalytic systems, extensive optimization is carried out with promising catalysts in order to best
understand their properties.
Many successful CO2 reduction catalysts with high (95%+) faradaic efficiency (FE) for a
given product have been developed throughout the years, but many of these catalysts exhibit
switchable selectivity where a change in the electrochemical setup, ligand design, or metal center
drastically shift the selectivity. Iron porphyrins are well known as selective CO 2-to-CO reduction
catalysts,
17-23
but the reactivity can be changed upon the addition of Lewis acid species (e.g. Mg
2+
)
to 30% formate production.
24, 25
Weak Brønsted acid 1-propanol also facilitates formate
production, yielding 35% formate with 60% CO.
26, 27
Cobalt porphyrins, while still able to perform
35

CO2 reduction, are largely studied as hydrogen evolution catalysts.
28-38
Nickel cyclams are
conventionally thought of as CO2 reduction catalysts,
39-46
but can be dependent on solvent choice
for their selectivity. Aprotic solvents favor formic acid (75%),
41
aqueous conditions from pH 4 to
5 and wet organic solvent show high selectivity for CO,
41-43
and highly acidic aqueous solutions
(pH < 2) begin to produce H2 in appreciable yields.
44

The installation of pendant proton groups is typically seen to enhance catalytic activity, but
changes to the ligand scaffold can have deeper impacts than intended. An iron carbonyl cluster has
been shown to reduce CO2 to formate with faradaic efficiencies near unity, even in aqueous
conditions.
47-49
Attempting to modify this catalyst with a pedant proton group in the form of
PPh2(CH2)2OH caused a complete shift to hydrogen evolution, producing almost no formate.
50

Further attempts to modify this catalyst with indirect proton donors resulted in similar HER
selectivity, with both pKa and group size affecting product formation.
51
Similarly, addition of a
2,6-dihydroxyphenyl substituent to the bipyridine backbone of [Mn(bpy)(CO)3Br] results in the
production of formate with faradaic efficiencies of 36-39% in the presence of strong acids in
addition to drastically reduced amounts of CO.
52

Our group has recently reported a cobalt aminopyridine macrocycle for the selective
reduction of CO2 to CO.
53, 54
This complex contains four pyridines which are bridged by four
amines, where each amine can be secondary or tertiary. In this system, the number of secondary
amines largely influences catalysis, with the kobs increasing from 20 s
-1
to 16,900 s
-1
and the
selectivity increasing from 36% FE to 98% for CO production when going from zero secondary
amines to four.
54
It was initially hypothesized that these secondary amines were behaving as
secondary sphere coordination groups due to the puckering of the cycle, but DFT analysis revealed
that these groups were actually engaging in hydrogen bonding interactions with added proton
sources which increased the local proton concentration around the substrate.
54
As shown by
previous reports of porphyrins containing pendant proton donor groups, slight changes in the
position of proton donor groups can have a large impact on the overall performance of the
catalyst.
55
By making use of the reactive secondary amines, the aminopyridine macrocycle can be
modified to position pendant groups in a more favorable position for direct ligand-substrate
interactions.
A modified version of the cobalt aminopyridine cycle was synthesized to have a pendant
nitro group as a method to study the effect it has on catalyst performance. The complex was
36

characterized, and its electrochemical behavior was investigated to determine its performance
relative to the parent cobalt complexes. Electrochemistry was performed with various proton
sources, in different solvents, and under different gaseous atmospheres to understand changes in
behavior caused by the introduction of the nitro group.

2.3 RESULTS AND DISCUSSION
2.3.1 Synthesis and Characterization
The macrocycle, (NH)1(NMe)3-Bridged Calix[4]pyridine (L
5
), was synthesized according
to literature precedent.
56
The nitrophenyl modified ligand, L
NO2
, was synthesized using previously
reported methods,
57
where L
5
, cesium carbonate, and 1-fluoro-2-nitrobenzene were heated in
DMSO before being purified by column chromatography with 2:1 ethyl acetate/hexanes to yield a
brown solid (Equation 2.1).

Equation 2.1. Synthesis of L
NO2
from L
5
and synthesis of CoL
NO2
from L
NO2
.

Cobalt(II) perchlorate hexahydrate and L
NO2
were dissolved in equal amounts of
acetonitrile and chloroform, respectively. The two solutions were mixed and stirred for 30 minutes
before removing the solvent to yield CoL
NO2
as a brown solid in quantitative yields. Following
metallation with cobalt, the peaks in the
1
H NMR spectrum appear broad (Figure 2.5) as would be
expected for a paramagnetic species. When using pyridine-d5 as the NMR solvent, four peaks
appear in the range of 31 to 22 ppm (Figure 2.6). These peaks are noticeably absent from the NMR
spectrum of cobalt-metallated L
5
(CoL
5
). A UV-Vis spectrum of CoL
NO2
taken in DMF solution
shows two main features as a large peak at 311 nm and a smaller peak at 660 nm, as well as
shoulders at 359 nm and 427 nm (Figure 2.7), which are consistent with previous reports.
57
FTIR
spectra for L
NO2
(Figure 2.8) and CoL
NO2
(Figure 2.9) are similar, and the peaks corresponding to
the nitro group appears at 1524 and 1360 cm
-1
for both (Figure 2.10).
37

Figure 2.1. Parent complex CoL
5
(left), CoL
6
(middle), and nitrophenyl lariat modified complex CoL
NO2
(right).
Counter anions removed for clarity.

X-ray quality crystals were grown by suspending CoL
NO2
, adding acetonitrile dropwise
until dissolution, and vapor diffusing in diethyl ether. Two perchlorate anions are present as
counterions, indicating a cobalt(II) species. The four pyridine groups of the macrocycle all
coordinate the cobalt metal center in a square planar geometry, with the bridging amines puckering
alternatively up and down in the same manner as the previously reported complexes (Figure 2.1
and 2.2).
53, 54, 57
A single acetonitrile solvent molecule is coordinated axially to the cobalt center
opposite the NO2 moiety. This matches CoL
5
and CoL
6
while the macrocycle with all secondary
amines has two solvents coordinated.
53, 54
The NO2 group is positioned over the Co metal center,
and the Co-O(2) distance is 2.83 Å. There is a contraction of the Co-Npy and Co-Nsolv bonds relative
to CoL
5
and the comparable cycle with four tertiary amines, CoL
6
(Table 2.1).
54
The average Co-
NPy lengths are 1.925(2) for CoL
NO2
, 1.943(2) for CoL
5
, and 1.944(7) for CoL
6
, suggesting that
the presence of the nitro group has some influence on the Co site. This contraction of bond lengths
could be detrimental for CO2 reduction as the methyl groups on the amines in CoL
5
and CoL
6
are
known to sterically hinder CO2 binding. By increasing the steric bulk around the metal center
further, CO2 binding may be difficult which may lead to competitive hydrogen evolution becoming
more significant.

38

Figure 2.2. Crystal structure of CoL
NO2
with side (A and B) and top (C) views. Noncoordinating counter ions,
unbound solvent species, and hydrogen atoms omitted for clarity.

Atoms Atomic Distance, Å
(CoL
NO2
)
Atomic Distance, Å
(CoL
5
)
Atomic Distance, Å
(CoL
6
)
Co-O 2.838

Co-N(CH3CN) 2.161 2.281 2.205
Co-NPy1 1.929 1.956 1.950
Co-NPy2 1.928 1.938 1.938
Co-NPy3 1.925 1.939 1.942
Co-NPy4 1.919 1.940 1.949
Table 2.1. Comparison of atomic distances in the metallated lariat and parent complexes as determined from the
crystal structures.

2.3.2 Cyclic Voltammetry in Acetonitrile
The ligand L
NO2
was analyzed electrochemically using cyclic voltammetry (CV) in 0.1 M
tetrabutylammonium hexafluorophosphate (TBAPF6) acetonitrile solution. Under an inert N2
atmosphere, weak reduction features can be seen at -1.65 V and -2.20 V vs Fc
+/0
along with
oxidative features at -1.62 V and -1.18 V (Figure 2.11). When the gaseous atmosphere is changed
to CO2, a large current increase occurs at the second reduction event. Upon titration with 2,2,2-
trifluoroethanol (TFE) as an added proton source, the current increases by a large amount initially
(0.09 M TFE) with subsequent additions only causing minimal increases until saturation at 0.63
M TFE (Figure 2.12).
CVs of the metal complex CoL
NO2
under a nitrogen atmosphere showed similar behavior
to the ligand and the previously reported parent cycle (Figure 2.3 A and 2.13).
53, 54
A plot of
log(scan rate) vs log(current) for the couple at -1.65 V vs Fc
+/0
yields a slope of 0.5, which is
indicative of a freely-diffusing species in solution according to the Randles-Sevcik equation
(Figures 2.14). Titrating with TFE resulted in a large increase in current, reaching a current density
39

of 82 mA/cm
2
at -2.90 V vs Fc
+/0
(Figure 2.3 B and 2.15). This large current increase at negative
potentials is consistent with hydrogen evolving polypyridyl systems, where large currents are
observed due to large overpotentials.
58

Figure 2.3. (A) CVs of 0.5 mM CoL
NO2
under N 2 (blue) and CO 2 (red) atmospheres in 0.1 M TBAPF 6 MeCN
solution. (B) CVs of 0.5 mM CoL
NO2
under an N 2 atmosphere with 0 M TFE (green) and 0.87 M TFE (purple)
added in 0.1 M TBAPF 6 MeCN solution. (C) CVs of 0.5 mM CoL
NO2
under a CO 2 atmosphere with 0 M TFE
(green) and 0.87 M TFE (purple) added in 0.1 M TBAPF 6 MeCN solution. (D) Controlled potential electrolysis of
0.5 mM CoL
NO2
under N 2 (blue) and CO 2 (red) with 0.87 M TFE in 0.1 M TBAPF 6 MeCN solution.

Under a CO2 atmosphere, CoL
NO2
CVs are similar to those under N2 in the absence of an
added proton source. Only a minor current increase is seen, as opposed to the behavior of the
ligand. This is inconsistent with the parent complexes, where a noticeable increase was seen when
changing atmospheres. However, current increases are seen upon titrating with TFE (Figure 2.3 C
and 2.16). The increase seen is notable in that the currents reached under CO2 are roughly 30 fold
less than under an N2 atmosphere. Under both CO2 and N2, the addition of TFE caused an anodic
shift in the reduction feature at -1.65 V which has been linked to the binding of a substrate to the
40

metal center.
59-62
Given that the shift takes place in N2 as well as CO2, this shift is most likely due
to the formation of a metal hydride.
Different acid sources were also tested. Titrating with water results in similar behavior,
with N2 scans reaching 15.2 mA/cm
2
at -2.68 V and CO2 scans reaching 3.76 mA/cm
2
at -2.55 V
upon the addition of 13.8 M H2O (Figures 2.17 and 2.18). Phenol also displays significant current
increases under N2 with a current of 70 mA/cm
2
at -2.70 V with 1.67 M (Figure 2.19). Under CO2,
a current of 31 mA/cm
2
is reached at -2.65 V under the same acid concentration (Figure 2.20 and
2.21). It is noted for the phenol titrations that different behavior is seen depending on if the
concentration of phenol in solution is lower or higher than 0.5 M.

2.3.3 Controlled Potential Electrolysis
Controlled potential electrolysis (CPE) studies were performed in acetonitrile at -2.70 V
vs Fc
+/0
under an N2 atmosphere with 1 M TFE as the proton source (Figure 2.3 D). As predicted
by CV studies, H2 was produced as the sole reduction product (Table 2.2). The production of H2
is unexpected as the parent complex showed no H2, demonstrating that the selectivity of the
catalyst has been altered by the introduction of the nitro group. When the atmosphere was switched
to CO2, the amount of charge passed in the CPE was much lower than under N2 despite a
reasonable amount of current passed in the CV. Three products were detected at the end of the 24
hour CPE: H2 (FE = 58), CO (FE = 2), and formate (FE < 1) (Table 2.3). The presence of formate
further supports the formation of a metal hydride, as CO2 insertion into a metal hydride is the
primary method of formate formation. The parent macrocyclic complexes (CoL
1
-CoL
6
) produced
exclusively CO from CPE under identical conditions. The addition of the pendant nitro group has
caused a shift in the selectivity of the metal complex, with CO no longer being a major product.
Switching the proton source to H2O under a CO2 atmosphere results in higher currents at -
2.70 V vs Fc
+/0
(Figure 2.4), which is consistent with the higher currents seen by CV. H2 was
present in even higher quantities under these conditions, reaching a maximum faradaic efficiency
of 71%, while CO decreased to 1% (Table 2.4). After 24 hours, 875 μmol of H2 were produced,
which is 2.5 times higher than the amount of H2 produced from TFE conditions. Using 1 M phenol
as the proton source resulted in the highest initial current. However, this current is not stable and
begins to rapidly decrease after 2 hours of bulk electrolysis. The faradaic efficiency for CO under
these conditions is similar to the FE under TFE conditions, reaching 4% at best (Table 2.5).
41

However, the H2 FE is the lowest achieved with a maximum FE of 46%. Additionally, trace
amounts of methane are seen by GC.

Figure 2.4. CPE of 0.5 mM CoL
NO2
under CO 2 in 0.1 M TBAPF 6 MeCN solution at -2.70 V vs Fc
+/0
with 1 M TFE
(red), water (blue), and phenol (green) as an added proton source.

As CoL
NO2
has all tertiary amines similar to CoL
6
, a loss of CO2 selectivity is not
unexpected; however, the extent of loss was unexpected as well as the high amount of H2 formed.
The steric crowding around the metal center could be attributing to this selectivity shift as protons
require less space than the relatively bulky CO2. In conjunction with the methyl groups which have
been shown to hinder CO2, the presence of the nitrophenyl group could further block the metal
center.
The need for polishing between CV scans and the unaccounted for faradaic efficiency
raised suspicions of the decomposition of the catalyst, possibly forming cobalt nanoparticles on
the electrode surface that may be active catalysts. A 2 hour controlled potential electrolysis
experiment was performed with 1 M phenol, as the high currents generated with phenol followed
by a large current drop suggested these conditions might be best suited for promoting catalyst
decomposition. Following the electrochemical experiment, part of the working electrode was
washed with acetonitrile and a separate part was washed with both acetonitrile and
dimethylformamide. Analysis of the electrode by X-ray photoelectron spectroscopy (XPS)
revealed the presence of cobalt in both washed regions (Figures 2.22), suggesting that a deposited
material is retained on the electrode. A controlled potential electrolysis experiment performed with
42

a washed electrode and 1 M phenol without any added CoL
NO2
produces only about one third of
the current seen when 1 M phenol and CoL
NO2
are used with a clean electrode (Figure 2.23). The
amount of hydrogen produced is also roughly one third of that produced with CoL
NO2
(Table 6),
indicating that catalyst decomposition may produce an active species on the electrode but the
activity of CoL
NO2
is still superior.

2.3.4 Cyclic Voltammetry Dependence Studies
Kinetic isotope effect measurements were performed with CoL
NO2
in acetonitrile solution
under N2 with H2O (Figure 2.17) and D2O (Figure 2.24) as the titrants, resulting in an H/D kinetic
isotope effect of 3.6 ± 1.4. This result suggests that protons are involved in the rate-determining
step, which matches with the selective H2 evolution seen under these conditions.
63
Similar kinetic
isotope effect studies were performed under CO2

(Figures 2.18 and 2.25), resulting in an H/D KIE
of 1.6 ± 0.5. This value is closer to what would be expected for CO2 binding to the metal center
though is still higher than is typically observed.
63

Solutions of CoL
NO2
under various conditions were titrated with TFE, H2O, and phenol in
stepwise increments to observe the shift in the couple at -1.65 V (Figures 2.27-2.29, Table 2.7). In
all cases, regardless of the identity of the proton source or gaseous atmosphere, the shift was
observed. TFE had the greatest rate of increase on the current shift, with phenol being second and
H2O having the weakest effect. However, phenol produced the largest overall shift in potential.
This would be expected for a hydride-forming species as phenol is the strongest acid under these
conditions.
64-68
Plotting the potential shift against the pKa of these acids in acetonitrile produces a
line with a slope of roughly 29 mV/pH unit (Figure 2.30), which suggests a 2 electron, 1 proton
event.
69
This is consistent with the size of the reduction event, but is inconsistent with the CV
characteristics in CoL
1
-CoL
6
.
The increase in the catalytic current with added protons was also investigated. While the
desired S-shaped curve could not be obtained during these studies, analysis of the relationship
between proton concentration and current was observed as a qualitative analysis. CoL
NO2
under
CO2 and N2 was titrated with TFE, H2O, and phenol, and the currents at -2.70 V vs Fc
+/0
were
plotted against the acid concentration (Figures 2.31-2.36). In all cases, a linear dependence on
proton concentration was observed. In the case of phenol, only concentrations up to 0.48 M were
considered.
43

An electrolyte solution under N2 with 0.5 M TFE was titrated with CoL
NO2
up to 0.64 mM.
The current initially increased from 0 mM to 0.24 mM before stagnating (Figure 2.37), suggesting
that catalyst saturation occurs at low concentrations. Plotting the currents for the lower
concentration data points against the concentration of the catalyst reveals a linear trend (Figure
2.38), suggesting that the catalytic process is first order in catalyst.

2.4 CONCLUSION
Modifications to a previously reported aminopyridine macrocycle produced a nitrophenyl-
modified ligand which was used to generate a cobalt complex. The ligand coordinates the metal in
a square planar fashion similar to the previously reported complexes, where the briding amines
pucker up and down in a saddle-like configuration. In the solid state, the nitro group is positioned
over the metal center with a Co-O distance of 2.83 Å, suggesting that the attachment of lariat arms
in this manner could be used as a means to explore direct proton donors. CV experiments showed
current increases under both N2 and CO2 with the addition of a proton source, with the N2 current
increases being an order of magnitude higher than those under a CO2 atmosphere. Controlled
potential electrolysis showed high faradaic efficiency for HER under both N2 and CO2, with small
amounts of CO2 reduction products also detected. This is in contrast to the previously reported
aminopyridine complexes which had 98% FE for CO2 reduction to CO, showing that the addition
of a single nitrophenyl group has a large impact on catalytic activity.

2.5 ACKNOWLEDGEMENTS
The authors are grateful to the University of Southern California (USC) for funding, the USC
Wrigley Institute for the Norma and Jerol Sonosky summer fellowship to EMJ, and to Thomas
Moulton and Ginny Dunn for the Harold and Lillian Moulton Graduate Fellowship to EMJ. The
studies of the molecular complex were supported by the National Science Foundation (NSF)
through the CAREER award (CHE-1555387). The authors are grateful to NSF (grant CRIF
1048807) and USC for their sponsorship of NMR spectrometers and X-ray diffractometer. The
authors would like to thank Dr. Alon Chapovetsky for discussion.

44

2.6 EXPERIMENTAL METHODS
General. Manipulations of air and moisture sensitive materials were carried out under nitrogen
either in a Vacuum Atmospheres drybox or on a dual-manifold Schlenk line. All solvents were
degassed with nitrogen and passed through activated alumina columns and stored over 4Å Linde-
type molecular sieves. The L
NO2
ligand was prepared according to literature precedent.
56, 57
All
other chemical reagents were purchased from chemical vendors and used without further
purification.

2.6.1 Physical Methods. NMR spectra were obtained using a Varian Mercury 400 MHz NMR
Spectrometer.

Elemental analyses were performed by Complete Analysis Laboratories, Inc., Parsippany, New
Jersey, 07054 or Robertson Microlit Laboratories, 1705 U.S. Highway 46, Suite 1D, Ledgewood,
New Jersey, 07852.

UV-Vis spectra were obtained using a Lambda 950 UV/Vis/NIR Spectrophotometer.

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.

The single-crystal X-ray diffraction data were collected on a Bruker SMART APEX DUO 3-circle
platform diffractometer, equipped with an APEX II CCD, using Mo Kα radiation (TRIUMPH
curved-crystal monochromator) from a fine-focus tube. The diffractometer was equipped with an
Oxford Cryosystems Cryostream 700 apparatus for low-temperature data collection. The frames
were integrated using the SAINT algorithm to give the hkl files corrected for Lp/decay.
70
The
absorption correction was performed using the SADABS program.
71
The structures were solved
by intrinsic phasing and refined on F
2
using the Bruker SHELXTL Software Package and
ShelXle.
72
All non-hydrogen atoms were refined anisotropically.

XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray source
was the Al K α line at 1486. 7 eV, and the hybrid lens and slot mode were used. Low-
45

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 regions
of interest. The sample chamber was maintained at < 9 × 10
–9
Torr. The XPS data were analyzed
using the CasaXPS software.

2.6.2 Electrochemical Methods. Electrochemical experiments were carried out using a Pine
potentiostat. The electrochemical experiments were carried out in three-electrode configuration
electrochemical cell further described below under nitrogen or carbon dioxide atmospheres using
a glassy carbon electrode as the working electrode.

2.6.3 Cyclic Voltammetry of L
NO2
and CoL
NO2
. Electrochemical studies of L
NO2
(0.5 mM) and
CoL
NO2
(0.5 mM) were performed in anhydrous dimethylformamide with tetrabutylammonium
hexafluorophosphate (TBAPF6) (0.1 M) as the supporting electrolyte. Platinum wire was used as
the counter electrode, silver wire as the reference electrode, and glassy carbon as the working
electrode. Decamethylferrocene was added as an internal standard to reference to ferrocene (Fc
+/0
),
and all potentials are given with reference to ferrocene. Samples were sparged with N2 or CO2 gas
for 10 minutes before experiments.

2.6.4 Controlled Potential Electrolysis. Controlled potential electrolysis (CPE) measurements to
determine Faradaic efficiency were conducted in a sealed two-chambered H cell where the first
chamber held the working and reference electrodes in 40 mL of electrolyte and the second chamber
held the auxiliary electrode in 20 mL of electrolyte. CPE experiments of were performed in 0.1 M
TBAPF6 DMF solution with 0.5 mM of analyte. 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 wire separated from the solution by a Vycor tip. The working compartment
was sparged with N2 or CO2 for 10 minutes before the experiment. The two chambers were both
under N2 or CO2 and separated by a fine porosity glass frit. 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 CO amount based on
46

total charge flowed is compared with the GC-detected CO produced from the controlled-potential
electrolysis experiment.

2.7 SYNTHETIC METHODS
The macrocycle (NH)1(NMe)3-bridged calix[4]pyridine was prepared according to literature
precedent.
56

2.7.1 Synthesis of L
NO2
. The ligand was synthesized using the following literature procedure.
57
A
dry flask under nitrogen was charged with (NH)1(NMe)3-bridged calix[4]pyridine (69.7 mg, 0.17
mmol), cesium carbonate (162.3 g, 0.50 mmol), and 2-fluoronitrobenzene (0.13 mL, 1.23 mmol).
DMF (14 mL) was added, and the solution was heated to 60 °C for 24 hours. The solution was the
extracted into ethyl acetate and washed with water. The organic phase was dried with Na2SO4 and
concentrated. The residue was purified via column chromatography with 2:1 hexanes/ethyl acetate
to obtain the product as a brown powder with 95% yield.

2.7.2 Synthesis of Co complex. A 1:1 ratio of L
NO2
and Co(ClO4)2·6H2O were dissolved in
chloroform and acetonitrile, respectively. The two solutions were then mixed and stirred for 30
minutes. The solvent was then removed by rotary evaporation to give CoL
NO2
as a brown powder
in quantitative yields. X-ray quality crystals were grown by suspending the powder in chloroform,
adding acetonitrile dropwise until dissolution, and vapor diffusing with diethyl ether. Anal. Calcd
for C33H31N11O10 Cl2Co: C, 42.83; H, 3.41; N, 15.00. Found: C, 44.84; H, 3.40; N, 16.87.

47

2.8 ADDITIONAL FIGURES

Figure 2.5.
1
H NMR spectrum (400 MHz) of CoL
NO2
in CD 3CN.

48

Figure 2.6.
1
H NMR spectra (400 MHz) overlay of CoL
NO2
(top) and CoL
5
(bottom) in deuterated pyridine (Py-d 5).

Figure 2.7. UV-Vis spectrum of CoL
NO2
in DMF solution.

CoL
NO2
CoL
5
49

Figure 2.8. Transmittance FTIR spectrum of L
NO2
.

Figure 2.9. Transmittance FTIR spectrum of CoL
NO2
.

50

Figure 2.10. Transmittance FTIR spectrum overlay of the NO 2 regions of L
NO2
and CoL
NO2
.

Figure 2.11. 0.5 mM L
NO2
in 0.1 M TBAPF 6 MeCN under N 2 (blue) and CO 2 (red) atmospheres. Scan rate = 100
mV/s.

51

Figure 2.12. 0.5 mM L
NO2
in 0.1 M TBAPF 6 MeCN under a CO 2 atmosphere titrated with TFE. Scan rate = 100
mV/s.

Figure 2.13. CV overlay of L
NO2
(blue), CoL
NO2
(red), and ZnL
NO2
(green) in 0.1 M TBAPF 6 acetonitrile solution
under N 2. Scan rate = 100 mV/s.
52

Figure 2.14. (A) Cyclic voltammogram of CoL
NO2
(0.5 mM) in MeCN with 0.1 M TBAPF 6 under N 2 with varied
scan rates with (B) a corresponding plot of log(current density) vs log(scan rate).

Figure 2.15. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under an N 2 atmosphere with
increasing amounts of TFE. Scan rate = 100 mV/s.

53

Figure 2.16. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with
increasing amounts of TFE. Scan rate = 100 mV/s.

Figure 2.17. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under an N 2 atmosphere with
increasing amounts of H 2O. Scan rate = 100 mV/s.

54

Figure 2.18. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with
increasing amounts of H 2O. Scan rate = 100 mV/s.

Figure 2.19. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under an N 2 atmosphere with
increasing amounts of Phenol. Scan rate = 100 mV/s.

55

Figure 2.20. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with
increasing amounts of Phenol. Scan rate = 100 mV/s.

Figure 2.21. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with
increasing amounts of Phenol. Low concentrations of Phenol only. Scan rate = 100 mV/s.

56

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr
599.8 76 0 0
4 hr
718.6 68 0 0
6 hr
796.7 64 0 0
8 hr
860.5 61 0 0
24 hr
1,127 45 0 0
Table 2.2. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile under N 2 with 1 M
TFE and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

H2 CO Formate
μ m ol FE (%) μ m ol FE (%) μ m ol FE (%)
2 hr
6.7 17 1.8 4
- -
4 hr
30.4 37 2.8 3
- -
6 hr
65.7 47 4.5 3
- -
8 hr
108.0 54 5.5 3
- -
24 hr
350.3 58 13.0 2
0.4 <1
Table 2.3. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile under CO 2 with 1 M
TFE and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

H2 CO Formate
μ m ol FE (%) μ m ol FE (%) μ m ol FE (%)
2 hr
155 67 4 1.5 - -
4 hr
265 69 5 1.26 - -
6 hr
348 71 6 1.18 - -
8 hr
406 69 6 1 - -
24 hr
875 58 9 <1 0 0
Table 2.4. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile under CO 2 with 1 M
H 2O and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

57

H2 CO Formate
μ m ol FE (%) μ m ol FE (%) μ m ol FE (%)
2 hr
443 46 37 4
- -
4 hr
679 38 55 3 - -
6 hr
681 36 67 3 - -
8 hr
708 34 73 4
- -
24 hr
484 18 85 3
0 0
Table 2.5. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile under CO 2 with 1 M
Phenol and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

Figure 2.22. XPS spectra of the Co 2p region for the acetonitrile (A) and acetonitrile and DMF (B) washed sections
of the glassy carbon working electrode following a two hour controlled potential electrolysis of CoL
NO2
with 1 M
phenol.

58

Figure 2.23. Controlled potential electrolysis of 1 M Phenol acetonitrile solution containing 0.5 mM CoL
NO2
and a
clean electrode (red) and a washed electrode in the absence of CoL
NO2
(black). 0.1 M TBAPF 6 acetonitrile solution
at -2.70 V vs Fc
+/0
.

μ m ol H2 FE (%) H2
Washed Electrode
139.5 55.1
Catalyst Solution
400.2 31.3
Table 2.6. Results of 2 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in acetonitrile under CO 2 with 1 M
Phenol and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
(catalyst solution) and the solvent-washed electrode with no added
CoL
NO2
and otherwise identical conditions (washed electrode).

59

Figure 2.24. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under an N 2 atmosphere with
increasing amounts of D 2O. Scan rate = 100 mV/s.

Figure 2.25. Titration of 0.5 mM CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with
increasing amounts of D 2O. Scan rate = 100 mV/s.

60

Figure 2.26. Titration of CoL
NO2
in 0.1 M TBAPF 6 MeCN solution under a CO 2 atmosphere with increasing
amounts of TFE focused on the reduction event at -1.65 V vs Fc
+/0
shifting anodically.

Figure 2.27. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the concentration of the
added proton source in solution under different conditions. Red circles = under CO 2 with titrated TFE. Blue circles =
under N 2 with titrated TFE.

61

Figure 2.28. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the concentration of the
added proton source in solution under different conditions. Red circles = under CO 2 with titrated H 2O. Blue circles =
under N 2 with titrated H 2O.

Figure 2.29. Plot of the shifted redox event potential (-1.65 V @ 0 M TFE) for CoL
NO2
vs the concentration of the
added proton source in solution under different conditions. Red circles = under CO 2 with titrated phenol. Blue
circles = under N 2 with titrated phenol.

62

Figure 2.30. Plot of potential vs pK a for CoL
NO2
under N 2 based on the shift of the couple at -1.65 V vs Fc
+/0
when
titrated with water, TFE, and phenol.

Gas Acid Slope (V/M)
CO2 TFE 0.1474
N2 TFE 0.1794
CO2 Phenol 0.0871
N2 Phenol 0.0880
CO2 H2O 0.0497
N2 H2O 0.0384
Table 2.7 The slope of the lines generated from plotting the shift in the -1.65 V vs Fc
+/0
couple vs added acid
concentration and their corresponding gas and acid environments.

63

Figure 2.31. Plot of current density vs TFE concentration for CoL
NO2
in MeCN solution under N 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

Figure 2.32. Plot of current density vs H 2O concentration for CoL
NO2
in MeCN solution under N 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

64

Figure 2.33. Plot of current density vs Phenol concentration for CoL
NO2
in MeCN solution under N 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

Figure 2.34. Plot of current density vs TFE concentration for CoL
NO2
in MeCN solution under CO 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

65

Figure 2.35. Plot of current density vs H 2O concentration for CoL
NO2
in MeCN solution under CO 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

Figure 2.36. Plot of current density vs Phenol concentration for CoL
NO2
in MeCN solution under CO 2 titrated with
corresponding acid source. Current was measured at -2.70 V vs Fc
+/0
for all concentrations.

66

Figure 2.37. Titrations of CoL
NO2
into 0.1 M TBAPF 6 acetonitrile solution under N 2 with 0.5 M TFE.

Figure 2.38. Current density vs catalyst concentration for CoL
NO2
.

67

Table 2.8. Sample and crystal data for CoL
NO2
.
Chemical formula C35H34Cl2CoN12O10
Formula weight 912.57 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.077 x 0.132 x 0.270 mm
Crystal habit yellow prism
Crystal system triclinic
Space group P -1
Unit cell dimensions a = 12.0282(14) Å α = 73.002(2)°

b = 12.5514(15) Å β = 70.196(2)°

c = 15.9811(19) Å γ = 65.433(2)°
Volume 2031.9(4) Å
3

Z 2
Density (calculated) 1.492 g/cm
3

Absorption coefficient 0.625 mm
-1

F(000) 938
Diffractometer Bruker APEX DUO
Radiation source fine-focus tube (MoKα , λ = 0.71073 Å)
Theta range for data
collection
1.81 to 26.37°
Index ranges -15<=h<=15, -15<=k<=15, -19<=l<=19
Reflections collected 35503
Independent reflections 8246 [R(int) = 0.0620]
Absorption correction multi-scan
Max. and min.
transmission
0.9530 and 0.8490
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 2018/3 (Bruker AXS, 2018)
Function minimized Σ w(Fo
2
- Fc
2
)
2

Data / restraints /
parameters
8246 / 542 / 634
Goodness-of-fit on F
2
1.026
Δ /σmax 0.001
68

Final R indices
6053 data;
I>2σ(I)
R1 = 0.0599, wR2 = 0.1097

all data R1 = 0.0881, wR2 = 0.1205
Weighting scheme
w=1/[σ
2
(Fo
2
)+(0.0388P)
2
+3.4275P]
where P=(Fo
2
+2Fc
2
)/3
Largest diff. peak and
hole
0.625 and -0.638 eÅ
-3

R.M.S. deviation from
mean
0.070 eÅ
-3

Table 2.9. Bond lengths (Å) for CoL
NO2
.
C1-N1 1.351(4) C1-C2 1.387(4)
C1-N8 1.393(4) C2-C3 1.372(5)
C2-H2 0.95 C3-C4 1.382(5)
C3-H3 0.95 C4-C5 1.381(4)
C4-H4 0.95 C5-N1 1.349(4)
C5-N2 1.409(4) C6-N3 1.351(4)
C6-C7 1.375(5) C6-N2 1.415(4)
C7-C8 1.385(5) C7-H7 0.95
C8-C9 1.374(5) C8-H8 0.95
C9-C10 1.379(5) C9-H9 0.95
C10-N3 1.351(4) C10-N4 1.405(4)
C11-N5 1.354(4) C11-N4 1.392(4)
C11-C12 1.396(5) C12-C13 1.365(5)
C12-H12 0.95 C13-C14 1.373(5)
C13-H13 0.95 C14-C15 1.385(4)
C14-H14 0.95 C15-N5 1.350(4)
C15-N6 1.397(4) C16-N7 1.357(4)
C16-C17 1.383(5) C16-N6 1.401(4)
C17-C18 1.379(4) C17-H17 0.95
C18-C19 1.371(5) C18-H18 0.95
C19-C20 1.385(4) C19-H19 0.95
C20-N7 1.353(4) C20-N8 1.393(4)
C21-C22 1.391(5) C21-C26 1.391(5)
C21-N2 1.428(4) C22-C23 1.381(5)
C22-H22 0.95 C23-C24 1.379(6)
C23-H23 0.95 C24-C25 1.381(5)
C24-H24 0.95 C25-C26 1.385(5)
69

C25-H25 0.95 C26-N10 1.474(4)
C27-N4 1.469(4) C27-H27A 0.98
C27-H27B 0.98 C27-H27C 0.98
C28-N6 1.465(4) C28-H28A 0.98
C28-H28B 0.98 C28-H28C 0.98
C29-N8 1.465(4) C29-H29A 0.98
C29-H29B 0.98 C29-H29C 0.98
C30-C31 1.452(5) C30-H30A 0.98
C30-H30B 0.98 C30-H30C 0.98
C31-N9 1.139(4) Co1-N7 1.919(3)
Co1-N5 1.925(2) Co1-N1 1.928(2)
Co1-N3 1.929(3) Co1-N9 2.161(3)
N10-O1' 1.142(8) N10-O2 1.172(4)
N10-O1 1.268(5) N10-O2' 1.344(7)
Cl1-O4 1.411(10) Cl1-O6 1.416(9)
Cl1-O5 1.424(9) Cl1-O3 1.433(9)
Cl1'-O6' 1.398(10) Cl1'-O5' 1.403(11)
Cl1'-O3' 1.413(10) Cl1'-O4' 1.429(11)
Cl2-O10 1.409(3) Cl2-O8 1.428(3)
Cl2-O7 1.431(2) Cl2-O9 1.453(3)
C32-C33 1.476(9) C32-H32A 0.98
C32-H32B 0.98 C32-H32C 0.98
C33-N11 1.126(7) C32'-C33' 1.478(15)
C32'-H32D 0.98 C32'-H32E 0.98
C32'-H32F 0.98 C33'-N11' 1.131(14)
C34-C35 1.542(10) C34-H34A 0.98
C34-H34B 0.98 C34-H34C 0.98
C35-N12 1.134(7) C34'-C35' 1.57(2)
C34'-H34D 0.98 C34'-H34E 0.98
C34'-H34F 0.98 C35'-N12' 1.109(18)

70

Table 2.10. Bond angles (°) for CoL
NO2
.
N1-C1-C2 121.9(3) N1-C1-N8 116.8(3)
C2-C1-N8 121.3(3) C3-C2-C1 118.6(3)
C3-C2-H2 120.7 C1-C2-H2 120.7
C2-C3-C4 120.2(3) C2-C3-H3 119.9
C4-C3-H3 119.9 C5-C4-C3 118.5(3)
C5-C4-H4 120.8 C3-C4-H4 120.8
N1-C5-C4 122.1(3) N1-C5-N2 116.5(3)
C4-C5-N2 121.4(3) N3-C6-C7 122.8(3)
N3-C6-N2 116.1(3) C7-C6-N2 121.0(3)
C6-C7-C8 117.9(3) C6-C7-H7 121.1
C8-C7-H7 121.1 C9-C8-C7 120.0(3)
C9-C8-H8 120.0 C7-C8-H8 120.0
C8-C9-C10 119.1(3) C8-C9-H9 120.4
C10-C9-H9 120.4 N3-C10-C9 121.6(3)
N3-C10-N4 117.2(3) C9-C10-N4 121.2(3)
N5-C11-N4 117.5(3) N5-C11-C12 121.2(3)
N4-C11-C12 121.2(3) C13-C12-C11 118.8(3)
C13-C12-H12 120.6 C11-C12-H12 120.6
C12-C13-C14 120.4(3) C12-C13-H13 119.8
C14-C13-H13 119.8 C13-C14-C15 118.6(4)
C13-C14-H14 120.7 C15-C14-H14 120.7
N5-C15-C14 122.0(3) N5-C15-N6 117.3(3)
C14-C15-N6 120.7(3) N7-C16-C17 122.2(3)
N7-C16-N6 116.6(3) C17-C16-N6 121.2(3)
C18-C17-C16 118.4(3) C18-C17-H17 120.8
C16-C17-H17 120.8 C19-C18-C17 120.4(3)
C19-C18-H18 119.8 C17-C18-H18 119.8
C18-C19-C20 118.4(3) C18-C19-H19 120.8
C20-C19-H19 120.8 N7-C20-C19 122.5(3)
N7-C20-N8 116.7(3) C19-C20-N8 120.8(3)
C22-C21-C26 117.9(3) C22-C21-N2 121.6(3)
C26-C21-N2 120.5(3) C23-C22-C21 120.5(4)
C23-C22-H22 119.7 C21-C22-H22 119.7
C24-C23-C22 120.7(4) C24-C23-H23 119.7
C22-C23-H23 119.7 C23-C24-C25 119.9(3)
C23-C24-H24 120.1 C25-C24-H24 120.1
C24-C25-C26 119.2(4) C24-C25-H25 120.4
71

C26-C25-H25 120.4 C25-C26-C21 121.8(3)
C25-C26-N10 116.5(3) C21-C26-N10 121.7(3)
N4-C27-H27A 109.5 N4-C27-H27B 109.5
H27A-C27-H27B 109.5 N4-C27-H27C 109.5
H27A-C27-H27C 109.5 H27B-C27-H27C 109.5
N6-C28-H28A 109.5 N6-C28-H28B 109.5
H28A-C28-H28B 109.5 N6-C28-H28C 109.5
H28A-C28-H28C 109.5 H28B-C28-H28C 109.5
N8-C29-H29A 109.5 N8-C29-H29B 109.5
H29A-C29-H29B 109.5 N8-C29-H29C 109.5
H29A-C29-H29C 109.5 H29B-C29-H29C 109.5
C31-C30-H30A 109.5 C31-C30-H30B 109.5
H30A-C30-H30B 109.5 C31-C30-H30C 109.5
H30A-C30-H30C 109.5 H30B-C30-H30C 109.5
N9-C31-C30 178.9(4) N7-Co1-N5 90.56(11)
N7-Co1-N1 88.26(10) N5-Co1-N1 174.98(11)
N7-Co1-N3 173.49(11) N5-Co1-N3 88.95(11)
N1-Co1-N3 91.66(11) N7-Co1-N9 96.16(10)
N5-Co1-N9 95.36(10) N1-Co1-N9 89.62(10)
N3-Co1-N9 90.34(10) C5-N1-C1 118.6(3)
C5-N1-Co1 120.89(19) C1-N1-Co1 120.4(2)
C5-N2-C6 120.3(3) C5-N2-C21 119.5(2)
C6-N2-C21 119.5(2) C10-N3-C6 118.4(3)
C10-N3-Co1 120.6(2) C6-N3-Co1 121.0(2)
C11-N4-C10 119.3(3) C11-N4-C27 117.6(3)
C10-N4-C27 116.8(3) C15-N5-C11 118.7(3)
C15-N5-Co1 120.8(2) C11-N5-Co1 120.4(2)
C15-N6-C16 119.4(3) C15-N6-C28 117.6(3)
C16-N6-C28 117.0(3) C20-N7-C16 117.9(3)
C20-N7-Co1 120.8(2) C16-N7-Co1 121.2(2)
C1-N8-C20 118.8(2) C1-N8-C29 118.5(3)
C20-N8-C29 117.1(3) C31-N9-Co1 162.2(3)
O2-N10-O1 124.3(4) O1'-N10-O2' 120.1(5)
O1'-N10-C26 124.7(5) O2-N10-C26 120.7(3)
O1-N10-C26 114.3(3) O2'-N10-C26 114.2(4)
O4-Cl1-O6 108.9(10) O4-Cl1-O5 112.1(10)
O6-Cl1-O5 107.8(10) O4-Cl1-O3 112.9(11)
O6-Cl1-O3 105.2(10) O5-Cl1-O3 109.6(11)
72

O6'-Cl1'-O5' 110.2(11) O6'-Cl1'-O3' 113.4(13)
O5'-Cl1'-O3' 104.3(12) O6'-Cl1'-O4' 110.5(11)
O5'-Cl1'-O4' 109.7(12) O3'-Cl1'-O4' 108.7(12)
O10-Cl2-O8 111.4(2) O10-Cl2-O7 108.9(2)
O8-Cl2-O7 109.16(16) O10-Cl2-O9 110.8(2)
O8-Cl2-O9 109.21(18) O7-Cl2-O9 107.28(18)
C33-C32-H32A 109.5 C33-C32-H32B 109.5
H32A-C32-H32B 109.5 C33-C32-H32C 109.5
H32A-C32-H32C 109.5 H32B-C32-H32C 109.5
N11-C33-C32 178.3(9) C33'-C32'-H32D 109.5
C33'-C32'-H32E 109.5 H32D-C32'-H32E 109.5
C33'-C32'-H32F 109.5 H32D-C32'-H32F 109.5
H32E-C32'-H32F 109.5 N11'-C33'-C32' 172.(2)
C35-C34-H34A 109.5 C35-C34-H34B 109.5
H34A-C34-H34B 109.5 C35-C34-H34C 109.5
H34A-C34-H34C 109.5 H34B-C34-H34C 109.5
N12-C35-C34 165.3(6) C35'-C34'-H34D 109.5
C35'-C34'-H34E 109.5 H34D-C34'-H34E 109.5
C35'-C34'-H34F 109.5 H34D-C34'-H34F 109.5
H34E-C34'-H34F 109.5 N12'-C35'-C34' 179.(6)

73

2.9 REFERENCES

1. Lewis, N. S.; Nocera, D. G., Powering the Planet: Chemical Challenges in Solar Energy
Utilization. Proceedings of the National Academy of Sciences 2006, 103 (43), 15729.
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78

CHAPTER 3. ELECTROCHEMISTRY OF A COBALT NITROPHENYL
SUBSTITUTED AMINOPYRIDINE COMPLEX IN DIMETHYLFORMAMIDE
SOLUTION

3.1 ABSTRACT
The cobalt aminopyridine system modified with a pendant nitro group was tested
electrochemically in DMF solution. The amount of measured CO in DMF under a CO2 atmosphere
exceeded the amounts seen in acetonitrile, reaching up to 14% faradaic efficiency (FE) with 2,2,2-
trifluoroethanol (TFE) while only 4% FE was seen in acetonitrile. However, CO was also seen
under an atmosphere of N2. After 24 hours of controlled potential electrolysis under both N2 and
CO2 atmospheres with added TFE, nearly identical amounts of CO were detected by gas
chromatography. Performing electrolysis at more positive potentials and using different proton
sources still produced CO. Since no CO was produced under N2 in acetonitrile, the decomposition
of the complex as the source of CO is unlikely. Therefore, the CO is hypothesized to be due to the
decarbonylation of DMF solution.

3.2 INTRODUCTION
The process of converting abundant small molecules into value added products has been
long attempted by chemists as a means to generate fuels and industrially relevant chemicals from
waste materials.
1-6
However, abundant molecules are typically highly stable, requiring a large
energy input to reach the desired product. In nature, enzymes are capable of performing these
complex and energetically costly processes at ambient pressures and temperatures.
2, 7-18
The study
of enzyme active sites has led to the identification of structural motifs present in multiple enzymes,
such as mutlimetallic cooperativity and hydrogen bonding interactions, which are believed to help
promote catalysis under desirable conditions. By incorporating these design elements into
synthetic catalysts, enhanced catalytic activity may be achieved.
Modification of successful catalysts has typically led to the improvement of catalytic
ability. Porphyrin-based catalysts have shown to be highly successful in the conversion of a variety
of small molecules.
19-38
Cobalt porphyrins have shown great promise as hydrogen evolution
catalysts,
30-32
with modifications of the ligand structure either through electronic
30
or steric
controls
31
allowing for enhancement of product formation. Similar cobalt porphyrins have also
79

been used for the other water splitting half reaction, performing both water oxidation
33
and oxygen
reduction.
35, 36
Iron porphyrins, known for being more selective for CO2 reduction than H2
evolution, have also made use of secondary sphere interactions, which both increase the local
proton environment and engage in direct hydrogen bonding interactions to stabilize the bound CO2
substrate.
24, 26, 37
Iron porphyrins also have seen success using charged substitutents to influence
the metal center.
19, 20, 28
Various other ligand architectures such as tripods
39-42
and cyclams
1, 3, 5, 43-
49
have also shown enhanced activity because of secondary sphere interactions.
While ligand modifications are often thought of as a simple way to enhance catalysis, it is
important to note that many catalytic reactions are in competition with one or more other processes,
and the introduction of new moieties may not always enhance the selectivity for the desired
process. An iron cluster was shown to have selectivity for CO2 reduction to formate near unity,
even in aqueous conditions.
50-52
When this catalyst was modified to contain pendant groups, the
selectivity for formate production plummeted and hydrogen evolution became the dominant
process.
53, 54
In the case where a pendant phenolic group was added, the faradaic efficiency for H2
was nearly 100%.
53
When weaker proton shuttles were used, a higher selectivity for H 2 was still
observed, but a relationship to the bulkiness of the added group was also considered important for
consideration, with more sterically bulk groups having better formate selectivity.
54
This shift in
selectivity is seen in many other catalysts as well, demonstrating that modifications meant for a
particular purpose may not always have the desired effect.
Our group has previously reported a family of cobalt aminopyridine macrocycle complexes
which selectively reduce carbon dioxide to carbon monoxide.
55, 56
These catalysts contain 4 amines
which can be either secondary or tertiary, with the substitution of the amines being critical for
catalysis. When all amines were secondary, a faradaic efficiency near unity was seen for CO 2 to
CO reduction with rates roughly four orders of magnitude higher than for the tertiary amines.
55, 56

The number of secondary and tertiary amines could also be controlled, with catalytic activity
scaling with the number of secondary amines.
56
DFT calculations revealed that this effect is due
to interspecies hydrogen bonding interactions, with the secondary amines stabilizing acid species
in solution in a position near the bound CO2 adduct.
56
In order to get more direct CO2-hydrogen
bonding interactions, ligand functionalization is required.
The installation of a lariat arm containing a nitro group has shown a significant shift in the
selectivity of the macrocycle in acetonitrile solution, producing ~75% H2 and only minimal
80

amounts of CO2 reduction products. In an attempt to explore different electrochemical conditions,
DMF was used as a solvent due to this being the optimized system for the parent complexes. The
lariat complex was studied at various potentials and with different acid groups.

3.3 RESULTS AND DISCUSSION
3.3.1 Cyclic Voltammetry
Using the species characterized in the previous work (L
NO2
and CoL
NO2
), cyclic
voltammetry (CV) was conducted in 0.1 M tetrabutylammonium hexafluorophosphate (TBAPF6)
dimethylformamide (DMF) solution. All compounds were 0.5 mM in solution and referenced to
the Fc
+/0
couple. L
NO2
showed little electrochemical response under a nitrogen atmosphere, but did
display a reversible couple at E1/2 = -1.66 V (Figure 3.1 A). When placed under a CO2 atmosphere,
the reversible couple becomes irreversible and a large current increase takes place at -2.36 V,
reaching a maximum current of c.a. 0.75 mA/cm
2
at -2.79 V. This type of current increase would
typically be representative of reactivity towards CO2. The solution was titrated with 2,2,2-
trifluoroethanol (TFE) as a proton source (Figure 3.1 B), but no further increase in current was
observed. These results are consistent with CVs conducted in acetonitrile. While acetonitrile can
be used as a proton source at highly negative potentials, DMF does not behave the same. The fact
that a large current increase occurs when switching from N2 to CO2 in both solutions and does not
change much with added TFE suggests that the interaction is largely proton independent.

Figure 3.1. Cyclic voltammograms of L
NO2
(0.5 mM) in DMF with 0.1 M TBAPF 6. (A) Overlay of scans under N 2
and CO 2. (B) Titrations with TFE under a CO 2 atmosphere. Scan rate = 100 mV/s.

CVs of CoL
NO2
under N2 are similar to those of the free ligand, though a current increase
does occur at more negative potentials (Figure 3.2 A). A plot of log(scan rate) vs log(current) for
(B)
(A)
81

the couple at -1.65 V vs Fc
+/0
yields a slope of 0.5, which is indicative of a freely-diffusing species
in solution according to the Randles-Sevcik equation (Figures 3.3 and 3.4). When the solution was
titrated with 2,2,2-trifluoroethanol (TFE), a large current enhancement occurred which reached a
maximum current denisty of 33 mA/cm
2
at -2.85 V vs Fc
+/0
(Figure 3.2 B). This large current
increase at negative potentials is consistent with hydrogen evolving polypyridyl systems.
57
This
same current enhancement is seen in acetonitrile solution, but the acetonitrile current increase is
much larger (c.a. 80 mA/cm
2
). When the gaseous atmosphere was switched to CO2, the peak at -
1.65 V becomes less reversible (Figure 3.2 A). Unlike the free ligand, the metal complex does not
show any enhanced current when subjected to a CO2 atmosphere. Again, this behavior is similar
to CVs conducted in acetonitrile. The previously reported aminopyridine macrocycles did not
display these behaviors: the ligands showed no unusual behavior and the Co complexes saw
current increases under CO2. While the current increases for the previously reported complexes
(CoL
1
-CoL
6
) may be in part due to the dissociation of their protons following dissolution
promoting catalysis,
56
protons which CoL
NO2
does not have, the full reason is likely more complex.
L
NO2
showing a current increase (in both acetonitrile and DMF) under CO2 while L
1
-L
6
and
CoL
NO2
show no response seems to suggest that the NO2 group may be responsible for this
behavior, but that this behavior is hindered when the ligand is metallated. Upon the addition of
TFE, a large current enhancement occurs which reaches 1 mA/cm
2
at -2.67 V vs Fc
+/0
(Figure 3.2
C). Like the increase seen under N2, a higher but otherwise similar current increase is seen in
acetonitrile. Upon the addition of TFE under both N2 and CO2 atmospheres, the reduction feature
at -1.65 V shifts more positively which suggests binding of a substrate to the metal center.
58-61

3.3.2 Controlled Potential Electrolysis
Controlled potential electrolysis (CPE) of CoL
NO2
at -2.70 V vs Fc
+/0
under CO2 in the
absence of TFE showed low current and no observable reduction products (i.e. H2, CO, formate)
(Figure 3.2 D) after 2 hours. With TFE added, CoL
NO2
produced high currents under both N2 and
CO2 atmospheres. In both cases, H2 was the dominant product with a 73% faradaic efficiency (FE)
under N2 and a 65% faradaic efficiency under CO2 (Table 3.1). These results differ from the
previously reported Co macrocycle complexes, which selectively reduced CO 2 to CO with 98%
FE.
56
Under CO2, small amounts of CO were detected (FE = 2%) along with trace amounts of
formate. CPE under N2 also produced trace amounts of CO, similar to the amount produced under
82

CO2, which could indicate that the CO produced is not actually a product of CO2 reduction. Wash
tests showed that no electrochemically active species was being deposited on the electrode surface
(Figure 3.5). Blank CPE experiments conducted with the free ligand and Co precursor (cobalt(II)
perchlorate hexahydrate) produced low currents, with only a small amount of H2 produced by the
Co precursor. The amount of H2 formed by the Co salt was less than that formed with CoL
NO2

(Table 3.2).

Figure 3.2. Electrochemical analysis of CoL
NO2
(0.5 mM) in DMF with 0.1 M TBAPF 6. (A) Overlay of cyclic
voltammetry scans under N 2 (blue) and CO 2 (red), (B) under an N 2 atmosphere with 0 M (green) and 1.02 M
(purple) TFE, and (C) under a CO 2 atmosphere with 0 M (green) and 0.93 M (purple) TFE. (D) Bulk electrolysis of
CoL
NO2
in DMF with 1 M TFE under N 2 (blue), 1 M TFE under CO 2 (red), and 0 M TFE under CO 2 (green).
[CoL
NO2
] = 0.5 mM. Scan rates = 100 mV/s.

H2 CO
μ m ol FE (%) μ m ol FE (%)
N2, 1 M TFE 311 73 ± 7 4 <1 ± 1
CO2, 1 M TFE 167 65 ± 6 5 2 ±2
CO2, 0 M TFE 0 0 0 0
Table 3.1. Results of 2 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF with 0.1 M TBAPF 6 and
with varied atmospheres and TFE concentrations.

(D)
(C)
(B)
(A)
83

3.3.3 Determination of CO Source
To better understand the product distribution of CoL
NO2
under both N2 and CO2, CPE
experiments were performed for 24 hours in order to generate larger amounts of gaseous products
for analysis (Figure 3.7). Under both N2 and CO2 atmospheres, the FE for H2 declined over time
while the FE for CO2 increased (Table 3.3 and 3.4). In these instances, the CO production was
similar between both gaseous atmospheres, further raising the question of whether or not the CO
produced under a CO2 atmosphere is due to CO2 reduction or if it is a product of a decomposition
pathway.
Six hour CPE experiments at more positive potentials (-2.25 V and -2.45 V) under CO2
with c.a. 1 M TFE produced less CO and H2 than those performed at -2.70 V, but the trend was
still apparent (Figure 3.8, Tables 3.5-3.7). Different proton sources were also tested to determine
if the pKa of the acid as well as its identity has an influence on the amount of CO produced. Acids
that produce either no CO under CO2 or produce a higher FE for CO may help to optimize catalyst
conditions. Water, methanol, and phenol were tested as proton sources under CO2 (Figure 3.9) and
showed similar behavior of decreasing H2 FE and increasing CO FE with time (Table 3.8-3.10).
In these studies, it was noted that the increase in CO FE seemed to be loosely related to pKa of the
proton source, with water having the smallest change and phenol having the largest. However, the
presence of such similar amounts of CO suggests that the identity of the acid does not influence
the side process which is generating the CO.
Given that no CO is seen in acetonitrile under an N2 atmosphere, ruling out decomposition
of CoL
NO2
as the source, the CO must be due to the decarbonylation of DMF solution. As the
solvent DMF contains a carbonyl and is the only possible source of CO in the N2 CPE experiments,
it is hypothesized that CoL
NO2
is predominantly a H2 evolution catalyst that, over time, produces
CO from DMF. Examples of group 7 metal complexes that perform decarbonylation reactions are
known in the literature.
62-66
Further investigation must be carried out to better grasp the behavior
of CoL
NO2
in DMF solution.

84

3.4 CONCLUSION
The modification of successful catalysts can often lead to unexpected changes in reaction
selectivity, sometimes leading to results wildly different from what was seen previously. In this
work, a lariat-modified variant of the cobalt aminopyridine macrocycle complexes behaves loses
the selectivity the parent complexes had for CO2 reduction when in DMF solution with TFE as an
added proton source. The results are similar to those seen under acetonitrile solution, with
hydrogen gas becoming the major product at highly negative overpotentials. However, under N 2,
the lariat also produces CO when in DMF solution. CO is not formed in acetonitrile solution under
N2, suggesting that the CO is reliant on the presence of DMF.

3.5 ACKNOWLEDGEMENTS
The authors are grateful to the University of Southern California (USC) for funding, the USC
Wrigley Institute for the Norma and Jerol Sonosky summer fellowship to EMJ, and to Mr. Thomas
Moulton and Ms. Ginny Dunn for the Harold and Lillian Moulton Graduate Fellowship for EMJ.
The studies of the molecular complex were supported by the National Science Foundation (NSF)
through the CAREER award (CHE-1555387). The authors are grateful to NSF (grant CRIF
1048807) and USC for their sponsorship of NMR spectrometers and X-ray diffractometer. The
authors thank Dr. Alon Chapovetsky for helpful discussion.

85

3.6 EXPERIMENTAL METHODS
General. Manipulations of air and moisture sensitive materials were carried out under nitrogen
either in a Vacuum Atmospheres drybox or on a dual-manifold Schlenk line. All solvents were
degassed with nitrogen and passed through activated alumina columns and stored over 4Å Linde-
type molecular sieves. The L
NO2
ligand was prepared according to literature precedent and CoL
NO2

was prepared as stated in Chapter 2. All other chemical reagents were purchased from chemical
vendors and used without further purification.

3.6.1 Physical Methods. NMR spectra were obtained using a Varian Mercury 400 MHz NMR
Spectrometer.

Elemental analyses were performed by Complete Analysis Laboratories, Inc., Parsippany, New
Jersey, 07054 or Robertson Microlit Laboratories, 1705 U.S. Highway 46, Suite 1D, Ledgewood,
New Jersey, 07852.

The single-crystal X-ray diffraction data were collected on a Bruker SMART APEX DUO 3-circle
platform diffractometer, equipped with an APEX II CCD, using Mo Kα radiation (TRIUMPH
curved-crystal monochromator) from a fine-focus tube. The diffractometer was equipped with an
Oxford Cryosystems Cryostream 700 apparatus for low-temperature data collection. The frames
were integrated using the SAINT algorithm to give the hkl files corrected for Lp/decay.
67
The
absorption correction was performed using the SADABS program.
68
The structures were solved
by intrinsic phasing and refined on F
2
using the Bruker SHELXTL Software Package and
ShelXle.
69
All non-hydrogen atoms were refined anisotropically.

3.6.2 Electrochemical Methods. Electrochemical experiments were carried out using a Pine
potentiostat. The electrochemical experiments were carried out in three-electrode configuration
electrochemical cell further described below under nitrogen or carbon dioxide atmospheres using
a glassy carbon electrode as the working electrode.

3.6.3 Cyclic Voltammetry of L
NO2
and CoL
NO2
. Electrochemical studies of L
NO2
(0.5 mM) and
CoL
NO2
(0.5 mM) were performed in anhydrous dimethylformamide with tetrabutylammonium
86

hexafluorophosphate (TBAPF6) (0.1 M) as the supporting electrolyte. Platinum wire was used as
the counter electrode, silver wire as the reference electrode, and glassy carbon as the working
electrode. Decamethylferrocene was added as an internal standard to reference to ferrocene (Fc
+/0
),
and all potentials are given with reference to ferrocene. Samples were sparged with N2 or CO2 gas
for 10 minutes before experiments.

3.6.4 Controlled Potential Electrolysis. Controlled potential electrolysis (CPE) measurements to
determine Faradaic efficiency were conducted in a sealed two-chambered H cell where the first
chamber held the working and reference electrodes in 40 mL of electrolyte and the second chamber
held the auxiliary electrode in 20 mL of electrolyte. CPE experiments of were performed in 0.1 M
TBAPF6 DMF solution with 0.5 mM of analyte. 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 wire separated from the solution by a Vycor tip. The working compartment
was sparged with N2 or CO2 for 10 minutes before the experiment. The two chambers were both
under N2 or CO2 and separated by a fine porosity glass frit. 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 CO amount based on
total charge flowed is compared with the GC-detected CO produced from the controlled-potential
electrolysis experiment.

87

3.7 ADDITIONAL FIGURES

Figure 3.3. Cyclic voltammogram of CoL
NO2
(0.5 mM) in DMF with 0.1 M TBAPF 6 under N 2 with varying scan
rates.

Figure 3.4. Plot of log(current) vs log(scan rate) for CoL
NO2
.

88

Figure 3.5. Cyclic voltammograms of CoL
NO2
(0.5 mM) under CO 2 in saturated DMF with 0.1 M TBAPF 6and 1 M
TFE after electrolysis (red) and of rinsed electrode (3 × 10 mL DMF) in a fresh DMF solution containing 0.1 M
TBAPF 6and 1 M TFE under CO 2 (blue). Scan rate = 100 mV/s. Measured in an H cell. Working electrode surface
area approximated as 8 cm
2
.

Figure 3.6. Controlled potential electrolysis of L
NO2
(red), Co(ClO 4)∙6H 2O (blue), and CoL
NO2
(green) under CO 2 in
DMF with 0.1 M TBAPF 6 and 1 M TFE. All concentrations were 0.5 mM. All electrolyses were performed at -2.70
V vs Fc
+/0
. Experiments were performed for 2 hours.

89

H2 CO
μmol FE
(%)
μmol FE
(%)
NO2L 0 0 0 0
Co(ClO4)2∙6H2O 90 84 ± 8 0 0
CoNO2L 167 65 ± 6 5 2 ± 1
Table 3.2. Results of 2 hour controlled potential electrolysis of 0.5 mM L
NO2
, Co(ClO 4)∙6H 2O, and CoL
NO2
in DMF
under CO 2 with 1 M TFE and 0.1 M TBAPF 6.

Figure 3.7. Controlled potential electrolysis of CoL
NO2
under N 2 (blue) and CO 2 (red) in DMF with 0.1 M TBAPF 6
and 1 M TFE. All concentrations were 0.5 mM. All electrolyses were performed at -2.70 V vs Fc
+/0
. Experiments
were performed for 24 hours.

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 311 73 ± 7 4 <1
4 hr 756 77 ± 8 35 4 ± 1
6 hr 1,026 69 ± 7 153 10 ± 1
8 hr 1,143 64 ± 6 245 14 ± 1
24 hr 1,350 53 ± 5 361 14 ± 1
Table 3.3. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under N 2 with 1 M TFE
and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

90

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 167 65 ± 6 5 2 ± 1
4 hr 297 62 ± 6 25 5 ± 1
6 hr 321 47 ± 5 49 7 ± 1
8 hr 566 64 ± 6 137 16 ± 1
24 hr 1,235 48 ± 5 210 8 ± 1
Table 3.4. Results of 24 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M TFE
and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

Figure 3.8. Controlled potential electrolysis of CoL
NO2
(0.5 mM) under CO 2 in DMF with 0.1 M TBAPF 6 and 1 M
TFE for 6 hours at -2.25 V (red), -2.45 V (blue), and -2.70 V (green) vs Fc
+/0
.

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 29 40 ± 4 0 0
4 hr 52 33 ± 3 0 0
6 hr 66 28 ± 3 0 0
Table 3.5. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M TFE
and 0.1 M TBAPF 6 at -2.25 V vs Fc
+/0
.

91

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 48 41 ± 4 0 0
4 hr 89 39 ± 4 7 3 ± 1
6 hr 139 40 ± 4 13 4 ± 1
Table 3.6. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M TFE
and 0.1 M TBAPF 6 at -2.45 V vs Fc
+/0
.

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 167 65 ± 6 5 2 ± 1
4 hr 297 62 ± 6 25 5 ± 1
6 hr 321 47 ± 5 49 7 ± 1
Table 3.7. Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M TFE
and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

Figure 3.9. Controlled potential electrolysis of CoL
NO2
(0.5 mM) under CO 2 in DMF with 0.1 M TBAPF 6 and 1 M
of H 2O (red), TFE (blue), phenol (green), or methanol (purple) as the proton source. All electrolyses were performed
at -2.70 V vs Fc
+/0
. Experiments were performed for 6 hours.

92

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 40 47 ± 5 0 0
4 hr 60 44 ± 4 3 2 ± 1
6 hr 78 43 ± 4 5 3 ± 1
Table 3.8 Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M H 2O
and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 178 60 ± 6 20 7 ± 1
4 hr 264 59 ± 6 45 10 ± 1
6 hr 290 58 ± 6 70 14 ± 1
Table 3.9 Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M phenol
and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

H2 CO
μ m ol FE (%) μ m ol FE (%)
2 hr 338 66 ± 6 4 <1
4 hr 601 63 ± 6 40 4 ± 1
6 hr 713 55 ± 5 105 8 ± 1
Table 3.10 Results of 6 hour controlled potential electrolysis of 0.5 mM CoL
NO2
in DMF under CO 2 with 1 M
methanol and 0.1 M TBAPF 6 at -2.70 V vs Fc
+/0
.

93

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Alters Product
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. Chemical
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Madison, WI, 2014.

98

CHAPTER 4. COVALENT ORGANIC FRAMEWORKS COMPOSED OF RHENIUM
BIPYRIDINE AND METAL PORPHYRINS: DESIGNING HETEROBIMETALLIC
FRAMEWORKS WITH TWO DISTINCT METAL SITES

4.1 ABSTRACT
The incorporation of homogeneous catalysts for CO2 reduction into extended frameworks has been
a successful strategy for increasing catalyst lifetime and activity, but the effects of the linkers on
catalysis are underexplored. In this work, a novel rhenium bipyridine complex was synthesized for
the purpose of designing a covalent-organic framework (COF) with both metalloporphyrin and
metal bipyridine moieties. Investigation of the rhenium complex as a homogeneous catalyst shows
a faradaic efficiency of 81(8)% for the electrocatalytic conversion of CO2 to CO upon the addition
of methanol as the proton source. Treatment of the rhenium complex with tetra(4-
aminophenyl)porphyrin (TAPP) under Schiff base conditions produces the desired COF, as
indicated by powder X-ray diffraction (PXRD) studies. Metallation of the porphyrins was
accomplished through post-synthetic modification with CoCl2 and FeCl3 metal precursors. The
retention of the PXRD peaks and appearance of new Co and Fe peaks in the corresponding X-ray
photoelectron spectroscopy (XPS) spectra suggest the successful incorporation of a secondary
metal site into the framework. Cyclic voltammetry measurements display increases in current
densities when the atmosphere is changed from N2 to CO2. Controlled potential electrolyses show
that the cobalt post metallated COF has the highest activity towards CO 2 reduction, reaching a
faradaic efficiency of 18(2)%.

4.2 INTRODUCTION
Metal- and covalent organic frameworks (MOFs and COFs, respectively) are classes of
crystalline coordination polymers that exhibit high porosities and surface areas as well as modular
structures.
1-7
Because of these properties, MOFs and COFs are attractive materials for applications
such as molecular sensing,
8-12
gas absorption and separation,
13-19
and catalysis.
7, 20-23

The use of MOFs and COFs for catalytic purposes has been extensively explored as the modular
nature of these materials has facilitated the incorporation of known molecular catalysts through
numerous available synthetic methods. The integration of molecular catalytic units into extended
frameworks has often resulted in improved efficiency, stability, and durability in comparison to
99

the corresponding molecular analogues.
24-26
Additionally, important catalytic properties such as
selectivity can be modulated by tuning the structural, electronic, and physical properties of the
frameworks through direct synthesis
27
or post synthetic modification.
28
This synthetic tunability
in conjunction with the improved durability of the molecular catalytic units following integration
into the framework environment grants the catalyst-modified frameworks benefits of both
homogeneous and heterogeneous catalysts.
Although MOFs exhibit several attractive qualities, the low electrical conductivity of the
MOFs has limited their ability to catalyze electrocatalytic reactions.
29, 30
The recent development
of two-dimensional frameworks with extensive -delocalization in the ab plane, and the
incorporation of redox-active ligands that can facilitate charge transport, has been a major
breakthrough for the field of electrocatalytic MOFs.
29, 30
MOFs and COFs have been recently
utilized for multiple types of electrocatalysis including hydrogen evolution, oxygen evolution,
oxygen reduction, and carbon dioxide reduction.
31
Porphyrins and bipyridines have proven to be
useful building blocks for these frameworks, performing a variety of transformations.
Electrocatalytically competent COFs composed of 5,10,15,20-tetra(4-
aminophenyl)porphyrincobalt(II) active sites were prepared using 1,4-benzenedicarboxaldehyde
(COF-366) and biphenyl-4,4'-dicarboxaldehyde (COF-367) as linkers.
32
These COFs were
reported to reduce CO2 to CO under aqueous conditions with faradaic efficiencies up to 90%,
turnover numbers up to 290,000, and an overpotential of 0.55 V.
32
Optimization of the
electrochemical set-up and introduction of linkers containing electron-withdrawing groups results
in an increase of current density up to 65 mA/mg.
33
Cobalt
34
and iron
35
porphyrin based MOFs
synthesized through atomic layer deposition and electrophoretic deposition, respectively, were
also identified as electrocatalysts for CO2 reduction. The cobalt porphyrin MOF produced CO
with a 76% faradaic efficiency and 1,400 turnovers per site under the same conditions as the cobalt
porphyrin COF.
34
The activity of the MOF was maintained for up to 7 hours, showing increased
lifetime relative to the molecular species.
34
The iron porphyrin MOF produced a mixture of CO
and H2 with a total faradaic efficiency of 100% at an overpotential of 0.65 V with a turnover
number of 1,520 upon the addition of a proton source.
35
Rhenium bipyridine-based MOF thin films
have been shown to catalyze the electrochemical reduction of CO2 to CO with faradaic efficiencies
around 93% and current densities up to 2 mA/cm
2
with an overpotential of 0.65 V.
36
MOFs based
on iron porphyrin
37
and nickel porphyrin
38
moieties have been used as oxygen evolution and
100

oxygen reduction catalysts, respectively, and cobalt bipyridine-based COFs
39
and ruthenium
bipyridine-based MOFs
40
have been investigated as oxygen evolution catalysts. These examples
highlight the synthetic ease with which well-known molecular porphyrin and biypridine-based
catalysts can be incorporated into a framework structure.
While selection of the appropriate catalytic active site to facilitate the desired
transformation is vital, the identity of the organic linker used for framework formation influences
several important properties that dictate the viability of the generated COFs as electrocatalysts. A
large library of organic linkers have been utilized in the syntheses of catalytically active COFs to
improve structural stability
41
in aqueous acidic/alkaline media and increase the accessible surface
area,
32, 41
which is of particular importance for substrate and product diffusion. Reasonably
conductive frameworks can make use of variations in the linkages to tune the active site, providing
increased control over activity and selectivity.
33
Issues of poor conductivity can be lessened in
certain cases by selecting linkers which orient metal centers in a manner which allows for charge
hopping,
37, 38
or by incorporating redox active linkages.
30
While this growing catalog of organic
linkers aims to maximize catalysis at the metal center, the linkers themselves are traditionally
catalytically inactive. An underexplored area of COF catalysis is the utilization of only
catalytically active building blocks, which will allow for the introduction of two unique metal sites.
Metalloporphyrin sites can be connected through metallobipyridine linkages to produce COFs with
two distinct metal sites. By substituting catalytically inactive linkages with active ones, COFs with
two distinct metal sites can be produced, which will be beneficial in formulating design principles
for the development of electrocatalytic MOFs.
Herein, 2,2'-bipyridine-5,5'-dicarbaldehyde and 2,2'-bipyridine-5,5'-
dicarbaldehydetricarbonylchlororhenium(I) are synthesized, electrochemically analyzed, and
integrated into a porphyrin-bipyridine bifunctional COF through a Schiff base reaction with
5,10,15,20-tetra(4-aminophenyl)porphyrin. Post synthetic modification with cobalt and iron
chloride salts results in the incorporation of the corresponding metal ions into the porphyrin
binding pocket, producing a COF with two distinct metal sites. Given the extensive literature on
the electrochemical CO2 reduction using rhenium bipyridine, cobalt porphyrin, and iron porphyrin
molecular species,
42-49
the electrocatalytic CO2 reduction activity was explored in aqueous media.

101

4.3 RESULTS AND DISCUSSION
4.3.1 Synthesis and Characterization of Rhenium Complex
The bipyridine-based linker was synthesized by radical bromination of 5,5'-dimethyl-2,2'-
bipyridine through the use of azobisisobutyronitrile (AIBN) and N-bromosuccinimide (NBS),
which produced the dibrominated product, 5,5'-bis(bromomethyl)-2,2'-bipyridine (Scheme 4.2).
Attempts to convert the 5,5'-bis(bromomethyl)-2,2'-bipyridine into the corresponding dialdehyde
through the reported procedure using Bredereck’s reagent followed by sodium periodate
50
were
unsuccessful. An alternative synthetic route using modified Sommelet reaction conditions (see
Experimental Methods) yielded 2,2'-bipyridine-5,5'-dicarbaldehyde (1) as a white solid.
51
The
1
H
NMR spectrum (Figure 4.13) of 1 in DMSO-d6 reveals peaks corresponding to the aldehyde proton
at δ 10.14 ppm (s) and the aromatic protons at δ 9.22 (d), 8.66 (d), and 8.43 (dd) ppm. Refluxing
1 with pentacarbonylchlororhenium(I) in toluene generates the rhenium complex 2 as a red solid
(Equation 4.1).
52
X-ray quality crystals of 2 were grown from a mixture of dichloromethane and
hexanes. The crystal structure of 2 (Figure 4.1) shows bidentate coordination of the bipyridine
ligand to the rhenium metal center with Re-N bond lengths of 2.172(4) and 2.184(4) Å. The
rhenium center maintains three of the five carbonyl groups, with two carbonyls oriented in the
equatorial plane and the third in an axial coordination. The chloride ligand occupies the other axial
position.

(4.1)

Figure 4.1. X-ray crystal structure of 2. Selected hydrogen atoms have been omitted from the bipyridine ligand for
clarity.
102

The
1
H NMR spectrum of 2 displays four peaks at δ 10.23 (s), 9.47 (d), 9.04 (d), and 8.73
(dd) ppm (Figure 4.15). The aromatic protons of the bipyridine ligand are deshielded due to the
presence of the rhenium center. The peaks in the
13
C NMR spectrum of 2 (Figures 4.16) are also
deshielded in comparison to the free bipyridine ligand (Figure 4.14) with the aromatic carbons
shifting to higher ppm values upon formation of 2. Additionally, two new peaks appear at δ 197
and 189 ppm, which correspond to the axial and equatorial carbonyl ligands.
The Fourier transform infrared (FT-IR) spectrum of 2 (Figure 4.2) reveals two intense
peaks at 2025 and 1897 cm
-1
, which are not present in 1 (Figure 4.17). These peaks are
characteristic of metal-carbonyls, with the 2025 cm
-1
peak corresponding to the fully-symmetric
stretching mode of the carbonyl ligands, while the broad peak at 1897 cm
-1
is due to the coalesced
in-phase and out-of-phase stretching modes. The positions of these peaks are consistent with
previously reported rhenium bipyridine complexes with carboxylate groups in the 5 and 5'
positions, Re(bpydc)(CO)3Cl (bpydc = 2,2'-bipyridine-5,5'-dicarboxylate), which displays three
peaks at 2022, 1920, and 1910 cm
-1
.
53

Figure 4.2. FT-IR spectrum of 2. The black box highlights the carbonyl stretching frequency region.

103

4.3.2 Electrochemical Studies of Molecular Species
Electrochemical experiments were conducted in dimethylformamide (DMF) solution with
0.5 mM of 1 or 2 and 0.1 M tetrabutylammonium hexafluorophosphate (TBAPF6) as the
supporting electrolyte. The cyclic voltammogram (CV) of the ligand 1 under a nitrogen
atmosphere shows two reversible one-electron reductions with E1/2 of -1.08 and -1.38 V versus
SCE (Figure 4.3A). Complex 2 also displays two reversible one-electron reductions under nitrogen
similar to the unmetallated ligand; however, these peaks are positively shifted to -0.56 and -0.90
V versus SCE. The peak current densities of these reversible couples increase linearly with respect
to the square root of the scan rate (Figures 4.18-4.21), as expected for freely diffusing species in
solution. The reduction potentials of 2 are positively shifted by roughly 800 mV in comparison to
previously reported reduction potentials for the well-established CO2 reduction catalysts, Re([2,2'-
bipyridine]-4,4'-tBu)(CO)3Cl (-1.45 V and -1.83 V versus SCE) and Re(bpy)(CO)3Cl (-1.34 V and
-1.73 V versus SCE).
54
Additionally, while 2 demonstrates two reversible couples that are ~340
mV apart, other known rhenium bipyridine complexes display a quasi-reversible first reduction
event, followed by an irreversible second reduction.
54

104

Figure 4.3. Electrochemical studies of 1 and 2 in DMF. (A) Overlay of the cyclic voltammograms of 1 (0.5 mM,
blue) and 2 (0.5 mM, red) under an N 2 atmosphere and (B) cyclic voltammograms of 2 (0.5 mM) under N 2 (purple)
and CO 2 (green) atmospheres. (C) Cyclic voltammogram of 2 (0.5 mM) under 1 atm of CO 2 with increasing
concentrations of MeOH (from 0 to 4 M). Cyclic voltammetry experiments were performed in DMF with 0.1 M
TBAPF 6 at a scan rate of 20 mV/s. (D) Controlled potential electrolysis of 1 (blue), 2 (red), and a blank solution
(dashed black) at -2.0 V versus SCE in DMF with 4 M MeOH under a CO 2 atmosphere.

N2 Atmosophere CO2 Atmosphere
H2 CO H2 CO
μmol FE (%) μmol FE (%) μmol FE (%) μmol FE (%)
Blank 35.0

73(7) 0 0 7.2

63(6) 0 0
1 2.4

5(1) 0 0 0 0 0 0
2 20.0

27(2) 0 0 0 0 73.7

81(8)
Table 4.1. Gaseous products measured using gas chromatography after 2 hours of bulk electrolysis at -2.0 V vs SCE
for a blank solution, 1, and 2 under N 2 and CO 2 atmospheres with 4 M MeOH.

105

When the solution of 2 was saturated with CO2, the reversible couples were unperturbed,
but a current increase was observed at a potential of approximately -1.60 V versus SCE (Figure
4.3B). Re([2,2'-bipyridine]-4,4'-tBu)(CO)3Cl and Re(bpy)(CO)3Cl display current increases near
similar potentials (-1.83 V and -1.73 V versus SCE, respectively) in the presence of CO2, which
was assigned to the onset of CO2 reduction with these rhenium complexes.
54
The current increase
observed in the CV experiments of 2 under CO2 is indicative of an interaction between 2 and CO2.
In comparison, ligand 1 displays negligible increase in current when the atmosphere is changed
from nitrogen to CO2 (Figures 4.22 and 4.23), indicating that the free ligand does not react with
CO2.
Acid titrations were conducted in DMF under CO2 to determine if the availability of excess
protons enhanced the catalytic activity of 2. When methanol is used as the proton source, current
increases are observed upon successive additions of MeOH with a five-fold increase occurring at
4 M MeOH (Figure 4.3C). A negligible current increase was observed for ligand 1 in the presence
of CO2 and 4 M MeOH (Figure 4.26). Additionally, species 1 and 2 display low current increases
under nitrogen and 4 M MeOH (Figures 4.24 and 4.25), indicating that CO 2 is required for the
current increases observed.
Controlled potential electrolysis (CPE) experiments were performed in 0.1 M TBAPF 6
DMF solution at -2.0 V versus SCE under an atmosphere of CO2, to identify and quantify the
reduction products. In the absence of an added proton source, 1 and 2 produced no quantifiable
amount of gaseous products after 4 hours of electrolysis. Upon the addition of 4 M MeOH,
complex 2 selectively produces 73.7 μmol of CO (17.5 C), corresponding to a faradaic efficiency
(FE) of 81(4)% (Figure 4.3D and Table 4.1). Other CO2 reduction products, such as formate/formic
acid, were not detected. Upon completion of electrolysis, the electrode was placed in a fresh
solution without the complex. The resulting cyclic voltammogram resembled the bare electrode,
which suggests no deposition of active species onto the electrode surface is occurring (Figure
4.27). CPE conducted under a N2 atmosphere and 4 M MeOH produced no CO (Figure 4.28 and
Table 4.1). To determine if 2 is responsible for the generated CO, control experiments were
performed at -2.0 V versus SCE with the free ligand 1 (0.5 mM) in the presence of CO2 and 4 M
MeOH (Figure 4.3D and Table 4.1). No CO was produced following these experiments, indicating
that the presence of the rhenium center in 2 is required for CO2 reduction. These control
106

experiments suggest that CO is only produced in the presence of 2, CO2, and an added proton
source (MeOH).
Rhenium bipyridine complexes are known to reduce CO2 to CO with faradaic efficiencies
near unity. The decrease in FE of complex 2 (81(8)%) could be due to the electrochemical
reduction of the aldehyde moiety. Aromatic aldehydes such as benzaldehyde were reported to
generate radicals under reductive conditions, which led to dimerized products with quinoid
structures.
55-58
Therefore, the incorporation of the rhenium-bipyridine catalyst 2 into an extended
framework is a viable method for eliminating the reactive aldehyde moieties through the formation
of imines by Schiff based reactions necessary to generate the covalent-organic frameworks of
interest.

4.3.3 Immobilization of Rhenium Complex via COFs
The incorporation of the rhenium bipyridine catalyst 2 into a covalent-organic framework
was accomplished by treating 2 with 5,10,15,20-tetra(4-aminophenyl)porphyrin (TAPP) in a
mixture of o-dichlorobenzene, 1-butanol, and 6 M acetic acid (1:1:0.2) in a sealed ampule under
vacuum, which led to the formation of COF-Re. An analogous derivative, whereby the
unmetallated bipyridine ligand 1 was used as the linker in the synthesis, yielded a second
framework denoted COF-Bpy. This framework is similar to the CuP-BPyPh COF reported
previously which contains a copper porphyrin and a bipyridine linker
59
and serves as a useful
structural comparison for COF-Re. This synthetic strategy allows for the preparation of a
framework containing a mixture of linkers 1 and 2. TAPP was treated with a 1:1 ratio of 1 and 2
to produce COF-Mix.

107

Scheme 4.1. Syntheses of COF-Re, COF-Re_Co, and COF-Re_Fe.

The crystalline structure of the synthesized frameworks was confirmed by powder X-ray
diffraction (PXRD) studies. The PXRD patterns of COF-Re and COF-Mix show major peaks at
2θ values of 3.1, 4.3, and 6.2° corresponding to 28.3, 20.6, and 14.2 Å, respectively (Figure 4.30
and Figure 4.31). These patterns are comparable to those of the 2-D cobalt porphyrin framework
constructed from cobalt-metallated TAPP and biphenyl-4,4'-dicarboxaldehyde (COF-367).
32
The
intensity of these peaks is slightly diminished in COF-Bpy, with only the peak at 3.1° being
observed in the PXRD pattern (Figure 4.29).
Attempts to directly synthesize a metalloporphyrin-based COF with either 1 or 2 as the
linker proved unsuccessful. Therefore, a post synthetic modification strategy was employed
whereby the free-base porphyrins were post metallated with either cobalt or iron precursors.
Insertion of metal ions into free-base porphyrin frameworks has been previously accomplished
through the use of post metallation techniques
35
and analogous conditions were utilized here.
COF-Re was heated in DMF at 80 °C in the presence of excess cobalt(II) or iron(III) chloride salts
for 24 hours generating COF-Re_Co and COF-Re_Fe, respectively.
The PXRD patterns of COF-Re_Co and COF-Re_Fe are identical to the pattern of COF-
Re (Figures 4.4, 4.32, and 4.33) confirming the preservation of the ordered structure following
post metallation. When the bipyridine sites in the framework are all fully occupied by rhenium as
in COF-Re, the resultant post metallated product retains its long-range order. However, when
COF-Bpy and COF-Mix are post metallated with iron or cobalt, the resulting product loses
crystallinity as evidenced by the loss of the diagnostic peaks in the PXRD pattern. Modelling was
108

performed with Materials Studio to understand the stacking of COF-Re and the post metallated
frameworks. Of the various stacking modes explored, the two space groups (P21212 and PCC2) that
more closely matched the experimental structure were both in an eclipsed conformation. In the
P21212 space group, the rhenium moieties are fully eclipsed, while in the P CC2 space group, the
bipyridine moieties are eclipsed, but the orientation of the rhenium tricarbonyl moieties alternate
between adjacent layers (Figures 4.5-4.8). These space groups display predicted peaks that align
well with the experimental major peak at 2θ of 3.1°, as well as with the peak at 6.2°. The peak at
4.3° is associated with the orientation of the rhenium moieties, as frameworks without rhenium
lack this peak in the predicted PXRD patterns, and moreover, this peak is absent in the
experimental PXRD pattern of COF-Bpy.

Figure 4.4. PXRD patterns of COF-Bpy (green), COF-Re (yellow), COF-Re_Co (light red), and COF-Re_Fe
(black).
109

Figure 4.5. P 21212 space group used to model COF-Re-Co.

Figure 4.6. Overlay of the P 21212 space group predicted (blue) and experimental (magenta) PXRD patterns.

110

Figure 4.7. P CC2 space group used to model COF-Re-Co.

Figure 4.8. Overlay of the P CC2 space group predicted (blue) and experimental (magenta) PXRD patterns.

111

The incorporation of cobalt and iron ions into COF-Re was confirmed by X-ray
photoelectron spectroscopy (XPS). Analysis of the COF-Re Re 4f spectrum (Figure 4.9A) showed
two peaks with binding energies of 44.2 and 41.9 eV, corresponding to Re 4f5/2 and 4f7/2. These
peaks and binding energies are similar to those of the previously reported rhenium(I) bipyridine
complexes.
36, 60, 61
There are also three nitrogen peaks present in the N 1s region at 400.5, 399.6,
and 398.1 eV (Figure 4.35) which are assigned to the three unique nitrogen environments in the
framework.

Figure 4.9. XPS analyses of COFs. (A) Re 4f XPS spectrum of COF-Re. (B) Co 2p XPS spectrum of COF-Re_Co.
(C) Fe 2p XPS spectrum of COF-Re_Fe.

XPS spectra of COF-Re_Co and COF-Re_Fe show retention of the rhenium peaks
(Figures 4.36 and 4.37). Additional peaks corresponding to Co 2p (Figure 4.5B) and Fe 2p (Figure
4.5C) are observed. The cobalt region in COF-Re_Co displays two peaks at 795.6 and 780.1 eV,
corresponding to Co 2p1/2 and 2p3/2 of a Co(II),
62
respectively. The iron region in the COF-Re_Fe
displays two peaks at 724.4 and 710.5 eV, corresponding to Fe 2p1/2 and 2p3/2 of Fe(III),
respectively.
63
These values are consistent with reported Co
62-64
and Fe
63, 65
porphyrin molecular
species. The retention of the Re peaks and the appearance of Co and Fe peaks is a strong indicator
of successful incorporation of two distinct metal centers in the COF. The N 1s region is unchanged
from COF-Re to COF-Re_Co (Figure 4.36A) and COF-Re_Fe (Figure 4.37A), respectively.
The vibrational stretching modes of the carbonyls in the COFs were analyzed by FT-IR
and compared to the stretches observed for 2 (Figures 4.2, 4.10, and 4.38-4.41). The FT-IR
spectrum of COF-Bpy shows no CO stretches as expected. On the other hand, the FT-IR spectrum
of COF-Re displays CO stretching frequencies at the same positions (2025 and 1897 cm
-1
) as in
complex 2, indicating the retention of the molecular structure of 2 upon incorporation into the
extended COF structure. Identical CO stretching frequencies are also present in COF-Re_Co and
COF-Re_Fe, further indicating that rhenium bipyridine moieties persist after post synthetic
112

modification. Additionally, other regions of the IR spectra show strong similarities, suggesting
structural retention even after post synthetic modification.

Figure 4.10. FT-IR spectra of COF-Bpy (green), COF-Re (yellow), COF-Re_Co (light red), and COF-Re_Fe
(dark red).

The accessible surface area of the activated COFs was measured using the Brunauer-
Emmett-Teller (BET) technique (see Experimental Methods). The frameworks COF-Re, COF-
Re_Co, and COF-Re_Fe display similar surface areas with values of 618.30, 615.97, and 723.12
m
2
/g, respectively (Table 4.3). These surface areas are comparable with those of other structurally
similar porphyrin-based COFs.
33, 59
Thermogravimetric analysis measurements indicate negligible
decreases in mass percent up to 300 °C, indicating that the materials have been freed of residual
solvent (Figure 4.42-4.44).
ICP analyses of the frameworks indicate a Re:Co ratio of 2 to 0.53 and a Re:Fe ratio of 2
to 0.31 for COF-Re_Co and COF-Re_Fe, respectively. The expected ratio is 2 rhenium centers
per 1 cobalt or iron center, suggesting that cobalt incorporation in COF-Re_Co is 53.4%, while
the iron incorporation in COF-Re_Fe is 30.7%. While previously reported post metallations of
porphyrin-based frameworks led to near unity incorporation of metal ions,
35
it has also been shown
that full incorporation is unnecessary as all metal sites may not be electrochemically active.
32

The activity of the COF-Re, COF-Re_Co, and COF-Re_Fe towards the electrocatalytic
reduction of CO2 was explored in pH 7.2 aqueous phosphate buffer solutions with 0.5 M KHCO3.
The conditions used here have been previously reported for successful CO2 reduction using cobalt
porphyrin containing frameworks (COF-366 and COF 367).
32
The COF-modified electrodes were
113

prepared by dropcasting the materials (0.5 mg) onto a 2 cm  1 cm strip of carbon fabric. Based
on the ICP-OES results, the estimated loading for COF-Re_Co is 4.66  10
-7
mol of Re and 1.24
 10
-7
mol of Co, while the loading for COF-Re_Fe is 5.15  10
-7
mol of Re and 7.90  10
-8
mol
of Fe.
Under a nitrogen atmosphere, none of the COFs displayed any observable redox features.
The CVs of COF-Re, COF-Re_Co, and COF-Re_Fe demonstrated a slight positive shift in onset
potential in the presence of CO2 with COF-Re_Fe exhibiting a current increase as well (Figure
4.11). These current increases are similar to those previously reported for similar porphyrin
frameworks.
32, 33

Figure 4.11. Cyclic voltammograms of (A) COF-Re, (B) COF-Re_Co, and (C) COF-Re_Fe in pH 7.2 aqueous
phosphate buffer solutions with 0.5 M KHCO 3 under N 2 (blue) and CO 2 (red) atmosphere. Scan rate = 100 mV/s.

Controlled potential electrolyses were conducted in pH 7.2 aqueous phosphate buffer
solutions with 0.5 M KHCO3 under a CO2 atmosphere for one hour at -1.1 V vs SHE to determine
if any gaseous products are formed (Figure 4.12, Table 4.2). No CO was generated for COF-Re
and the bare carbon fabric. COF-Re_Co produced 12.7 μmol of CO, corresponding to a faradaic
efficiency of 18(2)%. H2 was also formed during the electrolysis with a faradaic efficiency of
55(5)%. COF-Re_Fe produced 5.3 μmol of CO under the same conditions, corresponding to a
faradaic efficiency <2%. Again, H2 was formed in larger quantities with a faradaic efficiency of
65(6)%.
114

Figure 4.12. CPE of bare carbon fabric (dashed black), COF-Re (red), COF-Re_Co (blue), and COF-Re_Fe
(green) at -1.1 V vs SHE under CO 2 in pH 7.2 aqueous phosphate buffer with 0.5 M KHCO 3.

H2 CO
μmol FE (%) μmol FE (%)
Bare carbon fabric
72.1
67(6)
0
0
COF-Re
2.8
11(1)
0
0
COF-Re_Co
39.4
55(5)
12.7
18(2)
COF-Re_Fe
359.6
65(6)
5.3
<2
Table 4.2. Gaseous products measured by GC after 120 minutes of bulk electrolysis with bare carbon fabric, COF-
Re, COF-Re_Co, and COF-Re_Fe under CO 2 atmosphere.

Conducting CPE at -0.9 V, -1.0 V, and -1.3 V vs SHE showed no further enhancement of
CO production, which indicates that selectivity is not affected by changes in the reduction
potential. Doubling the amount of material dropcast on the electrode also had no impact on the
production of CO. The amount of metal salt used in the post metallation reactions was also
modified to as low as 1 equivalent in an attempt to maximize CO output, but that resulted in no
change. The post metallation procedure employed here has been reported for porphrin-based
frameworks to incorporate iron ions into available porphyrin sites with nearly 100% efficiency,
35

but in this report the maximum incorporation of the metal ions is 53.4% for Co and 30.7% for Fe.
These results may suggest that diffusion through the pores of the COFs is somewhat limited, which
may impede transport of CO2 and electrolyte to the active sites. Previous reports have also shown
that the electronics of the linkage can affect the activity of the porphyrin species.
33
The CVs of the
115

rhenium bipyridine complex 2 (Figure 4.3A) indicate that the rhenium moiety is reduced at
potentials more positive than those typically observed for other rhenium bipyridines, which may
suggest that the frameworks undergo multiple reductions events on different metal sites, and
therefore, compete for electrons. This competition could inhibit the generation of a doubly reduced
species required for CO2 reduction. This could lower the overall affinity for CO2 reduction,
resulting in the more thermodynamically favorable H2 evolution reaction taking place instead.

4.4 CONCLUSIONS
A novel rhenium bipyridine complex, 2,2'-bipyridine-5,5'-
dicarbaldehydetricarbonylchlororhenium(I) (2), was shown to be an active catalyst for
electrochemical CO2 reduction. The complex reduces CO2 to CO with 81(4)% FE and exhibits a
coordination geometry similar to that of the previously reported rhenium bipyridine complexes.
The immobilization of complex 2 in a COF was attempted in an effort to overcome the competing
reduction pathways involving terminal aldehydes. Treatment of 2 with 5,10,15,20-tetra(4-
aminophenyl)porphyrin generated the desired imine based framework, COF-Re, and subsequently
post metallation with iron and cobalt precursors led to COFs containing two distinct metal sites
(COF-Re_Co and COF-Re_Fe). XPS and FT-IR experiments confirmed the molecular nature of
the incorporated catalysts and PXRD studies established the retention of long-range order of the
frameworks after integration of two distinct metal sites. Following the successful synthesis of
heterobimetallic COFs, whose two metal sites are well-known molecular catalysts for the reduction
of CO2 to CO, electrochemical studies were performed to determine the ability of the frameworks
to reduce CO2. Slight positive shifts in the onset potentials and increases in current are observed
in the cyclic voltammograms of the frameworks when the atmosphere is changed from N2 to CO2.
This behavior is similar to that of previously reported cobalt porphyrin frameworks. COF-Re_Co
had the highest activity towards CO2 reduction, achieving a faradaic efficiency of 18(2)%. The
low selectivity for CO2 reduction is potentially linked to a competitive relationship between the
two metal centers as opposed to the desired synergistic relationship. Changes to the
metallobipyridine linker may be necessary to promote catalysis.

116

4.5 ACKNOWLEDGEMENTS
The authors are grateful to the University of Southern California (USC) for funding, the USC
Wrigley Institute for the Norma and Jerol Sonosky summer fellowship to EMJ, and to Mr. Thomas
Moulton and Ms. Ginny Dunn for the Harold and Lillian Moulton Graduate Fellowship for EMJ.
The studies of the molecular complex were supported by the National Science Foundation (NSF)
through the CAREER award (CHE-1555387), and the studies of the extended frameworks were
supported by the Nanoporous Materials Genome Center of the U.S. Department of Energy, Office
of Basic Energy Sciences, Division of Chemical Sciences, Geosciences and Biosciences under
Award DE-FG02-17ER16362. The authors are grateful to NSF (grant CRIF 1048807) and USC
for their sponsorship of NMR spectrometers and X-ray diffractometer. The authors thank Dr.
Patrick Cottingham of USC and Julia Oktawiec of Cal Tech for assistance with BET and Andrew
Clough for assistance with XPS. XPS data were collected and critical point CO2 washings were
conducted at the Center for Electron Microscopy and Microanalysis (CEMMA) at USC. The
authors would like to thank Dr. Courtney Downes for helpful discussion.

4.6 EXPERIMENTAL METHODS
General. Manipulations of air and moisture sensitive materials were carried out under nitrogen
either in a Vacuum Atmospheres drybox or on a dual-manifold Schlenk line. All solvents were
degassed with nitrogen and passed through activated alumina columns and stored over 4Å Linde-
type molecular sieves. TAPP
66
and 5,5'-bis(bromomethyl)-2,2'-bipyridine
67
were prepared
according to literature procedures. All other chemical reagents were purchased from chemical
vendors and used without further purification. NMR spectra were obtained using a Varian Mercury
400 MHz NMR Spectrometer. Elemental analyses and ICP-OES were performed by Complete
Analysis Laboratories, Inc., Parsippany, New Jersey, 07054 or Robertson Microlit Laboratories,
1705 U.S. Highway 46, Suite 1D, Ledgewood, New Jersey, 07852. Critical point CO2 washings
were performed with a Tousimis Autosamdri®-815, Series B sample drier.

4.6.1 Physical Methods. UV-Vis spectra were obtained using a Lambda 950 UV/Vis/NIR
Spectrophotometer.

117

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.

Powder X-ray diffraction (PXRD) studies were performed on a Rigaku Ultima IV X-Ray
diffractometer in reflectance parallel beam/parallel slit alignment geometry. The measurements
employed Cu Kα line focused radiation at 1760 W (40 kV, 44 mA) power and a Ge crystal detector
fitted with a 10 mm radiation entrance slit. Samples were mounted on zero-background sample
holders. Samples were observed using a 0.01° 2θ step scan from 2.0 – 30.0° with an exposure time
of 0.60 s per step.

The single-crystal X-ray diffraction data were collected on a Bruker SMART APEX DUO 3-circle
platform diffractometer, equipped with an APEX II CCD, using Mo Kα radiation (TRIUMPH
curved-crystal monochromator) from a fine-focus tube. The diffractometer was equipped with an
Oxford Cryosystems Cryostream 700 apparatus for low-temperature data collection. The frames
were integrated using the SAINT algorithm to give the hkl files corrected for Lp/decay.
68
The
absorption correction was performed using the SADABS program.
69
The structures were solved
by intrinsic phasing and refined on F
2
using the Bruker SHELXTL Software Package and
ShelXle.
70
All non-hydrogen atoms were refined anisotropically.

Modelling was performed in Materials Studio version 8. The chosen space group was P 21212 with
lattice parameters a = 35.4985 Å, b = 47.041 Å, c = 6.6542 Å and α = β = γ = 90°. Overlaying the
predicted PXRD pattern with the experimental pattern places the two largest theoretical peaks in
agreement with the experimental ones. This result indicates that the structure is largely eclipsed in
a manner similar to that of COF-367, though with a slight lattice expansion to accommodate the
presence of the rhenium centers. The broad experimental peak at 4.12° 2θ is only observed for the
rhenium-containing frameworks, and not for COF-Bpy, suggesting that this is a rhenium related
peak. Two different stacking orientations of the rhenium moieties where explored. The first
orientation in the P21212 space group has a fully eclipsed geometry. A PCC2 space group was also
tested (c = 13.2542 Å, all other parameters unchanged), in which the bipyridine moieties are
eclipsed, but the orientation of the rhenium tricarbonyl moieties alternate between adjacent layers
to relieve some steric strain (Figures S22-S25). These two space groups display predicted peaks
118

that align well with the major experimental peak at 2θ of 3.1°, as well as with the peak at 6.2°. The
peak at 4.12° 2θ is broad and only present in the rhenium containing frameworks, whereas the
peaks at 3.1° and 6.2° 2θ are present in all COFs, suggesting that the frameworks are generally in
the eclipsed orientation.

XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray source
was the Al K α line at 1486. 7 eV, and the hybrid lens and slot mode were used. 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 regions
of interest. The sample chamber was maintained at < 9 × 10
–9
Torr. The XPS data were analyzed
using the CasaXPS software.

Thermogravimetric analysis was performed on the activated COFs with a Netzsch TG 209 F3
Tarsus under N2 with a flow rate of 40 mL/min in alumina crucibles. The temperature was
controlled by the furnace, and ranged from 30 °C to 900 °C at a ramp rate of 5 °C/min.

Gas adsorption isotherms for pressures in the range 0-1.2 bar were measured by a volumetric
method using a Micromeritics ASAP 2420 instrument. A typical sample of ca. 30 mg of material
was transferred to a pre-weighed analysis tube, which was capped with a Transeal and evacuated
by heating under dynamic vacuum until an outgas rate of less than 2.6 μbar/min was achieved. The
evacuated analysis tube containing the degassed sample was then carefully transferred to an
electronic balance and weighed again to determine the mass of sample. The tube was then
transferred back to the analysis port of the gas adsorption instrument. The outgas rate was again
confirmed to be less than 2.6 μbar/min. For all isotherms, warm and cold free space correction
measurements were performed using ultra-high purity He gas (UHP grade 5.0, 99.999% purity);
N2 isotherms at 77 K were measured in liquid nitrogen baths, using UHP-grade gas sources. Oil-
free vacuum pumps and oil-free pressure regulators were used for all measurements to prevent
contamination of the samples during the evacuation process or of the feed gases during the
isotherm measurements.

119

4.6.2 Electrochemical Methods. Electrochemical experiments were carried out using a Pine
potentiostat. Carbon fabric (99.5% carbon) was purchased from Fuel Cell Earth LLC. The fabric
was cut into 1 × 2.5 cm
2
pieces, treated with 6 M HCl overnight to remove trace metal impurity,
rinsed thoroughly with Milli-Q water, and dried in air before use. The electrochemical experiments
were carried out in three-electrode configuration electrochemical cell further described below
under nitrogen or carbon dioxide atmospheres using carbon fabric or glassy carbon electrodes as
the working electrode.

4.6.3 Preparation of Carbon Fabric COF-modified Working Electrodes. 0.5 mg of COF
catalyst was weighed out on a balance, suspended in 80 μL of DMF, and sonicated for 20 minutes.
The cleaned carbon fabric (1 × 2.5 cm
2
) was placed on a glass plate, and the suspended material
was dropcast over the bottom 1 × 2 cm
2
area and left to dry overnight. A silver wire was then
threaded through the top 1 × 0.5 cm
2
portion of the carbon fabric for electrical contact.

4.6.4 Cyclic Voltammetry of 1 and 2. Electrochemical studies of 1 (0.5 mM) and 2 (0.5 mM)
were performed in anhydrous dimethylformamide with tetrabutylammonium hexafluorophosphate
(TBAPF6) (0.1 M) as the supporting electrolyte. Platinum wire was used as the counter electrode,
silver wire as the reference electrode, and glassy carbon as the working electrode.
Decamethylferrocene was added as an internal standard to reference to ferrocene (Fc
+/0
), and all
potentials are given with reference to the saturated calomel electrode (SCE) by adding 0.40 V.
Samples were sparged with N2 or CO2 gas for 10 minutes before experiments.

4.6.5 Cyclic Voltammetry of COFs. Electrochemical experiments for the prepared COF materials
were performed in a pH 7.2 phosphate buffer containing 0.5 M KHCO3. For the pH 7.2 solution,
NaH2PO4 (0.468 g) and Na2HPO4 (1.637 g) were dissolved in water (100 mL). KHCO3 (5.00 g)
was then added to the solution. The pH was measured using a benchtop Mettler Toledo pH meter.
The working electrode was carbon fabric that had been treated with dropcast COF material and
connected to the potentiostat through a silver wire. The counter electrode was a clean carbon fabric
connected through a silver wire. The reference electrode was a Ag/AgCl/saturated KCl (aq)
electrode separated from the solution by a Vycor tip. Samples were sparged with N2 or CO2 gas
for 10 minutes before experiments. Electrolyses were first conducted at -1.30 V vs
120

Ag/AgCl/saturated KCl (aq) for 10 minutes to activate the surface of the carbon fabric 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. Potentials were
converted to the standard hydrogen electrode (SHE) by adding a value of 0.205 V.

4.6.6 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
electrolyte and the second chamber held the auxiliary electrode in 25 mL of electrolyte. The two
chambers were both under N2 or CO2 and separated by a fine porosity glass frit. 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 CO
amount based on total charge flowed is compared with the GC-detected CO produced from the
controlled-potential electrolysis experiment.

4.6.7 Controlled Potential Electrolysis of 1 and 2. CPE experiments of 1 (0.5 mM) and 2 (0.5
mM) were performed in 0.1 M TBAPF6 DMF solution 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 wire separated from the solution by a Vycor tip. The working compartment
was sparged with N2 or CO2 for 10 minutes before the experiment.

4.6.8 Controlled Potential Electrolysis of COFs. CPE experiments of COFs were performed in
0.5 M KHCO3 (aq) solution. The solution was prepared by dissolving KHCO3 (5.00 g) in water
(100 mL). The working electrode was carbon fabric that had been treated with dropcast COF
material connected to the potentiostat through a silver wire. The counter electrode was a cleaned
carbon fabric connected through a silver wire. The reference electrode was a Ag/AgCl/saturated
KCl (aq) electrode separated from the solution by a Vycor tip. The working compartment was
sparged with N2 or CO2 for 10 minutes before the experiment.

121

4.7 SYNTHETIC METHODS
2,2'-Bipyridine-5,5'-dicarbaldehyde (1).
51
Dry glassware was charged with 5,5'-
bis(bromomethyl)-2,2'-bipyridine (0.7630 g, 2.34 mmol) and hexamethylenetetraamine (1.4695 g,
10.49 mmol). The solids were dissolved in 40 mL of dry dichloromethane, and the solution was
brought to reflux in air. After approximately 30 minutes, a precipitate began forming. The reaction
mixture was continued to reflux for 12 hours, then allowed to cool to room temperature. The
solution was filtered and the white solid was taken up in 40 mL of 1 M acetic acid. The solution
was brought to reflux for 10 hours, cooled, and briefly placed in the freezer to precipitate the
organic product. The white solid was collected by filtration (55% yield). The product was
crystallized by slow evaporation of an acetone solution.
1
H NMR (400 MHz, DMSO-d6) δ 10.14
(s, 2H), 9.11 (q, 2H), 8.66 (q, 2H), 8.27 (q, 2H).
13
C NMR (100 MHz, DMSO-d6) δ 192.11, 158.02,
151.55, 137.57, 131.67, 121.87.

2,2'-Bipyridine-5,5'-dicarbaldehydetricarbonylchlororhenium(I) (2).
52
Dry toluene (50 mL)
was warmed in a dry reflux apparatus under nitrogen. Pentacarbonylchlororhenium(I) (146.9 mg,
0.40 mmol) and 1 (68.4 mg, 0.32 mmol) were added to the apparatus, and the solution was brought
to reflux. After 18 hours, the solution was cooled, and a red solid appeared. The solid was collected
by filtration and washed with hexanes to remove excess toluene. Yield = 86%.
1
H NMR (400 MHz,
DMSO-d6) δ 10.23 (s, 2H), 9.47 (q, 2H), 9.04 (d, 2H), 8.73 (q, 2H).
13
C NMR (100 MHz, DMSO-
d6) δ 197.21, 190.31, 189.24, 157.62, 154.81, 139.25, 133.86, 125.91. Anal. Calcd for
C15H8ClN2O5Re: C, 34.79; H, 1.56; N, 5.41. Found: C, 35.02; H, 1.56; N, 5.03.

COF-Bpy Synthesis.
32
An ampule was charged with TAPP (15.7 mg, 0.023 mmol) and 1 (8.6 mg,
0.041 mmol). A combination of o-dichlorobenzene, 1-butanol, and 6 M acetic acid (0.5:0.5:0.1
mL) was added, and the ampule was sonicated for 20-30 minutes. The mixture was degassed once
by freeze-pump-thaw method, and the ampoule was flame sealed before being placed in an oven
at 120°C for two days. The ampules were opened while still hot, and the contents were filtered and
washed briefly with acetone. The powder was washed via Soxhlet extraction with dioxane (24
hours) and acetone (24 hours). The powder was then washed 5 times with supercritical CO2 before
being placed under vacuum at 100° C for 18 hours.A purple solid was collected in 89% yield.

122

COF-Re Synthesis.
32
An ampule was charged with TAPP (30.2 mg, 0.045 mmol) and 2 (43.5 mg,
0.090 mmol). A combination of o-dichlorobenzene, 1-butanol, and 6 M acetic acid (1:1:0.2 mL)
was added, and the ampule was sonicated for 20-30 minutes. The mixture was degassed once by
freeze-pump-thaw method, and the ampoule was flame sealed before being placed in an oven at
120°C for two days. The ampules were opened while still hot, and the contents were filtered and
washed briefly with acetone. The powder was washed via Soxhlet extraction with dioxane (24
hours) and acetone (24 hours). The powder was then washed 5 times with supercritical CO2 before
being placed under vacuum at 100° C for 18 hours.A red-purple solid was collected in 85% yield.
Anal. Calcd for C74H42Cl2N12O6Re2: C, 54.24; H, 2.58; N, 10.26. Found: C, 52.01; H, 2.97; N,
9.48.

COF-Mix Synthesis.
32
An ampule was charged with TAPP (40.4 mg, 0.060 mmol), 1 (12.3 mg,
0.060 mmol), and 2 (22.8 mg, 0.050 mmol). A combination of o-dichlorobenzene, 1-butanol, and
6 M acetic acid (1:1:0.2 mL) was added, and the ampule was sonicated for 20-30 minutes. The
mixture was degassed once by freeze-pump-thaw method, and the ampoule was flame sealed
before being placed in an oven at 120°C for two days. The ampules were opened while still hot,
and the contents were filtered and washed briefly with acetone. The powder was washed via
Soxhlet extraction with dioxane (24 hours) and acetone (24 hours). The powder was then washed
5 times with supercritical CO2 before being placed under vacuum at 100° C for 18 hours. A red-
purple solid was collected in 83% yield.

Synthesis of COF-Re_Co.
35
COF-Re (50 mg) was added to 10 mL of DMF in a vial. A large
excess of cobalt (II) chloride hexahydrate (300 mg) was added to the vial, and the vial was heated
to 80 °C in an oven for 24 hours. The COF powder was collected via filtration and washed briefly
with acetone. The powder was washed via Soxhlet extraction with dimethylformamide (24 hours)
and acetone (24 hours). The powder was then washed 5 times with supercritical CO2 before being
placed under vacuum at 100° C for 18 hours. The product was a red-purple solid after workup and
was recovered in quantitative yields. ICP-OES: Rhenium = 17.34%, Cobalt = 1.47%. Anal. Calcd
for C74H 40Cl2CoN 12O6Re2: C, 52.42; H, 2.38; N, 9.91. Found: C, 52.34; H, 3.28; N, 9.37.

123

Synthesis of COF-Re_Fe.
35
COF-Re (50 mg) was added to 10 mL of DMF in a vial. A large
excess of iron (III) chloride (300 mg) or iron (II) chloride tetrahydrate (300 mg) was added to the
vial, and the vial was heated to 80 °C in an oven for 24 hours. The COF powder was collected via
filtration and washed briefly with acetone. The powder was washed via Soxhlet extraction with
dimethylformamide (24 hours) and acetone (24 hours). The powder was then washed 5 times with
supercritical CO2 before being placed under vacuum at 100° C for 18 hours. The product was a
red-purple solid after workup and was recovered in quantitative yields. ICP-OES: Rhenium =
19.17%, Iron = 0.88%. Anal. Calcd for C74H40Cl3FeN12O6Re2: C, 52.52; H, 2.38; N, 9.93. Found:
C, 58.15; H, 3.67; N, 10.93.

Scheme 4.2. Synthesis of 1 and 2.

124

Figure 4.13. 400 MHz
1
H NMR spectrum of 1 in DMSO-d 6.
125

Figure 4.14. 100 MHz
13
C{
1
H} NMR spectrum of 1 in DMSO-d 6.
126

Figure 4.15. 400 MHz
1
H NMR spectrum of 2 in DMSO-d 6.

127

Figure 4.16. 100 MHz
13
C{
1
H} NMR spectrum of 2 in DMSO-d 6.

128

Figure 4.17. FTIR spectrum of 1.

Figure 4.18. Variable scan rate of the first redox couple of 2. Scan rates are all mV/s.
129

Figure 4.19. Plot of current density vs square root of the scan rate for the first redox couple of 2. Current densities
were measured at -0.585 V and -0.542 V vs SCE.

Figure 4.20. Variable scan rate of both redox couple of 2. Scan rates are all mV/s.
130

Figure 4.21. Plot of current density vs square root of the scan rate for the second redox couple of 2. Current densities
were measured at -0.931 V and -0.869 V vs SCE.

Figure 4.22. Cyclic voltammograms of 0.5 mM 1 under N 2 and CO 2 environment in DMF with 0.1 M TBAPF 6.

131

Figure 4.23. Cyclic voltammograms of 1 and 2 (0.5 mM) under CO 2 in DMF with 0.1 M TBAPF 6.

Figure 4.24. Cyclic voltammograms of 1 (blue) and 2 (red) (0.5 mM) under N 2 in DMF with 4 M MeOH and 0.1 M
TBAPF 6.
132

Figure 4.25. Cyclic voltammograms of 2 (0.5 mM) under N 2 in DMF and 0.1 M TBAPF 6 with 0 M (purple) and 4 M
(green) MeOH.

Figure 4.26. Cyclic voltammogram of bare glassy carbon (black dashed), 1 (blue), and 2 (red) (0.5 mM) in CO 2
saturated DMF with 0.1 M TBAPF 6 and 4 M MeOH. Scan rate = 100 mV/s. Measured in an H cell. Working
electrode surface area = 8 cm
2
.

133

Figure 4.27. Cyclic voltammograms of 2 (0.5 mM) under CO 2 in saturated DMF with 0.1 M TBAPF 6 and 4 M
MeOH before CPE (green) and of rinsed electrode (3  10 mL DMF) in a fresh DMF solution containing 0.1 M
TBAPF 6 and 4 M MeOH under CO 2 (purple). Scan rate = 100 mV/s. Measured in an H cell. Working electrode
surface area = 8 cm
2
.

Figure 4.28. Controlled potential electrolyses of 1 (blue), 2 (red), and a blank solution (dashed black) at -2.0 V
versus SCE in DMF with 4 M MeOH under a N 2 atmosphere.

134

Figure 4.29. PXRD pattern of COF-Bpy.

Figure 4.30. PXRD pattern of COF-Re.

135

Figure 4.31. PXRD pattern of COF-Mix.

Figure 4.32. PXRD pattern of COF-Re_Co.

136

Figure 4.33. PXRD pattern of COF-Re_Fe.

Figure 4.34. N 1s XPS spectrum of COF-Bpy.

137

Figure 4.35. N 1s XPS spectrum of COF-Re.

Figure 4.36. XPS analyses of COF-Re_Co. (A) N 1s XPS spectrum. (B) Re 4f XPS spectrum.

Figure 4.37. XPS analyses of COF-Re_Fe. (A) N 1s XPS spectrum. (B) Re 4f XPS spectrum.
138

Figure 4.38. IR spectrum of COF-Bpy.

Figure 4.39. IR spectrum of COF-Re.

139

Figure 4.40. IR spectrum of COF-Re_Co.

Figure 4.41. IR spectrum of COF-Re_Fe.

140

Sample Surface Area (m
2
/g)
COF-Re 618.30
COF-Re_Co 615.97
COF-Re_Fe 723.12
Table 4.3. BET surface areas of the synthesized COFs.

Figure 4.42. N 2 sorption isotherm of COF-Re at 77K. Red = adsorption, Blue = desorption.

Figure 4.43. N 2 sorption isotherm of COF-Re_Co at 77K. Red = adsorption, Blue = desorption.

141

Figure 4.44. N 2 sorption isotherm of COF-Re_Fe at 77K. Red = adsorption, Blue = desorption.

Figure 4.45. TGA trace of the COF-Re activated sample under N 2.
142

Figure 4.46. TGA trace of the COF-Re_Co activated sample under N 2.

Figure 4.47. TGA trace of the COF-Re_Fe activated sample under N 2.

143

Table 4.4. Crystal data and structure refinement for 2.
Chemical formula C15H8ClN2O5Re
Formula weight 517.88 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.065 × 0.066 × 0.078 mm
Crystal habit clear orange-red prism
Crystal system monoclinic
Space group P 1 21/c 1
Unit cell dimensions a = 14.9714(17) Å α = 90°
b = 11.0724(13) Å β = 102.998(2)°
c = 9.5523(11) Å γ = 90°
Volume 1542.9(3) Å
3

Z 4
Density (calculated) 2.229 g/cm
3

Absorption coefficient 8.078 mm
-1

F(000) 976
Diffractometer Bruker APEX DUO
Radiation source fine-focus tube, MoKα
Theta range for data
collection
2.31 to 30.90°
Index ranges -20 ≤ h ≤ 21, -15 ≤ k ≤ 15, -13 ≤ l ≤ 13
Reflections collected 37876
Independent reflections 4773 [R(int) = 0.0520]
Coverage of independent
reflections
98.0%
Absorption correction multi-scan
144

Structure solution
technique
direct methods
Structure solution
program
SHELXTL XT 2013/1 (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(Fo
2
- Fc
2
)
2

Data / restraints /
parameters
4773 / 0 / 217
Goodness-of-fit on F2 1.040
Δ /σm ax 0.001
Final R indices 3611 data; I > 2σ(I)
R1 = 0.0347,
wR2 = 0.0745

all data
R1 = 0.0561,
wR2 = 0.0818

Weighting scheme
w = 1/[σ
2
(Fo
2
) + (0.0380P)
2
+ 4.0159P]
where P = (Fo
2
+2Fc
2
)/3

Largest diff. peak and hole 3.265 and -0.926 eÅ
-3

R.M.S. deviation from mean 0.176 eÅ
-3

Table 4.5. Bond lengths (Å) for 2.
C1-O1 1.155(6) C1-Re1 1.922(5)
C2-O2 1.131(7) C2-Re1 1.936(5)
C3-O3 1.129(6) C3-Re1 1.922(6)
C4-N1 1.341(6) C4-C5 1.409(8)
C4-H4 0.95 C5-C6 1.381(10)
C5-C14 1.502(8) C6-C7 1.369(8)
C6-H6 0.95 C7-C8 1.401(7)
145

C7-H7 0.95 C8-N1 1.373(6)
C8-C9 1.472(7) C9-N2 1.358(6)
C9-C10 1.396(7) C10-C11 1.385(8)
C10-H10 0.95 C11-C12 1.369(7)
C11-H11 0.95 C12-C13 1.401(7)
C12-C15 1.499(7) C13-N2 1.339(6)
C13-H13 0.95 C14-O4 1.201(8)
C14-H14 0.95 C15-O5 1.170(7)
C15-H15 0.95 Cl1-Re1 2.4948(14)
N1-Re1 2.172(4) N2-Re1 2.184(4)

Table 4.6. Bond angles (°) for 2.
O1-C1-Re1 177.4(5) O2-C2-Re1 177.3(4)
O3-C3-Re1 178.7(4) N1-C4-C5 121.1(5)
N1-C4-H4 119.4 C5-C4-H4 119.4
C6-C5-C4 119.3(5) C6-C5-C14 122.6(5)
C4-C5-C14 118.0(6) C7-C6-C5 119.7(5)
C7-C6-H6 120.2 C5-C6-H6 120.2
C6-C7-C8 119.6(5) C6-C7-H7 120.2
C8-C7-H7 120.2 N1-C8-C7 120.7(5)
N1-C8-C9 115.6(4) C7-C8-C9 123.7(5)
N2-C9-C10 120.8(5) N2-C9-C8 115.7(4)
C10-C9-C8 123.5(5) C11-C10-C9 119.3(5)
C11-C10-H10 120.3 C9-C10-H10 120.3
C12-C11-C10 120.1(5) C12-C11-H11 120.0
C10-C11-H11 120.0 C11-C12-C13 118.0(5)
C11-C12-C15 123.1(5) C13-C12-C15 118.9(5)
146

N2-C13-C12 122.8(5) N2-C13-H13 118.6
C12-C13-H13 118.6 O4-C14-C5 120.8(8)
O4-C14-H14 119.6 C5-C14-H14 119.6
O5-C15-C12 121.9(5) O5-C15-H15 119.0
C12-C15-H15 119.0 C4-N1-C8 119.4(4)
C4-N1-Re1 124.0(4) C8-N1-Re1 116.5(3)
C13-N2-C9 119.0(4) C13-N2-Re1 124.2(3)
C9-N2-Re1 116.8(3) C1-Re1-C3 87.1(2)
C1-Re1-C2 89.5(2) C3-Re1-C2 88.8(2)
C1-Re1-N1 99.21(18) C3-Re1-N1 93.79(19)
C2-Re1-N1 171.04(18) C1-Re1-N2 174.31(18)
C3-Re1-N2 92.52(18) C2-Re1-N2 96.17(17)
N1-Re1-N2 75.15(15) C1-Re1-Cl1 95.96(16)
C3-Re1-Cl1 175.85(15) C2-Re1-Cl1 94.09(19)
N1-Re1-Cl1 82.94(12) N2-Re1-Cl1 84.18(11)

147

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152

CHAPTER 5. A COBALT TRIPHENYLENEHEXATHIOLATE MOF FOR OXYGEN
REDUCTION

5.1 ABSTRACT
The attainment of cheap, robust materials capable of replacing platinum and platinum group metals
for the oxygen reduction half of hydrogen-based fuel cells is of growing interest. Triphenylene-
based MOF cobalt triphenylene-2,3,6,7,10,11-hexathiol (CoTHT) has previously been reported as
a hydrogen evolution catalyst, and other triphenylene-based MOFs have been reported for oxygen
reduction. The introduction of CoTHT to an O2-saturated KOH solution showed large current
enhancements relative to the same setup in the absence of O2. Weak physisorption resulted in loss
of MOF material during electrochemical experiments, so the use of a Nafion and carbon black ink
mixture was used to generate an electrically conductive, immobilized film. This ink mixture still
showed loss of current and an increase in onset with successive scans, which was determined to
be due to the irreversible oxidation of the MOF.

5.2 INTRODUCTION
As concerns over the continued use of fossil fuels as an energy source grow, alternative
and carbon-free methods of producing energy have become more attractive. One such method is
using hydrogen gas in fuel cells, where H2 is split into its component protons and electrons in order
to generate an electrical current.
1-5
While progress has been made in enhancing the efficiency of
this hydrogen oxidation step, the limiting factor of the overall process is the oxygen reduction
reaction (ORR) at the cathode.
1-5
The reduction of oxygen to water requires four electrons and four
protons as well as a potential of 1.23 V vs SHE, making oxygen reduction more
thermodynamically and kinetically demanding than the hydrogen oxidation counter reaction.
Current fuel cells perform both half reactions with platinum group metals (PGMs), making use of
their high efficiencies and low overpotentials.
6
However, PGMs are expensive and prone to
poisoning from the CO generated via steam reforming.
4
These high costs are exacerbated due to
ORR requiring more PGMs to competently perform as the counter to hydrogen oxidation.
4
To
reduce the high cost of fuel cells enough for them to serve as viable alternatives to fossil fuel
generators, cheap but efficient catalysts for ORR must be developed.
153

Metal organic frameworks (MOFs) are a class of materials known for their high surface
area, permanent porosity, and regular structure.
7-10
MOFs are extremely attractive as catalysts due
to their modular structure, which allows for molecular catalysts to be incorporated into a
heterogeneous framework.
7-10
Integrating homogeneous catalysts in this manner retains or
enhances efficiency while also imparting the stability and durability of heterogeneous catalysts.
7-
10
However, MOFs often suffer from low conductivity, resulting in poor charge transfer
capabilities which are prohibitive for catalysis.
11-13
As a result, many MOFs either rely on graphitic
additives to achieve the necessary conductivity for electrocatalysis
14-16
or are used as scaffolds and
pyrolyzed.
5, 6, 16

The use of pristine MOFs for ORR has been explored, but successful catalysts are few in
number.
16-18
A Cu-BTC (BTC = benzenetricarboxylate) based MOF was the first reported use of
a MOF for ORR, with an onset of potential of 0.46 V vs RHE.
19
The MOF lacked stability under
the aqueous conditions used for catalysis and decomposed rapidly, but still served as proof that
ORR with pure MOFs was feasible. An iron porphyrin-based MOF, PCN-223-Fe, was reported to
reduce oxygen in organic media at -0.3 V vs NHE.
20
Water was the main reduction product, with
the formation of the two electron peroxide product decreasing to less than 6% as the potential was
made more negative at -0.6 V vs NHE and with acetic acid as a proton source.
20
A nickel
hexaiminotriphenylene (NiHITP) MOF performs ORR at 0.82 V vs RHE in pH 13 aqueous
solution, which corresponds to a 0.18 V overpotential compared to platinum (1.00 V vs RHE).
21

Ni3HITP2 mainly produced the two electron peroxide product with a fadaic efficiency of 100% at
the onset and 63% at more negative potentials as H2O becomes more favorable.
21
The MOF was
stable under basic conditions for over 7 hours, though was less stable under acidic conditions.
21

Further studies on the NiHITP MOF revealed that the active site for catalysis was not on
the metal center, but rather on the β-carbon of the MOF.
22
The ability of the MOF to delocalize
charge was determined to be a critical component as the molecular version of the catalyst
performed poorly.
22
In situ XANES showed no change in the Ni oxidation state during catalysis,
and DFT studies showed the β-carbon was the most favorable binding site.
22
Subsequent studies
sought to tune the metal center and chelating atoms, resulting in CoHHTP (HHTP =
hexahydroxytriphenylene), NiHHTP, CuHHTP, and CuHITP structural analogues being tested for
ORR.
23
CuHITP was the most active catalyst, generating twice as much current density as NiHITP,
but rapidly deactivated with successive scans.
23
The NiHHTP was less active and deactivated more
154

quickly than NiHITP, which was attributed to better overlap between the imine moieties and the
metal center.
23
All MOFs had similar Tafel slopes with the exception of CoHHTP, and all were
first order with respect to O2 which suggests that changes in the metal center and chelating atoms
do not largely affect the mechanism of oxygen reduction.
23
The success of these MOFs is further
supported by other groups.
19, 24, 25

Our group has previously reported an electrically conductive cobalt
triphenylenehexathiolate (CoTHT) MOF, which is structurally analogous to the NiHITP system.
26,
27
Initial studies have shown that CoTHT is a competent electrocatalyst for the hydrogen evolution
reaction,
26
with other triphenylene MOFs also proving to be potent electrocatalysts.
28, 29

Additionally, CoTHT displays unique conductivity results, transitioning from a semiconductor to
a metal at low temperatures.
27
Herein, this work aims to utilize CoTHT as an electrocatalyst for
the oxygen reduction reaction. Current responses under both N2 and O2 atmospheres are observed,
and the stability of the MOF species is tested. Various deposition methods are explored in order to
optimize catalytic performance. The practicality of the use of CoTHT for ORR is discussed in
regards to the results.

5.3 RESULTS AND DISCUSSION
5.3.1 Cyclic Voltammetry of CoTHT
CoTHT was prepared as a film and characterized as previously reported (Scheme 5.1).
26, 27

The glassy carbon disk working electrode was pressed through the film with the working face
facing upwards (Scheme 5.2). The electrode was then raised through the film such that a portion
of the film lay evenly across the face of the electrode. Excess solvent was allowed to evaporate
before use. The electrode was then placed in a pH 13 (0.1 M KOH) solution.

155

Scheme 5.1. Synthesis of CoTHT from cobalt and triphenylene hexathiol.

Scheme 5.2. Deposition of CoTHT film onto glassy carbon through physical adsorption.

Cyclic voltammetry (CV) was conducted under an atmosphere of N2 with the glassy carbon
insert placed into a rotating ring disk set up. Ag/AgCl was used as the reference electrode and Pt
wire was used as the counter. Little electrochemical response was seen for a range of 1 to 0.2 V vs
RHE (pH = 13) (Figure 5.1). When the gaseous atmosphere was switched to O2, a large current
enhancement was seen at 0.8 V vs RHE. Onsets at this potential are similar to onsets seem for the
NiHITP system.
21

156

Figure 5.1. CV of CoTHT in pH 13 aqueous solution under N 2 (blue) and O 2 (red). Scan rate = 10 mV/s. Rotation
rate = static.

CVs of the bare glassy carbon electrode, dropcasted triphenylenehexathiolate ligand, and
dissolved cobalt(II) chloride were shown to generate lower amounts of current density than the
CoTHT MOF (Figure 5.2). They also showed more negative onsets, demonstrating that the MOF
is the cause of the current increase at 0.8 V vs RHE and not some trace impurity or exposed glassy
carbon electrode.

Figure 5.2. CVs of the bare electrode (black dashed), CoCl 2 metal precursor (blue), THT ligand precursor (green),
and CoTHT MOF (red) in pH 13 solution under O 2. Scan rate = 10 mV/s. Rotation rate = static.

In order to better evaluate the electrochemical behavior, the working electrode was rotated
at rates ranging from 100 rpm to 2500 rpm. Under these conditions, the diffusion of substrate to
157

the electrode is no longer limiting which allows for determination of the number of electrons
passed through the use of the Koutecky-Levich equation.
21
The current increased with increasing
rotation rate (Figure 5.3), but the onset of catalysis shifted more negative before stopping at
roughly the same onset as the glassy carbon electrode as seen in Figure 5.2. Throughout the
experiment, black flakes could be seen floating in the solution, and analysis of the working
electrode after the CVs showed a noticeable loss of MOF material from the electrode surface. The
physisorption of the MOF material to the electrode surface is likely too weak to keep the MOF
attached throughout the CVs, especially at higher rotation rates. In order to maintain the MOF-
modified surface, the use of Nafion as a polymeric binder was investigated as success had been
seen for other MOFs.
23

Figure 5.3. CVs of CoTHT at various rotation rates ranging from 0 (red) to 2500 (purple) rpm in pH 13 solution
under O 2. Scan rate = 10 mV/s.

5.3.2 Cyclic Voltammetry of CoTHT-Nafion-Carbon Black Ink Mixture
To better immobilize the MOF onto the glassy carbon surface, the material was sonicated
with Nafion, ethanol, and water according to previously reported procedures.
30
The material was
then dropcast onto the glassy carbon and allowed to dry before use. CVs of the Nafion-bound
catalyst showed worse performance than the MOF by itself, as demonstrated by a decrease in
current density and an increase in overpotential (Figure 5.4 A). Given that many MOFs have poor
electrical conductivity even before being broken into fragments due to sonication and that Nafion
can be insulating, charge transfer may be limiting in this mixture. To circumvent this issue, carbon
158

black was added to the mixture as a conductive support to increase electron transport. With the
addition of carbon black, the currents achieved increased substantially and the onset was returned
to where it was previously (Figure 5.4 B). While the addition of Nafion and carbon black
strengthened the electrode-MOF surface interaction and enhanced charge transport, the current
decreased with successive scans. The loss of material from the electrode surface has been mitigated
and therefore does not explain this drastic loss in current.

Figure 5.4. CVs of CoTHT films prepared with Nafion (A) and Nafion and carbon black (B) at pH 13 under O 2.
Scan rate = 10 mV/s. Rotation rate = static.

The stability of the MOF under such basic conditions was questioned, thus weaker base
solution (pH 10, Figure 5.10) was tested. The MOF far outperformed the bare glassy carbon
electrode and carbon black ink mixture alone under these conditions, proving that the MOF was
still active even at lower pHs. However, subsequent scans again resulted in a loss of activity
(Figure 5.11). As the catalyst is known to be stable under acidic conditions from the studies
performed for HER, acidic pH solutions of 0.28 (Figures 5.12 and 5.13) and 1.16 (Figures 5.14
and 5.15) were tested as well. A significant decrease was seen in catalyst activity as the pH became
more acidic. These studies failed to show any improvement in the stability of the MOF, suggesting
that the MOF may be deactivating due to the O2 atmosphere.
To further confirm the deactivation of the MOF ink mixture, 25 consecutive CV sweeps were
conducted under O2. As the scans progressed, a steady decrease in current and increase in
overpotential could be seen (Figure 5.5). No material was seen floating in the solution, which
suggests some transformation or decomposition must be occurring within the ink mixture. Further
structural characterization experiments were conducted to analyze the MOF material post-
electrochemistry and compare to the structural data from before electrochemical experiments.
A B
159

Figure 5.5. 25 consecutive CVs of CoTHT ink mixture under O 2 in pH 13 KOH solution, where the current
decreases and the overpotential increases after each scan. Rotation rate = 1600 rpm. Scan rate = 10 mV/s.

5.3.3 X-Ray Photoelectron Spectroscopy of CoTHT-Nafion-Carbon Black Ink Mixture
Drastic changes could be seen in the X-ray photoelectron spectroscopy (XPS) spectra from
before CVs to after. The sulfur 2s and 2p (Figure 5.6 B and D) and cobalt 2p (Figure 5.7 B) regions
of the ink mixture used in the 25 scan experiment were analyzed. Both sulfur regions show the
most distinct change, with features growing in at higher binding energies. While these features are
somewhat present in the 2s and 2p regions before electrolysis (Figure 5.6 A and C), they come to
dominate the spectra following CV experiments. The growth of these features is consistent with
the oxidation of the sulfur environments as seen for FeTHT materials.
31
Changes are also seen in
the cobalt region, where the before scan seems to show two major features with shoulder peaks
(Figure 5.7 A) and the after scan is closer to just two main peaks (Figure 5.7 B). The consistency
with the oxidized FeTHT suggests that CoTHT may be oxidized at the sulfur and possibly cobalt
sites.

160

Figure 5.6. XPS spectra of the S 2s and S 2p regions before (A and C) and after (B and D) CVs.

Figure 5.7. XPS spectra of the Co 2p region before (A) and after (B) CVs.

To test the hypothesis that CoTHT may be oxidizing under O2, the MOF was exposed to
air for 10 straight days and the XPS spectra for Co and S were obtained (Figure 5.8). While the Co
2p region (Figure 5.8 A) resembles the Co spectrum before CV, the S 2s and 2p regions (Figure
5.8 B and C) are much more consistent with the post-electrochemistry spectra. This evidence
further supports that the CoTHT MOF is oxidizing during the CV experiments. This behavior is
unexpected as the NiHITP MOF and its family of triphenylene compounds are all predicted by
DFT to undergo O2 insertion at the β-carbon of the triphenylene, which would leave the metal
(A)
(B)
(C)
(D)
(A) (B)
161

center and binding atom (Co and S in our case) unaffected. If the Co and S sites are oxidizing, this
may suggest that the active site of catalysis in our case is Co or S, but that turnover is hindered
which results in irreversible oxidation.

Figure 5.8. Co 2p (a), S 2s (b), and S 2p (c) XPS spectra of CoTHT following 10 days of O 2 exposure.

5.3.4 Cyclic Voltammetry of the Oxidized CoTHT-Nafion-Carbon Black Ink Mixture
To further confirm that the oxidation of the CoTHT MOF is the reason behind the loss in
activity, the air-oxidized MOF was used to prepare a similar ink mixture as before. Again, 25
consecutive CV scans were performed under O2. The oxidized MOF ink showed much more
consistency scan to scan, with the onset taking place at more negative potentials than the
unexposed CoTHT (Figure 5.9). The current was also lower, though not as much as the fresh ink
at the end of the 25 scans. This may be due to the increased alteration of the Co sites in the fresh
ink post CV as seen by XPS being related to a co-oxidation with the S, whereas the preoxidized
MOF is incapable of this due to the sulfurs already being heavily oxidized. Further studies would
be needed to better understand this behavior.

Figure 5.9. 25 consecutive CVs of CoTHT ink prepared air free (red) and pre-oxidized CoTHT ink (black) in pH 13
solution under O 2. Scan rate = 10 mV/s. Rotation rate = 1600 rpm.
162

5.3.5 Theoretical ORR Activity of CoTHT
The failure of CoTHT to effectively reduce O2 is at odds with many theoretical predictions
stating that CoTHT and other similar Co-based systems should be excellent catalysts for ORR.
32-
34
These predictions screened transition metals from groups 8, 9, and 10 for their ORR and OER
capabilities. In these cases, Co showed ideal properties with O2 binding strongly enough for
catalysis to occur but not so strongly that the catalyst would become poisoned. This was in
combination with the low predicted overpotential for ORR predicted for CoTHT (0.30 V) which
is similar to the experimentally observed overpotential. Of note is the prediction for the reaction
mechanism; the initial O2 binding is predicted to occur at the cobalt metal center, and following
intermediates OOH and OH are also localized on the Co.
33
However, the O intermediate is
predicted to be shared across the Co-S bond.
33
This could explain the experimental behavior if the
O was then transferred to the S atom and could not be released. The sulfur would be oxidized, and
O2 reduction would be slowed or stopped as the sharing of the O between Co and S would be more
difficult if S has no free orbitals. The O would then be stuck on Co, explaining the possible
oxidation seen for the cobalt site. Therefore, it is likely that the catalyst can initially perform ORR
but the sulfur atom becomes irreversibly oxidized, preventing further turnovers.

163

Scheme 5.3. Proposed mechanism for the reaction between CoTHT (truncated for clarity) and O 2 based on DFT
calculations.

5.4 CONCLUSION
CoTHT, a successful hydrogen evolution catalyst, was explored for its use as an oxygen
reduction catalyst. While current increases were seen in the presence of O2 by CV, the weak
physisorption of the catalyst required enhanced immobilization via Nafion which further required
the addition of carbon black as a conductive support. The ink mixture was better able to keep the
material on the electrode, but the loss of current and increase in the overpotential was still present.
XPS studies revealed that the Co and S atoms were undergoing a change in their oxidation state
during electrochemical experiments, with that change seemingly being oxidation. XPS of a
purposefully oxidized CoTHT supported this theory. CVs of an ink mixture containing the
purposefully oxidized MOF showed CV character similar to the fresh MOF following multiple
164

consecutive CVs. While previously reported triphenyle-based MOFs for oxygen reduction have
proposed a β-carbon active site with no metal or chelating atom involvement, M3THT2 MOFs are
predicted to have largely metal-based redox chemistry. It is believed that the cooperativity between
the Co and S atoms is what leads to catalyst deactivation, and thus different metals may need to be
used to optimize oxygen reduction with THT.

5.5 ACKNOWLEDGEMENTS
The authors are grateful to the University of Southern California (USC) for funding, the USC
Wrigley Institute for the Norma and Jerol Sonosky summer fellowship to EMJ, and to Thomas
Moulton and Ginny Dunn for the Harold and Lillian Moulton Graduate Fellowship to EMJ. The
studies of the extended frameworks were supported by the Nanoporous Materials Genome Center
of the U.S. Department of Energy, Office of Basic Energy Sciences, Division of Chemical
Sciences, Geosciences and Biosciences under Award DE-FG02-17ER16362. The authors are
grateful to NSF (grant CRIF 1048807) and USC for their sponsorship of NMR spectrometers and
X-ray diffractometer. The authors thank Keying Chen for assistance with sample preparation and
Andrew Clough for assistance with XPS. XPS data were collected at the Center for Electron
Microscopy and Microanalysis (CEMMA) at USC. The authors would like to thank Dr Courtney
Downes for discussion.

165

5.6 EXPERIMENTAL METHODS
General. 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). All the solvents used were degassed under vacuum and refilled with nitrogen (10 ×).
Triphenylene-2,3,6,7,10,11-hexathiol ligand (THT) was synthesized according to literature
procedure.
35
All other chemical reagents were purchased from commercial vendors and used
without further purification.

5.6.1 Electrochemical Methods
Electrochemistry experiments were carried out using a Pine WaveDriver 40 potentiostat in a three
electrode configuration electrochemical cell under an inert atmosphere. A glassy carbon electrode
(GCE, 0.07065 cm
2
surface area) or a rotating disk electrode (RDE, glassy carbon insert, 0.196
cm
2
surface area) was used as the working electrode. GCE and RDE were polished with 0.05 μm
Al2O3 polish powder and sonicated in water prior to use. A graphite rod, purchased from Graphite
Machining, Inc. (Grade NAC-500 Purified, < 10 ppm ash level), was used as the counter electrode.
The reference electrode, placed in a separate compartment and connected by a Vycor tip, was
based on an aqueous Ag/AgCl/saturated 3.5 M 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.

Rotating disk and rotating ring disk electrode experiments were conducted using a Pine
WaveVortex 10 rotator.

5.6.2 Physical Characterization Methods
X-ray photoelectron spectroscopy (XPS) data were collected using a Kratos AXIS Ultra
instrument. The monochromatic X-ray source was the Al K α line at 1486.7 eV, and the hybrid
lens and slot mode were used. 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
166

collected on individual XPS regions of interest. The sample chamber was maintained at < 9× 10
−9

Torr. The XPS data were analyzed using the CasaXPS software.

High resolution synchrotron powder X-ray diffraction data was collected using the 11-BM
beamline mail-in program at the Advanced Photon Source (APS), Argonne National Laboratory,
with an average wavelength of 0.412750 Å. Discrete detectors covering an angular range from 0.5
to 30
o
2θ are scanned over a 34
o
2θ range, with data points collected every 0.001
o
2θ and scan
speed of 0.01
o
/s.

5.6.3 Koutecký-Levich Equation
21
The number of electrons passed at a given potential can be determined by plotting 1/current density
vs 1/√rpm
B = n ∙ 0.62FDO2
2/3
ν
-1/6
cO2(
2π
/60)
1/2

Where,
B = 1/slope of 1/Current density vs 1/√rpm
n = number of electrons
F = Faraday constant
DO2 = O2 diffusion coefficient in the electrolyte
ν = kinematic viscosity of the electrolyte
cO2 = saturation concentration of O2 in the electrolyte at 1 atm O2 pressure

B and n are variables and the rest are constants
F = 96,485 C
DO2 = 1.9 ∙ 10
-5
cm/s
ν = 0.1 m
2
/s
cO2 = 1.26 ∙ 10
-6
mol/cm
3

The equation can therefore be condensed to:
B = n ∙ 2.55 × 10
-5

167

5.7 SYNTHETIC METHODS
5.7.1 Synthesis of CoTHT
The synthesis of CoTHT was previously reported by our laboratory.
26, 27
A 120 mL jar was
charged with a solution of CoCl2∙6H2O (40.0 mg, 0.168 mmol) in water (40 mL). Separately, a
suspension of THT (2.5 mg, 0.006 mmol) in N-methyl-2-pyrrolidone (NMP) (0.1 mL) was diluted
with ethyl acetate until the total volume of the suspension reached 5 mL, sealed, and briefly
sonicated to form a uniform suspension. Ethyl acetate (35 mL) was gently layered on top of the
aqueous solution to create a liquid-liquid interface; the suspension of THT was then gently added
to the ethyl acetate layer and the jar was sealed and allowed to stand still. A black film appeared
at the liquid-liquid interface over 5 days, which was then collected as powder, solvent exchanged
with methanol (3 × 20 mL), and dried under vacuum.
5.7.2 Deposition of CoTHT for Electrochemical Study
Preparation of ink composite (1): 2 mg of CoTHT was mixed with a desired amount of carbon
black (Vulcan XC-72R), 20 μL of Nafion solution (0.5 wt%), 45 μL of water, and 135 μL of
ethanol, followed by sonication for 1 hour to form a uniformly dispersed suspension. A desired
amount of suspension was then drop casted onto a glassy carbon electrode (GCE) or a rotating disk
electrode (RDE, glassy carbon insert) using a microsyringe, and dried in nitrogen atmosphere at
room temperature.

Direct deposition of CoTHT: Deposition was carried out by drop casting the film formed at the
liquid-liquid interface to the electrode substrate using a glass pipette. Following deposition, the
electrode was washed with water and methanol.

Preparation of CoTHT/Nafion mixture: 2 mg of CoTHT was mixed with 20 μL of Nafion solution
(0.5 wt%), 45 μL of water, and 135 μL of ethanol, followed by sonication for 1 hour to form a
uniformly dispersed suspension. 10 μL of suspension was then drop casted onto a GCE and dried
in nitrogen atmosphere at room temperature.

168

5.8 ADDITIONAL FIGURES

Figure 5.10. CV of CoTHT ink (red), bare glassy carbon electrode (black dashed), and carbon black ink mixture
(purple dotted) in pH 10 aqueous solution under O 2 (red). Scan rate = 10 mV/s. Rotation rate = static.

Figure 5.11. Subsequent CVs of CoTHT ink in pH 10 aqueous solution under O 2. Scan rate = 10 mV/s. Rotation
rate = 1600 rpm.
169

Figure 5.12. CV of glassy carbon in pH 0.28 aqueous solution under N 2 (blue) and O 2 (red). Scan rate = 10 mV/s.
Rotation rate = 1600 rpm.

Figure 5.13. CV of CoTHT in pH 0.28 aqueous solution under N 2 (blue) and O 2 (red). Scan rate = 10 mV/s.
Rotation rate = 1600 rpm.
170

Figure 5.14. CV of glassy carbon in pH 1.16 aqueous solution under N 2 (blue) and O 2 (red). Scan rate = 10 mV/s.
Rotation rate = 1600 rpm.

Figure 5.15. CV of CoTHT in pH 1.16 aqueous solution under N 2 (blue) and O 2 (red). Scan rate = 10 mV/s.
Rotation rate = 1600 rpm.

171

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CHAPTER 6. FUTURE DIRECTIONS AND OUTLOOK

The continuation of these projects is important for providing starting points for future
graduate students. The ideation of the future directions of these projects should therefore be the
purview of the graduating student. In this section, the previously discussed projects will be
considered through the lens of not only what has been done but what can reasonably be
accomplished in the future with the information currently available.

6.1 Lariat Macrocycle for Solar Energy Storage
The NO2 lariat has produced unexpected reactivity for the macrocycle which should be
explored more in depth. The shift in selectivity from solely CO is a result of the formation of a
cobalt hydride, but experimental proof (i.e. a crystal structure of the hydride) for the formation of
the hydride and on which face of the cobalt metal center it forms is still lacking. While it is most
likely that the hydride forms opposite the lariat arm, resembling pentapyridyl HER catalysts, it is
theoretically possible that the nitro group could stabilize the hydride, and the Co-O distance is
sufficiently long to fit a hydride. Density functional theory should also be employed to predict and
explain the behavior of CoL
NO2
under catalytic conditions. These data should then be supported
with stoichiometric transformations for the purpose of isolating intermediates in order to support
a mechanism for product formation. Previous attempts using potassium graphite (KC8) have shown
some promise with the 4 peaks in the paramagnetic region in deuterated pyridine becoming more
shielded, but more easily controlled reducing agents such as sodium naphthalide should be
explored in order to better manipulate the reduction process.
The production of formate is valuable and should be explored in more detail in order to
maximize the amount of formate produced. Weaker proton sources such as isopropanol may be
necessary to increase the selectivity for formate as strong acids will favor CO formation and HER.
Solvent choice may also be an important consideration as has already been seen in the case of
DMF. More non-conventional solvents may be explored in order to test if they influence formate
production. Trace amounts of methane have been detected in the presence of stronger acids (phenol
and large excess of water). The formation of methane through electroreduction of CO 2 is rare, so
if this process could be better understood and optimized, it may lead to interesting chemistry.
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The original intent of the lariat cycle was to explore the effect of the positioning of NH
moieties on the cycle reactivity as the parent cycle amines are not direct proton donors but rather
more like docks and guides for proton sources in solution. The NO2 lariat was pursued as it was
air stable, while the NH2 version of the lariat oxidized back to the NO2 lariat in air. As the
understanding of this project has grown, it may now be possible to isolate a metallated NH2 lariat.
This process will likely involve purification in the glovebox. Following the reduction of the NO2
with H2 and Pd/C, the solution can be filtered quickly, poured back into a Schlenk flask, and
concentrated before being brought in the box as the reported synthesis of the NH2 species suggests
that oxidation of the ligand is not immediate. Alternatively, the reaction could be worked up
entirely in the box. Filtering through celite to remove the Pd/C can easily be accomplished in the
box, and removing the solvent is possible though likely time consuming. While much of the
workups proposed in the paper suggest column chromatography, recent results show that silica
plugs are sufficient for purification by washing out impurities with dichloromethane before
removing the purified product with ethyl acetate. Incorporation of a metal center may be sufficient
to protect the complex from oxidation, but if not, electrochemical experiments may be set up or
performed in the glove boxes.
Along with the phenyl linker, a benzyl version of the lariat chain was initially explored.
This project was initially favored because the NH2 was air stable, but crystallizing the metal
complex was difficult and often failed to provide pure product. However, through studying the
phenyl NO2 lariat, techniques that were optimized for the phenyl NO2 worked well for the benzyl
NH2, also. The benzyl lariat can be purified through the same silica plug procedures mentioned
above, and metal complexes formed with Co and Ni have been crystallized from vapor diffusion
of ether into a mixture of acetonitrile and chloroform, though X-ray structures have never been
generated. In addition to having NH2 groups that can possibly participate as secondary
coordination sphere H donors, Chris Chang has shown that the positioning of the proton donor can
have a significant influence on catalysis. A family of lariats can be envisioned with the number of
carbons between the lariat phenyl ring and macrocycle ranging from 0 to at least 2, with each
potentially having different capabilities for CO2 reduction. The number of lariat chains on a
particular macrocycle can also be manipulated through the use of an appropriate parent
macrocycle. While it may not be possible to form a tetra-substituted macrocycle, a cycle with two
lariat arms may be possible. It may also be possible to form a macrocycle containing one lariat
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chain and unreacted NH groups. This could be useful for tuning selectivity, as the presence of the
secondary amines stabilizes a CO2 adduct but the lariat seems to shift favorability to metal
hydrides. This combination could help maintain the selectivity of CO2 to CO reduction but alter
other aspects of catalysis.

Scheme 6.1. Synthesis of longer lariat arm macrocycles which are more air stable.

If the presence of NH2 groups can be shown to have reactivity that favors the formation of
CO2 reduction products, different hydrogen bonding/donating substituents can be introduced as
lariat chains. Hydroxyl groups have been heavily studied in the literature and could be installed
here through the use of the appropriate lariat chain. Since the attachment of the lariat to the
macrocycle involves the use of base, the hydroxyl will likely need to be protected during this step.
Methoxy groups may also be installed through a similar method in order to give a better
understanding of the effect of direct proton donors in this position as methoxy groups can behave
as indirect donors, engaging in H-bonding interactions with acid species in solution. This approach
would be more similar to the original macrocycles, but with indirect donors in different positions
relative to the metal center and CO2 adduct. To this effect, the NH2 species can also be methylated
with methyl iodide to form and NMe2 or NHMe species to study the effect methylated indirect
donors have relative to protonated direct donors with both chalcogens and pnictogens.
Furthermore, the NH2 can be tuned through reactions with aldehydes or esters to form more
complex architectures. The use of row 3 moieties such as thiols and phosphines may also be
synthesized, and may even impart different binding.
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Figure 6.1. Different functional groups that can be explored as direct proton relays to bound CO 2.

The addition of the lariat may also be used in other ways; a moiety which can attach to a
surface may be installed in order to immobilize the macrocycle onto an electrode surface.
Heterogenizing species can improve their overall catalytic lifetime and performance. However,
previous methods of heterogenizing the macrocycle were of unsuccessful or otherwise
underexplored. A diazonium can be formed with the linker via an NH2, which our group has
successfully accomplished previously with rhenium bipyridine systems. This will likely require a
lariat with the NH2 in the para position, which has not been explored but the synthetic conditions
should be largely similar to those for the ortho substituted lariats. The question will remain as to
whether or not the diazonium synthesis conditions will perform side reactions. If the macrocycle
is unmetallated, the exposed pyridines may react with the NOBF4, which could then hinder metal
coordination. Alternatively, metallating before performing the NOBF4 reaction could protect the
pyridines but the metal could then undergo irreversible reactions. There is also the potential for
the amines to be reactive which may decompose the cycle. A method that may be useful could be
the attachment of a para substituted lariat arm to the surface before attaching it to the macrocycle.
By attaching the arm first, a macrocycle with at least one secondary amine should be able to bind
covalently to the modified surface provided the lariat arm maintains a reactive leaving group. If
diazonium synthesis is not possible, other methods of covalent surface attachment may be
explored.

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Scheme 6.2. Possible general methods for covalently attaching the macrocycle to an electrode surface via diazonium
chemistry.

With the development of the parent cycle alongside the lariat cycle, revisiting approaches
that were explored for the parent may be more fruitful with the lariat. The use of Fe and Ni metal
centers proved to be a poor choice for the parent macrocycles. The resulting molecular orbitals for
these complexes as determined by DFT show that the cobalt species has the most favorable energy
map for reducing CO2, while Fe and Ni have unfavorable orbital energies and rapidly decompose.
As the lariat seems to have different selectivity than the parent complex, the ligand orbitals likely
have different energies that may better overlap with Fe or Ni. DFT predictions of the orbital energy
levels should be run beforehand to check the best option for a metal center, but if a different metal
is suggested to be better than Co, the complex could easily be made and tested.

179

6.2 Decarbonylation Reactions using Lariat
The production of CO under an inert atmosphere in DMF solution was small but
significant. While this result was abandoned in favor of performing electrochemistry in
acetonitrile, there is still value in exploring this further. The first experimental approach should be
to determine if this procedure can be accomplished stoichiometrically. The complex can be
reduced with an excess of KC8 in a DMF solution, and a proton source such as TFE or phenol can
be added to see if CO is generated. Through this method, an intermediate may also be trapped to
observe if DMF is coordinating to the metal complex during the reduction process. Performing the
reduction in the absence of a proton source may help this.
If the decarbonylation of DMF is confirmed, different carbonyl containing species should
be explored. DMF is an amide species which is a decent foundation for screening activity. Included
in the class of amide containing species are amino acids, and the cleavage of these bonds may be
beneficial for biological purposes. In addition to amides, there are aldehyde, ketone, and anhydride
species which may be cleaved with this complex. A large screening effort will need to be
undertaken to determine what will be the best species to use for cleavage.
Perhaps the most interesting avenue to explore would be the dehydrogenation of formic
acid to generate CO and H2. The catalyst is known to reduce protons to H2, so if formic acid could
be cleaved, it could create an energetic carbon cycle where CO is reduced to formate/formic acid
which can then be cleaved later to produce H2 as fuel. This idea is stringent on the ability of the
catalyst to decarbonylate formic acid. This can be tested by added formate to a solution of catalyst
in acetonitrile with reducing species such as KC8. Formic acid may be considered as well but the
acid strength of formic acid may be strong enough to promote HER, which might complicate the
interpretation of results. Therefore, formic acid should be explored first and tested for the
production of both H2 and CO. If this occurs, then formic acid should also be tested.

6.3 Bimetallic COF
The first method that can be attempted to improve the bimetallic COFs does not involve
the COF itself but rather the equipment and methods used for its synthesis and purification. For
the initial synthesis of the porphyrin unit, Soxhlet extractors should be used for the purification as
reported by White. After metallating the porphyrin, the material should be recrystallized instead
of simply washed. Through this approach, more pure porphyrins may be used in synthesizing the
180

COF which may improve crystallinity. Due to the low solubility of the metalloporphyrins, this
may not be feasible. In this case, other purification methods such as Soxhlet extractor washing
may be useful to obtain the purest product possible. The synthesis of the COF itself can be
performed under better conditions as well. The ovens used possibly do not reach nor maintain the
proper temperature, and there is no method by which to monitor the temperature over the course
of the two to three day reaction. This may be attributing to the weak crystallinity, where other
groups report much sharper XRD peaks for similar structures. A more idealized setup would
include an oven with digital temperature regulation that is also not general use. The use of an oil
bath instead of an oven gives nearly the same intensity peaks by PXRD, suggesting that the current
oven use is not ideal. The opening of ampules can potentially also be improved by using better
glass cutting equipment. The sealed ampules can take up to 10 minutes to open, with the reverse
reaction becoming more and more favored as the temperature lowers. A true glass file would make
opening the ampules much faster, once again potentially resulting in more crystalline material.
Finally, the last step of the purification of the frameworks involves washing with CO 2, but the
equipment used is not a true supercritical CO2 setup. While this may not influence much, it can at
least be attempted in order to have more active surface area available within the COF and therefore
more metal sites.
The design of the COFs may be modified through more complex synthetic approaches. The
first change should be to determine the possibility of reducing the number of rhenium bipyridine
sites in the COF. The bulk of the rhenium centers may be hindering mass transport through the
framework, and reducing the number of sites may promote higher activity. It has been shown that
synthesis containing a 1:1 ratio of bipyridine to rhenium bipyridine forms a crystalline framework,
but it is unknown if the final product maintains this ratio. Given the differences in reactivity of the
aldehyde groups as noted by the crystallinity of the purely bipyridine or purely rhenium bipyridine
COFs, rhenium bipyridine linkers will likely react first and may possibly crash smaller portions of
the 2D framework out of solution before much of the bipyridine has time to react. This issue is
compounded by the fact that the bipyridine would not be the desired linker, but rather the biphenyl
variant to avoid issues related to post metallation when the bipyridine species are unmetallated.
While the bipyridine-based COF typically displayed weak diffraction peaks, the biphenyl COF
often was worse despite it reportedly having high crystallinity in the work of Yaghi and coworkers.
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With an optimized experimental setup, it may be possible to form mixed biphenyl/rhenium
bipyridine COFs to form optimized frameworks for CO2 reduction.
In addition to optimized electrocatalysis, photocatalysis may be a route to explore for these
frameworks. Rhenium bipyridine is known to be a single component photocatalyst, and porphyrins
have been used in photocatalysis as well. The combination of these two may act as a highly
photosensitive material. MOFs containing rhenium bipyridine moieties have been successfully
used for photocatalytic CO2 reduction, therefore COFs behaving as photocatalysts can be inferred.
The COF can be suspended in a solution containing the appropriate components for photocatalysis
(e.g. sacrificial donor, CO2, etc.). The optimized catalyst for electrocatalysis may not be the
optimal catalyst for photocatalysis, thus a screening of conditions may be necessary for ideal
catalytic performance. Furthermore, the substitution of photosensitizers based on ruthenium and
iridium may be possible, particularly if the mixtures of rhenium bipyridine and biphenyl can be
determined.

Figure 6.2. General photocatalysis setup for COFs using a sacrificial donor species and lamp illumination.

While the initial concept for this project was purely CO2 reduction, other types of catalysis
may be possible with these COFs. It has already been reported that oxygen reduction is possible
with porphyrin-based COFs, thus this framework may also be applicable. The use of bipyridine-
based frameworks for various types of catalysis, such as water oxidation, has also been reported.
Modifications to the metal centers of both the bipyridine and the porphyrin will allow for access
to a wide range of catalysis including oxygen evolution, oxygen reduction, hydrogen evolution,
hydrogen oxidation, CO2 reduction, nitrogen reduction, and ammonia oxidation. The combination
182

of two metal centers could provide a pathway to a COF that can do more complex reactions. For
example, a porphyrin species known for hydrogen evolution combined with a bipyrdine complex
capable of water oxidation could be used on the cathode and anode for complete water splitting.
Cascade catalysis is another possible route to explore, relying on one metal center to perform the
first part of a reaction and using the second metal to perform the following step. Through this
method, complex reactions such as CO2 to methanol can be attempted.

6.4 Cobalt Triphenylenehexathiolate MOF
Given that the MOF oxidizes under the conditions necessary to perform ORR, there is not
much left that can be accomplished with this material. The biggest consideration for the future of
this project is that experimental evidence contradicts previous theoretical predictions. New DFT
should be performed to better understand how and why the material is deactivated, as well as what
steps forward can be undertaken to prevent oxidation. The DFT of the NiHITP complex may also
need to be considered, as this class of MOFs states that the active site for ORR is the β carbon of
the triphenylene. The differences between these successful catalysts and CoTHT should be further
studied to determine if future endeavors in this project are worthwhile.
One possible future approach to using THT MOFs for ORR is to use different metal centers
to modulate the electronic properties. The HHTP and HITP complexes have shown that metal-
ligand orbital overlap and the identity of the metal center can play a large role in the success of the
catalyst. In fact, the Co species explored in these studies also performed poorly in part due to the
material crystallizing in a trigonal configuration instead of hexagonal. Changing from Co to Ni or
Cu may aid in stabilizing the MOF or making it more resistant to oxidation. Studies on CoTHT
have suggested that the orbital energies of Co and S are similar, and therefore may be more
susceptible to negative interactions. However, Ni and Cu have different molecular orbital
structures and may therefore be more optimal for ORR. The difficulty then comes from the method
of growing the MOF films as previous attempts to make crystalline NiTHT have been unsuccessful
due to the rapid rate at which Ni coordinates. Additionally, literature examples show that Cu forms
more complicated structures where multiple Cu atoms can coordinate to the same sulfur.
Optimization of this process or use of different MOF synthesis techniques may be necessary to
perform ORR effectively with THT based frameworks.

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Abstract (if available)

Abstract The use of renewable energy in place of current fossil fuel technology has grown rapidly in its practicality, but is still limited by insufficient storage methods. Due to the inability to store energy long term, energy generated from solar panels must be used in the area and time frame within which it is generated. This hinders the availability of energy at night, which is when a large amount of energy is used. By developing catalysts which can use this energy to generate chemical fuels or valuable industrial products, the reliance on fossil fuels can still be minimized. Success has been seen with both solubilized homogeneous catalysts and surface-bound heterogeneous catalysts. In this work, the development of novel homogeneous catalysts and the heterogenization of well-known homogeneous catalysts into covalent organic and metal-organic frameworks will be presented. These species will be applied for CO₂ reduction and O₂ reduction.

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Asset Metadata

Creator Johnson, Eric Michael (author)

Core Title Design of catalysts for the transformation of abundant small molecules into solar fuels

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Publication Date 04/21/2020

Defense Date 04/21/2020

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

Tag catalysis,CO₂ reduction,electrocatalysis,O₂ reduction,OAI-PMH Harvest,solar fuels

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor Marinescu, Smaranda (committee chair), Feakins, Sarah (committee member), Williams, Travis (committee member)

Creator Email ericmjohnson135@gmail.com,johnsoem@usc.edu

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c89-286226

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Source University of Southern California (contributing entity), University of Southern California Dissertations and Theses (collection)

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