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Rhenium bipyridine catalysts with pendant amines: substituent and positional effects on the electrocatalytic reduction of CO₂ to CO

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Rhenium bipyridine catalysts with pendant amines: substituent and positional effects on the electrocatalytic reduction of CO₂ to CO

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Rhenium Bipyridine Catalysts with Pendant Amines:
Substituent and Positional Effects on the Electrocatalytic
Reduction of CO
2
to CO

by

Ashley N. Hellman

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 2021

ii

“Perhaps...the Earth was made round so that we would not see too far down the road.”
-Karen Blixen, “Out of Africa”
iii

ACKNOWLEDGEMENTS

Well, USC, it’s certainly been a wild ride. I never actually saw myself as a “Doctor,” so
sitting here writing this today is still a sort of surreal experience. Grad school has pushed my
boundaries in ways I never could have imagined, and I look back at a very different person when
I started at age 23. While the past five and a half years have been fraught with plenty of moments
of doubt and depression, they have largely been filled with incredible memories, and I have so
many people to thank for helping me on this journey.
First, I want to thank my PI, Professor Smaranda C. Marinescu, for allowing me this
opportunity for personal and professional growth. I have enjoyed the ability to direct my projects
and investigate my interests within the realm of topics studied in our lab. I am the chemist I am
today because I joined our group, and I look forward to taking this skillset into my future
endeavors. I would also like to thank Professor Travis Williams for allowing me the opportunity
to collaborate on electrochemical experiments for his group, which culminated in a journal
publication – this experience early on in my grad school career boosted my confidence and showed
me that others truly cared about my progress and success. Thank you as well to thank the other
members of my thesis and qualifying exam committees, Professors Mark Thompson, Noah
Malmstadt, and Ralf Haiges for their guidance and support throughout these stressful processes.
Going a little further back, I also need to thank my Austin College PI, Professor Stephanie
Gould, for always encouraging me to apply myself to my full potential and not choose a boring
job – you still have the right to say “I told you so.” Thank you for preparing me for grad school
and for all of your help along the way. Thank you as well to my undergrad academic advisor,
Professor David Baker, for the guidance when I was (for the first of several times) trying to figure
iv

out my life’s path. Also thank you for that awesome JanTerm trip to Guatemala, Belize, and
Mexico.
A big shoutout obviously goes to the rest of the MarinesCrew, both past and present. The
friendliness and support of everyone in the group was a big part of the reason I joined, and I truly
would not have made it through grad school without you guys. In particular, thanks to Alon
Chapovetsky for the training and guidance when I first joined the group, and to Eric Johnson for
always being super helpful (and down to gossip). Thanks to my lab bestie Nick, my first friend in
LA from visitation weekend debauchery. From all the DTLA train rides to Christmas decorating
and nights out in WeHo, I have truly treasured your friendship throughout the years. (Although
it’s probably a good thing we never became roommates because seeing each other 24/7 might’ve
been a bit too much.) Thanks to Geo (and Carlos, MarinesCrew member-adjacent) for finding Zeke
and being awesome FUN-cles, watching him when I was out of town. Keying, I’m so glad we got
to bond during our COVID-era two-person shift – I really enjoyed all of the matcha, coffee, and
boba trips, cat stories, and hot department gossip updates. Jeff and David, thanks for all of your
help with the IR that I never got trained on, and Jeremy, thanks for helping with XPS. Finally,
thanks to Jeremiah for your interest in taking over my project – I’m excited to see my “babies”
live on and find out what cool new things you do with these complexes.
One of the most important acknowledgements is for my grad school girls, Savannah, Abby,
and Caroline. You guys have been some of the best friends I could have asked for, and your support
throughout grad school has gotten me through so many hard times, both personal and professional.
Thank you for the wine and movie nights, the takeout, the celebrations, and even for listening to
my quals practice talks. Savannah, thanks for cooking with me, watching our favorite sitcoms,
helping me get used to dog(s), and always being down to vent. Also, I miss our girls’ trips, so we
v

should really start that back up once we all have better-paying jobs and COVID is over. I can’t
wait to see what incredible things you guys do in the future.
Thank you to all of my other grad school friends, for study sessions, practice talks, Netflix
parties, and general good times. I’m a true believer that the people in our lives make the
experiences so much better, and that has been an incredibly important part of my life at USC.
Thank you as well to my LA friends outside grad school – as much as I need my grad school
friends, sometimes it’s just really nice to not have to think about it for a while. Shoutouts to my
gym families from both The Phoenix Effect and Depot Fitness. Having that escape from grad
school to relieve stress and work on my physical strength has been so central to my life here in
LA, and it would not have been the same without all of your support and encouragement, both for
lifting and for life. Many thanks to my travel friends – my Aussie and Kiwi friends, especially
Georgie and Justine, who I’m so glad I got to visit before grad school got particularly crazy, and
my Iceland girls, Jenny, Louise, and Marina, with whom I’ve had so many great adventures
throughout Europe and America. And a massive thank you goes to my OC cousins Matt, Leesel,
Jenn, and David (and Evie and Major) for hosting me during quarantine and providing me a refuge
from the city. It’s meant so much to me to have family close by and to get to know you all better.
A big acknowledgement goes to my therapist Kim – I was always skeptical about therapy
until my doctor made me find someone to talk to. I almost quit after a few weeks, but I’m so glad
I stuck it out, because it has been one of the greatest experiences of my life. Thank you for
encouraging me to push my personal boundaries and for teaching me to set boundaries with others.
Thank you for spending 25% of our sessions giving me the skills to manage my mental health and
the other 75% just giving me a safe space to vent about literally anything going on in my life.
Thank you for reminding me to be kind to myself and that I have actually accomplished a lot,
vi

despite the serious imposter syndrome. I am happy with the person I am today, and I would not
have become that person without your help; I truly appreciate everything you’ve done for me.
Another set of important “thank yous” is for my friends from TX. Thanks to my amazing
Omega Zeta sisters for representing our pillars to the fullest, always setting great examples and
upholding a community of strong, supportive women. Thank you to Stephanie L. for
commiserating about grad school and being an awesome chemistry friend, and thank you to
Stephanie C. for flying to Vancouver with me just to run a half marathon and picking up our
friendship so easily. I’m going to throw some PA acknowledgements in here too and send a big
thank you to my fellow grad students from Penn State for the support when I was undergoing my
quarter-life crisis. Thanks specifically to Laura and Katelyn – I’m so glad we’ve kept in contact,
and that I’ve gotten to see and chat with both of you since starting my PhD program.
Indu, remember when I always refused to tell you you were awesome? Well, I’ll tell you a
million times now, because you truly are awesome. I can’t believe we’ve now been friends for
over half of our lives, IBFFWA – that we’ve remained such good friends over the years, despite
only living in the same place for a total of three months since high school, is one of my favorite
achievements to date. Thank you for visiting me in LA not once, but twice, and for becoming an
even better friend (I literally didn’t even know that was possible) since the start of COVID. Thank
you for the FaceTimes and phone calls, for the birthday gifts, and for DoorDashing me wine and/or
dessert on some of my bad days. Thank you for watching cheesy Netflix Christmas movies with
me, and for getting into The Mentalist so we can commiserate about Jane being stupid together.
Also, one day we will watch that Psych movie. All in all, I can’t even express how much your
support has meant to me – your constant encouragement and celebration for each little goal and
vii

big milestone has been so, so appreciated. I can’t wait to be your bridesmaid in a few months, and
I will cherish our friendship forever.
Amy, I want to thank you as well for all of the incredible support you’ve shown over the
years. I am so thankful we can pick up right where we left off, no matter how long it’s been since
we last talked to each other. You have been one of the best friends I could have imagined in my
life, and it’s crazy that even after more than ten years (!!!) of friendship, we are still basically the
same person. I can’t wait until we’ve both figured out our lives, and I really hope that someday we
can be next door neighbors (or at least live in the same state again). Thank you for binging shows
on Netflix with me during COVID, for our three-hour conversations about life, and for regularly
checking in during my thesis writing process. I am so thankful we met in Dave’s CI class and
found out that we were also both on the tennis team, studying the same subject, and joining the
same sorority (well, eventually). You’re a forever kind of friend, and I am truly looking forward
to more adventures with you in the future.
Most importantly, the biggest imaginable thank you goes to my parents, Bob and Debbie
Hellman. Your unending support, encouragement, and willingness to have serious discussions
about life (or just let me vent) have helped me end up where I am today. Thank you for everything
you have sacrificed to ensure that I can live my best possible life, for always welcoming me home
with open arms even though I decided to move across the country (twice), and for pushing me to
be a strong, successful woman. I can’t imagine going through life without you as my parents, and
I am so thankful that my relationships with both of you have gotten even stronger over the last
several years. Dad, I’m sorry I’ve now forgotten your birthday twice – I’m blaming it on grad
school brain, but I still feel terrible. Thank you for the financial and business advice and for helping
me with a ridiculous number of issues over the years. Mom, thanks for the genes for all of those
viii

premature gray hairs (well, maybe some of those are from grad school...), for the chats about life,
and for always sending cute cat pictures. I appreciate you both putting up with me in my times of
extreme stress, and hopefully I’ll be less unpleasant once I have a job and am not so anxious all
the time. Seriously, though, thank you for everything.
ix

TABLE OF CONTENTS
Epigraph .........................................................................................................................................ii
Acknowledgements..........................................................................................................................iii
Table of Contents............................................................................................................................ix
List of Figures...............................................................................................................................xiii
List of Schemes............................................................................................................................xviii
List of Tables.................................................................................................................................xix
Chapter 1: General Introduction....................................................................................................1
1.1 Global Energy Consumption..........................................................................................2
1.2 Climate Change and Renewable Energy Sources ..........................................................3
1.3 Carbon Dioxide Reduction Reaction..............................................................................5
1.4 Homogeneous Electrocatalysts for the Carbon Dioxide Reduction...............................6
1.4.1 Homogeneous Catalysts with Secondary Sphere Interactions........................8
1.4.1 Rhenium and Manganese Bipyridine Catalysts with Secondary Sphere
Interactions.................................................................................................10
1.4.2.1 Rhenium Bipyridine Catalysts........................................................10
1.4.2.2 Manganese Bipyridine Catalysts....................................................13
1.5 Heterogenization of Molecular Electrocatalysts for the Carbon Dioxide Reduction
Reaction.................................................................................................................15
1.5.1 Heterogeneous Rhenium Bipyridine Catalysts.............................................16
1.5.1.1 Covalent Organic Frameworks......................................................16
1.5.1.2 Non-Covalent Immobilization.......................................................16
1.5.1.3 Covalent Immobilization................................................................17
1.6 Outline of Thesis..........................................................................................................19
1.7 References....................................................................................................................22

Chapter 2: Rhenium bipyridine catalysts with hydrogen bonding pendant amines for CO2
reduction........................................................................................................................................30
x

2.1 Abstract........................................................................................................................31
2.2 Introduction..................................................................................................................31
2.3 Results and Discussion ................................................................................................33
3.2.1 Synthesis and Characterization......................................................................33
3.2.2 Cyclic Voltammetry......................................................................................43
3.2.3 Controlled Potential Electrolysis...................................................................49
2.4 Conclusion....................................................................................................................53
2.5 Experimental Methods and Additional Figures............................................................54
2.5.1 Materials and Synthesis ................................................................................54
2.5.2 Electrochemistry...........................................................................................54
2.5.3 Synthesis of N
6
,N
6
'-dimethyl-[2,2'-bipyridine]-6,6'-diamine (L
1
).................55
2.5.4 Synthesis of N-methyl-[2,2'-bipyridine]-6-amine (L
2
)..................................56
2.5.5 Synthesis of N
6
,N
6
,N
6'
,N
6'
-tetramethyl-[2,2'-bipyridine]-6,6'-diamine (L
3
)...57
2.5.6 Synthesis of N,N-dimethyl-[2,2'-bipyridin]-6-amine (L
4
).............................57
2.5.7 Synthesis of Re(L
1
)(CO)3Cl (1).....................................................................58
2.5.8 Synthesis of Re(L
2
)(CO)3Cl (2).....................................................................58
2.5.9 Synthesis of Re(L
3
)(CO)3Cl (3).....................................................................59
2.5.10 Synthesis of Re(L
4
)(CO)3Cl (4)...................................................................59
2.5.11 Additional Figures.......................................................................................61
2.6 References....................................................................................................................73

Chapter 3: Influence of intermolecular hydrogen bonding interactions on the electrocatalytic
reduction of CO2 to CO by 6,6'-amine substituted rhenium bipyridine complexes.........................76
3.1 Abstract........................................................................................................................77
3.2 Introduction..................................................................................................................77
3.3 Results and Discussion ................................................................................................81
3.3.1 Synthesis and Characterization......................................................................81
3.3.2 Cyclic Voltammetry......................................................................................86
3.3.3 Controlled Potential Electrolysis...................................................................94
xi

3.4 Conclusions..................................................................................................................99
3.5 Experimental Methods and Additional Figures..........................................................100
3.5.1 Materials and Synthesis ..............................................................................100
3.5.2 Electrochemistry.........................................................................................101
3.5.3 Synthesis of 6,6'-NH2(bpy)[Re(CO)3Cl] (NH2)...........................................102
3.5.4 Additional Figures.......................................................................................103
3.6 References..................................................................................................................111

Chapter 4: Primary and secondary sphere effects of amine substituent position on rhenium
bipyridine electrocatalysts for CO2 reduction.............................................................................116
4.1 Abstract......................................................................................................................117
4.2 Introduction................................................................................................................117
4.3 Results and Discussion ..............................................................................................119
4.3.1 Synthesis and Characterization....................................................................119
4.3.2 Cyclic Voltammetry....................................................................................121
4.3.3 Controlled Potential Electrolysis.................................................................128
4.4 Conclusions................................................................................................................129
4.5 Experimental Methods and Additional Figures..........................................................130
4.5.1 Materials and Synthesis ..............................................................................130
4.5.2 Electrochemistry.........................................................................................131
4.5.3 Synthesis of 4,4’-NH2..................................................................................132
4.5.4 Additional Figures.......................................................................................133
4.6 References..................................................................................................................138

Chapter 5: Immobilization of a Molecular Rhenium Bipyridine Electrocatalyst for CO2
Reduction.....................................................................................................................................142
5.1 Abstract......................................................................................................................143
5.2 Introduction................................................................................................................143
5.3 Results and Discussion ..............................................................................................146
xii

5.3.1 Synthesis and Characterization....................................................................146
5.3.2 Electropolymerization.................................................................................149
5.3.3 Characterization of Electropolymerized Films............................................152
5.3.3.1 XPS..............................................................................................152
5.3.4 Electrochemistry of Films...........................................................................157
5.3.5 Electrocatalytic Studies...............................................................................161
5.3.5.1 Cyclic Voltammetry.....................................................................161
5.4 Conclusions................................................................................................................163
5.5 Experimental Methods...............................................................................................164
5.5.1 Materials and Synthesis ..............................................................................164
5.5.2 Electrochemistry.........................................................................................164
5.5.3 Synthesis of 4,4’-N2
+
-Re.............................................................................166
5.5.4 Physical Methods........................................................................................166
5.5.4.1 XPS..............................................................................................166
5.5.4.2 FTIR.............................................................................................167
5.5.4.3 UV-vis..........................................................................................167
5.6 References..................................................................................................................168
Bibliography................................................................................................................................171

xiii

LIST OF FIGURES
Figure 1.1 Global energy consumption by type..............................................................................3
Figure 1.2 United States coal and renewable energy consumption by source................................4
Figure 1.3 Depiction of homogeneous and heterogeneous catalyst activity...................................6
Figure 1.4 Structure of the active site of [NiFe]-CO dehydrogenase with a bound CO2
molecule...............................................................................................................................8
Figure 2.1
1
H NMR spectra of ligands L
1
–L
4
in CDCl3...............................................................32
Figure 2.2
1
H NMR spectra of complexes 1–4 in CDCl3.............................................................34
Figure 2.3 Solid state structures of complexes 1–4......................................................................34
Figure 2.4 Solid state structures of complexes 1 and 2 with N-H bonds included.......................35
Figure 2.5 Solid state structures of complexes 1, 2, and 4 showing N–C–C–N torsion angles....36
Figure 2.6 Solid state structures of complexes 1, 2, and 4 showing N–C–C–N torsion angles....37
Figure 2.7 FTIR spectra of complexes 1–4...................................................................................38
Figure 2.8 Cyclic voltammograms of complexes 1–4 and Re(bpy)(CO)3Cl in an acetonitrile
solution under N2 atmosphere............................................................................................40
Figure 2.9 Randles-Sevcik plots of the first and second reduction features of complexes 1–4....40
Figure 2.10 Variable scan rate studies of complexes 1–4.............................................................42
Figure 2.11 Cyclic voltammograms of complex 2 in a DMF solution under N2 atmosphere......42
Figure 2.12 Cyclic voltammograms complexes 1–4 under N2, CO2, N2 + 2 M TFE, and CO2
+ 2 M TFE..........................................................................................................................43
Figure 2.13 Cyclic voltammograms of complexes 1–4 under CO2 with added PhOH, TFE, and
H2O....................................................................................................................................43
Figure 2.14 Current versus time over one hour of controlled potential electrolysis for complexes
1–4......................................................................................................................................44
Figure 2.15 UV-vis before and after CPE of complexes 3 and 4....................................................45
Figure 2.16 Cyclic voltammograms of a blank cell and complex 4 before controlled potential
electrolysis with no added TFE, 1 M TFE, and 2 M TFE....................................................45
Figure 2.17 Cyclic voltammograms of 1 under CO2 with increasing concentrations of H2O.........61
Figure 2.18 Cyclic voltammograms of 1 under CO2 with increasing concentrations of MeOH.....62
Figure 2.19 Cyclic voltammograms of 1 under CO2 with increasing concentrations of TFE.........62
Figure 2.20 Cyclic voltammograms of 1 under CO2 with increasing concentrations of PhOH......63
Figure 2.21 Cyclic voltammograms of 2 under CO2 with increasing concentrations of H2O.........63
xiv

Figure 2.22 Cyclic voltammograms of 2 under CO2 with increasing concentrations of MeOH.....64
Figure 2.23 Cyclic voltammograms of 2 under CO2 with increasing concentrations of TFE.........64
Figure 2.24 Cyclic voltammograms of 2 under CO2 with increasing concentrations of PhOH......65
Figure 2.25 Cyclic voltammograms of 3 under CO2 with increasing concentrations of H2O.........65
Figure 2.26 Cyclic voltammograms of 3 under CO2 with increasing concentrations of MeOH.....66
Figure 2.27 Cyclic voltammograms of 3 under CO2 with increasing concentrations of TFE.........66
Figure 2.28 Cyclic voltammograms of 3 under CO2 with increasing concentrations of PhOH......67
Figure 2.29 Cyclic voltammograms of 4 under CO2 with increasing concentrations of H2O.........67
Figure 2.30 Cyclic voltammograms of 4 under CO2 with increasing concentrations of MeOH.....68
Figure 2.31 Cyclic voltammograms of 4 under CO2 with increasing concentrations of TFE.........68
Figure 2.32 Cyclic voltammograms of 4 under CO2 with increasing concentrations of PhOH......69
Figure 2.33 Cyclic voltammograms of 1 under N2 with increasing amounts of TFE......................69
Figure 2.34 Cyclic voltammograms of 2 under N2 with increasing amounts of TFE......................70
Figure 2.35 Cyclic voltammograms of 3 under N2 with increasing amounts of TFE......................70
Figure 2.36 Cyclic voltammograms of 4 under N2 with increasing amounts of TFE......................71
Figure 2.37 Cyclic voltammogram of 2 in DMF under N2 compared with that under CO2............71
Figure 2.38 Cyclic voltammograms of 2 under CO2 with increasing concentrations of TFE.........72
Figure 3.1
1
H NMR spectrum of NH2-Re in CDCl3.......................................................................82
Figure 3.2 UV-vis spectra of NH2-Re, NHMe-Re, and NMe2-Re in MeCN..................................82
Figure 3.3 Solid state structure of 6,6'-NH2-Re..............................................................................83
Figure 3.4 FT-IR spectrum of NH2-Re...........................................................................................85
Figure 3.5 Cyclic voltammograms of 1 mM NH2-Re, NHMe-Re, and NMe2-Re under N2 in a
MeCN solution...................................................................................................................86
Figure 3.6 Variable scan rate studies of NH2-Re under N2...........................................................88
Figure 3.7 Randles-Sevcik plots for NH2-Re................................................................................88
Figure 3.8 Cyclic voltammograms of NH2-Re, NHMe-Re, and NMe2-Re under N2 in MeCN
and DMF...........................................................................................................................89
Figure 3.9 Variable scan rate studies of NH2-Re, NHMe-Re, and NMe2-Re in DMF under
N2.......................................................................................................................................90
Figure 3.10 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of
DMF and TBACl..............................................................................................................91
xv

Figure 3.11 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of
triethylamine......................................................................................................................92
Figure 3.12 Cyclic voltammograms of NH2-Re and NHMe-Re in MeCN under N2 with the
addition of no acid, H2O, and TFE......................................................................................93
Figure 3.13 Cyclic voltammograms of NH2-Re, NHMe-Re, and NMe2-Re under N2, CO2, 1 M
PhOH, 1 M TFE, and 1 M H2O...........................................................................................94
Figure 3.14 Current versus time over one hour of controlled potential electrolysis of NH2-Re,
NHMe-Re, and NMe2-Re at -2.00 V, -2.10 V, -2.20 V, and -2.30 V...................................95
Figure 3.15 Plot of Faradaic efficiency for CO (%) versus potential (V) for CPE experiments
of NH2-Re...........................................................................................................................97
Figure 3.16 CPE cell showing NHMe-Re after 1 hour electrolysis under N2 at -1.96 V.................98
Figure 3.17 Cyclic Voltammograms of NHMe-Re in a CV cell following 1 hour of CPE under
N2 at -1.96 V and overlay of CVs with added DMF and NEt3.............................................98
Figure 3.18 First and second reduction features of NH2-Re in DMF............................................103
Figure 3.19 First and second reduction features of NHMe-Re in DMF........................................104
Figure 3.20 First and second reduction features of NMe2-Re in DMF.........................................104
Figure 3.21 Variable scan rate studies of NH2-Re under N2 in DMF............................................105
Figure 3.22 Variable scan rate studies of NHMe-Re under N2 in DMF........................................105
Figure 3.23 Variable scan rate studies of NMe2-Re under N2 in DMF.........................................106
Figure 3.24 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of
H2O..................................................................................................................................106
Figure 3.25 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of
TFE..................................................................................................................................107
Figure 3.26 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of
PhOH...............................................................................................................................107
Figure 3.27 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of
H2O..................................................................................................................................108
Figure 3.28 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of
TFE..................................................................................................................................108
Figure 3.29 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of
H2O..................................................................................................................................109
Figure 3.30 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of
TFE..................................................................................................................................109
Figure 3.31 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of
PhOH...............................................................................................................................110
xvi

Figure 4.1
1
H NMR of 4,4’-NH2-Re in DMSO-d6.......................................................................120
Figure 4.2 UV-vis spectrum of 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re in MeCN............120
Figure 4.3 FT-IR spectrum of 4,4’-NH2-Re.................................................................................121
Figure 4.4 Variable scan rate CVs and Randles-Sevcik plot of 4,4’-NH2-Re under N2................122
Figure 4.5 Overlay of cyclic voltammograms of 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re
under N2 in MeCN and DMF............................................................................................124
Figure 4.6 Variable scan rate studies of 4,4’-NH2-Re and 5,5’-NH2-Re under N2.......................124
Figure 4.7 CVs of 4,4’-NH2-Re under N2, CO2, N2 + 1 M TFE, and CO2 + 1 M TFE...................125
Figure 4.8 CVs of 4,4’-NH2-Re and 5,5’-NH2-Re under CO2 with 1 M PhOH, 1 M TFE, and
1 M H2O...........................................................................................................................126
Figure 4.9 CVs of 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re under CO2 with 1 M TFE........127
Figure 4.10 CVs of 4,4’-NH2-Re and 5,5’-NH2-Re under N2 with 1 M PhOH, 1 M TFE, and
1 M H2O...........................................................................................................................128
Figure 4.11 Current versus time over one hour of controlled potential electrolysis of 4,4’-
NH2-Re at -2.10 V, -2.20 V, -2.30 V, and -2.40 V.............................................................129
Figure 4.12 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations
of H2O..............................................................................................................................133
Figure 4.13 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations
of TFE..............................................................................................................................133
Figure 4.14 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations
of PhOH...........................................................................................................................134
Figure 4.15 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations
of H2O..............................................................................................................................134
Figure 4.16 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations
of TFE..............................................................................................................................135
Figure 4.17 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations
of PhOH...........................................................................................................................135
Figure 4.18 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations
of H2O..............................................................................................................................136
Figure 4.19 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations
of TFE..............................................................................................................................136
Figure 4.20 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations
of PhOH...........................................................................................................................137
Figure 5.1
1
H NMR and
19
F NMR of 4,4’-N2
+
-Re in MeCN-d3...................................................147
xvii

Figure 5.2 UV-vis spectrum of 4,4’-N2
+
-Re in MeCN.................................................................148
Figure 5.3 FTIR spectrum of 4,4’-N2
+
-Re....................................................................................148
Figure 5.4 Cyclic voltammograms of electropolymerization of 4,4’-N2
+
-Re on a glassy
carbon stick electrode.......................................................................................................150
Figure 5.5 Cyclic voltammograms of electropolymerization of 4,4’-N2
+
-Re on an FTO
electrode...........................................................................................................................150
Figure 5.6 XPS survey scans for modified FTO electrodes..........................................................153
Figure 5.7 High-resolution XPS of the Re 4f region for modified FTO electrodes......................154
Figure 5.8 High-resolution XPS of the Sn 2p region for modified FTO electrodes......................154
Figure 5.9 High-resolution XPS of the Cl 2p region for a modified FTO electrode.....................155
Figure 5.10 High-resolution XPS of the P 2p region for a modified FTO electrode.....................155
Figure 5.11 High-resolution XPS of the F 1s region for a modified FTO electrode.....................156
Figure 5.12 High-resolution XPS of the N 1s region for a modified FTO electrode.....................156
Figure 5.13 High-resolution XPS of the C 1s region for a modified FTO electrode.....................157
Figure 5.14 Cyclic voltammograms of modified GC stick electrodes..........................................157
Figure 5.15 Variable scan rate studies of modified glassy carbon electrodes...............................159
Figure 5.16 Randles-Sevcik plots for modified glassy carbon electrodes....................................160
Figure 5.17 Double-layer charging current density (ΔJ = Janodic-Jcathodic) at the open-
circuit potential for modified glassy carbon electrodes.....................................................161
Figure 5.18 Cyclic Voltammograms of modified glassy carbon electrodes under N2, CO2,
and CO2 + 0.5 M TFE.......................................................................................................162
Figure 5.19 Cyclic Voltammograms of modified FTO electrodes under N2, CO2, and CO2 +
0.5 M TFE........................................................................................................................163

xviii

LIST OF SCHEMES
Scheme 1.1 CO2 reduction reaction equations with their associated standard reduction
potentials............................................................................................................................11
Scheme 2.1 Synthetic scheme for complexes 1–4..........................................................................33
Scheme 2.2 Synthetic schemes for 1, 2, 3, and 4.............................................................................61
Scheme 3.1 Synthetic scheme for complexes NH2, NHMe, and NMe2........................................81
Scheme 3.2 Synthetic scheme for the formation of NH2-Re.........................................................103
Scheme 4.1 Synthesis of complex 4,4’-NH2-Re...........................................................................119
Scheme 5.1 Synthesis of 4,4’-N2
+
-Re from 4,4’-NH2-Re.............................................................146
Scheme 5.2 Scheme demonstrating electrochemical grafting of 4,4’-N2
+
-Re onto various
electrode surfaces by cyclic voltammetry and subsequent electropolymerization............149

xix

LIST OF TABLES
Table 2.1 Crystal data and structure refinement for complex 1......................................................38
Table 2.2 Crystal data and structure refinement for complex 2......................................................39
Table 2.3 Crystal data and structure refinement for complex 3......................................................40
Table 2.4 Crystal data and structure refinement for complex 4......................................................41
Table 2.5 Nitrogen–rhenium bond lengths (in Ångstroms) for complexes 1, 2, and 4....................42
Table 2.6 Pertinent IR spectroscopy wavenumbers for complexes 1–4.........................................43
Table 2.7 Reduction potentials for complexes 1–4 under N2..........................................................46
Table 2.8 Catalytic parameters for complexes 1–4 and Re(bpy)(CO)3Cl under CO2 in MeCN
with added TFE..................................................................................................................51
Table 3.1 Crystal data and structure refinement for NH2-Re..........................................................83
Table 3.2 Pertinent IR spectroscopy wavenumbers for NH2-Re, NHMe-Re, and NMe2-Re........85
Table 3.3 Reduction potentials for NH2-Re, NHMe-Re, and NMe2-Re in MeCN and DMF
under N2............................................................................................................................86
Table 3.4 CPE results for NMe2-Re in MeCN under CO2 with 1 M TFE......................................96
Table 3.5 CPE results for NHMe-Re in MeCN under CO2 with 1 M TFE......................................96
Table 3.6 CPE results for NH2-Re in MeCN under CO2 with 1 M TFE..........................................96
Table 4.1 Reduction potentials for 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re in MeCN and
DMF under N2..................................................................................................................122
Table 4.2 CPE results of 4,4’-NH2-Re in MeCN under CO2 with 1 M TFE..................................129

1

CHAPTER 1

General Introduction

2

1.1 GLOBAL ENERGY CONSUMPTION
2
Figure 1.1 Global energy consumption by type. Reprinted with permission from Ref 1.
With the world’s continually increasing population and an expanding global economy
comes the necessity for an increase in energy consumption to match. The world today relies heavily
on fossil fuel consumption to provide energy for electricity and many daily activities. In fact, as
of 2019, a staggering 84% of the world’s energy sources was made up of various fossil fuels,
including oil, natural gas, and coal, while only 5% came from renewable sources (Figure 1.1).
1

Further, with only about 115 years left of coal production and approximately 50 years of both oil
and natural gas, the increased use of renewable resources is necessary for both the transition to
clean energy and to keep up with world population increases and the resulting energy
3

consumption.
2,3
In 2018, the global electricity demand grew by 4% to more than 23,000 TWh,
making electricity usage nearly 20% of the total energy consumption.
4
The greatest increases in
energy demand that year occurred in China (3.5%, 107 Mtoe) and the United States (3.7%, 80
Mtoe), with expanded usage of coal as well as solar and wind generation to meet this outstanding
electricity demand. Further increases in global energy demand are expected to occur for at least
the next few decades, as countries such as India, Africa, and China continue to grow in population
and industrialize.
4

1.2 CLIMATE CHANGE AND RENEWABLE ENERGY SOURCES
Carbon emissions, in particular carbon dioxide (CO2), are responsible for the majority of
the climate change that has occurred thus far, and this process is not slowing down.
5–7
As of 2019,
coal-fired power generation accounted for 30% of global energy-related CO2 emissions; through
fossil fuel combustion processes, over 34 gigatons of CO2 were released into the atmosphere.
4

While the 0.5% increase in CO2 emissions in 2019 was an improvement from the 2.1% increase
in 2018, significant changes need to be made – in order to achieve net zero carbon emissions by
2050, CO2 emissions need to decrease by nearly 2.6 gigatons annually, the same as the estimated
decrease during the COVID-19 pandemic. Due to these carbon emissions, average global
temperatures have increased 1°C from pre-industrial levels and are expected to reach a 1.5°C
increase within the next two decades, leading to extreme weather events, coral bleaching, and other
impacts on both the ecosystem and society.
8,9
Combining the need to combat detrimental climate
change and carbon emissions while maintaining the ability to meet growing global energy demands
requires a significant increase in renewable resource consumption.
In 2019, United States annual renewable energy consumption surpassed coal for the first
time in 130 years, with the total renewable energy consumption reaching a record high of 11.5
4

quadrillion Btu, showing a promising trend in the use of renewable energy sources (Figure 1.2).
10

While renewable resource consumption has increased in recent years (1.5% from 2019 to 2020),
it still makes up only about 11% of the total energy consumption.
11,12
The majority of renewable
resource consumption is generated by wind and hydro power, with less than 1% attributed to
sustainable solar energy due to high costs and low efficiencies. The sun delivers 4.6 x 10
20
joules
of energy to the earth in just one hour, enough to power human activity for an entire year.
13

However, solar power is both diffuse and intermittent, only able to be harvested during the day
and in clear weather, preventing humans from fully harnessing the power of the sun. In order to
maximize solar energy consumption, it is therefore necessary to find a method for storing this
energy for later use.
14–16
Solar and other renewable energy storage will allow for more sustainable
pathways for producing fuels and chemicals, including hydrogen, hydrocarbons, oxygenates, and
ammonia, that can be used for further energy production.
17
By converting some of the most
abundant chemicals in the atmosphere such as water, nitrogen, and carbon dioxide into these
products via electrochemical processes, more sustainable and efficient energy production can be
realized.

5

Figure 1.2 United States coal and renewable energy consumption by source from 1950 to 2019.
Reprinted with permission from Ref 10.

1.3 CARBON DIOXIDE REDUCTION REACTION
The carbon dioxide reduction reaction (CO2RR) is a useful method for recycling harmful
CO2 while also allowing for energy storage for use in both renewable electricity and industrial
chemicals.
18–20
A variety of reduction products can be formed through this reaction, including
carbon monoxide (CO), formic acid (HCOOH), methanol (MeOH), methane (CH4), ethanol
(EtOH), and ethylene (C2H4). However, carbon dioxide is a linear molecule, requiring significant
energy to reorganize into a different geometry and bond order; thus, the thermodynamic activation
barrier and high overpotential necessary for this reaction prevents most of these products from
being easily accessible (Scheme 1.1). While processes involving more protons and electrons
ultimately reduce the thermodynamic barrier, product formation is instead limited by kinetic
constraints.
19
To overcome these barriers, it is necessary to use a catalyst.
CO2 + e
-
→ CO2˙
-
E°’ = -1.90 V
CO2 + 2H
+
+ 2e
-
→ CO + H2O E°’ = -0.53 V
CO2 + 2H
+
+ 2e
-
→ HCOOH E°’ = -0.61 V
CO2 + 4H
+
+ 4e
-
→ HCOH + H2O E°’ = -0.48 V
CO2 + 6H
+
+ 6e
-
→ CH3OH + H2O E°’ = -0.38 V
CO2 + 8H
+
+ 8e
-
→ CH4 + H2O E°’ = -0.24 V
Scheme 1.1 CO2 reduction reaction equations with their associated standard reduction potentials.
19

6

One of the most common CO2RR products is carbon monoxide (CO). Requiring an
overpotential of 0.53 V, only two protons and two electrons are needed to form CO from CO2,
making both the thermodynamic and kinetic barriers comparatively low for this reaction. Since
CO is a gas, it is easy to separate and can then be used to make hydrocarbons and other useful
fuels, such as gasoline, via the Fischer-Tropsch process.
21,22
Formation of other CO2RR products
such as formic acid as well as hydrogen from the competing hydrogen evolution reaction (HER)
can lead to issues with selectivity, making catalyst selectivity an important consideration. The
specific product(s) formed during the CO2RR depend on the mechanism with the lowest energy
barriers and whether a CO2 molecule or a hydride forms an adduct with the metal center of the
catalyst.
23
Catalysts can thus be designed with ideal properties for each scenario, allowing for
optimization of selectivity as well as other catalytic metrics.

1.4 HOMOGENEOUS ELECTROCATALYSTS FOR THE CARBON DIOXIDE
REDUCTION REACTION
There are two main classes of electrocatalysts that facilitate the CO2RR: homogeneous
catalysts and heterogeneous catalysts (Figure 1.3). Homogeneous catalysts are molecular
complexes that operate in organic or aqueous solution and typically are defined by diffusion at the
electrode surface. Some key benefits for homogeneous catalysts include their well-defined
structures, tunable coordination environments, and relative ease of functional group modification
for improving catalytic performance.
24
These properties allow for insight into the CO2RR
mechanisms for unique classes of homogeneous catalysts, providing a better understanding of how
different modifications affect the various catalytic parameters. However, these catalysts have
pitfalls as well, including low overall stability and poor recyclability, preventing large-scale
7

industrial applications.
25
Additionally, catalyst solubility, concentration, diffusion, and product
separation can all be limiting factors. Heterogeneous catalysts will be discussed further in section
1.5.

Figure 1.3 Depiction of homogeneous and heterogeneous catalyst activity. Reproduced with
permission from Ref 24.

A number of catalytic metrics are typically used in the assessment of catalytic activity and
performance: overpotential (η), the standard reduction potential subtracted from the applied
potential; Faradaic efficiency (FE), the ratio of moles of product divided by half the charge passed
during controlled potential electrolysis (CPE); turnover number (TON), the total moles of product
per mole of catalyst from CPE; turnover frequency (TOF), the TON per unit of time; and the
catalytic rate constant (kobs), which can be used as a proxy for TOF under pseudo first-order
conditions.
26–28
These parameters are evaluated for each electrocatalyst and used for comparison
to others throughout the literature; they will be used for further discussion of various CO2RR
catalysts in the following sections.
8

1.4.1 HOMOGENEOUS CATALYSTS WITH SECONDARY SPHERE
INTERACTIONS
In nature, the enzyme [NiFe]-CO dehydrogenase (CODH) catalyzes the reduction of CO2
to CO and water with high rate and efficiency.
29–32
The active site of CODH from
Carboxydothermus hydrogenoformans contains a nickel center bound to three sulfur ligands and
an iron center coordinated by amino acid residues (His and Cys), a sulfido ligand, and a water
ligand.
31
During CO2RR, the NiFe cluster binds a CO2 adduct, which is stabilized and facilitated
by hydrogen bonding secondary sphere interactions with perfectly positioned amino acid residues
within the binding pocket (Figure 1.4).

Figure 1.4 Structure of the active site of [NiFe]-CO dehydrogenase with a bound CO2 molecule.
Due to the exceptional activity achieved by billions of years of evolution, chemists have
often used nature as inspiration for the design of homogeneous catalysts for CO2RR, introducing
pendant protons capable of this CO2 adduct stabilization. While emulating these beneficial
properties with synthetic compounds is not trivial, a vast number of CO2RR homogeneous
electrocatalysts with wide-ranging capabilities for CO2 reduction have been reported to date; only
selected examples with second-sphere interactions will be discussed herein.
19,25,33
Nickel cyclams
9

and related nickel macrocyclic compounds are a well-studied class of CO2RR electrocatalysts.
34–
39
With macrocyclic backbones and secondary amine substituents, they have often been used as a
pseudo-standard for the electrocatalytic reduction of CO2 in both aqueous and organic media.
Modification of these catalysts has led to a wide range of activities for the selective formation of
CO due to electronic effects, and mechanistic studies have shown evidence of Ni-CO2 adducts
during catalysis. However, they often require the use of environmentally toxic mercury electrodes.
Iron porphyrins have also been extensively studied as CO2RR electrocatalysts.
40
In
particular, iron porphyrins with pendant protons have been shown to have excellent activity for
the selective reduction of CO2 to CO. Iron porphyrin electrocatalysts with phenol substituents
induced CO2-adduct stabilization through hydrogen bonding interactions, achieving catalytic
current increases in cyclic voltammetry (CV) studies at lower overpotentials than the unsubstituted
iron porphyrin catalyst due to the presence of an internal proton source.
41
Porphyrins with pendant
amide groups were shown to have enhanced catalytic rates due to through-space positional
effects.
42
By varying the location of the amide N-H group (ortho- and para- as well as proximal
and distal to the porphyrin plane), the results illustrated the significance of positioning for ideal
second-sphere hydrogen bonding interactions, as the para-functionalized isomers did not display
notable through-space effects on catalytic activity. Iron hangman porphyrins with pendant phenol,
guanidinium, and sulfonic acid groups reduce CO2 to CO with >93% FE, with computational
studies showing that the positioning of the pendant protons allows for 2.1−6.6 kcal/mol
stabilization of the CO2 adduct at the metal center during catalysis.
43

Several cobalt macrocycle complexes have also been shown to possess high activity for
CO2RR. Cobalt aminopyridine macrocycles with pendant protons demonstrated a dependence of
catalytic activity on the number of available proton donors, with a decrease in FE from 98% to
10

36% and an overall reaction rate three orders of magnitude lower when the four secondary amine
moieties are replaced by tertiary methylamines.
44
A further look at this series of aminopyridine
macrocycles with varying numbers of pendant protons demonstrate a first order dependence of
catalytic reaction rate on number of pendant secondary amines.
45
Density functional theory (DFT)
studies indicate that the pendant amines bind acid molecules in the catalytic solution, forming an
extended hydrogen bonding network that shuttles protons from the acid source to the activated
CO2 molecule. Altering the macrocyclic pyridine backbone with electron donating -NMe2 groups
increases the negative charge on the CO2 adduct and enhances the protonation steps of the catalytic
mechanism; this provides additional support for the conclusion that the pendant amines facilitate
proton transfer from the acid source to the CO2 adduct during electrocatalysis.
46
Cobalt
macrocyclic catalysts with bipyridyl-N-heterocyclic carbene frameworks have also demonstrated
high activity for the selective reduction of CO2 to CO in aqueous media, with FECO ranging from
78% to 98% depending on the rigidity of the macrocyclic framework.
47

1.4.2 RHENIUM AND MANGANESE BIPYRIDINE CATALYSTS WITH
SECONDARY SPHERE INTERACTIONS
1.4.2.1 RHENIUM BIPYRIDINE CATALYSTS
Rhenium and manganese bipyridine complexes (Re(bpy)(CO)3Cl and Mn(bpy)(CO)3Br
[bpy = 2,2'-bipyridine]) are some of the most well-studied classes of CO2RR electrocatalysts.
48–52

Extensive work has also been done investigating Re and Mn(bpy) electrocatalysts with pendant
protons capable of undergoing secondary sphere interactions.
53
Several rhenium bipyridine
CO2RR electrocatalysts with pendant -OH groups have been discussed in the literature. A Re(bpy)
catalyst with 6,6’-hydroxy groups displayed a moderate FECO of 70% with TON <1, suggesting
11

slow product release and decomposition of intermediates at the negative potentials required for
catalysis.
54
This activity is attributed to electrochemical reductive cleavage of the O-H bond
leading to deprotonation. The 4,4’-anologue performed significantly better, producing 450 µL CO
with FE = 95%. Another study of the catalytic activity of a Re(bpy) catalyst with 3,3’-OH groups
showed that it also undergoes a 1 e
−
reductive cleavage of the O-H bond, forming a 5-coordinate
[Re(3,3′-DHBPY-2H
+
)(CO)3]
3−
species.
55
CPE experiments yielded a FECO of 83% after one hour,
falling in between the values observed for the 4,4’- and 6,6’-OH catalysts. Two rhenium complexes
with local proton sources in the form of phenyl diol and triol substituents exhibit moderate current
enhancements at the first and second reduction potentials under CO2 with added water but undergo
electrode passivation during catalysis, yielding low to moderate TONCO between 2 and 14.
56
This
behavior suggests the amount of local proton source does not play a significant role in catalytic
activity for these complexes.
Several biologically-inspired Re(bpy) catalysts with pendant amide and amino acid groups
have also been reported. The introduction of methyl acetamidomethyl groups at the 4,4’-positions
of a Re(bpy) complex was found to enhance the rate of a bimolecular catalytic mechanism
involving the formation of a hydrogen bonded dimer.
57
This species was catalytically active for
the reductive disproportionation of CO2 to CO and CO3
2-
at an approximately 250 mV lower
overpotential than the corresponding formation of CO and H2O with typical Re(bpy) catalysts,
although with lower FE and TOF (54-56% FECO). Further modification with tyrosine amino acid
residues led to an increase in the rate of catalysis compared to a phenylalinine-substituted analogue
without pendant protons, undergoing the same bimolecular reductive disproportionation
mechanism as the amide-substituted catalyst.
58
The increased rate was attributed to the positioning
of pendant phenol groups of the tyrosine residues near the rhenium metal center, allowing for a
12

similar supramolecular assembly of hydrogen bonding interactions as that observed for the amide-
substituted version. Additional amino acid residues, peptides, and hydrogen bonding groups were
also appended to this family of rhenium bipyridine catalysts to introduce proton relays for the
cooperative reduction of CO2 to CO.
59
These studies demonstrated that peptide backbones at a
length of five amino acids can adopt conformations with observable intramolecular interactions
even on the nuclear magnetic resonance (NMR) timescale. This data further suggests that both the
unimolecular and bimolecular mechanisms can operate in tandem on the cyclic voltammetry (CV)
timescale, allowing for optimization of catalytic rate and overpotential based on substituent
modification.
A rhenium bipyridine catalyst with a thiourea tether in the second coordination sphere
demonstrated the selective reduction of CO2 to CO with a TOF of 3040 s
-1
.
60
This fast catalytic
rate was attributed to stabilization of the Re-CO2 adduct, which was observed by
1
H NMR, and
additional experiments indicated that the pendant thiourea acted as a local proton source as well,
further enhancing catalytic activity. A charged imidazolium functionalized Re(bpy) catalyst with
an available pendant proton exhibited a slightly greater current increase under CO2 than the
unsubstituted complex at a potential about 170 mV more positive, induced by hydrogen bonding
second-sphere interactions.
61
Controlled potential electrolysis (CPE) experiments with added
water yielded moderate FECO, likely due to consumption of the proton source. A series of rhenium
bipyridine catalysts with pendant arylamine groups, with the primary amine positioned in the
ortho-, meta-, and para- sites of the aniline substituent, were investigated for both primary and
secondary sphere effects on electrocatalytic reduction of CO2.
62
In the presence of an additional
proton source, these catalysts produced CO with FE ≥ 89% and TOFs ranging from 109 to 239 s
-1

compared to 73 s
-1
for the unsubstituted Re(bpy) catalyst under analogous conditions. This study
13

demonstrated the importance of positioning of the pendant proton relative to the catalytic active
site at the metal center, as catalytic activity changed substantially when the position of the primary
amine was altered.

1.4.2.2 MANGANESE BIPYRIDINE CATALYSTS
Several manganese bipyridine catalysts with pendant amines have also been studied for the
electrocatalytic reduction of CO2. Manganese catalysts typically have different activity than their
rhenium counterparts, often more readily undergoing dimerization and other deactivation
processes and showing lower overall activity and selectivity for CO2 reduction. The difference in
behavior can be observed for catalysts for which both manganese and rhenium analogues have
been studied. Interestingly, a manganese bipyridine catalyst with a charged imidazolium group
and an available pendant proton displayed a slightly higher FECO (78%) than that observed for the
rhenium analogue, still at very mild reduction potentials.
63
Based on computational studies, the
catalytic activity at these potentials was attributed to the formation of a local hydration shell that
facilitated protonation of CO2 reduction intermediates during catalysis. Additionally, a series of
arylamine-substituted Mn(bpy) catalysts was investigated for the effects of pendant proton
positioning on catalytic activity, similar to the studies performed for the family of rhenium
catalysts.
64
As was observed for rhenium, the catalyst with a pendant amine in closest proximity
to the metal center displayed the greatest overall activity for CO2 reduction, displaying a nine-fold
increase in TOF compared to the unsubstituted manganese complex (901 and 102 s
-1
, respectively)
at an overpotential lower by 150 mV. While the catalytic rates observed were significantly higher
than those of the rhenium complexes, the FE values were overall lower (≥ 70%). A manganese
bipyridine catalyst with a local proton source in the form of phenyl diol substituents demonstrated
14

moderate CO2RR activity with a FECO of 70% with no added acid.
65
However, replacing the
pendant protons with methyl groups resulted in no catalytic activity without added acid, illustrating
the importance of the built-in local proton source on CO2 reduction, unlike the minimal effect
observed for the rhenium analogues. Further CPE studies resulted in a mixture of CO and formic
acid products, which was also not observed using rhenium, highlighting the difference in
selectivity by changing the metal center.
Manganese bipyridine complexes modified with tertiary amines demonstrated a
dependence of product selectivity on proximity of the amines in the secondary coordination sphere
to the metal center.
66
While these complexes do not contain pendant protons, the amines undergo
in-situ protonation, facilitating the formation of a manganese hydride intermediate that leads to
formic acid production. With a maximum TOF of 5500 s
-1
, high FEs, and mild overpotentials,
these catalysts show high activity for the conversion of CO2 to formic acid. A similar Mn(bpy)
complex with pendant methoxy groups demonstrates both hindered Mn-Mn dimerization due to
steric bulk and a second-sphere hydrogen bonding interaction that lowers the activation barrier for
C-OH bond cleavage during CO2RR.
67
Although this catalyst also does not contain pendant
protons, the methoxy groups are capable of forming weak hydrogen bonding interactions with both
the CO2 adduct and the added proton source, reducing the required overpotential by going through
a “protonation-first” reduction pathway while maintaining high FECO and selectivity over proton
reduction. A manganese catalyst with a pendant phenolic group also displays stabilization of the
CO2 adduct during catalysis, allowing for proton-assisted C-O bond cleavage with a lower
activation barrier.
68
While the observed catalytic current increase is seven times that of the
unsubstituted manganese complex, the overpotential is roughly the same, and FECO remains
moderate to high at an average of 76%.
15

1.5 HETEROGENIZATION OF MOLECULAR ELECTROCATALYSTS FOR THE
CARBON DIOXIDE REDUCTION REACTION
Heterogeneous catalysts are catalysts that are confined to an electrode surface rather than
involving the homogeneous diffusion of catalyst in solution. While typical heterogeneous catalysts
also have several disadvantages, including ill-defined active sites, limited surface area for the
catalytic reaction to take place, and issues with mass transport, heterogenization of molecular
species allows for the combination of the advantages of both homogeneous and heterogeneous
catalysts. Heterogeneous molecular catalysts come with their own set of advantages: often a lower
concentration of the molecular catalyst is necessary once immobilized on a surface; solubility in
organic solvents is not required for maintaining high catalytic activity; recyclability and separation
after catalysis are greatly increased; and steric confinement of the immobilized species reduces the
chances of deactivation via dimerization or aggregation.
69
Heterogenization can occur through a
variety of methods, including both covalent and non-covalent attachment of the molecular complex
to an electrode surface or periodic assembly into porous materials such as metal organic
frameworks (MOFs) or covalent organic frameworks (COFs).
70–72
Non-covalent immobilization
involves the interaction of a catalyst with a substrate through π-π interactions or space
confinement, while covalent immobilization involves surface grafting or polymerization via either
chemical or electrochemical methods. A variety of molecular catalysts for electrocatalytic CO2
reduction, such as porphyrins, phthalocyanines, cyclams, and bipyridines, have been
heterogenized using a range of substrates, but only the most pertinent involving heterogenization
of bipyridine-based complexes will be discussed herein.
69,73

16

1.5.1 HETEROGENEOUS RHENIUM BIPYRIDINE CATALYSTS
1.5.1.1 COVALENT ORGANIC FRAMEWORKS
There are a few examples in the literature of covalent organic frameworks (COFs)
incorporating rhenium bipyridine complexes. A rhenium 5,5’-diamino-2,2’-bipyridine complex
was integrated into a COF via a modified Schiff base condensation and post-synthetic metalation;
it was subsequently deposited onto glassy carbon electrodes as a composite with carbon black and
polyvinylidene fluoride (PVDF).
74
At an applied potential of -2.8 V vs Fc
+/0
, CO2 electrolysis
resulted in formation of CO with 81% FE. However, after 30 minutes of CPE, electrocatalytic
activity decreased due to inhibited mass transport and substrate diffusion. In another report, a
heterobimetallic COF containing both rhenium bipyridine and metalloporphyrin (Co and Fe)
moieties was also formed using a Schiff base reaction and post-metalation of the porphyrins.
75

Electrocatalytic CO2RR studies demonstrated higher activity for COFs with cobalt porphyrins,
although production of CO was low, with an FE of only 18%. This low activity was linked to a
competitive, rather than synergistic, catalytic relationship between the two metal centers.

1.5.1.2 NON-COVALENT IMMOBILIZATION
Non-covalent surface attachment mainly uses weak π–π interactions, which can lead to
catalyst leaching and loss of activity during long-term electrochemical experiments.
69
A rhenium
bipyridine complex modified with pyrene linkers was non-covalently attached to graphitic carbon
electrode surfaces.
76
Electrocatalytic reduction of CO2 resulted in selective CO production with a
FE of 70% and TON of 58. However, these modified electrodes did not maintain long-term
17

stability, and the current decreased to background levels within one hour of electrolysis due to loss
of rhenium from the surface.

1.5.1.3 COVALENT IMMOBILIZATION
Various techniques of covalent immobilization, including both chemical and
electrochemical methods, have been used to generate modified electrodes with rhenium bipyridine
moieties. Condensation of fac-Re(5,6-diamino-1,10-phenanthroline)(CO)3Cl to o-quinone edge
defects on glassy carbon electrodes generated graphite-conjugated rhenium (GCC-Re) catalysts
that demonstrated high activity for the reduction of CO2 to CO in acetonitrile. During electrolysis,
catalytic currents reached 50 mA/cm
2
, and CO was produced with 96% FE. However, the
requirement of defect sites within the graphitic carbon greatly limits the substrate scope as well as
the catalyst loading on the electrode surface.
Another useful method for immobilizing molecular catalysts to electrode surfaces is via
covalent attachment in the form of electropolymerization. A notable example of this is the
preparation of catalyst films by the electrochemical reduction of rhenium bipyridine catalysts with
polymerizable vinyl functional groups [(vbpy)Re(CO)3Cl; vbpy = 4-Me-4′-vinylbpy].
77–80
By
applying an electrochemical potential negative enough to reduce the ligand, vinyl−vinyl coupling
led to the formation of vinyl-linked surface-confined polymers. Though electrons are conducted
through the polymer in these types of films, catalytic activity is generally comparable to that of
the molecular complexes; however, due to the high concentration of active sites at the electrode
surface, catalytic reactivity is often even higher. For electropolymerized films of [(vbpy)Re-
(CO)3(CH3CN)]
+
with 0.1 M TBAH in acetonitrile (MeCN), CO2RR catalysis yielded CO as the
18

main product with 90% FE, although oxalate was also produced with 9% FE.
78
In co-polymerized
films with 30% [(vbpy)Re-(CO)3(CH3CN)]
+
and 70% Ru(II)(bpy) complex, no oxalate was
detected, suggesting that consecutive rhenium sites were necessary to drive the C-C coupling that
forms oxalate. This vinyl(bpy) complex was also electropolymerized onto nanocrystalline TiO2 on
glass, allowing for an increased number of redox active sites and enhanced catalytic activity due
to the nano-porous nature of the substrate.
81
However, one downside to this method of
electropolymerization is the inherent flexibility of the films due to the methylene spacers as well
as the unintended side-products, such as Re−Re dimers, formed from the vinyl radicals during
grafting, leading to instability and film degradation.
The catalytic activity of glassy carbon electrodes functionalized with fac-
Re(apbpy)(CO)3Cl [apbpy = 4-(4-aminophenyl)-2,2’-bipyridine] was investigated.
82
This
complex was electropolymerized via two approaches: oxidation of a terminal amino group,
forming C-N bonds, and reduction of a diazonium salt, forming C-C bonds. These modified
electrodes generated CO with TON ranging from 7 to 12 and FECO near unity, retaining their
properties for 24 hours in acetonitrile and with one week of air exposure. Unlike the amino
oxidation grafting method, the diazonium reduction method allowed for multilayers via successive
grafting scans and attained a longer overall lifetime. Graphene oxide was modified with fac-Re(4-
amino-bpy)(CO)3Cl by diazonium grafting and deposited onto glassy carbon electrodes for
electrocatalytic CO2 reduction.
83
Electrolysis yielded TOFs up to 4.44 s
-1
and syngas production
with a CO/H2 ratio of 7:5. Production of H2 was attributed to inherent catalytic activity of the
glassy carbon electrode. A study using thiophene-substituted Re(bpy) complexes for oxidative
electropolymerization onto glassy carbon electrodes for the CO2RR resulted in FECO values
ranging from 34 to 85% with TON of 489 to 519.
84
The lower activity within the series is explained
19

by the insertion of C≡C, which may react with CO2 when reduced or push the rhenium metal center
toward the solution, leading to behavior more like the molecular complex.
Several studies involving more rigid electropolymerized films with the aim of increasing
long-term stability were also reported. Electropolymerized films of an alkyne-substituted Re(bpy)
complex were generated on platinum-plate electrodes via radical coupling.
85
While counteracting
some of the previous issues with side product generation and lack of control with grafting, CO2
electrolysis resulted in a FECO of only 33%, although no film degradation was observed. In another
report, a [2,2′-bipyridine]-5,5′-bis(diazonium) rhenium complex was electropolymerized onto a
variety of electrode supports, yielding rigid, conjugated, structurally ordered film growth.
86

Electrochemical studies with high surface area modified graphite rod electrodes demonstrated the
dependence of catalytic activity on catalyst loading, up to an optimal value that occurred between
5 and 10 grafting scans. For the electrode modified with 5 scans, CO was produced with an FE of
99% and TON of 3606, and no dimer formation was observed during catalysis. Analogous films
were grown via electropolymerization on carbon cloth electrodes and studied for CO2RR.
87
These
carbon cloth films exhibited an 80-fold increase in catalytic activity over the modified graphite
electrodes, reaching TON per rhenium site of 290,000 with FECO >99%. This high catalytic activity
was attributed to a mechanism initiated by electrochemical charging of the conjugated backbone
followed by anion dissociation, CO2 coordination, and protonation.

1.6 OUTLINE OF THESIS
While much progress has been made in both homogeneous and heterogeneous catalyst
generation for CO2RR, there are still gaps in our knowledge concerning both function and
20

mechanism of catalysis. In particular, both molecular and immobilized rhenium bipyridine
electrocatalysts have shown great promise for selective production of CO with high activities. This
thesis will present further insights into this class of catalysts for CO2RR. Synthesis,
characterization, and electrochemical studies of novel rhenium bipyridine catalysts will be
discussed and compared, as well as the integration of these catalysts into heterogeneous systems.
Chapter two will discuss the development of a new family of pendant amine-substituted
rhenium bipyridine catalysts. These four catalysts are either mono- or bis-substituted in the 6- and
6’-positions with secondary or tertiary amines (methylamine and dimethylamine, respectively).
The difference in electrochemical behavior with the availability of a pendant proton for secondary
amine substituents is explored for a better understanding of the effect of these bio-inspired groups
on catalytic activity. Further differences in behavior based on mono- or bis-substitution are also
examined.
Chapter three delves deeper into the effects of available pendant protons on electrocatalytic
activity. The previously discussed family of 6,6’-amine substituted rhenium bipyridine catalysts is
expanded to include bis-substituted primary, secondary, and tertiary amines. The similarities
between the NH2- and NHMe-substituted complexes with pendant proton donors are discussed,
and additional electrochemical studies with the NHMe-substituted complex are presented for
evidence of the formation of hydrogen bonding networks. The NH2-substituted complex displays
an interesting dependence of FECO on electrochemical potential, which is not observed for the
other catalysts within the family.
In chapter four, another new primary amine-substituted rhenium bipyridine catalyst is
presented, and its electrocatalytic behavior is explored. Along with additional studies of the
previously reported 5,5’-NH2 complex, comparisons are made between the three NH2-substituted
21

catalysts to determine the effect of substituent and pendant proton position on catalytic activity.
Evidence of a more positive catalytic onset potential supports the favorable positioning of the NH2
group near the rhenium metal center, although the 4,4’-substitued complex also displays an
interesting dependence of FECO on electrochemical potential, suggesting that the pendant protons
do still play a role in catalysis.
Finally, chapter five discusses the heterogenization of the 4,4’-NH2 rhenium bipyridine
catalyst, based on previous studies involving the 5,5’-NH2 analogue. A new 4,4’-diazonium
complex is presented and utilized for electropolymerization onto a variety of electrode substrates.
Heterogeneous characterization techniques and electrochemical studies are used for a comparison
of the resulting films to the previously reported molecular wires. While these studies are still
preliminary, the data suggests that grafting with the 4,4’-substituted complex retains a lower
amount of control on film thickness and thus non-linear trends in electrochemical behavior based
on the number of grafting scans.

22

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30

CHAPTER 2

Rhenium bipyridine catalysts with hydrogen bonding pendant amines for CO2 reduction

A portion of this chapter has appeared in print and is reproduced with permission from The Royal
Society of Chemistry:
Hellman, A. N.; Haiges, R.; Marinescu, S. C.* Rhenium bipyridine catalysts with hydrogen
bonding pendant amines for CO2 reduction. Dalton Transactions. 2019, 48, 14251-14255.
31

2.1 ABSTRACT
Rhenium tricarbonyl bipyridine complexes modified with pendant secondary and tertiary
amines in the 6- and 6'- positions were synthesized and characterized. Electrocatalytic studies
performed under CO2 with 2,2,2,-trifluoroethanol display large current densities, corresponding
to the reduction of CO2 to CO with moderate Faradaic efficiencies (51-73%).

2.2 INTRODUCTION
The electrochemical conversion of CO2 to valuable chemicals and fuels is an important
step toward creating a carbon-neutral infrastructure.
1
In nature, the enzyme CO-dehydrogenase
(CODH) catalyzes the selective conversion of CO2 to CO with high activity.
2–4
The CO2 activation
is facilitated by hydrogen bonding interactions between the bound substrate and the amino acid
residues present in the secondary coordination sphere.
2–4
Due to the high thermodynamic barrier
associated with the reduction of CO2 to CO, a catalyst capable of reducing CO2 via proton-coupled
electron transfer (PCET) steps is attractive for lowering the required energy.
5–7
A variety of bio-
inspired CO2-reduction electrocatalysts with pendant proton donors have been reported, such as
nickel cyclams with pendant amines
8–10
, iron porphyrins with pendant phenolic,
11

trimethylanilinium,
12
and amide
13
groups, and cobalt aminopyridine macrocycles with pendant
amines.
14,15

Metal bipyridine complexes, including Re(bpy)(CO)3Cl and Mn(bpy)(CO)3Cl (bpy = 2,2'-
bipyridine), are some of the most well-studied classes of electrocatalysts for CO2 reduction.
16–19

In particular, Re(bpy)(CO)3Cl has been shown to have good selectivity for the reduction of CO2
to CO, although with low catalytic rate constants and deactivation via dimerization. However, the
performance of Mn and Re bipyridine catalysts can be modulated through the addition of pendant
32

groups capable of hydrogen bonding interactions. Manganese bipyridine complexes modified with
pendant phenolic,
20
methoxy,
21
and imidazolium
22
groups exhibit increased current enhancements
at more positive potentials than Mn(bpy)(CO)3Cl. We have previously reported that a rhenium
bipyridine catalyst modified with pendant amines in the 5- and 5'- positions converts CO2 to CO
with a 99% Faradaic efficiency (FE), although at more negative potentials than that of other
reported rhenium bipyridine catalysts.
23
Bio-inspired rhenium bipyridine catalysts modified with
amino acid substituents in the 4- and 4'- positions undergo both inter- and intramolecular hydrogen
bonding interactions, leading to current increases at approximately 250 mV more positive
potentials than that of Re(bpy)(CO)3Cl, due to the promotion of an alternative bimolecular
mechanism.
24
Rhenium bipyridine complexes bearing di- and tri-phenolic groups display current
enhancements at both the first and second reduction potentials under CO2 with added H2O, but
undergo electrode passivation during catalysis, yielding turnover numbers (TON) between 2 and
14.
25
An aminophenethyl-modified rhenium bipyridine catalyst exhibits a current increase near the
second reduction potential under CO2, with near unity FE and selectivity for CO production and a
TON of 6.
26
Additionally, a rhenium bipyridine complex featuring an imidazolium group shows
current increases under CO2 at a potential 170 mV more positive than that of Re(bpy)(CO)3Cl.
27

This increased activity was attributed to hydrogen bonding second-sphere interactions present in
the imidazolium catalyst.
27
Motivated by the prior work on bio-inspired rhenium bipyridine
catalysts, we report here a family of substituted rhenium bipyridine catalysts with pendant
secondary and tertiary amines and explore their electrochemistry for CO2 reduction.

33

2.3 RESULTS AND DISCUSSION
2.3.1 Synthesis and Characterization

Scheme 2.1 Synthetic scheme for complexes 1–4.

Amine-substituted bipyridine ligands L
1
–L
4
were synthesized using a modified Ullmann
coupling procedure (Scheme 2.1). Ligands L
1
and L
2
display a broad resonance at  4.55 and 4.58
ppm, respectively, corresponding to the NH moieties (Figure 2.1). Rhenium complexes 1–4 were
generated by refluxing L
1
–L
4
with Re(CO)5Cl in anhydrous toluene for 12 h. The
1
H NMR
resonances of complexes 1–4 are shifted downfield in comparison to the corresponding resonances
of ligands L
1
–L
4
, as expected upon addition of the electron withdrawing rhenium metal center
(Figure 2.2). Complexes 1 and 2 each display a broad resonance at  6.23 and 6.09 ppm,
respectively, corresponding to the NH moieties.
34

Figure 2.1
1
H NMR spectra in CDCl3 of (a) ligand L
1
(400 MHz), (b) ligand L
2
(400 MHz), (c)
ligand L
3
(500 MHz), and (d) ligand L
4
(400 MHz).
35

Figure 2.2
1
H NMR spectra in CDCl3 of (a) complex 1 (400 MHz), (b) complex 2 (500 MHz),
(c) complex 3 (400 MHz), and (d) complex 4 (500 MHz).

36

Single crystal X-ray diffraction studies of complexes 1–4 reveal equatorial coordination of
the bipyridine ligand and a facial arrangement of the three carbonyl moieties, typical of other
Re(bpy)(CO)3 complexes (Figures 2.3-2.4). Due to low quality diffraction data for complex 3,
bond lengths and angles will only be discussed for complexes 1, 2, and 4. The Re–N(pyridine)
bond lengths in complexes 1 and 2 range between 2.18 and 2.20 Å, and are analogous to those
observed for the Re(bpy)(CO)3Cl (2.17 Å).
28
However, the Re–N bond lengths in complex 4 are
de-symmetrized: one is slightly shorter (2.162(2) Å), and one is elongated to 2.226(2) Å, due to
the presence of the bulky dimethylamino group. Further, while the bipyridine ligand in complexes
1 and 2 are essentially planar, with bipyridine N—C—C—N torsion angles of only 0.6(3)° and
1.3(3)°, respectively, complex 4 has a torsion angle of 12.2(2)° (Figures 2.5-2.6). Crystal structure
refinement data and select bond lengths are presented in Tables 2.1-2.5.

Figure 2.3 Solid state structures of (a) 1, (b) 2, (c) 3, and (d) 4. Color legend of the atoms: gray –
C; blue – N; red – O; green – Cl; pink – Re. Solvent molecules and hydrogen atoms are excluded
for clarity.

37

Figure 2.4 Solid state structures of (a) 1 and (b) 2 with N-H bonds included. Color legend of the
atoms: gray – C; blue – N; red – O; green – Cl; pink – Re; purple – H. Solvent molecules and
additional hydrogen atoms are excluded for clarity.

Figure 2.5 Solid state structures of (a) 1, (b) 2, and (c) 4 showing labels for the nitrogen and carbon
atoms involved in the N—C—C—N torsion angles. Color legend of the atoms: gray – C; blue –
N; red – O; green – Cl; pink – Re. Solvent molecules and additional hydrogen atoms are excluded
for clarity.

38

Figure 2.6 Solid state structures of (a) 1, (b) 2, and (c) 4 showing torsion angles of the N—C—
C—N bonds. Color legend of the atoms: gray – C; blue – N; red – O; green – Cl; pink – Re.
Solvent molecules and hydrogen atoms are excluded for clarity.

Table 2.1 Crystal data and structure refinement for 1.

Chemical formula C18H21ClN5O4Re
Formula weight 593.05 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.085  0.163  0.168 mm
Crystal habit clear light yellow prism
Crystal system triclinic
Space group
P1
¯

Unit cell dimensions a = 8.953(3) Å α = 80.373(4)°

b = 9.085(3) Å β = 76.566(4)°

c = 13.399(4) Å γ = 78.668(4)°
Volume 1030.9(5) Å
3

Z 2
Density (calculated) 1.911 g/cm
3

Absorption coefficient 6.058 mm
-1

F(000) 576
Diffractometer Bruker APEX DUO
Radiation source fine-focus tube, MoKα
Theta range for data collection 1.58 to 30.59°
39

Index ranges -12 ≤ h ≤ 12, -12 ≤ k ≤ 12, -19 ≤ l ≤ 19
Reflections collected 25548
Independent reflections 6210 [R(int) = 0.0470]
Coverage of independent reflections 98.0%
Absorption correction multi-scan
Max. and min. transmission 0.6270 and 0.4290
Structure solution technique direct methods
Structure solution program SHELXTL XT 2014/4 (Bruker AXS, 2014)
Refinement method Full-matrix least-squares on F
2

Refinement program SHELXL-2014/6 (Sheldrick, 2014)
Function minimized Σ w(Fo
2
- Fc
2
)
2

Data / restraints / parameters 6210 / 2 / 272
Goodness-of-fit on F
2
1.068
Δ/σmax 0.002
Final R indices 5716 data; I > 2σ(I) R1 = 0.0219, wR2 =
0.0488

All data R1 = 0.0259, wR2 =
0.0501
Weighting scheme w=1/[σ
2
(Fo
2
)+(0.0147P)
2
+0.5921P]
where P = (Fo
2
+2Fc
2
)/3
Largest diff. peak and hole 0.972 and -1.081 eÅ
-3

R.M.S. deviation from mean 0.130 eÅ
-3

Table 2.2 Crystal data and structure refinement for 2.
Identification code Ashley012918
Chemical formula C17H18ClN4O4Re
Formula weight 564.00 g/mol
Temperature 103(2) K
Wavelength 0.71073 Å
Crystal size 0.089  0.258  0.322 mm
Crystal habit clear light yellow prism
Crystal system triclinic
Space group
P1
¯

Unit cell dimensions a = 8.912(3) Å α = 84.241(5)°

b = 8.957(3) Å β = 87.294(5)°

c = 12.030(4) Å γ = 79.875(5)°
Volume 940.1(5) Å
3

Z 2
Density (calculated) 1.992 g/cm
3

Absorption coefficient 6.636 mm
-1

F(000) 544
40

Diffractometer Bruker APEX DUO
Radiation source fine-focus tube, MoKα
Theta range for data collection 1.70 to 30.60°
Index ranges -12 ≤ h ≤ 12, -12 ≤ k ≤ 12, -16 ≤ l ≤ 17
Reflections collected 23077
Independent reflections 5677 [R(int) = 0.0488]
Coverage of independent reflections 97.9%
Absorption correction multi-scan
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 2016/6 (Bruker AXS, 2016)
Function minimized Σ w(Fo
2
- Fc
2
)
2

Data / restraints / parameters 5677 / 0 / 250
Goodness-of-fit on F
2
1.072
Δ/σmax 0.003
Final R indices 5501 data; I >
2σ(I)
R1 = 0.0198, wR2 = 0.0495

all data R1 = 0.0206, wR2 = 0.0499
Weighting scheme w=1/[σ
2
(Fo
2
)+(0.0150P)
2
+0.7497P]
where P=(Fo
2
+2Fc
2
)/3
Largest diff. peak and hole 0.903 and -1.492 eÅ
-3

R.M.S. deviation from mean 0.135 eÅ
-3

Table 2.3 Crystal data and structure refinement for 3.

Identification code Ashley_bisNMe2
Chemical formula C18H19Cl4N4O3Re
Formula weight 667.37 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.020  0.100  0.130 mm
Crystal system triclinic
Space group
P1
¯

Unit cell dimensions a = 9.368(5) Å α = 75.978(9)°

b = 11.275(6) Å β = 76.488(8)°

c = 11.587(6) Å γ = 81.595(8)°
Volume 1149.3(11) Å
3

Z 2
Density (calculated) 1.928 g/cm
3

Absorption coefficient 5.778 mm
-1

F(000) 644
41

Diffractometer Bruker APEX DUO
Radiation source fine-focus tube (MoKα , λ = 0.71073 Å)
Theta range for data collection 1.85 to 28.07°
Reflections collected 7978
Coverage of independent reflections 94.9%
Absorption correction multi-scan
Max. and min. transmission 0.8890 and 0.5090
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 7978 / 261 / 277
Goodness-of-fit on F
2
1.028
Final R indices 5841 data; I >
2σ(I)
R1 = 0.1016, wR2 = 0.2579

all data R1 = 0.1429, wR2 = 0.2963
Weighting scheme w=1/[σ
2
(Fo
2
)+(0.1896P)
2
+29.3759P]
where P=(Fo
2
+2Fc
2
)/3
Largest diff. peak and hole 11.085 and -6.164 eÅ
-3

R.M.S. deviation from mean 0.526 eÅ
-3

Table 2.4 Crystal data and structure refinement for 4.

Identification code Ashley082218
Chemical formula C15H13ClN3O3Re
Formula weight 504.93 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.200  0.242  0.352 mm
Crystal habit clear orange prism
Crystal system triclinic
Space group
P1
¯

Unit cell dimensions a = 7.333(2) Å α = 91.690(4)°

b = 8.821(3) Å β = 99.041(4)°

c = 13.059(4) Å γ = 107.903(4)°
Volume 791.1(4) Å
3

Z 2
Density (calculated) 2.120 g/cm
3

Absorption coefficient 7.866 mm
-1

F(000) 480
42

Table 2.5 Nitrogen–rhenium bond lengths (in Ångstroms) for complexes 1, 2, and 4.

Complex Re—N1(pyridine) Re—N2(pyridine)
1 2.194(2) 2.202(2)
2 2.175(2) 2.200(2)
4 2.1624(17) 2.2255(17)

Complexes 1–4 were further characterized by FT-IR spectroscopy. Three carbonyl
stretches are observed for each complex: one high-energy mode (a1') and two lower-energy modes
(a" and a2'), as expected for fac-Re(CO)3 complexes (Figure 2.7 and Table 2.6).
29,30
The a1' mode
displays a slight red-shift from 2011 to 2013 cm
–1
in complexes 1 and 3, respectively, and from
2014 to 2016 cm
–1
in complexes 2 and 4, respectively, as expected upon increasing the electron
Diffractometer Bruker APEX DUO
Radiation source fine-focus tube, MoKα
Theta range for data collection 1.58 to 30.64°
Index ranges -10 ≤ h ≤ 10, -12 ≤ k ≤ 12, -18 ≤ l ≤ 18
Reflections collected 19445
Independent reflections 4782 [R(int) = 0.0340]
Coverage of independent reflections 97.8%
Absorption correction multi-scan
Max. and min. transmission 0.3020 and 0.1680
Structure solution technique direct methods
Structure solution program SHELXTL XT 2014/4 (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 4782 / 0 / 210
Goodness-of-fit on F
2
1.087
Δ/σmax 0.003
Final R indices 4686 data; I > 2σ(I) R1 = 0.0157, wR2 = 0.0381

all data R1 = 0.0162, wR2 = 0.0384
Weighting scheme w=1/[σ
2
(Fo
2
)+(0.0114P)
2
+0.5504P]
where P=(Fo
2
+2Fc
2
)/3
Largest diff. peak and hole 0.947 and -1.237 eÅ
-3

R.M.S. deviation from mean 0.117 eÅ
-3

43

density of the bipyridine substituents. This difference, although minimal, indicates a slightly
greater electron density at the rhenium center for the bis-substituted complexes compared to the
mono-substituted complexes. Complexes 1 and 2 also display bands at 3416 and 3391 cm
–1
,
respectively, corresponding to the N-H stretches (Table 2.6).

Figure 2.7 FT-IR spectra of complex 1 (red), complex 2 (orange), complex 3 (green), and
complex 4 (blue).

Table 2.6 Pertinent IR spectroscopy wavenumbers for complexes 1–4.

Complex N-H C=O (a1') C=O (a") C=O (a2')
1 3416 2011 1897 1857
2 3391 2014 1915 1869
3 – 2013 1903 1876
4 – 2016 1899 1880

2.3.2 Cyclic Voltammetry
Cyclic voltammetry (CV) experiments of complexes 1–4 under N2 display either a
reversible or quasi-reversible first reduction feature followed by an irreversible second reduction
feature (Figure 2.8). All potentials are listed versus Fc
+/0
. Complexes 1 and 3 exhibit the first
44

reduction feature at -1.96 V and -1.92 V, respectively, and the second reduction feature at -2.25 V
and -2.26 V, respectively (Table 2.7). Complexes 2 and 4 display the first reduction feature at -
1.87 V and -1.84 V, respectively, and the second reduction feature at -2.16 V and -2.19 V,
respectively. The reduction potentials of complexes 1 and 3 are slightly more negative in
comparison to those of complexes 2 and 4 due to an increased number of electron-donating groups
(NMe). All four complexes display reduction features at more negative potentials than the
unsubstituted complex, Re(bpy)(CO)3Cl, which has reduction features at -1.75 V and -2.12 V
under identical conditions. Randles-Sevcik plots indicate that complexes 1–4 are all freely
diffusing in solution (Figure 2.9).

Figure 2.8 CVs of 1 mM 1 (red), 2 (orange), 3 (green), 4 (blue), and Re(bpy)(CO)3Cl (gray) under
N2 in a 0.1 M TBAPF6 MeCN solution. Dashed lines illustrate the first reduction feature, and dotted
45

lines (1 and 2) show the presence of an additional intermediate reduction feature. Scan rate: 100
mV/s.

46

Figure 2.9 Randles-Sevcik plots for the first (left) and second (right) reductions of (a) 1, (b) 2,
(c) 3, and (d) 4; the linear correlation indicates that the species is freely diffusing based on the
Randles-Sevcik equation.

Table 2.7 Reduction potentials (V vs. Fc
+/0
) for complexes 1 through 4 under N2. Scan rate: 100
mV/s.
Catalyst 1
st
reduction Add’l reduction 2
nd
reduction
1 -1.96 -2.05 -2.25
2 -1.87 -2.01 -2.16
3 -1.92 – -2.26
4 -1.84 – -2.19

Complexes 1 and 2 both exhibit an additional reduction feature at -2.05 and -2.01 V,
respectively, at 100 mV/s (Figure 2.8). Variable scan rate studies of complexes 1 and 2 show that
this additional reduction feature disappears at faster scan rates (above 400 mV/s), suggesting that
this reduction event is coupled to a chemical process which proceeds too slowly to be observed on
the CV timescale (Figure 2.10). This chemical process is assigned to Re—Re dimerization, which
has been reported previously in analogous rhenium bipyridine catalysts.
31
In contrast, complexes
3 and 4, which lack NH moieties, do not display this additional feature even at low scan rates. This
behavior suggests that dimerization is facilitated by the presence of pendant NH moieties, which
may engage in hydrogen bonding interactions. Because MeCN is a poor hydrogen bond donor and
acceptor, the solute—solute hydrogen bonding interactions that facilitate Re—Re dimerization are
expected to be more prevalent in MeCN than in DMF.
32
To further probe the proposed hydrogen
bonding interactions, the electrochemistry of 2 was explored in DMF. CVs of 2 in DMF under N2
exhibit reduction features at -1.91 and -2.28 V, consistent with the two reduction features typically
observed for rhenium bipyridine complexes (Figure 2.11). A third irreversible feature at -2.45 V
is observed and assigned to the dimerization. Interestingly, this reduction appears at a potential
47

more negative than the additional feature in MeCN, suggesting that DMF can disrupt the hydrogen
bonding interactions and reduce the extent of hydrogen bonding-facilitated dimerization.

Figure 2.10 Variable scan rate studies of (a) 1, (b) 2, (c) 3, and (d) 4. Conditions: 1 mM catalyst
in MeCN with 0.1 M TBAPF6.

Figure 2.11 First and second reduction potentials of 2 in DMF. 1 mM catalyst in DMF with 0.1
M TBAPF6, scan rate: 100 mV/s.
48

Upon switching the gas from N2 to CO2 (1 atm), complexes 1–4 exhibit current
enhancements at either the same or more negative potentials than their second reduction features
(Figure 2.12). Bis-substituted complexes 1 and 3 exhibit smaller current increases under CO2 than
mono-substituted complexes 2 and 4. Notably, complex 4 displays a drastic increase in current
under CO2 saturation, reaching 9 mA/cm
2
. Additionally, complex 4 displays trace crossing of the
forward and reverse scans, indicating substrate depletion in the diffusion layer on the CV
timescale.
33,34

Figure 2.12 Cyclic voltammograms of (a) 1, (b) 2, (c) 3, and (d) 4 under N2 (black), CO2 (red),
N2 + 2 M TFE (blue), and CO2 + 2 M TFE (green). Conditions: 1 mM catalyst in MeCN with 0.1
M TBAPF6. Scan rate: 100 mV/s.

The current response upon addition of a Brønsted acid (water, 2,2,2-trifluoroethanol, or
phenol) was analyzed for each complex (Figure 2.12 and 2.13). While titrations with water (H2O)
49

led to current increases near the potentials of the second reduction features, addition of 2,2,2-
trifluoroethanol (TFE) or phenol (PhOH) caused substantial current enhancements at potentials
near the first reduction features. Moreover, TFE provided slightly greater current responses than
PhOH, so this was chosen as the optimal acid source for further studies (Figure 2.13). CVs
performed in the presence of TFE under N2 show only minimal increases in current densities,
indicating that the observed current responses are not due to proton reduction (Figure 2.12).

Figure 2.13 Cyclic voltammograms of (a) 1, (b) 2, (c) 3, and (d) 4 under CO2 with added PhOH
(orange), TFE (green), and H2O (blue). Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6.
Scan rate: 100 mV/s.

2.3.3 Controlled Potential Electrolysis

Controlled potential electrolysis (CPE) experiments were performed for 1 hour under CO2
at the potentials of maximum current density for complexes 1–4 (Figure 2.14). Gas
50

chromatography (GC) analyses of the electrolysis cell headspace confirmed production of CO for
complexes 1–4 (Table 2.8). CPE experiments with mono-substituted complexes 2 and 4 performed
at -2.06 and -2.32 V produced 157 and 190 μmol CO with Faradaic Efficiencies (FE) of 73 and
58%, respectively. The amounts of CO produced by complexes 2 and 4 are comparable to that of
Re(bpy)(CO)3Cl (183 μmol CO and 86% FE at -2.20 V) under identical conditions. However,
complex 2 attains these values at a slightly more positive potential (-2.06 V) than complex 4 and
Re(bpy)(CO)3Cl (-2.32 and -2.20 V, respectively). CPE experiments with bis-substituted
complexes 1 and 3 at -2.10 and -2.32 V produce lower amounts of CO (55 and 27 μmol CO) with
51 and 53% FE, respectively. No other products, such as hydrogen or formic acid, were detected
for any of these catalysts, and wash tests indicate that no species active for CO2 reduction deposit
on the electrode during CPE. Total turnover numbers (TON) from CPE were calculated for
complexes 1–4 and Re(bpy)(CO)3Cl according to Eq. 2.1:
𝑚𝑜𝑙 𝐶𝑂 𝑝𝑟𝑜𝑑𝑢𝑐𝑒𝑑 𝑚𝑜𝑙 𝑐𝑎𝑡𝑎𝑙𝑦𝑠𝑡 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜 𝑛

where mol catalyst in solution is 1 mM = 0.00004 mol in 40 mL solution. Although the values
are low, all catalysts except complex 3 produce greater than one equivalent of CO per mole of
catalyst in the electrolysis solution (Table 2.8). Electrolyses of complexes 1–4 and TFE under N2
display low current densities and no CO production; additionally, UV-vis spectra taken before
and after electrolysis under CO2 support retention of the original structures, suggesting that
decomposition and loss of carbonyl ligands is not a contributing factor toward the production of
CO (Figure 2.15).
51

Figure 2.14 Current versus time over one hour of controlled potential electrolysis for complexes
1 (red), 2 (orange), 3 (green), and 4 (blue). BE studies were performed using 1 mM catalyst in
MeCN with 0.1 M TBAPF6.

Table 2.8 Catalytic parameters for complexes 1–4 and Re(bpy)(CO)3Cl (1 mM) under CO2 in
MeCN with 0.1 M TBAPF6 and TFE (2 M for complexes 1, 2, 4, and Re(bpy)(CO)3Cl and 0.8 M
for complex 3).

Complex Ecat (V vs. Fc
+/0
) FECO (%) μmol CO TON icat/ip (100 mV/s)
1 -2.10 51 55 1.4 8.2
2 -2.06 73 157 3.9 18.3
3 -2.32 53 27 0.7 22.3
4 -2.32 58 190 4.8 29.5
Re(bpy)(CO)3Cl -2.20 86 183 4.6 29.1

In the absence of TFE as an acid source, only complex 4 produced CO (44 μmol with a FE
of 18%), and no other products were detected for any of the catalysts (Figure 2.16). Use of PhOH
as the Brønsted acid for complex 2 led to steady current decreases and no formation of CO2
reduction products. Further, use of DMF as the solvent with TFE as the acid source for complex 2
52

led to a stable current response during CPE and a FE for CO analogous to that with MeCN as the
solvent (71%), with 134 μmol CO produced. CPE experiments performed in the absence of the
catalysts display negligible current responses and CO production, indicating that the glassy carbon
working electrode is not responsible for the observed current responses during catalysis.

Figure 2.15 UV-vis before and after CPE of the working solution of the electrolysis cell for (a)
complex 3 and (b) complex 4.

Figure 2.16 Cyclic voltammograms of a blank cell and 4 before controlled potential electrolysis
with no added TFE, 1 M TFE, and 2 M TFE, showing a dependence of catalytic onset and current
response on acid concentration.

As CPE studies confirmed the reduction of CO2 to CO by complexes 1–4, CV data recorded
with added TFE were used to determine their normalized peak current values (icat/ip, where icat =
53

the peak current density under catalytic conditions, and ip = the peak current density of the catalysts
under N2). These icat/ip values were used to estimate the activity of the rhenium catalysts.
Complexes 1–4 display normalized peak current values of 8.2, 18.3, 22.3, and 29.5, respectively
(Table 2.8). The icat/ip and value for complex 4 is comparable to that of Re(bpy)(CO)3Cl (29.1)
under identical conditions, suggesting similar catalytic activity. However, the calculated icat/ip
values are overall lower than those reported for other substituted rhenium bipyridine catalysts,
30

suggesting that these complexes undergo slower catalysis.

2.4 CONCLUSION
In summary, a series of rhenium bipyridine complexes modified in the 6- and 6'- positions
with secondary and tertiary amines were synthesized and characterized by
1
H NMR, FT-IR, and
single crystal X-ray diffraction. Single crystal X-ray diffraction studies of complexes 1–4 confirm
the facial arrangement of the three carbonyl moieties, typical of other Re(bpy)(CO)3 complexes.
Cyclic voltammograms under N2 display an additional reduction feature at slow scan rates for
complexes 1 and 2, which contain NH moieties. These additional features were attributed to
hydrogen-bonding dimerization due to their disappearance in DMF. Under CO2, complexes 1–4
display irreversible current enhancements at the second reduction feature. Titrations with TFE lead
to greater current enhancements at the first reduction feature, attaining current densities up to 12
mA/cm
2
for complex 4. During CPE experiments, mono-substituted complexes 2 and 4 performed
analogously to Re(bpy)(CO)3Cl under identical conditions, producing similar volumes of CO,
whereas bis-substituted complexes 1 and 3 produced only small amounts of CO. Under optimized
conditions, the Faradaic efficiencies for all four complexes were moderate, ranging from 51 to
73% CO produced, with no other quantifiable products observed.
54

2.5 EXPERIMENTAL METHODS AND ADDITIONAL FIGURES
2.5.1 Materials and Synthesis
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres drybox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. All solvents were degassed with nitrogen and passed through activated
alumina columns and stored over 4Å Linde-type molecular sieves. Deuterated solvents were dried
over 4Å Linde-type molecular sieves prior to use. Proton NMR spectra were acquired at room
temperature using Varian (Mercury 400 2-Channel, VNMRS-500 2-Channel, VNMRS- 600 3-
Channel, and 400-MR 2-Channel) spectrometers and referenced to the residual
1
H resonances of
the deuterated solvent (
1
H: CDCl3) and are reported as parts per million relative to
tetramethylsilane. Elemental analyses were performed using Thermo Scientific™ FLASH 2000
CHNS/O Analyzers. All the chemical reagents were purchased from commercial vendors and used
without further purification.

2.5.2 Electrochemistry
Electrochemistry experiments were carried out using a Pine potentiostat. The experiments
were performed in a single compartment electrochemical cell under nitrogen or CO 2 atmosphere
using a 3 mm diameter glassy carbon electrode as the working electrode, a platinum wire as
auxiliary electrode and a silver wire as the reference electrode. Ohmic drop was compensated using
the postive feedback compensation implemented in the instrument. All electrochemical
experiments were performed with iR compensation using the current interrupt (RUCI) method in
AfterMath. Typical values for the cell resistance were around 0.16-0.17 ohms. All experiments in
this paper were referenced relative to ferrocene (Fc) with the Fe
3+/2+
couple at 0.0 V. Alternatively,
55

in cases when the redox couple of ferrocene overlapped with other features of interest,
decamethylferrocene (Fc*) was used as an internal standard with the Fe*
3+/2+
couple at –0.48 V.
All electrochemical experiments were performed with 0.1 M tetrabutylammonium
hexafluorophosphate as supporting electrolyte. The concentrations of the rhenium catalysts were
generally at 1 mM and experiments with CO2 were performed at gas saturation in acetonitrile
(MeCN) or dimethylformamide (DMF).
Controlled-potential electrolysis measurements were conducted in a two-chamber H cell.
The first chamber held the working and reference electrodes in 40 mL of 0.1 M
tetrabutylammonium hexafluorophosphate and TFE in MeCN or DMF. The second chamber held
the auxiliary electrode in 25 mL of 0.1 M tetrabutylammonium hexafluorophosphate in MeCN or
DMF. The two chambers were separated by a fine porosity glass frit. The reference electrode was
placed in a separate compartment and connected by a Vycor tip. Glassy carbon plate electrodes (6
cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and auxiliary electrodes. 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. Faradaic efficiencies were determined by diving the
measured CO produced by the amount of CO expected based on the charge passed during the bulk
electrolysis experiment. For each species the controlled-potential electrolysis measurements were
performed at least twice, leading to similar behavior. The reported Faradaic efficiencies and mmol
of CO produced are average values.
2.5.3 Synthesis of N
6
,N
6
'-dimethyl-[2,2'-bipyridine]-6,6'-diamine (L
1
)
A high pressure flask was charged with 6,6'-dibromo-2,2'-bipyridine (105 mg, 0.33 mmol), CuI
(6.3 mg, 0.033 mmol), L-proline (7.8 mg. 0.067 mmol), and Na2CO3 (110 mg, 1 mmol).
56

Subsequently, 1 mL DMSO and 0.1 mL water were added, and the solution was purged under N 2
for five minutes while stirring. Methylamine (2 mL, 50 mM) was added via syringe, and the flask
was sealed. The reaction mixture was heated to 130 °C and allowed to stir for 24 hours. After
cooling, the organic phase was extracted with ethyl acetate and washed with copious amounts of
water. The organic phase was dried over Na2SO4, and the solvent was removed by rotary
evaporation. The product was obtained as a yellow-brown solid in approximately 70% yield.
1
H
NMR (CDCl3):  7.63 (d, 1H, NC5H3), 7.54 (t, 1H, p-NC5H3), 6.40 (d, 1H, NC5H3), 4.55 (s, 1H,
NH), 2.98 (d, 3H, NCH3) ppm.

2.5.4 Synthesis of N-methyl-[2,2'-bipyridine]-6-amine (L
2
)
A high pressure flask was charged with 6-bromo-2,2'-bipyridine (160 mg, 0.67 mmol), CuI (6.3
mg, 0.033 mmol), L-proline (7.8 mg, 0.067 mmol), and Na2CO3 (110 mg, 1 mmol). Subsequently,
1 mL DMSO and 0.1 mL water were added, and the solution was purged under N2 for five minutes
while stirring. Methylamine (2 mL, 50 mM) was added via syringe, and the flask was sealed. The
reaction mixture was heated to 130 °C and allowed to stir for 48 hours. After cooling, the organic
phase was extracted with ethyl acetate and washed with copious amounts of water. The organic
phase was dried over Na2SO4, and the solvent was removed by rotary evaporation. The product
was obtained as a yellow-brown solid in approximately 72% yield.
1
H NMR (CDCl3):  8.65 (d,
1H, NC5H3), 8.33 (d, 1H, NC5H3), 7.78 (t, 1H, p-NC5H3), 7.68 (d, 1H, NC5H3), 7.58 (t, 1H, p-
NC5H3), 6.45 (d, 1H, NC5H3), 4.58 (s, 1H, NH), 3.01 (d, 3H, NCH3) ppm.

57

2.5.5 Synthesis of N
6
,N
6
,N
6'
,N
6'
-tetramethyl-[2,2'-bipyridine]-6,6'-diamine (L
3
)
A high pressure flask was charged with 6,6'-dibromo-2,2'-bipyridine (105 mg, 0.33 mmol), CuI
(6.3 mg, 0.033 mmol), L-proline (7.8 mg, 0.067 mmol), and Na2CO3 (110 mg, 1 mmol).
Subsequently, 1 mL DMSO and 0.1 mL water were added, and the solution was purged under N 2
for five minutes while stirring. Dimethylamine (2 mL, 30 mM) was added via syringe, and the
flask was sealed. The reaction mixture was heated to 130 °C and allowed to stir for 48 hours. After
cooling, the organic phase was extracted with ethyl acetate and washed with copious amounts of
water. The organic phase was dried over Na2SO4, and the solvent was removed by rotary
evaporation. The product was obtained as an orange-brown solid in approximately 56% yield.
1
H
NMR (CDCl3):  7.71 (d, 1H, NC5H3), 7.56 (t, 1H, p-NC5H3), 6.53 (d, 1H, NC5H3), 3.16 (s, 6H,
N(CH3)2) ppm.

2.5.6 Synthesis of N,N-dimethyl-[2,2'-bipyridine]-6-amine (L
4
)
A high pressure flask was charged with 6-bromo-2,2'-bipyridine (160 mg, 0.67 mmol), CuI (6.3
mg, 0.033 mmol), L-proline (7.8 mg, 0.067 mmol), and Na2CO3 (110 mg, 1 mmol). Subsequently,
1 mL DMSO and 0.1 mL water were added, and the solution was purged under N2 for five minutes
while stirring. Dimethylamine (2 mL, 30 mM) was added via syringe, and the flask was sealed.
The reaction mixture was heated to 130 °C and allowed to stir for 48 hours. After cooling, the
organic phase was extracted with ethyl acetate and washed with copious amounts of water. The
organic phase was dried over Na2SO4, and the solvent was removed by rotary evaporation. The
product was obtained as an orange-brown solid in approximately 78% yield.
1
H NMR (CDCl3): 
8.64 (d, 1H, NC5H3), 8.41 (d, 1H, NC5H3), 7.77 (t, 1H, p-NC5H3), 7.69 (d, 1H, NC5H3), 7.59 (t,
1H, p-NC5H3), 7.25 (d, 1H, NC5H3), 6.57 (d, 1H, NC5H3), 3.16 (s, 6H, N(CH3)2) ppm.
58

2.5.7 Synthesis of Re(L
1
)(CO)3Cl (1)
A three neck flask was charged under N2 with ligand L
1
(102 mg) and rhenium pentacarbonyl
chloride (165 mg, 1 eq.). While stirring, dry toluene (20 mL) was added via syringe. The reaction
mixture was heated to reflux and allowed to stir overnight. After cooling, yellow solid precipitated,
which was collected by vacuum filtration and washed with diethyl ether. The filtrate was
transferred to a jar that was capped and placed in the freezer to allow more solid to precipitate
from solution. The total yield of bright yellow powder (complex 1) was approximately 95%.
Yellow crystals suitable for X-ray diffraction were grown from diffusion of diethyl ether into a
dimethylformamide (DMF) solution of 1.
1
H NMR (CDCl3):  7.70 (t, 1H, p-NC5H3), 7.32 (d, 1H,
NC5H3), 6.66 (d, 1H, NC5H3), 6.23 (s, 1H, NH), 3.07 (s, 3H, NCH3) ppm. Anal. Calcd. for 1: C,
34.65; H, 2.71; N, 10.78. Found: C, 34.68; H, 2.63; N, 10.63.

2.5.8 Synthesis of Re(L
2
)(CO)3Cl (2)
A three neck flask was charged under N2 with ligand L
2
(168 mg) and rhenium pentacarbonyl
chloride (235 mg, 1 eq.). While stirring, dry toluene (35 mL) was added via syringe. The reaction
mixture was heated to reflux and allowed to stir overnight. After cooling, yellow solid precipitated,
which was collected by vacuum filtration and washed with diethyl ether. The filtrate was
transferred to a jar that was capped and placed in the freezer to allow more solid to precipitate
from solution. The total yield of bright yellow powder (complex 2) was approximately 97%.
Yellow crystals suitable for X-ray diffraction were grown from diffusion of diethyl ether into a
dimethylformamide (DMF) solution of 2.
1
H NMR (CDCl3):  9.04 (d, 1H, NC5H3), 8.06 (d, 1H,
NC5H3), 7.99 (t, 1H, p-NC5H3), 7.75 (t, 1H, p-NC5H3), 7.45 (m, 2H, NC5H3), 6.74 (d, 1H, NC5H3),
59

6.09 (s, 1H, NH), 3.10 (d, 3H, NCH3). Anal. Calcd. for 2: C, 34.25; H, 2.26; N, 8.56. Found: C,
34.11; H, 2.28; N, 8.38.

2.5.9 Synthesis of Re(L
3
)(CO)3Cl (3)
A three neck flask was charged under N2 with ligand L
3
(149 mg) and rhenium pentacarbonyl
chloride (271 mg, 1 eq.). While stirring, dry toluene (40 mL) was added via syringe. The reaction
mixture was heated to reflux and allowed to stir overnight. After cooling, orange solid precipitated,
which was collected by vacuum filtration and washed with diethyl ether. The filtrate was
transferred to a jar that was capped and placed in the freezer to allow more solid to precipitate
from solution. The total yield of bright orange powder (complex 3) was approximately 94%.
Orange crystals suitable for X-ray diffraction were grown from diffusion of diethyl ether into a
dimethylformamide (DMF) and chloroform solution of 3.
1
H NMR (CDCl3):  7.89 (t, 1H, p-
NC5H3), 7.65 (d, 1H, NC5H3), 7.18 (d, 1H, NC5H3), 3.07 (s, 6H, N(CH3)2). Anal. Calcd. for 3: C,
37.26; H, 3.31; N, 10.22. Found: C, 37.07; H, 3.22; N, 9.87.

2.5.10 Synthesis of Re(L
4
)(CO)3Cl (4)
A three neck flask was charged under N2 with ligand L
4
(105 mg) and rhenium pentacarbonyl
chloride (158 mg, 1 eq.). While stirring, dry toluene (20 mL) was added via syringe. The reaction
mixture was heated to reflux and allowed to stir overnight. After cooling, orange solid precipitated
and was collected by vacuum filtration and washed with diethyl ether. The filtrate was transferred
to a jar that was capped and placed in the freezer to allow more solid to precipitate from solution.
The total yield of bright orange powder (complex 4) was approximately 99%. Orange crystals
60

suitable for X-ray diffraction were grown from diffusion of diethyl ether into a dimethylformamide
(DMF) solution of 4.
1
H NMR (CDCl3):  9.08 (d, 1H, NC5H3), 8.03 (m, 2H, NC5H3), 7.89 (t, 1H,
p-NC5H3), 7.67 (d, 1H, NC5H3), 7.48 (t, 1H, p-NC5H3), 7.18 (d, 1H, NC5H3), 6.57 (d, 1H, NC5H3),
3.12 (s, 6H, N(CH3)2). Anal. Calcd. for 4: C, 35.68; H, 2.60; N, 8.32. Found: C, 35.79; H, 2.67; N,
8.30.

61

2.5.11 Additional Figures

Scheme 2.2 Synthetic schemes for (a) 1, (b) 2, (c) 3, and (d) 4.

62

Figure 2.17 Cyclic voltammograms of 1 under CO2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.18 Cyclic voltammograms of 1 under CO2 with increasing concentrations of MeOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.19 Cyclic voltammograms of 1 under CO2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

63

Figure 2.20 Cyclic voltammograms of 1 under CO2 with increasing concentrations of PhOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.21 Cyclic voltammograms of 2 under CO2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

64

Figure 2.22 Cyclic voltammograms of 2 under CO2 with increasing concentrations of MeOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.23 Cyclic voltammograms of 2 under CO2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

65

Figure 2.24 Cyclic voltammograms of 2 under CO2 with increasing concentrations of PhOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.25 Cyclic voltammograms of 3 under CO2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

66

Figure 2.26 Cyclic voltammograms of 3 under CO2 with increasing concentrations of MeOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.27 Cyclic voltammograms of 3 under CO2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

67

Figure 2.28 Cyclic voltammograms of 3 under CO2 with increasing concentrations of PhOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.29 Cyclic voltammograms of 4 under CO2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

68

Figure 2.30 Cyclic voltammograms of 4 under CO2 with increasing concentrations of MeOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.31 Cyclic voltammograms of 4 under CO2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

69

Figure 2.32 Cyclic voltammograms of 4 under CO2 with increasing concentrations of PhOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.33 Cyclic voltammograms of 1 under N2 with increasing amounts of TFE. Conditions:
1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

70

Figure 2.34 Cyclic voltammograms of 2 under N2 with increasing amounts of TFE. Conditions:
1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.35 Cyclic voltammograms of 3 under N2 with increasing amounts of TFE. Conditions:
1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

71

Figure 2.36 Cyclic voltammograms of 4 under N2 with increasing amounts of TFE. Conditions:
1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 2.37 Cyclic voltammogram of 2 in DMF under N2 compared with that under CO2. 1 mM
catalyst in DMF with 0.1 M TBAPF6, scan rate: 100 mV/s.

72

Figure 2.38 Cyclic voltammograms of 2 under CO2 with increasing concentrations of TFE. 1
mM catalyst in DMF with 0.1 M TBAPF6, scan rate: 100 mV/s.

73

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76

CHAPTER 3

Influence of intermolecular hydrogen bonding interactions on the electrocatalytic reduction of
CO2 to CO by 6,6'-amine substituted rhenium bipyridine complexes

77

3.1 ABSTRACT
The introduction of biologically inspired motifs into electrocatalysts is an attractive
strategy for efficiently transforming harmful fossil fuel combustion products such as carbon
dioxide into useful chemical fuels. Herein, we present a series of Re(bpy)(CO)3Cl electrocatalysts
with pendant primary, secondary, and tertiary amines with the aim of determining the effect of
available pendant protons on the reduction of CO2 to CO. Cyclic voltammetry studies indicate that
the availability of pendant protons leads to intermolecular hydrogen bonding interactions, altering
the electrochemical behavior of the complexes. Further, controlled potential electrolysis studies
show a clear trend in the catalytic activity of these complexes based on the availability of pendant
protons. For NMe2-Re complex with no pendant protons, the Faradaic efficiency (FE) remains
quite stable with changing potential (41-65% FECO), but for NH2-Re with maximum available
pendant protons, the FECO increases with more negative potentials, peaking at 83% FECO. Together
with the formation of H2 by the NH2-substituted Re(bpy)(CO)3Cl complex, this suggests a change
in electrocatalytic behavior due to the intermolecular hydrogen bonding interactions that occur
with additional pendant protons.

3.2 INTRODUCTION
With the world’s increasing levels of atmospheric CO2 and detrimental climate change, it
is vital for the global infrastructure to transition from fossil fuel consumption toward renewable
energy sources. While solar energy is an attractive, abundant, and clean alternative to harmful
fossil fuel combustion, its intermittent nature makes energy storage a necessity. Therefore, several
strategies need to be developed to efficiently capture, store, transport, and convert the solar energy
into useful fuels. The electrolytic transformation of abundant small molecules into value-added
78

chemical feedstocks is an attractive strategy that would help mitigate this problem.
1
In particular,
the electrochemical conversion of carbon dioxide (CO2) to chemical fuels is an efficient method
for recycling harmful fossil fuel combustion products and for creating a more carbon-balanced
infrastructure.
2
Following the efficient and selective conversion of CO2 to CO, the Fischer-Tropsch
process can be utilized to produce useful hydrocarbon fuels.
3

In nature, the enzyme CO-dehydrogenase (CODH) efficiently catalyzes the selective
conversion of CO2 to CO, wherein the bound CO2 adduct is stabilized via second-coordination-
sphere hydrogen bonding interactions of the appropriately positioned amino-acid residues.
4–6

While CO2 reduction is a challenging process due to the high thermodynamic barrier, the use of a
catalyst emulating the beneficial properties found in enzymes increases the efficiency of this
transformation by lowering the energy input.
7,8
Although second-sphere interactions can be
difficult to control in synthetic catalysts, many biologically inspired CO2 reduction catalysts with
pendant proton donors for second-sphere interactions have been reported in an effort to harness
these intramolecular interactions for enhanced catalytic activity.
9
Iron porphyrins with pendant
phenolic,
10
trimethylanilinium,
11
amide,
12
and hangman groups,
13
nickel cyclams with pendant
amines,
14–17
polypyridyl iron complexes with pendant protic groups,
18
and cobalt aminopyridine
macrocycles with pendant amines
19,20
are only a select few examples of catalysts inspired by
nature.
Metal bipyridine complexes, such as Re(bpy)(CO)3Cl and Mn(bpy)(CO)3Br (bpy = 2,2'-
bipyridine), have been well-studied as electrocatalysts for the reduction of CO2 to CO.
21–25
While
this class of catalysts is known to be selective for CO production, dimerization and other side
processes can lead to low turnovers and catalytic deactivation.
26,27
Mn(bpy)(CO)3Br complexes
tend to undergo more rapid dimerization, with lower current densities, stabilities, and overall
79

Faradaic efficiencies (FEs).
28
While modifications of the ligand scaffold can influence electronic
and steric properties and drastically affect catalytic performance, introduction of pendant proton
donors to both Re(bpy)(CO)3Cl and Mn(bpy)(CO)3Br complexes were reported to introduce
second-sphere hydrogen bonding interactions that stabilize the CO2 adduct during catalysis and
increase catalytic activity.
29–33
Pendant phenolic groups installed on Mn(bpy)(CO)3Br complexes
have led to CO2 reduction with relatively low overpotentials
29
and production of CO and formic
acid without addition of a Brønsted acid
30
due to the inherent local proton source in close proximity
to the CO2 binding site. Incorporation of pendant secondary and tertiary amines,
31

dimethoxyphenyl,
32
and imidazolium
33
groups have also allowed for tuning of product selectivity
and decreased overpotentials, respectively.
Recently, the number of studies of Re(bpy)(CO)3Cl complexes with pendant proton donors
has rapidly increased.
34–43
Positioning of pendant hydroxy groups on the bipyridine backbone led
to deprotonation of the pendant protons, and a catalytic reductive disproportionation pathway was
observed, with slow CO production.
34
However, when the OH group is moved from the 4,4'-
positions to the 3,3'-positions, intra-ligand hydrogen bonding interactions allowed for a positive
shift in the reduction potential.
35
Re-bpy complexes with pendant di- and tri-phenolic groups
installed in the 6-position displayed small current enhancements even at the first reduction
potential, enhanced further by the addition of methanol (MeOH) or water.
36
An imidazolium-
functionalized Re(bpy)(CO)3Cl catalyst with an available pendant proton yielded greater (although
still moderate) FEs for CO compared to an analogous complex without intramolecular hydrogen
bonding capabilities and the unsubstituted Re(bpy)(CO)3Cl complex.
37
A thiourea group
positioned in the secondary coordination sphere acted as a local proton source, resulting in a
catalyst with optimal performance and an overall high FE for CO, in the absence of a Brønsted
80

acid.
38
Additionally, catalysts with ortho-, meta-, and para-aniline groups in the 6-position of the
bipyridine ring all outperformed unsubstituted Re(bpy)(CO)3Cl, with high turnover frequencies
(TOFs) and FEs for CO; the meta-NH2 allowed for the greatest catalytic activity due to the ideal
positioning of the pendant proton donor.
39
Modification of the bipyridine backbone with bio-
inspired methyl acetamidomethyl
40
groups and amino acids
41
promoted supramolecular assembly
and Re—Re dimerization that induced an alternative bimolecular mechanism for reductive
disproportionation. Further studies with the amino acid-modified complexes allowed for tuning of
the catalytic pathway and product selectivity via proton relays and hydrogen bonding
interactions.
42
We have also reported a rhenium bipyridine catalyst with 5,5'-NH2 substituents that
is selective for CO2 to CO conversion, although with a high overpotential.
43

We recently reported a study on Re(bpy)(CO)3Cl catalysts modified with secondary and
tertiary amines in the 6- and 6'- positions, wherein the mono-substituted complexes outperformed
the bis-substituted ones, with selective CO production but moderate FEs.
44
Inspired by the
literature reports on enhancement of catalytic activity through second-sphere interactions and
available pendant proton donors, we sought to more directly compare three 6,6'-substituted
Re(bpy)(CO)3Cl complexes with primary, secondary, and tertiary amines to determine the role of
available pendant protons in catalysis. Herein, we report the synthesis, characterization, and
electrochemistry of a new primary amine-substituted Re(bpy)CO3Cl catalyst (NH2-Re, Scheme 1)
with hydrogen bonding properties for the reduction of CO2 to CO. Additional electrochemical
studies were also performed on the previously reported secondary and tertiary amine-substituted
catalysts (NHMe-Re and NMe2-Re), and compared to the NH2-Re, to determine the effect of
available pendant protons capable of creating extended hydrogen bonding networks on
electrochemical behavior and catalyst performance.
81

3.3 RESULTS AND DISCUSSION
3.3.1 Synthesis and Characterization

Scheme 3.1 Synthetic scheme of complexes NH2, NHMe, and NMe2 from ligands NH2-L, NHMe-
L, and NMe2-L.

The rhenium tricarbonyl complexes NHMe-Re and NMe2-Re were prepared as previously
reported.
44
The synthesis of NH2-L was adapted from literature procedures,
45,46
and the rhenium
complex (NH2-Re) was subsequently generated by refluxing the NH2-L ligand with Re(CO)5Cl in
anhydrous toluene for 18 hours (Scheme 3.1). The
1
H nuclear magnetic resonance (NMR)
spectrum of NH2-Re in CDCl3 displays three aryl peaks: one triplet at  7.64 ppm and two doublets
at  7.37 and 6.74 ppm in a 1:1:1 ratio (Figure 3.1), assigned to the protons on the bipyridine
backbone. Additionally, a broad singlet appears at  5.61 ppm, which integrates in a 2:1 ratio
relative to each of the aryl peaks and is assigned to the symmetric NH2 moieties in the 6- and 6'-
positions. All three complexes were further characterized by UV-vis, displaying the characteristic
Re(bpy)(CO)3Cl peaks between 250 and 450 nm (Figure 3.2). The NH2 complex was also
characterized by elemental analysis (EA), and its elemental composition matches the expected
values.
82

Figure 3.1 400 MHz
1
H NMR spectrum of 6,6'-NH2(bpy)[Re(CO)3Cl] (NH2-Re) in CDCl3.

Figure 3.2 Overlay of the UV-vis spectra of NH2-Re (purple), NHMe-Re (blue), and NMe2-Re
(pink) in MeCN.

Colorless crystals of NH2-Re were grown via slow evaporation of DMF. Single crystal X-
ray diffraction studies display a six coordinate Re(I) metal center and facial arrangement of three
carbonyl moieties, analogous to other Re(bpy)(CO)3 complexes.
47
The bipyridine ligand occupies
a quasi-equatorial coordination, with Re–N–C–N torsion angles of 16.9(3)° and 14.1(2)° (Figure
3.3, Table 3.1). These angles are similar to those previously reported for NHMe-Re of 16.3(3)°
and 18.2(3)°. However, Re(bpy)(CO)3 complexes with substituents in other positions are typically
much more planar due to less steric bulk around the metal center. Additionally, the Re–N(bpy)
83

bond lengths (2.178(2) Å and 2.190(2) Å) in NH2-Re are analogous to, although slightly shorter
than, those of NHMe-Re (2.194(2) Å and 2.202(2) Å), consistent with less steric bulk and a lower
degree of ligand tilting.

Figure 3.3 Solid state structure of 6,6'-NH2-Re, (a) front and (b) side views. Color legend of the
atoms: gray – C; blue – N; red – O; green – Cl; pink – Re; white – H. Solvent molecules and
additional hydrogen atoms are excluded for clarity.

Table 3.1 Crystal data and structure refinement for NH2-Re.
Identification code Ashley102919
Chemical formula C16H17ClN5O4Re
Formula weight 564.99 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.216 x 0.300 x 0.389 mm
Crystal habit clear colourless prism
Crystal system monoclinic
Space group P 1 21/c 1
Unit cell dimensions a = 12.1840(16) Å α = 90°

b = 8.4095(11) Å β = 103.408(2)°

c = 18.769(3) Å γ = 90°
Volume 1870.7(4) Å
3

Z 4
Density (calculated) 2.006 g/cm
3

Absorption coefficient 6.672 mm
-1

F(000) 1088
Diffractometer Bruker APEX II CCD Bruker APEX DUO
Radiation source fine-focus tube (MoKα, λ = 0.71073 Å)
Theta range for data collection 1.72 to 30.48°
Index ranges -17 ≤ h ≤ 17, -11 ≤ k ≤ 11, -26 ≤ l ≤ 26
84

Reflections collected 35952
Independent reflections 5654 [R(int) = 0.0424]
Coverage of independent reflections 99.4%
Absorption correction multi-scan
Structure solution technique direct methods
Structure solution program SHELXT 2014/5 (Sheldrick, 2014)
Refinement method Full-matrix least-squares on F2
Refinement program SHELXL-2018/3 (Sheldrick, 2018)
Function minimized Σ w(Fo
2
- Fc
2
)
2

Data / restraints / parameters 5654 / 3 / 258
Goodness-of-fit on F
2
1.016
Δ/σmax 0.008
Final R indices 5323 data; I >
2σ(I)
R1 = 0.0163, wR2 = 0.0399

all data R1 = 0.0181, wR2 = 0.0407
Weighting scheme w = 1/[σ
2
(Fo
2
) + (0.0158P)
2
+ 1.5571P]
where P = (Fo
2
+2Fc
2
)/3
Largest diff. peak and hole 0.661 and -0.842 eÅ
-3

R.M.S. deviation from mean 0.105 eÅ
-3

NH2-Re was further characterized by FT-IR spectroscopy and compared to the previously
reported spectra of NHMe-Re and NMe2-Re.
44
Three carbonyl stretches were observed for NH2-
Re at 2027, 1911, and 1897 cm
-1
, characteristic of the one high-energy mode (a'1) and two low-
energy modes (a'' and a'2), typically observed for fac-Re(CO)3 complexes (Figure 3.4, Table
3.2).
48,49
NH2-Re also displays a broad N-H stretch centered at 3234 cm
-1
, consistent with a primary
amine with intermolecular hydrogen bonding interactions. This IR band is much broader than the
sharp N-H peak observed for NHMe-Re at 3416 cm
-1
, suggesting that NH2-Re has a greater affinity
for intermolecular hydrogen bonding interactions. A sharp peak at 1655 cm
-1
is also visible for
NH2-Re, assigned to N-H bending modes.
85

Figure 3.4 FT-IR spectrum of NH2-Re.

Table 3.2 Pertinent IR spectroscopy wavenumbers for NH2-Re, NHMe-Re, and NMe2-Re.

Complex N-H C=O (a1') C=O (a") C=O (a2') Reference
NH2-Re 3234 2027 1911 1897 This study
NHMe-Re 3416 2011 1897 1857 57
NMe2-Re – 2013 1903 1876 57

86

3.3.2 Cyclic Voltammetry

Figure 3.5 CVs of 1 mM NH2-Re (purple), NHMe-Re (blue), and NMe2-Re (pink) under N2 in a
0.1 M TBAPF6 MeCN solution. Dashed lines illustrate the first reduction feature and dotted lines
(for NH2-Re and NHMe-Re) show the presence of an additional intermediate reduction feature.
Scan rate: 100 mV/s.

Table 3.3 Reduction potentials (V vs. Fc
+/0
) for NH2-Re, NHMe-Re, and NMe2-Re in MeCN and
DMF under N2. Scan rate: 100 mV/s.
Catalyst
1
st
reduction Intermediate reduction 2
nd
reduction
MeCN DMF MeCN DMF MeCN DMF
NH 2-Re -2.00 -2.06 -2.11 -- -2.23 -2.42
NHMe-Re -1.96
44
-2.00 -2.05
44
-- -2.25
44
-2.36
87

NMe 2-Re -1.92
44
-1.94 -- -- -2.26
44
-2.46

For comparison of the electrochemical behavior of NH2-Re to our previously reported
NHMe-Re and NMe2-Re complexes, cyclic voltammetry (CV) studies were first performed under
N2 in acetonitrile (MeCN) solutions (1 mM catalyst with 0.1 M tetrabutylammonium
hexafluorophosphate [TBAPF6] as supporting electrolyte). All electrochemical potentials
discussed are referenced versus Fc
+/0
. CVs of NH2-Re under N2 display three irreversible reduction
events at -2.00 V, -2.11 V, and -2.23 V (Figure 3.5, Table 3.3). The first reduction feature gains
quasi-reversibility with increased scan rates (Figure 3.6), suggesting that at high scan rates, the
electrochemical process can outcompete the irreversible chemical one. Randles-Sevcik plots
indicate that the catalyst is freely diffusing in solution, as expected for a molecular species (Figure
3.7). The three observed reduction features are similar to those reported for NHMe-Re,
44
which
also exhibits a third reduction feature in the same potential range as the two one-electron reduction
events typically observed for Re(bpy)(CO)3 complexes.
21,24,47
We have previously assigned the
second reduction feature for NHMe-Re to the formation of a species with hydrogen bonding
interactions, as the behavior is similar to the four reduction features observed previously for a
Re(bpy)(CO)3 complex with hydrogen bonding methyl acetamidomethyl groups.
40
Moreover,
NMe2-Re, which has tertiary amine substituents, does not display this intermediate feature,
suggesting that this hydrogen bonding process is related to the availability of pendant protons, and
supports our previous assignment of hydrogen bonding interactions. Further, variable scan rate
studies of NH2-Re demonstrate that the intermediate (the 2
nd
reduction) feature disappears at scan
rates higher than 200 mV/s, indicating that this feature is related to a chemical process that occurs
too slowly to be observed on these timescales (Figure 3.6).
88

Figure 3.6 Variable scan rate studies of (a) first and (b) first and second reductions of NH2-Re
under N2. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6.

Figure 3.7 Randles-Sevcik plots for the first (a) and second (b) reductions of NH2-Re at -2.00 and
-2.23 V, respectively; the linear correlation and slope of c.a. 0.5 indicates that the species is freely
diffusing based on the Randles-Sevcik equation.

To further probe the hydrogen bonding interactions of NH2-Re and NHMe-Re, CVs of all
three complexes were performed in DMF due to DMF’s ability to disrupt intermolecular hydrogen
bonding interactions.
40
As expected, both NH2-Re and NHMe-Re, which possess pendant protons
capable of hydrogen bonding, exhibit very different electrochemical behavior in DMF than in
MeCN. Under N2, both complexes display only one quasi-reversible reduction feature at -2.06 V
(for NH2-Re) and -2.00 V (for NHMe-Re), followed by an irreversible feature at -2.42 V (for NH2-
Re) and -2.36 V (for NHMe-Re), resembling the behavior of NMe2-Re and Re(CO)3Cl complexes
without hydrogen bonding capabilities (Figure 3.8, Table 3.3).
24,40
The pendant protons of the
89

amines allow for intermolecular hydrogen bonding interactions in MeCN, which is neither a good
hydrogen bond donor nor acceptor. However, DMF, which is a much better hydrogen bond donor
and acceptor than MeCN, disrupts these intermolecular interactions and inhibits the formation of
hydrogen bonding networks, leading to an overall change in shape and the disappearance of the
additional electrochemical reduction feature. NMe2-Re, which possesses no pendant protons,
exhibits similar electrochemical behavior in DMF as in MeCN, with a quasi-reversible reduction
feature at -1.94 V and irreversible feature -2.46 V (Figure 3.8, Table 3.3). Variable scan rate studies
in DMF for all the three complexes show minimal changes in the reduction features (Figure 3.9).

Figure 3.8 Overlay of the cyclic voltammograms of NH2-Re (purple), NHMe-Re (blue), and
NMe2-Re (pink) under N2 in MeCN (dashed) and DMF (solid). Conditions: 1 mM catalyst with
0.1 M TBAPF6, scan rate: 100 mV/s.
90

Figure 3.9 Variable scan rate studies of (a) NH2-Re, (b) NHMe-Re, and (c) NMe2-Re under N2.
Conditions: 1 mM catalyst in DMF with 0.1 M TBAPF6.

Two additional experiments were conducted to corroborate the assignment of the observed
reduction events for NHMe-Re. First, CVs were performed on NHMe-Re in MeCN and titrated
with increasing amounts of DMF (Figure 3.10a). The first reduction feature at -1.95 V gradually
disappears, corresponding to the interruption of intermolecular hydrogen bond formation due to
competing hydrogen bonding between the NHMe group and DMF. Thus, this feature can
confidently be assigned to the formation of hydrogen bonding networks. The second experiment
performed was the titration of tetrabutylammonium chloride (TBACl) into a MeCN solution of
NHMe-Re (Figure 3.10b). The addition of as little as 5 mM TBACl leads to the disappearance of
the -1.95 V reduction feature, and a current increase for the -2.05 V reduction feature. CVs with
added TBACl in the 5 mM to 100 mM concentration range exhibit a current decrease in the
reduction feature at -2.05 V. As the additional supply of Cl
-
in the solution inhibits irreversible
91

chloride loss from Re(CO)3Cl complexes, this feature can be attributed to loss of the chloride
ligand.

Figure 3.10 Cyclic voltammogram of NHMe-Re under N2 with increasing concentrations of (a)
DMF and (b) TBACl. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100
mV/s.

The electrochemical behavior of NHMe-Re in the presence of bases was also evaluated to
determine the impact of available pendant protons on the observed redox features. Triethylamine
(NEt3) was titrated into the solution during cyclic voltammetry under an N 2 atmosphere. Similar
to titrations with DMF, the first reduction feature at -1.95 V gradually disappeared with increasing
concentrations of NEt3. At concentrations higher than 2 M, a reduction feature at -2.10 V, slightly
more negative than the original intermediate feature (-2.05 V), grows in more strongly (Figure
3.11). With 4 M NEt3, a single reversible feature was observed, similar to a bulky manganese
bipyridine complex that does not undergo dimerization or hydrogen bonding during
electrochemical studies,
51
for which the redox feature was assigned to two overlapping one-
electron reduction events combined with MeCN ligand loss. These results suggest that the presence
of the pendant protons plays an important role in the observed electrochemical behavior.
92

Figure 3.11 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of
triethylamine. 1s mM catalyst in MeCN with 0.1 M TBAPF6, scan rate: 100 mV/s.

The addition of weak Brønsted acids, such as water and 2,2,2-trifluoroethanol (TFE) under
N2, causes a notable transformation for both NH2-Re and NHMe-Re (Figure 3.12). For NH2-Re,
the addition of a large excess of acid – 1 M TFE or 2 M H2O – changes the shape of the CV to two
irreversible reduction events at -2.03 V and -2.56 V with TFE and -2.00 V and -2.39 V with H2O
(Figure 3.12a). For NHMe-Re, the addition of only 0.1 M TFE produces two irreversible reduction
events at -1.97 V and -2.27 V; however, 3 M H2O was necessary to observe the appearance of two
distinct reduction features, at the same potentials as when TFE was used (Figure 3.12b). The
distinct shift in electrochemical behavior with the addition of these proton sources suggests that
the weak acids are capable of disrupting the intermolecular hydrogen bonding interactions between
adjacent catalyst molecules, thus inhibiting the formation of an extended hydrogen bonding
network, similar to the effect observed when the solvent is switched from MeCN to DMF.
93

Figure 3.12 Cyclic voltammograms of (a) NH2-Re and (b) NHMe-Re in MeCN under N2 with the
addition of no acid (grey), H2O (blue, 2 M for NH2-Re and 3 M for NHMe-Re), and TFE (orange,
1 M for NH2-Re and 0.1 M for NHMe-Re). 1 mM catalyst with 0.1 M TBAPF6, scan rate: 100
mV/s.

Upon saturation with CO2, CVs of NH2-Re display a small increase in current beginning
at the first reduction event, similar to the behavior observed for NHMe-Re and NMe2-Re under
CO2 (Figure 3.13). However, the maximum current densities of NH2-Re and NHMe-Re (each <1
mA/cm
2
) are lower than that of NMe2-Re, which displays a peak-shaped curve achieving a current
density of 1.7 mA/cm
2
at -2.38 V as previously reported.
44
Various Brønsted acids (water, 2,2,2-
trifluoroethanol, and phenol) were evaluated for the reduction of CO2 by NH2-Re (Figures 3.13a).
As with NHMe-Re and NMe2-Re, TFE was found to provide the greatest current enhancements at
acid concentrations up to 2 M. Additionally, current enhancement with TFE occurred at a more
positive potential than with water. Although the current enhancement with phenol (PhOH)
occurred at a potential similar to that with TFE, the peak current density with PhOH was lower
than that with TFE. Comparison of the three 6,6'-substituted catalysts shows that NMe2-Re
obtained the highest current density of 5.0 mA/cm
2
in the presence of 1 M TFE, followed by NH2-
Re and NHMe-Re (3.0 and 1.8 mA/cm
2
, respectively) (Figure 3.13). Additionally, current densities
were the largest for all three catalysts under TFE conditions relative to 1 M PhOH or water, so
TFE was chosen as the optimal proton source for further studies. Titration with Brønsted acids
94

under N2 led to negligible current enhancements, indicating that the current responses under CO2
are likely not due to reduction of protons from the acid sources (Figure 3.12).

Figure 3.13 CVs of 1 mM (a) NH2-Re, (b) NHMe-Re, and (c) NMe2-Re under N2 (black), CO2
(red), 1 M PhOH (orange), 1 M TFE (green), and 1 M H2O (blue) (0.1 M TBAPF6 in MeCN; scan
rate: 100 mV/s).

3.3.3 Controlled Potential Electrolysis
Controlled potential electrolysis (CPE) experiments were performed for one hour under
CO2 at various potentials for all three catalysts. Gas chromatography (GC) was used to analyze the
gaseous products formed during electrolysis. CPE experiments for all three complexes without
added acid produced no detectable gaseous products, indicating that an additional acid source is
necessary for catalytic turnover. As TFE provided the greatest current densities during CV studies,
1 M TFE was added to each CPE solution. CPE for each complex was performed at 100 mV
95

increments from -2.00 V to -2.30 V to explore the impact of catalytic activity and potential on
product distribution. For NMe2-Re, the Faradaic efficiency (FE) for CO varied less with the
potential at which the CPE was performed, ranging from 41 to 65% (Table 3.4 and Figure 3.14c),
and no catalytic products other than CO and trace H2 were detected. CPE experiments of NHMe-
Re showed a slight dependence of FE for CO on the potential applied, such that the greatest FE
(76%) was obtained at -2.1 V, near the potential of maximum current density in the CV
experiments (Table 3.5 and Figure 3.14b). Additionally, the amount of CO produced at -2.00 V
(55 µmol) was comparable to that at -2.10 and -2.20 V (74 and 55 µmol, respectively), but the FE
for CO (17%) at -2.00 V, was significantly lower. As no other products, such as H2 or formate,
were detected, this suggests that alternative Faradaic processes, such as the likely formation of the
proposed hydrogen bonding networks, are more dominant at -2.00 V.

Figure 3.14 Current versus time over one hour of controlled potential electrolysis of (a) NH2-Re,
(b) NHMe-Re, and (c) NMe2-Re at -2.00 V (red), -2.10 V (orange), -2.20 V (green), and -2.30 V
96

(blue). BE studies were performed using 1 mM catalyst in MeCN with 1M TFE and 0.1 M
TBAPF6.

Table 3.4 CPE results for NMe2-Re (1 mM) in MeCN under CO2 with 0.1 M TBAPF6 and 1 M
TFE.
Potential (V) FECO (%) FEH2 (%) μmol CO TON
-2.00 55 ± 13.4 0 21 0.5
-2.10 41 ± 9.9 0 66 1.7
-2.20 65 ± 1.5 3 ± 0.0 77 1.9
-2.30 50 ± 11.4 4 ± 0.1 124 3.1

Table 3.5 CPE results for NHMe-Re (1 mM) in MeCN under CO2 with 0.1 M TBAPF6 and 1 M
TFE.
Potential (V) FECO (%) FEH2 (%) μmol CO TON
-2.00 17 ± 0.4 0 55 1.4
-2.10 76 ± 5.5 0 74 1.9
-2.20 51 ± 1.2 0 55 1.4
-2.30 55 ± 1.3 0 127 3.2

Table 3.6 CPE results for NH2-Re (1 mM) in MeCN under CO2 with 0.1 M TBAPF6 and 1 M
TFE.
Potential (V) FECO (%) FEH2 (%) μmol CO TON
-2.00 0 0 0 0
-2.10 32 ± 0.7 43 ± 0.5 11 0.3
-2.20 63 ± 4.5 15 ± 0.9 69 1.8
-2.30 83 ± 1.9 9 ± 0.1 125 3.1
-2.35 71 ± 1.6 0 127 3.2

CPE experiments of NH2-Re display a unique linear relationship between applied potential
and FECO (Figures 3.14a and 3.15). All CPE results of NH2-Re are presented in Table 3.6. At -2.00
V, no gaseous products were detected by GC, but both the amount of CO produced and the FE
increase with more negative potentials, peaking at -2.30 V. CPE at -2.35 V yielded a comparable
amount of CO to that at -2.30 V, but with a drop in FE. Substantial amounts of H2 were also
97

produced under these conditions, decreasing with more negative potentials; the greatest amount of
H2 (43% FEH2, 15 µmol) was produced at -2.10 V, near the potential with maximum current density
observed in CV studies (-2.15 V).

Figure 3.15 Plot of Faradaic efficiency for CO (%) versus potential (V) for CPE experiments of
NH2-Re.

Further CPE experiments of NHMe-Re under N2 were performed to independently observe
the proposed hydrogen bonding network and confirm the results observed by CV studies. Holding
the potential at the first reduction event (-1.96 V) for one hour led to a gradual color change from
bright yellow to brownish-orange (Figure 3.16). Following CPE, 15 mL of the solution from the
working compartment of the electrolysis cell were syringed into a CV cell. Post-CPE CVs showed
a single reversible reduction event at -2.04 V, analogous with the behavior observed in CV studies
collected in the presence of increasing amounts of NEt3 or DMF (Figure 3.17). Peak-to-peak
separation of the reversible -2.04 V feature is 48 mV, compared to 88 mV observed for Fc
+/0
under
the same conditions, similar to a bulky manganese bipyridine complex that does not undergo
dimerization of hydrogen bonding during electrochemical studies (Table 3.6).
51
These results
98

indicate that the reversible -2.04 V feature can be assigned as a two-electron reduction event,
compared to two subsequent one-electron reduction events typically observed for this class of
compounds.
51
The change in CV behavior is consistent with the inhibition of intermolecular
hydrogen bonding interactions. Further, GC analysis of the headspace yielded a large production
of CO (20 µmol), indicating loss of carbonyl ligands during electrolysis. Analogous CPE
experiments performed with NH2-Re showed minimal change in CV behavior, and no gaseous
products were observed by GC.

Figure 3.16 CPE cell showing NHMe-Re after 1 hour electrolysis under N2 at -1.96 V, showing
color change from yellow to brown/orange.

99

Figure 3.17 CV of (a) NHMe-Re in a CV cell following 1 hour of CPE under N2 at -1.96 V vs
Fc
+/0
and (b) overlay of CVs with added DMF (blue) and NEt3 (green) (0.1 M TBAPF6 in MeCN;
scan rate: 100 mV/s).

Although NH2-Re and NHMe-Re both have available pendant protons to undergo hydrogen
bonding interactions during electrochemical studies, the different trends revealed by CPE indicate
that the number of pendant protons plays an important role in catalytic activity. NMe2-Re, with no
available pendant protons, displays a relatively consistent FECO regardless of the potential applied
during CPE. However, for NHMe-Re and NH2-Re, which both have available pendant protons,
the FECO does change with potential, more noticeably for NH2-Re than for NHMe-Re. These two
complexes exhibit higher FE values at potentials corresponding to the peak current densities
observed by CV. FE values at more positive potentials are lower, suggesting that a smaller portion
of the charge produced is actually contributing to the conversion of CO2 to CO; at these potentials,
it is likely that NH2-Re and NHMe-Re are undergoing intermolecular hydrogen bonding
interactions that interfere with the desired catalytic conversion. The additional production of H2
by NH2-Re, which decreases with more negative potentials, further indicates that an alternative
process is occurring more readily at more positive potentials, possibly encouraged by these
hydrogen bonding interactions. This indicates that NH2-Re has an optimal potential for the
conversion of CO2 to CO and allows for control over the selectivity by changing the potential
during CPE.

3.4 CONCLUSION
In conclusion, we have reported here the synthesis, characterization, via
1
H NMR, FT-IR
and single crystal X-Ray diffraction, of a new Re(bpy)(CO)3Cl complex modified with primary
100

amines in the 6- and 6'-positions. Cyclic voltammograms of NH2-Re under N2 display an additional
intermediate reduction event, analogous to our previously reported NHMe complex, which
indicates the formation of hydrogen bonding networks. Switching the electrochemical solvent
from MeCN to DMF, which can interrupt the formation of these hydrogen bonds, led to a
noticeable change in behavior for both NH2-Re and NHMe-Re, with only the two typical
Re(bpy)(CO)3Cl reduction features observed. Further CV experiments of NHMe-Re with
increasing amounts of DMF, triethylamine, and weak Brønsted acids water and TFE further
support the interruption of hydrogen bonding networks, with CVs displaying either two subsequent
or two overlapping reduction events, without the additional feature observed in MeCN alone.
Under CO2, NH2-Re displayed a moderate increase in current, which was further increased with
the addition of TFE, yielding a maximum current density of 3.0 mA/cm
2
. CPE experiments of
NH2-Re showed a linear trend between FECO and electrochemical potential, with the greatest
selectivity for CO occurring at -2.30 V (83%). Varying amounts of hydrogen were also formed
from -2.10 V to -2.30 V. These results differ from those of NHMe-Re and NMe2-Re, which
produce negligible H2 and vary less in FECO with potential, suggesting that the greater number of
available pendant protons do play a role in enhancing catalytic activity.

3.5 EXPERIMENTAL METHODS AND ADDITIONAL FIGURES
3.5.1 Materials and Synthesis
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres drybox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. All solvents were degassed with nitrogen and passed through activated
101

alumina columns and stored over 4Å Linde-type molecular sieves. Deuterated solvents were dried
over 4Å Linde-type molecular sieves prior to use. Proton NMR spectra were acquired at room
temperature using Varian (Mercury 400 2-Channel, VNMRS-500 2-Channel, VNMRS- 600 3-
Channel, and 400- MR 2-Channel) spectrometers and referenced to the residual
1
H resonances of
the deuterated solvent (
1
H: CDCl3) and are reported as parts per million relative to
tetramethylsilane. Elemental analyses were performed using Thermo Scientific™ FLASH 2000
CHNS/O Analyzers. All the chemical reagents were purchased from commercial vendors and used
without further purification.

3.5.2 Electrochemistry
Electrochemistry experiments were carried out using a Pine potentiostat. The experiments
were performed in a single compartment electrochemical cell under nitrogen or CO 2 atmosphere
using a 3 mm diameter glassy carbon electrode as the working electrode, a platinum wire as the
auxiliary electrode and a silver wire as the reference electrode. Ohmic drop was compensated using
the positive feedback compensation implemented in the instrument. All electrochemical
experiments were performed with iR compensation using the current interrupt (RUCI) method in
AfterMath. Typical values for the cell resistance were around 160-170 ohms. All potentials in this
paper were referenced relative to ferrocene (Fc) with the Fe
3+/2+
couple at 0.0 V. Alternatively, in
cases when the redox couple of ferrocene overlapped with other features of interest,
decamethylferrocene (Fc*) was used as an internal standard with the Fe*
3+/2+
couple at –0.48 V.
All electrochemical experiments were performed with 0.1 M tetrabutylammonium
hexafluorophosphate (TBAPF6) as supporting electrolyte. The concentrations of the rhenium
102

catalysts were generally at 1 mM and experiments with CO2 were performed at gas saturation in
acetonitrile (MeCN) or dimethylformamide (DMF).
Controlled-potential electrolysis (CPE) measurements were conducted in a two-chamber
H cell. The first chamber held the working and reference electrodes in 40 mL of 0.1 M TBAPF6
and TFE in MeCN or DMF. The second chamber held the auxiliary electrode in 25 mL of 0.1 M
TBAPF6 in MeCN or DMF. The two chambers were separated by a fine porosity glass frit. The
reference electrode was placed in a separate compartment and connected by a Vycor tip. Glassy
carbon plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and
auxiliary electrodes. 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. Faradaic
efficiencies were determined by dividing the measured CO produced by the amount of CO
calculated based on the charge passed during the CPE experiment. The reported Faradaic
efficiencies and mol of CO produced are average values.
3.5.3 Synthesis of 6,6'-NH2(bpy)[Re(CO)3Cl] (NH2)
Under N2, 6,6'-NH2(bpy) (0.24 g, 1.3 mmol) and Re(CO)5Cl (0.48 g, 1.3 mmol) were
refluxed in 30 mL anhydrous toluene for 20 hours. After cooling, the resulting yellow precipitate
was collected by vacuum filtration and washed with ether, yielding a bright yellow solid in
approximately 92% yield. Anal. Calcd. for C13H10ClN4O3Re: C, 31.74; H, 2.05; N, 11.39. Found:
C, 32.36; H, 1.47; N, 11.60.

103

3.5.4 Additional Figures

Scheme 3.2 Synthetic scheme for the formation of 6,6'-NH2(bpy)[Re(CO)3Cl] (NH2-Re).

Figure 3.18 First (gray) and second (black) reduction features of NH2-Re. Conditions: 1 mM
catalyst in DMF with 0.1 M TBAPF6. Scan rate: 100 mV/s.

104

Figure 3.19 First (gray) and second (black) reduction features of NHMe-Re. Conditions: 1 mM
catalyst in DMF with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 3.20 First (gray) and second (black) reduction features of NMe2-Re. Conditions: 1 mM
catalyst in DMF with 0.1 M TBAPF6. Scan rate: 100 mV/s.

105

Figure 3.21 Variable scan rate studies of NH2-Re under N2. Conditions: 1 mM catalyst in DMF
with 0.1 M TBAPF6.

Figure 3.22 Variable scan rate studies of NHMe-Re under N2. Conditions: 1 mM catalyst in DMF
with 0.1 M TBAPF6.

106

Figure 3.23 Variable scan rate studies of NMe2-Re under N2. Conditions: 1 mM catalyst in DMF
with 0.1 M TBAPF6.

Figure 3.24 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

107

Figure 3.25 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 3.26 Cyclic voltammograms of NH2-Re under N2 with increasing concentrations of PhOH.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

108

Figure 3.27 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 3.28 Cyclic voltammograms of NHMe-Re under N2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

109

Figure 3.29 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 3.30 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

110

Figure 3.31 Cyclic voltammograms of NH2-Re under CO2 with increasing concentrations of
PhOH. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

111

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116

CHAPTER 4

Primary and secondary sphere effects of amine substituent position on rhenium bipyridine
electrocatalysts for CO2 reduction
117

4.1 ABSTRACT
It is well-established that electrocatalysts with biologically inspired moieties provide
advantageous qualities for the efficient transformation of fossil fuel combustion products, such as
carbon dioxide, into value-added products. Herein, we present a family of Re(bpy)(CO)3Cl
electrocatalysts with pendant NH2 groups in the 4,4’-, 5,5’-, and 6,6’-positions of the bipyridine
backbone with the objective of determining the effects of combined primary and secondary sphere
interactions on the reduction of CO2 to CO. Cyclic voltammetry studies under CO2 show that the
catalytic onset potential is more positive for 6,6’-NH2-Re, indicating that proximity of the available
pendant protons to the metal center is important during catalysis. Controlled potential electrolysis
studies demonstrate that, similar to 6,6’-NH2-Re, 4,4’-NH2-Re displays a dependence of Faradaic
efficiency (FE) on catalytic potential, reaching a maximum FECO of 91% at -2.40 V; this FE is
greater than that observed for 6,6’-NH2-Re but occurs at a more negative potential. Ultimately, it
is important that the various catalytic parameters are properly balanced to provide the greatest
overall catalytic activity.

4.2 INTRODUCTION
The effects of climate change are detrimental to both the environment and society. With a
significant portion of infrastructure based on the consumption of fossil fuels, it is paramount to
develop the technology to convert small molecules into renewable fuel sources.
1
Specifically, the
conversion of carbon dioxide (CO2) into useful products is beneficial in that it allows for recycling
of harmful combustion byproducts while working toward a carbon-neutral infrastructure.
2
In order
to overcome the high thermodynamic barrier associated with the conversion of CO2 to carbon
monoxide (CO), catalysts must be designed and employed.
3
Throughout its evolutionary history,
118

the CO-dehydrogenase (CODH) enzyme in nature has adapted to efficiently and selectively reduce
CO2 to CO.
4–6
In recent years, chemists have aimed to emulate nature by tuning both the electronic
effects of the primary coordination sphere and the through-space effects of the secondary
coordination sphere. Specifically, altering the position and electron-withdrawing nature of a
functional group on the ligand of the catalyst can influence such catalytic parameters as the
overpotential, and adding and properly positioning pendant proton donors can assist with
stabilization of the CO2 adduct that forms during catalysis, enhancing overall catalytic activity.
A variety of bio-inspired CO2 reduction electrocatalysts have been designed to date,
including (but not limited to) nickel cyclams with pendant amines,
7–9
iron porphyrins with pendant
phenolic,
10
trimethylanilinium,
11
and amide groups,
12
and cobalt aminopyridine macrocycles with
pendant amines.
13–15
Further, many rhenium
16–26
and manganese
27–32
bipyridine [Re(bpy)(CO)3Cl
and Mn(bpy)(CO)3Br, (bpy = 2,2′-bipyridine)] catalysts with pendant proton donors and resulting
second-sphere interactions have also been reported.
33
Positional effects of methyl substituents on
the bipyridine backbone have been studied,
34
as well as the electronic effects from altering the
electron-donating or -withdrawing nature of substituents within the same position on the ligand.
35–
37
These studies demonstrate that both sterics and electronics influence electrocatalytic behavior.
However, a small number of CO2 reduction electrocatalysts exist that investigate tandem primary
and secondary coordination sphere effects.
We previously studied a family of 6,6’-substituted rhenium bipyridine complexes with
primary, secondary, and tertiary amines to determine the effect of available pendant protons on
electrocatalytic activity, wherein the -NH2 substituted complex (6,6’-NH2-Re) underwent
hydrogen bonding interactions and outperformed 6,6’-NHMe-Re and 6,6’-NMe2-Re. Building
upon these studies, we report here the synthesis, characterization and electrochemistry of a novel
119

4,4’-amino substituted Re(bpy)(CO)3Cl complex (4,4’-NH2-Re, Scheme 4.1), along with further
electrochemical studies of the previously reported 5,5’-substituted analogue (5,5’-NH2-Re).
25
We
further sought to more directly compare 4,4’-NH2-Re with our previously studied 5,5’-NH2-Re
and 6,6’-NH2-Re to determine the electronic and steric positional effects of the -NH2 substituent
on electrocatalytic activity for the reduction of CO2 to CO.

4.3 RESULTS AND DISCUSSION
4.3.1 Synthesis and Characterization

Scheme 4.1 Synthesis of complex 4,4’-NH2-Re.

Synthesis of ligand 4,4’-NH2-L was adapted from literature procedures.
38,39
Upon isolation
of the ligand, complex 4,4’-NH2-Re was synthesized via reflux of 4,4’-NH2-L and Re(CO)5Cl in
anhydrous toluene overnight (Scheme 4.1). Following collection of the precipitate,
1
H nuclear
magnetic resonance (NMR) of 4,4’-NH2-Re in DMSO-d6 displays three aromatic peaks: a doublet
at  8.22 ppm, a singlet at  7.21 ppm, and another doublet at  6.65 ppm in a 1:1:1 ratio, assigned
to the protons along the backbone of bipyridine ring (Figure 4.1). An additional 2H singlet at 
7.12 ppm is assigned to the symmetric NH2 substituents in the 4- and 4’-positions of the bipyridine
ring. The small singlet at  7.95 ppm represents the protons of co-crystallized N,N-
120

dimethylformamide (DMF). 4,4’-NH2-Re was further characterized by UV-vis, displaying a
characteristic Re(bpy)(CO)3Cl peak at 350 nm (Figure 4.2). Unfortunately, crystallization of 4,4’-
NH2-Re for single crystal X-ray diffraction studies was unsuccessful.

Figure 4.1
1
H NMR of 4,4’-NH2-Re in DMSO-d6 (400 MHz). The peak at δ 7.95 ppm represents
co-crystallized DMF.

Figure 4.2 UV-vis spectrum of 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re in MeCN.

4,4’-NH2-Re was further characterized by FT-IR spectroscopy (Figure 4.3). Three carbonyl
stretches are observed at 2014, 1917, and 1869 cm
-1
, corresponding to one high-energy mode (a1')
and two lower-energy modes (a" and a2'), as expected for fac-Re(CO)3 complexes and similar to
121

those observed for 6,6’-NH2-Re. There are also several primary amine asymmetric and symmetric
N-H stretches at 3493, 3405, 3308, and 3200 cm
-1
, as well as a strong -NH2 scissoring absorption
at 1640 cm
-1
. The sharper -NH2 stretches compared to the broad ones observed for 6,6’-NH2-Re
suggest that hydrogen bonding interactions are less favorable for this complex.

Figure 4.3 FT-IR spectrum of 4,4’-NH2-Re.

4.3.2 Cyclic Voltammetry

For comparison of the electrochemical behavior of 4,4’-NH2-Re to our previously reported
5,5’-NH2-Re and 6,6’-NH2-Re, cyclic voltammetry (CV) studies were first performed under N2 in
acetonitrile (MeCN) solutions (1 mM catalyst with 0.1 M tetrabutylammonium
hexafluorophosphate [TBAPF6] as supporting electrolyte). All electrochemical potentials
discussed are referenced versus Fc
+/0
. CVs of 4,4’-NH2-Re display an irreversible reduction event
at -2.14 V followed by a second broad irreversible reduction event near -2.46 V, both of which do
not gain reversibility with faster scan rates (Figure 4.4a, Table 4.1). The Randles-Sevcik plot
shows a slope of 0.5, indicating that this complex is freely diffusing in solution (Figure 4.4b). This
122

electrochemical behavior is analogous to that of 5,5’-NH2-Re, which displays irreversible
reduction features at -2.11 and -2.47 V.
25
The electrochemical behavior of both 4,4’-NH2-Re and
5,5’-NH2-Re, however, differs from that of 6,6’-NH2-Re, which displays three irreversible
reduction events under N2 at -2.00 V, -2.11 V, and -2.23 V. The more negative reduction potentials
of 4,4’-NH2-Re and 5,5’-NH2-Re compared to those of 6,6’-NH2-Re suggest that these complexes
are more difficult to reduce, potentially due to the positions of the NH2 substituents on the
bipyridine backbone. Further, the lack of the additional intermediate feature observed in CVs of
6,6’-NH2-Re, which was attributed to intermolecular hydrogen bonding interactions, is not
apparent in CVs of 4,4’-NH2-Re and 5,5’-NH2-Re, suggesting that the position of the NH2 groups
on the bipyridine backbone influences the hydrogen bonding capabilities of the amine-substituted
Re(bpy) complexes.

Figure 4.4 (a) Variable scan rate CVs and (b) Randles-Sevcik plot of 4,4’-NH2-Re under N2.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6.

Table 4.1 Reduction potentials (V vs. Fc
+/0
) for 4,4’-NH2-Re, 5,5’-NH2-Re, and 6,6’-NH2-Re in
MeCN and DMF under N2. Scan rate: 100 mV/s.
Catalyst
1
st
reduction Intermediate reduction 2
nd
reduction
MeCN DMF MeCN DMF MeCN DMF
123

4,4’-NH 2-Re -2.14 -2.17 -- -- -2.46 -2.44
5,5’-NH 2-Re
25
-2.11 -2.21 -- -- -2.47 --
6,6’-NH 2-Re -2.00 -2.06 -2.11 -- -2.23 -2.42

Further electrochemical studies of 4,4’-NH2-Re and 5,5’-NH2-Re were performed in DMF
to investigate how the position of the NH2 substituent influences intermolecular hydrogen bonding
interactions in these complexes. Since DMF is a significantly better hydrogen bond donor and
acceptor than MeCN, it can disrupt intermolecular hydrogen bonding interactions, leading to an
observable change in electrochemical behavior. CVs of 4,4’-NH2-Re and 5,5’-NH2-Re in DMF
display only moderately different behavior than in MeCN, compared to 6,6’-NH2-Re, which differs
more significantly upon change in solvent (Figure 4.5, Table 4.1). Under N2, 4,4’-NH2-Re exhibits
two irreversible reduction events at -2.17 V and -2.44 V, analogous to the features observed in
MeCN. Similarly, 5,5’-NH2-Re exhibits a quasi-reversible reduction event at -2.21 V, slightly
more negative than the corresponding feature in MeCN; however, the lack of a second reduction
event and appearance of a small oxidation feature suggest that the solvent does slightly influence
the behavior of 5,5’-NH2-Re. Upon increasing the scan rate, the behavior of both 4,4’-NH2-Re and
5,5’-NH2-Re changes more substantially; the second reduction feature of 4,4’-NH2-Re grows in
more strongly, and the reduction features of both complexes become more quasi-reversible,
suggesting that at faster scan rates, these electrochemical processes can outcompete the irreversible
chemical ones that may influence hydrogen bonding (Figure 4.6). The position of the NH2 groups
on the bipyridine backbone does apparently influence the extent to which these amine-substituted
complexes can undergo hydrogen bonding interactions, based on the electrochemical differences
observed for each of these complexes upon switching from MeCN to DMF.
124

Figure 4.5 Overlay of cyclic voltammograms of 4,4’-NH2-Re (green), 5,5’-NH2-Re (purple), and
6,6’-NH2-Re (pink) under N2 in MeCN (dashed) and DMF (solid). Conditions: 1 mM catalyst with
0.1 M TBAPF6, scan rate: 100 mV/s.

125

Figure 4.6 Variable scan rate studies of (a) 4,4’-NH2-Re and (b) 5,5’-NH2-Re under N2.
Conditions: 1 mM catalyst in DMF with 0.1 M TBAPF6.

Upon saturation of the electrochemical solution with CO2, CVs of 4,4’-NH2-Re display a
moderate increase in current density with an onset more positive than the first reduction event
under N2 (Figure 4.7). The maximum current density reaches nearly 2 mA/cm2, greater than the
maximum current densities of both 5,5’-NH2-Re and 6,6’-NH2-Re (each <1 mA/cm
2
). Various
Brønsted acids (water, 2,2,2-trifluoroethanol, and phenol) were evaluated as proton sources for the
reduction of CO2 for both 4,4’-NH2-Re and 5,5’-NH2-Re (Figure 4.8). With the addition of 1 M
acid, 4,4’-NH2-Re displayed a trend similar to 6,6’-NH2-Re, with TFE providing the greatest
current densities at -2.40 V; while this is >200 mV more negative than the potential at which 6,6’-
NH2-Re achieved peak current density, the current increase is approximately three times greater.
5,5’-NH2-Re, however, reached the greatest current density with 1 M PhOH, also at -2.40 V.
Addition of water to both complexes leads to a negligible increase in current; for consistency, TFE
was used for all further electrochemical studies, since it provided the greatest current increases for
the majority of catalysts in the series.

Figure 4.7 CVs of 4,4’-NH2-Re under N2 (black), CO2 (red), N2 + 1 M TFE (blue), and CO2 + 1
M TFE (green). Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6; scan rate 100 mV/s.
126

Figure 4.8 CVs of 1 mM (a) 4,4’-NH2-Re and (b) 5,5’-NH2-Re under CO2 with 1 M PhOH
(orange), 1 M TFE (green), and 1 M H2O (blue) (0.1 M TBAPF6 in MeCN; scan rate: 100 mV/s).

Interestingly, 4,4’-NH2-Re and 5,5’-NH2-Re have catalytic onset potentials and peak
current density potentials that are more negative than those of 6,6’-NH2-Re (Figure 4.9). This
suggests that the position of the NH2 group likely plays an important role in the catalytic
mechanism. As has been observed for other CO2 reduction catalysts, available pendant proton
donors in close proximity to the metal center can help stabilize the CO2 adduct that forms during
catalysis, allowing the reduction of CO2 to occur with a lower overpotential. In the NH2-substituted
catalysts, this trend is only observed with addition of Brønsted acids and is not seen under CO2
alone, indicating that the acid source itself may participate in hydrogen bonding interactions that
influence catalysis, similar to our previous observations and calculations for a cobalt macrocycle
CO2 reduction catalyst.
13,14
The NH2 moieties in 4,4’-NH2-Re and 5,5’-NH2-Re are likely
positioned too far from the rhenium metal center to have a significant effect on the catalytic
potentials.
127

Figure 4.9 CVs of 4,4’-NH2-Re (green), 5,5’-NH2-Re (purple), and 6,6’-NH2-Re (pink) under CO2
with 1 M TFE. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6; scan rate 100 mV/s.

Addition of Brønsted acids under N2 was also performed for both 4,4’-NH2-Re and 5,5’-
NH2-Re (Figure 4.10). While water provides no current increase for 4,4’-NH2-Re under N2,
addition of 1 M TFE does lead to a slight increase in current density. Addition of 1 M PhOH leads
to a significant increase in current density, reaching approximately 4 mA/cm
2
, which is not
substantially lower than the current density achieved under CO2 (5.7 mA/cm
2
). Without the
presence of CO2 for catalytic CO production, these results imply that some other process, such as
proton reduction, may be occurring due to the presence of the added proton source. PhOH has a
smaller pKa in MeCN than TFE and water, making it a stronger proton donor; this likely explains
why the increase in current density is more substantial with the addition of PhOH compared to the
other Brønsted acids.
35
Interestingly, 5,5’-NH2-Re displays a similar increase in current density
with the addition of PhOH, reaching >7 mA/cm
2
, which is slightly greater than the current density
with 1 M PhOH under CO2. Since the maximum current density for 5,5’-NH2-Re occurs at a
potential more negative than that under CO2, this process may not compete with CO2 reduction at
less negative potentials.
128

Figure 4.10 CVs of 1 mM (a) 4,4’-NH2-Re and (b) 5,5’-NH2-Re under N2 with 1 M PhOH
(orange), 1 M TFE (green), and 1 M H2O (blue) (0.1 M TBAPF6 in MeCN; scan rate: 100 mV/s).

4.3.3 Controlled Potential Electrolysis

Controlled potential electrolysis (CPE) studies of 4,4’-NH2-Re were performed for one
hour under CO2 with added TFE (1 M) at various electrochemical potentials (Figure 4.11). Gas
chromatography (GC) was used for analysis of the headspace of the CPE cell after electrolysis.
All CPE results for 4,4’-NH2-Re are presented in Table 4.2. Current versus time plots display a
decrease in current over time at each potential, indicating long-term instability of this catalyst. The
highest catalytic activity (FECO = 91%, 65 µmol CO) is observed at -2.40 V, near the potential of
maximum current density observed in CV studies with added TFE. At each potential, only CO was
produced, with no other gaseous products such as H2 observed by GC. Similar to CPE studies of
6,6’-NH2-Re, 4,4’-NH2-Re displays a dependence of FECO on electrochemical potential, with
greater amounts of CO produced and higher FE values at more negative potentials. Since this trend
is only observed for the NH2-substituted rhenium bipyridine catalysts, it indicates that the
availability of pendant protons does play a role in catalytic activity, potentially as a proton shuttle
for the Brønsted acid. Additionally, the total amounts of CO produced by 4,4’-NH2-Re are
approximately two-fold lower than those produced by 6,6’-NH2-Re, with lower overall TON,
129

substantiating the beneficial positioning of the pendant amine near the rhenium metal center in
6,6’-NH2-Re.

Figure 4.11 Current versus time over one hour of controlled potential electrolysis of 4,4’-NH2-Re
at -2.10 V (orange), -2.20 V (green), -2.30 V (blue), and -2.40 V (purple). BE studies were
performed using 1 mM catalyst in MeCN with 1M TFE and 0.1 M TBAPF6.

Table 4.2 CPE results of 4,4’-NH2-Re in MeCN under CO2 with 1 M TFE.
Potential (V) FECO (%) FEH2 (%) μmol CO TON
-2.10 27 0 14 0.4
-2.20 45 0 19 0.5
-2.30 61 0 32 0.8
-2.40 91 0 65 1.6

4.4 CONCLUSION
In summary, a novel 4,4’-NH2 substituted rhenium bipyridine catalyst was synthesized and
characterized, and its electrochemical behavior was investigated and compared to the previously
studied 5,5’- and 6,6’-analogues. Electrochemical studies under both N2 and CO2 display key
differences between 4,4’-NH2-Re and 6,6’-NH2-Re, indicating that the position of the primary
amine on the bipyridine backbone influences the formation of hydrogen bonding interactions.
130

Additionally, catalytic studies demonstrate a more negative onset potential and catalytic
overpotential for 4,4’-NH2-Re and 5,5’-NH2-Re compared to 6,6’-NH2-Re. Although 4,4’-NH2-
Re shows a similar trend to 6,6’-NH2-Re of FECO dependence on electrochemical potential, the
lower overall production of CO and more negative optimized catalytic potential indicate that
positioning of the primary amine farther from the metal center negatively impacts electrocatalytic
activity.

4.5 EXPERIMENTAL METHODS AND ADDITIONAL FIGURES
4.5.1 Materials and Synthesis
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres drybox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. All solvents were degassed with nitrogen and passed through activated
alumina columns and stored over 4Å Linde-type molecular sieves. Deuterated solvents were dried
over 4Å Linde-type molecular sieves prior to use. Proton NMR spectra were acquired at room
temperature using Varian (Mercury 400 2-Channel, VNMRS-500 2-Channel, VNMRS- 600 3-
Channel, and 400- MR 2-Channel) spectrometers and referenced to the residual
1
H resonances of
the deuterated solvent (
1
H: CDCl3) and are reported as parts per million relative to
tetramethylsilane. Elemental analyses were performed using Thermo Scientific™ FLASH 2000
CHNS/O Analyzers. All the chemical reagents were purchased from commercial vendors and used
without further purification.

131

4.5.2 Electrochemistry
Electrochemistry experiments were carried out using a Pine potentiostat. The experiments
were performed in a single compartment electrochemical cell under nitrogen or CO 2 atmosphere
using a 3 mm diameter glassy carbon electrode as the working electrode, a platinum wire as the
auxiliary electrode and a silver wire as the reference electrode. Ohmic drop was compensated using
the positive feedback compensation implemented in the instrument. All electrochemical
experiments were performed with iR compensation using the current interrupt (RUCI) method in
AfterMath. Typical values for the cell resistance were around 160-170 ohms. All potentials in this
paper were referenced relative to ferrocene (Fc) with the Fe
3+/2+
couple at 0.0 V. Alternatively, in
cases when the redox couple of ferrocene overlapped with other features of interest,
decamethylferrocene (Fc*) was used as an internal standard with the Fe*
3+/2+
couple at –0.48 V.
All electrochemical experiments were performed with 0.1 M tetrabutylammonium
hexafluorophosphate (TBAPF6) as supporting electrolyte. The concentrations of the rhenium
catalysts were generally at 1 mM and experiments with CO2 were performed at gas saturation in
acetonitrile (MeCN) or dimethylformamide (DMF).
Controlled-potential electrolysis (CPE) measurements were conducted in a two-chamber
H cell. The first chamber held the working and reference electrodes in 40 mL of 0.1 M TBAPF6
and TFE in MeCN or DMF. The second chamber held the auxiliary electrode in 25 mL of 0.1 M
TBAPF6 in MeCN or DMF. The two chambers were separated by a fine porosity glass frit. The
reference electrode was placed in a separate compartment and connected by a Vycor tip. Glassy
carbon plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and
auxiliary electrodes. 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)
132

equipped with a BID detector and a Restek ShinCarbon ST Micropacked column. Faradaic
efficiencies were determined by dividing the measured CO produced by the amount of CO
calculated based on the charge passed during the CPE experiment. The reported Faradaic
efficiencies and mol of CO produced are average values.

4.5.2 Synthesis of 4,4’-NH2-Re
Under N2, 4,4’-NH2(bpy) (102 mg, 0.55 mmol) and Re(CO)5Cl (199 mg, 0.55 mmol)
were refluxed in 30 mL anhydrous toluene for 18 hours. After cooling, the resulting precipitate
was collected by vacuum filtration and washed with ether, yielding a yellow solid in
approximately 94% yield.

133

4.5.3 Additional Figures

Figure 4.12 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations of
H2O. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 4.13 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations of
TFE. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

134

Figure 4.14 Cyclic voltammograms of 4,4’-NH2 under CO2 with increasing concentrations of
PhOH. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 4.15 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations of H2O.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

135

Figure 4.16 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations of TFE.
Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 4.17 Cyclic voltammograms of 4,4’-NH2 under N2 with increasing concentrations of
PhOH. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

136

Figure 4.18 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations of
H2O. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

Figure 4.19 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations of
TFE. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

137

Figure 4.20 Cyclic voltammograms of 5,5’-NH2 under CO2 with increasing concentrations of
PhOH. Conditions: 1 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

138

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142

CHAPTER 5

Immobilization of a Molecular Rhenium Bipyridine Electrocatalyst for CO2 Reduction

143

5.1 ABSTRACT
Immobilization of well-defined molecular complexes is an attractive method for combining
the advantages of homogeneous and heterogeneous electrocatalysis. Herein, we present
modification of a variety of electrode substrates by a [2,2'-bipyridine]-4,4'-bis(diazonium) rhenium
complex (4,4’-N2
+
-Re) via radical electropolymerization of the diazonium substituent. Post-
modification characterization and electrochemical studies demonstrate the successful attachment
of the rhenium bipyridine moieties, although with variance in the quantity and electrochemical
behavior depending on number of grafting scans. Further studies, including additional
characterization and electrocatalysis, will be done to determine the extent of control over grafting
and film thickness and the overall electrocatalytic activity for CO2 reduction of these films.

5.2 INTRODUCTION
Today’s extenuating global energy demands put an enormous pressure on finding suitable
alternative energy sources. While harnessing the energy of the sun, which can provide enough
energy in one hour to power human activity for an entire year, would be ideal, the diffuse,
intermittent nature and storage hindrances provide challenges.
1
It is thus necessary to find a way
to store solar energy for use at night and during poor weather conditions, when sunlight is not
readily available.
2,3
Solar-powered electrocatalysis is an attractive method for utilizing the sun as
a clean energy source while simultaneously recycling harmful byproducts due to fossil fuel use,
such as carbon dioxide (CO2); this process allows for the conversion of toxic CO2 into carbon
monoxide (CO), which can then be used for the production of other chemical feedstocks.
4

Two main classes of electrocatalysts – homogeneous and heterogeneous – have been
developed to successfully undergo the transformation of CO2 to CO.
5
In general, homogeneous
144

(molecular) catalysts tend to possess well-defined active sites and tunability via the ligand scaffold,
but they can suffer from long-term instability and limited recyclability.
4
Heterogeneous (surface-
bound) catalysts, on the other hand, often have more ill-defined active sites and limited tunability
due to synthetic challenges, but they have the advantage of increased recyclability and stability.
6

One method for maximizing the benefits of these catalysts is by attaching a molecular complex to
an electrode via heterogenization, allowing for the more well-defined active sites of molecular
catalysts while maintaining the increased stability and recyclability of surface-bound catalysts.
7

Rhenium bipyridine catalysts [Re(bpy)(CO)3Cl, (bpy = 2,2′-bipyridine)] have been well-
studied thus far in the literature, proving to be efficient and selective electrocatalysts for the
reduction of CO2 to CO.
8–10
However, these catalysts can suffer from dimerization, deactivation,
and poor long-term stability. Re(bpy) catalysts have also been employed in heterogenization
efforts via both covalent and non-covalent methods, with fairly good success.
11–13
However, one
important limitation with these methods is that the type of electrode to which the catalysts can be
bound is limited due to the method of immobilization.
Electropolymerization is an enticing alternative to other covalent and non-covalent
methods for immobilization of a molecular complex.
14
This method has been employed previously
for heterogenization of a rhenium 4-vinyl-4'-methyl-2,2'-bipyridine complex (Re(vbpy)(CO)3Cl)
onto platinum disk electrodes and for semiconductor materials, resulting in high activity for the
reduction of CO2 to CO.
15–18
However, the methylene spacers and vinyl radicals generated via
electropolymerization of this complex led to side products due to radical-radical coupling
reactions, and the films were thus limited in stability and prone to decomposition. However, this
issue can be counteracted by altering overall structure and the functional group used for grafting;
in particular, introducing rigidity and conjugation into the polymer chains allows for increased
145

structural control and reduced side reactions.
19
An alkyne-substituted rhenium bipyridine complex
with increased rigidity was shown to be active for CO2 reduction, although with a low Faradaic
efficiency (FE) for CO (33%).
20
Our group has demonstrated that altering the radical-forming
functional group by introducing diazonium substituents allows for both a more rigid polymer
structure and increased activity and stability.
21,22
Further advantages include a wide substrate
scope, long-range order, and control over structure and film thickness via number of grafting scans
and electrochemical potential window.
Based on our previous work on immobilized rhenium bipyridine conjugated polymers and
our investigations into positional effects of molecular NH2-substituted Re(bpy)(CO)3Cl catalysts,
we sought to explore the effect of diazonium (and thus grafting) position along the bipyridine
backbone.
21,22
We present here the synthesis and characterization of a new diazonium-substituted
rhenium bipyridine complex (4,4’-N2
+
-Re, Scheme 5.1), which was subsequently employed for
immobilization on a variety of substrates. Modification of the ligand structure allows for variance
in electropolymerization directionality and polymer chain flexibility. Following substrate
modification, each electrode was studied by a variety of surface characterization methods for
structure determination and comparison to the 5,5’-analogue, and the electrochemical properties
and catalytic activity and stability were thoroughly investigated.

146

5.3 RESULTS AND DISCUSSION
5.3.1 Synthesis and Characterization

Scheme 5.1 Synthesis of 4,4’-N2
+
-Re from 4,4’-NH2-Re.

The diazonium complex 4,4’-N2
+
-Re was synthesized from 4,4’-NH2-Re according to
literature precedent for the synthesis of 5,5’-N2
+
-Re (Scheme 5.1).
21
Under N2, a solution of
nitrosonium tetrafluoroborate (NOBF4) in acetonitrile (MeCN) was added dropwise to a stirring
suspension of 4,4’-NH2-Re in MeCN at -40°C, leading to an immediate color change from yellow
to dark purplish-black. Upon addition of diethyl ether, purplish-black solid precipitated from
solution and was isolated by vacuum filtration.
1
H nuclear magnetic resonance (NMR) of the
product in MeCN-d3 displays three aromatic signals at  9.71 ppm,  9.34 ppm, and  8.66 ppm,
corresponding to the protons on the bipyridine backbone (Figure 5.1). These peaks are shifted
downfield from those of 4,4’-NH2-Re due to the electron withdrawing nature of the diazonium
substituent. Further, the absence of the broad amine signal observed for 4,4’-NH2-Re indicates the
successful conversion to 4,4’-N2
+
-Re. The
19
F NMR spectrum of 4,4’-N2
+
-Re displays two signals
at  -150.40 and -150.96 ppm, representing the BF4
-
counterion (Figure 5.1b).
147

Figure 5.1 (a)
1
H NMR of 4,4’-N2
+
-Re in MeCN-d3 (400 MHz) and (b)
19
F NMR of 4,4’-N2
+
-Re
in MeCN-d3 (500 MHz).

4,4’-N2
+
-Re was further characterized by UV-vis spectroscopy and Fourier Transform
Infrared (FTIR) spectroscopy. The UV-vis spectrum shows two broad features at 450 and 570 nm
(Figure 5.2). Additionally, a feature at 354 nm corresponds to the MLCT band for the rhenium
bipyridine moiety, as was originally observed for 4,4’-NH2-Re. The FTIR spectrum displays three
characteristic carbonyl stretches at 2023, 1917, and 1894 cm
-1
(Figure 5.3). An additional
diazonium stretch appears at 2115 cm
-1
.

148

Figure 5.2 UV-vis spectrum of 4,4’-N2
+
-Re in MeCN.

Figure 5.3 FTIR spectrum of 4,4’-N2
+
-Re.

149

5.3.2 Electropolymerization of 4,4’-N2
+
-Re

Scheme 5.2 Scheme demonstrating electrochemical grafting of 4,4’-N2
+
-Re onto various electrode
surfaces by cyclic voltammetry and subsequent electropolymerization. Conditions: 0.5 mM
complex in MeCN with 0.1 M TBAPF6 supporting electrolyte.

A glassy carbon (GC) stick electrode was immersed in a solution of 4,4’-N2
+
-Re in MeCN
(0.5 mM in 10 mL with 0.1 M TBAPF6 supporting electrolyte). Consecutive cyclic voltammetry
(CV) scans were performed at 100 mV/s to facilitate electropolymerization of 4,4’-N2
+
-Re on the
electrode surface (Figure 5.4a, Scheme 5.2). Upon the initial reduction scan, a broad irreversible
reduction event occurs at -1.13 V versus Fc
+/0
(all subsequent potentials are referenced versus the
ferrocenium/ferrocene couple). This reduction feature appears at a potential 42 mV more negative
than the analogous feature observed for 5,5’-N2
+
-Re under similar conditions; however, the feature
is still at a significantly more positive potential than the features typically observed for molecular
rhenium bipyridine complexes, including the precursor 4,4’-NH2-Re (-2.14 V). This broad
reduction feature can be assigned to the one-electron reduction of the diazonium group, as was
previously reported for 5,5’-N2
+
-Re.
21
Similar reduction behavior has also been observed for other
reported diazonium complexes, demonstrating that, upon the addition of an electron, dinitrogen is
released and an aryl radical is formed near the electrode surface, promoting
electropolymerization.
23
Subsequent CV scans show a negative shift in the potential of this
reduction event, similar to consecutive scans reported for 5,5’-N2
+
-Re; this behavior indicates an
150

impedance in charge transport between the solution and the electrode surface, as is expected for
the formation of a molecular film.
24
Further, the increase in current density with each scan indicates
the successful electropolymerization, as a larger number of rhenium bipyridine species are present
at the electrode surface with each consecutive CV.
25
Interestingly, applying the same grafting
parameters for electropolymerization of a fluorine-doped tin oxide (FTO) electrode yields no
discernible reduction events (Figure 5.5).

Figure 5.4 Cyclic voltammograms of electropolymerization of 4,4’-N2
+
-Re on a glassy carbon
stick electrode showing five consecutive scans from (a) -0.6 V to -1.6 V and (b) -0.6 V to -2.6 V.
Conditions: 0.5 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate: 100 mV/s.

151

Figure 5.5 Cyclic voltammograms of electropolymerization of 4,4’-N2
+
-Re on an FTO electrode
showing five consecutive scans from -0.6 V to -1.6 V. Conditions: 0.5 mM catalyst in MeCN with
0.1 M TBAPF6. Scan rate: 100 mV/s.

If the switching potential (Ps) is limited to -1.6 V, a single reduction event is observed for
each CV; however, if the electrochemical window is expanded and P s is increased to -2.6 V,
additional features are observed (Figure 5.4b). In the initial CV scan, two additional irreversible
reduction events are observed at -2.18 V and -2.33 V, nearer the potential at which the reduction
events of the molecular complex appear. Only the first of the two features is present on the second
CV scan, and with subsequent scans, the overall current density decreases substantially and the
reduction events disappear altogether. This behavior at more negative potentials is markedly
different from 5,5’-N2
+
-Re, for which the broad feature is still present and current density decreases
only slightly by the ninth CV when Ps is -2.6 V.
21
Since the disappearance of these reduction events
can be attributed to hindered electrode kinetics with film growth, this divergence in behavior
indicates that the direction of film growth is likely very different between the 4,4’- and 5,5’-
analogues due to the position of the diazonium group on the bipyridine backbone.
26
Performing
CVs with a Ps of -2.6 V rapidly increases the rate of electropolymerization for 4,4’-N2
+
-Re, likely
occurring too quickly for ideal film growth. Further, the more negative potentials observed for the
diazonium reduction events of 4,4’-N2
+
-Re compared to those observed for 5,5’-N2
+
-Re, even with
a less negative Ps of -1.6 V, substantiates the difference in film growth between the two analogues.
4,4’-N2
+
-Re appears to undergo more rapid film growth than does 5,5’-N2
+
-Re, likely also
hindering electrode kinetics more quickly.
A standardized grafting methodology for use with a variety of electrodes was replicated
from the previous work on 5,5’-N2
+
-Re.
21
CV scans were performed using an initial potential (Pi)
152

of -0.6 V and a Ps of -1.6 V, with a consistent scan rate of 1 V/s to minimize the Helmholtz layer
and trapping of electrolyte during electropolymerization, yielding a more uniform film.
27

Additionally, subsequent CVs, rather than controlled potential electrolysis (CPE), were employed
to further minimize trapped charges throughout the film. All electrodes, including glassy carbon,
FTO, gold, and graphite rod electrodes, were modified using consecutive CV scans (n = total
number of grafting scans) and used for further electrochemical and characterization studies. After
n consecutive grafting scans were performed at 1 V/s, each electrode was rinsed with MeCN and
acetone to remove any physisorbed diazonium species or other side products. As opposed to 5,5’-
N2
+
-Re, electrodes grafted with 4,4’-N2
+
-Re displayed no color change when the scan rate used
was 1 V/s; electrodes grafted using a scan rate of 100 mV/s displayed a slight coloration when n ≥
10. The difference in observable color change upon grafting suggests qualitatively that the films
produced by electropolymerization of 4,4’-N2
+
-Re are of a lower thickness than those produced
using 5,5’-N2
+
-Re.

5.3.3 Characterization of Electropolymerized Films
5.3.3.1 XPS
Modified FTO electrodes were analyzed by high-resolution X-ray photoelectron
spectroscopy (XPS) (Figure 5.6). Rhenium 4f features are present for all four electrodes after
solvent washing and sonication, suggesting robust covalent attachment of the diazonium complex
to the electrode surface (Figure 5.7). However, there is not a clear trend in the intensity of these
features with increasing number of grafting scans, indicating that there may not be as much order
in the films made with 4,4’-N2
+
-Re compared to the linear wires generated using 5,5’-N2
+
-Re. For
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the n = 1 electrode, high-resolution Sn 2p features are clearly present, whereas these features
disappear for films with n ≥ 5 (Figure 5.8). These Sn features correspond to the underlying FTO
substrate; passivation with increasing scans suggests that film thickness has exceeded the XPS
sampling depth (5-10 nm), indicating at least minimal control over film thickness with grafting
scans. Interestingly, high-resolution Cl 2p spectra show the absence of Cl for each grafted film,
indicating chloride dissociation as part of the electropolymerization mechanism for this complex
(Figure 5.9). Additionally, P 2p spectra and F 1s spectra display features representing trapped PF6
-

anion from the electrolyte during electropolymerization and BF4
-
anion from the diazonium
species, respectively (Figures 5.10-5.11). High-resolution N 1s and C 1s spectra also indicate the
presence of these elements on the electrode surface (Figures 5.12-5.13).

Figure 5.6 XPS survey scans for modified FTO electrodes (a) n = 1, (b) n= 5, (c) n = 10, and (d)
n = 20.

154

Figure 5.7 High-resolution XPS of the Re 4f region for modified FTO electrodes (a) n = 1, (b) n
= 5, (c) n = 10, and (d) n = 20.

Figure 5.8 High-resolution XPS of the Sn 2p region for modified FTO electrodes (a) n = 1 and (b)
n = 20.

155

Figure 5.9 High-resolution XPS of the Cl 2p region for a modified FTO electrode with n = 20.

Figure 5.10 High-resolution XPS of the P 2p region for a modified FTO electrode with n = 20.

156

Figure 5.11 High-resolution XPS of the F 1s region for a modified FTO electrode with n = 20.

Figure 5.12 High-resolution XPS of the N 1s region for a modified FTO electrode with n = 20.

157

Figure 5.13 High-resolution XPS of the C 1s region for a modified FTO electrode with n = 20.

5.3.4 Electrochemistry of Films

Figure 5.14 Cyclic voltammograms of modified GC stick electrodes with n = 1 (red), 5 (orange),
10 (green), and 20 (blue). Conditions: 0.5 mM catalyst in MeCN with 0.1 M TBAPF6. Scan rate:
100 mV/s.

After grafting, all modified electrodes were analyzed by cyclic voltammetry under N2 in
MeCN with 0.1 M TBAPF6 supporting electrolyte. For each electrode type (GC stick, FTO, and
158

gold), four individual electrodes were modified, with n = 1, 5, 10, and 20, for consistent
comparison with 5,5’-N2
+
-Re. CVs of modified GC stick electrodes are shown in Figure 5.14. For
n = 20, an irreversible reduction event occurs at -1.13 V, similar to the behavior observed for 5,5’-
N2
+
-Re; this feature has been previously ascribed to a buildup and subsequent discharge of excess
electrolyte within the polymerized film.
16,25,28
The observance of this feature only for n = 20 is
again indicative of thinner film production using 4,4’-N2
+
-Re, as the charge transport hindrance
and electrolyte buildup are not evident by CV for n ≤ 10. An additional broad quasi-reversible
reduction event is observed at -2.09 V for n = 10 and 20, assigned to a two-electron reduction of
the electropolymerized films.
21
The convalescence of the two previously observed reduction
events into a single broad feature is likely due to hindered electron kinetics with thicker films.
29

The peak separation (ΔE = 80 mV) at n = 20 suggests that films of this thickness behave more like
a diffusion-controlled species due to the flexibility of longer polymer chains.
21
Additionally, the
lack of this redox event with lower film thicknesses is likely due to the more ill-defined polymer
chains of 4,4’-N2
+
-Re compared to 5,5’-N2
+
-Re. Randles Sevcik analyses based on variable scan
rate studies were performed for each electrode using variable scan rate studies (Figures 5.15-5.16).
Interestingly, these plots show a trend of increasing slope corresponding to the number of CV
scans: for n = 1, the slope is near 0.5, whereas for n = 20, the slope is about 1. This trend is not
observed for the polymer chains of 5,5’-N2
+
-Re, which display the ideal value near 1 for all n >
1.
21

159

Figure 5.15 Variable scan rate studies of modified glassy carbon electrodes with (a) n = 1, (b) n =
5, (c) n = 10, and (d) n = 20. Conditions: MeCN with 0.1 M TBAPF6.

160

Figure 5.16 Randles-Sevcik plots for modified glassy carbon electrodes with (a) n = 1, (b) n = 5,
(c) n = 10, and (d) n = 20.

Due to the lack of distinct redox features, CVs could not be used to determine the
electroactive rhenium sites.
30
However, the electroactive surface coverage for each modified
electrode was estimated using double layer capacitance (Cdl) measurements. Variable scan rate
CVs were performed in a 100 mV window around the open circuit potential (OCP), yielding ΔJ
versus scan rate plots with linear slopes (Figure 5.17). Unlike films of electropolymerized 5,5’-
N2
+
-Re, films formed using 4,4’-N2
+
-Re do not display a linear correlation between Cdl and number
of grafting scans; this suggests that there is less control over film growth for the 4,4’-analogue,
likely due to the lack of a single directionality for bond formation. This behavior is observed for
modified FTO electrodes as well.
161

Figure 5.17 Double-layer charging current density (ΔJ = Janodic-Jcathodic) at the open-circuit
potential for modified glassy carbon electrodes with (a) n = 1, (b) n = 5, (c) n = 10, and (d) n = 20
as a function of scan rate.

5.3.5 Electrocatalytic Studies
5.3.5.1 Cyclic Voltammetry
Cyclic voltammograms of modified glassy carbon electrodes were performed under
saturated CO2 conditions to determine electrocatalytic activity for CO2 reduction. Upon switching
the atmosphere from N2 to CO2, each electrode displays a slight increase in current density (except
for n = 5, which displays a slight drop in current) (Figure 5.18). Similar to the Cdl measurements,
the increase in current density does not correlate with the total number of grafting scans, supporting
lesser control over electropolymerization and film growth for 4,4’-N2
+
-Re compared to 5,5’-N2
+
-
162

Re. The addition of 0.5 M of the Brønsted acid 2,2,2-trifluoroethanol (TFE) under CO2 led to a
more substantial increase in current density for n = 1, 10, and 20 (Figure 5.18). In particular, for n
= 20, the current density achieved >3 mA/cm2, more than three times the current achieved with the
other films, even though the Cdl measurement resulted in the smallest value. This suggests that
thicker films are not optimal for electrocatalytic activity, although a certain threshold is likely
necessary. CVs of modified FTO electrodes under the same conditions yielded similar behavior
for n = 1, 5, 10, and 20, with moderate enhancement in current density upon saturation with CO2
and subsequent addition of 0.5 M TFE (Figure 5.19).

Figure 5.18 CVs of modified glassy carbon electrodes with (a) n = 1, (b) n = 5, (c) n = 10, and (d)
n = 20 under N2 (blue), CO2 (red), and CO2 + 0.5 M TFE (green). Conditions: 0.1 M TBAPF6 in
MeCN; scan rate 100 mV/s.

163

Figure 5.19 CVs of modified FTO electrodes with (a) n = 1, (b) n = 5, (c) n = 10, and (d) n = 20
under N2 (blue), CO2 (red), and CO2 + 0.5 M TFE (green). Conditions: 0.1 M TBAPF6 in MeCN;
scan rate 100 mV/s.

5.4 CONCLUSION
While the studies presented here are still preliminary, we have demonstrated the successful
electropolymerization of a novel 4,4’-diazonium bipyridine complex onto a variety of electrode
surfaces, including glassy carbon and FTO. XPS and electrochemical studies indicate a lower
degree of film order than that observed for electropolymerized films of the previously reported
5,5’-analogue, likely due to the positioning of the diazonium groups on the bipyridine backbone.
However, further characterization studies, including infrared reflection absorption spectroscopy
(IRRAS), scanning electron microscopy (SEM), atomic force microscopy (AFM), and inductively
coupled plasma (ICP) will need to be performed to analyze catalyst loading and morphology and
to determine the extent of control of this grafting method on both film thickness and film order.
164

Additionally, electrocatalytic studies using graphite rod electrodes for higher surface area coverage
will be completed to analyze the catalytic activity of these films for CO2RR.

5.5 EXPERIMENTAL METHODS
All manipulations of air and moisture sensitive materials were conducted under a nitrogen
atmosphere in a Vacuum Atmospheres drybox or on a dual manifold Schlenk line. The glassware
was oven-dried prior to use. All solvents were degassed with nitrogen and passed through activated
alumina columns and stored over 4Å Linde-type molecular sieves. Deuterated solvents were dried
over 4Å Linde-type molecular sieves prior to use. Proton NMR spectra were acquired at room
temperature using Varian (Mercury 400 2-Channel, VNMRS-500 2-Channel, VNMRS- 600 3-
Channel, and 400- MR 2-Channel) spectrometers and referenced to the residual
1
H resonances of
the deuterated solvent (
1
H: MeCN-d3) and are reported as parts per million relative to
tetramethylsilane. Elemental analyses were performed using Thermo Scientific™ FLASH 2000
CHNS/O Analyzers. All the chemical reagents were purchased from commercial vendors and used
without further purification.
Electrochemistry experiments were carried out using a Pine potentiostat. The experiments
were performed in a single compartment electrochemical cell under nitrogen or CO 2 atmosphere
using a 3 mm diameter glassy carbon electrode as the working electrode, a platinum wire as the
auxiliary electrode and a silver wire as the reference electrode. Ohmic drop was compensated using
the positive feedback compensation implemented in the instrument. All electrochemical
experiments were performed with iR compensation using the current interrupt (RUCI) method in
AfterMath. Typical values for the cell resistance were around 160-170 ohms. All potentials in this
paper were referenced relative to ferrocene (Fc) with the Fe
3+/2+
couple at 0.0 V. Alternatively, in
165

cases when the redox couple of ferrocene overlapped with other features of interest,
decamethylferrocene (Fc*) was used as an internal standard with the Fe*
3+/2+
couple at –0.48 V.
All electrochemical experiments were performed with 0.1 M tetrabutylammonium
hexafluorophosphate (TBAPF6) as supporting electrolyte. The concentrations of the rhenium
catalysts were generally at 1 mM and experiments with CO2 were performed at gas saturation in
acetonitrile (MeCN) or dimethylformamide (DMF).
Controlled-potential electrolysis (CPE) measurements were conducted in a two-chamber
H cell. The first chamber held the working and reference electrodes in 40 mL of 0.1 M TBAPF6
and TFE in MeCN or DMF. The second chamber held the auxiliary electrode in 25 mL of 0.1 M
TBAPF6 in MeCN or DMF. The two chambers were separated by a fine porosity glass frit. The
reference electrode was placed in a separate compartment and connected by a Vycor tip. Glassy
carbon plate electrodes (6 cm × 1 cm × 0.3 cm; Tokai Carbon USA) were used as the working and
auxiliary electrodes. 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. Faradaic
efficiencies were determined by dividing the measured CO produced by the amount of CO
calculated based on the charge passed during the CPE experiment. The reported Faradaic
efficiencies and mol of CO produced are average values.

Synthesis of 4,4’-N2
+
-Re: 4,4’-NH2-Re (45 mg) was suspended in acetonitrile (1.9 mL) under N2.
Separately, a solution of nitrosonium tetrafluoroborate (NOBF4) (24 mg) was dissolved in a
minimal amount of acetonitrile (0.9 mL). Both solutions were cooled to -40 °C. Once cooled, the
suspension of 4,4’-NH2-Re was added dropwise to the NOBF4 solution, leading to an immediate
166

color change from yellow to dark purple/black. Addition of diethyl ether (~6 mL) resulted in the
formation of a dark purple precipitate, which was collected by vacuum filtration and stored in the
dark at -27 °C (
1
H in CD3CN: δ 9.71, 9.34 and 8.66 ppm).

Physical 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 1486.6 eV. Low-resolution survey spectra were acquired between
binding energies of 1–1200 eV. Higher-resolution detailed scans, with a resolution of ~0.1 eV,
were collected on individual XPS lines of interest. The sample chamber was maintained at < 2 ×
10–9 Torr. The XPS data were analyzed using the CasaXPS software.
FT-IR
FT-IR spectra were acquired using a Bruker Vertex 80v spectrometer. Reflectance spectra
were collected using a VeeMAX III specular reflectance accessory from Pike Instruments.
Samples were positioned face-down over an aperture (3/8” diameter). All IRRAS measurements
were collected with a 56° angle of incidence under vacuum pressure with 1 cm-1 resolution.
Polarization studies were performed with a ZnSe polarizing lens purchased from Pike Instruments.
The spectra measured for unmodified substrates under identical experimental parameters (angle,
polarization, and resolution) were subtracted as background. Studies were performed at 1 cm-1
resolution with 128 scans. ATR-FTIR measurements for complex 2 were performed using a Bruker
Optics Alpha FTIR spectrometer in the ATR mode.
167

UV-Vis
UV-Vis spectra were collected using a UV-1800 Shimadzu UV spectrophotometer. FTO
samples were studied in transmittance mode and the spectrum measured for an unmodified FTO
substrate was subtracted as background.

168

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Abstract (if available)

Abstract The conversion of harmful carbon dioxide into value added products is an advantageous method for creating a carbon neutral infrastructure and promoting renewably energy sources. In nature, the enzyme CO dehydrogenase performs this process with high rate and efficiency due to hydrogen bonding stabilization of the CO₂ adduct by neighboring amino acids. We emulated this process by introducing pendant amines onto highly active rhenium bipyridine moieties, which have been found to selectively produce CO. The inclusion of pendant protons led to the formation of intermolecular hydrogen bonding interactions, which was studied further with electrochemical studies on 6,6'-substituted rhenium bipyridine complexes with primary, secondary, and tertiary amines. As catalysts with 6,6'-primary amine substituents demonstrated a dependence of product selectivity on electrocatalytic potential, the substituent position on the bipyridine backbone was altered to explore the effects of the hydrogen bonding interactions and position of pendant protons in relation to the active metal center. Finally, the 4,4'-substituted complex was electropolymerized onto various electrode surfaces to determine the effect of substituent position on film growth and electrocatalytic activity.

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Creator Hellman, Ashley N. (author)

Core Title Rhenium bipyridine catalysts with pendant amines: substituent and positional effects on the electrocatalytic reduction of CO₂ to CO

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Publication Date 03/18/2021

Defense Date 03/03/2021

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

Tag amines,carbon dioxide,carbon neutral,catalysis,CO production,CO₂ reduction,electrocatalysis,electrochemistry,electropolymerization,grafting,heterogeneous,homogeneous,hydrogen bonding,inorganic chemistry,OAI-PMH Harvest,pendant protons,renewable energy,rhenium bipyridine

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor Marinescu, Smaranda C. (committee chair), Malmstadt, Noah (committee member), Thompson, Mark (committee member)

Creator Email ashley892@sbcglobal.net,hellmana@usc.edu

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