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Studying iron and nickel analogues of an efficient carbon dioxide reduction electrocatalyst
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Studying iron and nickel analogues of an efficient carbon dioxide reduction electrocatalyst
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i
STUDYING IRON AND NICKEL ANALOGUES OF AN EFFICIENT
CARBON DIOXIDE REDUCTION ELECTROCATALYST
by
Geovanni M. Rangel
A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirements for the Degree
MASTER OF SCIENCE
(CHEMISTRY)
August 2019
ii
ACKNOWLEDGEMENTS
Thank you to Prof. Smaranda C. Marinescu and the Marinescu group. A special thank you to my
close friends Carlos, Damir, and Keying.
iii
TABLE OF CONTENTS
Acknowledgements .......................................................................................................................... ii
Table of Contents ........................................................................................................................... iii
List of Figures ................................................................................................................................ iv
List of Schemes ............................................................................................................................... vi
List of Tables ................................................................................................................................. vii
Chapter 1. General Introduction ....................................................................................................1
1.1 Global Energy Perspective .............................................................................................2
1.2 Carbon Dioxide Conversion ..........................................................................................2
1.3 Homogeneous Catalysts for CO2 Conversion ................................................................3
1.4 Co Complexes Bearing Azacalix[4](2,6)pyridine Ligands ...........................................5
1.5 References ......................................................................................................................9
Chapter 2. Physical and Electrochemical Characterization of Fe and Ni Analogues of an Efficient
Co-based CO2-to-CO Electrocatalyst bearing an Azacalix[4](2,6)pyridine Macrocycle .................13
2.1 Introduction ..................................................................................................................14
2.2 Results and Discussion ................................................................................................16
2.3 Conclusion ...................................................................................................................22
2.4 References ....................................................................................................................23
2.5 Supplemental Information for Chapter 2 .....................................................................25
2.6 Supplemental Information References .........................................................................47
iv
LIST OF FIGURES
Figure 1.1 Heterogeneous and Homogeneous Electrochemical CO2 Reduction ............................4
Figure 1.2 Co Complexes Bearing Azacalix[4](2,6)pyridine Ligands ...........................................5
Figure 1.3 Electrochemical Studies of CoL
1
..................................................................................6
Figure 1.4 kobs vs Number of Pendant NH Groups of CoL
1
...........................................................7
Figure 1.5 EECC Catalytic Cycle of CoL
1
.....................................................................................8
Figure 2.1 Crystal Structures of FeL
1
and NiL
1
..........................................................................16
Figure 2.2 UV-Vis Spectrum of FeL
1
, CoL
1
, and NiL
1
...............................................................18
Figure 2.3 CVs of FeL
1
, CoL
1
, and NiL
1
.....................................................................................18
Figure 2.4 TFE CV Titrations of FeL
1
, CoL
1
, and NiL
1
..............................................................20
Figure 2.5 DFT-determined Molecular Orbital Diagrams of the ML
1
Series ..............................22
Figure S1. FTIR Spectrum of FeL
1
..............................................................................................29
Figure S2. FTIR Spectrum of CoL
1
..............................................................................................29
Figure S3. FTIR Spectrum of NiL
1
...............................................................................................30
Figure S4. Beer-Lambert Plot of FeL
1
..........................................................................................30
Figure S5. Beer-Lambert Plot of CoL
1
.........................................................................................31
Figure S6. Beer-Lambert Plot of FeL
1
..........................................................................................31
Figure S7. CVs of NiL
1
at Varying Scan Rates ............................................................................32
Figure S8. Plot of Peak Current Density vs Scan Rate of NiL
1
....................................................32
Figure S9. CVs of FeL
1
at Varying Scan Rates ............................................................................33
Figure S10. Plot of Peak Current Density vs Scan Rate of FeL
1
..................................................34
Figure S11. Multiple CV Scans of FeL
1
under CO2 .....................................................................35
Figure S12. Multiple CV Scans of NiL
1
under CO2 .....................................................................36
Figure S13. Multiple CV Scans of FeL
1
under N2 .......................................................................37
v
Figure S14. Multiple CV Scans of NiL
1
under N2 ........................................................................38
Figure S15. CPE of FeL
1
at -2.68 V .............................................................................................38
Figure S16. CPE of FeL
1
at -2.40 V .............................................................................................39
Figure S17.CPE of NiL
1
...............................................................................................................39
Figure S18. CV and CPE of FeL
1
-GCE .......................................................................................40
Figure S19. CV and CPE of NiL
1
-GCE .......................................................................................40
Figure S20. SEM Images of FeL
1
-GCE .......................................................................................40
Figure S21. SEM Images of CoL
1
-GCE ......................................................................................41
Figure S22. SEM Images of NiL
1
-GCE .......................................................................................41
Figure S23. XPS Spectra of ML
1
-GCE Series .............................................................................42
vi
LIST OF SCHEMES
Scheme 2.1 Synthesis of ML
1
complexes .....................................................................................18
vii
LIST OF TABLES
Table 1.1. Selected CO2 Reduction Processes ................................................................................3
Table 2.1 Average relevant bond distances of FeL
1
, CoL
1
, and NiL
1
.........................................17
Table 2.2 Summary of Electrochemical Properties of the ML
1
series..........................................19
Table S1. Crystal data and structure refinement for FeL
1
............................................................42
Table S1. Crystal data and structure refinement for NiL
1
.............................................................44
1
CHAPTER 1
General Introduction
2
1.1 Global Energy Perspective
Supplying the world with reliable, clean, and sustainable energy remains a scientific challenge of
upmost importance. The global energy consumption is expected to at least double by the middle
of the century due to population and economic growth.
1
Even though the global average energy
intensity has declined in the past 100 yr from improvements in energy related technology, energy
consumption is expected to increase from 13.5 TW in 2001 to 27 TW by 2050.
2
Studies included
in the World Energy Assessment Report estimate that with the world’s total fossil fuel reserves,
which includes coal, oil, natural gas, and shales, would be able to support a 25 to 30 TW energy
consumption rate for at least several centuries.
1
However, consuming energy on such a scale with
our current energy system that is dominated by fossil fuels incurs detrimental consequences to the
environment.
The global dependence on fossil fuels has led to a rapid increase in carbon dioxide (CO2) emissions
since the industrial revolution, surpassing an atmospheric concentration of 400 ppm in the year
2015.
3
Given that high concentrations of atmospheric CO2 result in a stronger greenhouse effect,
contributing largely the rapid changing climate of our planet, there is increasing discussion among
the scientific community of CO2 mitigation in our energy system.
4
1.2 Carbon Dioxide Conversion
Until renewable energy technologies have advanced enough to compete with fossil fuels, the
global CO2 emissions will continue to increase.
1
This has lead to a dramatic growth of research
interest in the catalytic conversion of CO2 in recent years, which tackles the CO2 problem through
“CO2 utilization” which ultimately leads to “CO2 mitigation”. CO2 conversion is a pathway to a
sustainable future for two reasons. 1) CO2 as a feedstock chemical can reduce the carbon footprint
of the chemical industry and 2) CO2 can serve as an energy vector for harvesting renewable
electricity into the mobility and transport sector.
4
Using CO2 as an alternative feedstock aims to minimize the carbon footprint of the petrochemical
industry. This can be achieved either by incorporating CO2 into the current chemical industry or
by developing novel synthetic routes with other reagents whereby CO2 can serve as a C1 building
block.
5
Target products that have gained considerable attention included organic carbonates,
synthesis gas, carboxylic acids and carboxylation products, and methanol. The benefits of these
3
routes are not limited to just mitigating the climate change impact but also reducing fossil fuel
depletion and providing milder production pathways for chemical markets worth hundreds of
millions of dollars.
4
While using CO2 as a feedstock chemical directly impacts the chemical sector, conversion of CO2
to chemical fuels aims to aid the mobility and transport sectors. One avenue to this end is the
reaction of CO2 with hydrogen (generated from renewable sources) for the production of various
synthetic fuels; such as, liquified natural gas or Fischer-Tropsch products like gasoline and diesel.
This allows for the reduction of global warming effects by “closing” the CO2 cycle and mitigating
local emissions.
6
Unfortunately, the kinetic and thermodynamic barriers associated with CO2
conversion continues to limit the efficiency of these processes. Thus, it is of great desire to design
catalysts which can assist in the conversion of CO2 into valuable products.
1.3 Homogeneous Catalysts for CO2 Conversion
One promising approach to the conversion of CO2 is via electrocatalytic reduction.
4,5,7–9
However,
this pathway remains challenging, due to the thermodynamic stability of CO2. As seen in Table
1.1, the processes for electrochemical CO2 conversion operate at very negative potentials.
10
Additionally, electroreduction of CO2 can occur via one-, two-, four-, six-, and eight-electron
reduction pathways. This leads to a variety of possible products including but not limited to CO,
HCOOH, oxalate, formaldehyde, methanol, and methane. Due to these many possible reaction
pathways, electrochemical conversion of CO2 usually produces a mixture of products. Thus,
designing novel and optimal electrocatalysts is essential for achieving electrochemical conversion
of CO2 with high rates and selectivity at minimal overpotentials.
Table 1.1. Selected CO2 Reduction Processes in pH 7 aqueous solution vs NGE at 25
o
C.
Reaction E
0
CO2 + 2 H
+
+ 2 e
-
→ CO + H2O -0.53 V
CO2 + 2 H
+
+ 2 e
-
→ HCO2H -0.61 V
CO2 + 4 H
+
+ 4 e
-
→ HCHO + H2O -0.48 V
CO2 + 6 H
+
+ 6 e
-
→ CH3OH + H2O -0.38 V
CO2 + 8 H
+
+ 8 e
-
→ CH4 + 2 H2O -0.24 V
CO2 + e
-
→ CO2
•-
-1.90 V
4
Figure 1.1 Electrochemical CO2 Reduction via Heterogeneous (top) and Homogeneous
Catalysis (bottom).
5
Table 1.1 illustrates that the direct uncatalyzed single-electron reduction of CO2 generates a highly
reactive radical. The largely negative potential required to generate this radical is due to the high
reorganization energy difference between the linear CO2 molecule and the bent radical.
10
By
having the homogeneous catalyst in solution, the catalyst can act as an electron shuttle, whereby
the electrode reduces the catalyst which in turn reduces CO2. This allows one to bypass the energy
intensive direct reduction of CO2 and perform reduction at the potential of the catalyst.
Additionally, molecular catalysts are desirable due to their tunability and well-understood
structures, compared to their heterogeneous counterparts. It is for these reasons that molecular
based catalysts for the homogeneous electrochemical reduction of CO2 has gained considerable
interest in the past several decades.
5,10
Some of the most well-known homogeneous catalysts are Mn and Re complexes of the type fac-
[M(bpy)(CO)3X] (with X = Br, Cl).
11
The fac-[Re(bpy)(CO)3Cl] and bpy derivative complexes
were first studied by Lehn et al.
12
and Meyer et al.
13
and were shown to be highly active for CO2-
to-CO conversion. However, to make electrochemical CO2 reduction more sustainable, focus has
shifted to incorporate more earth abundant metals such as Fe and Ni.
5
One such example includes the Fe porphyrin systems extensively studied by Saveant and
coworkers. These studies began with the tetraphenyl porphyrin complex [(tpp)Fe
III
]Cl which is
able to convert CO2 to CO, but suffers from deactivation when in aprotic conditions.
14
However,
it was shown that the incorporation of weak Brønsted acids dramatically enhances the current
generated.
15
Furthermore, it was found that the selectivity of the products could be tuned by the
identity of the Brønsted acid used: where the use of 1-propanol results in a mixture of CO and
formate, the use of 2,2,2-trifluoroethanol (TFE) yields CO exclusively.
16
Catalytic performance is
enhanced further by the incorporation of phenolic groups in the ortho- and ortho’- positions of the
phenyl rings.
17
It was found that this enhancement is due to the stabilization of the Fe
II
-CO2 adduct
from the high local concentration of protons. This system was further optimized by the
incorporation of trimethylammonium groups at those same ortho- and ortho’ positions.
18
To date,
the trimethylammonium substituted Fe porphyrin complexes represent the most efficient
homogeneous catalyst for the electroconversion of CO2-to-CO.
5
Other CO2 conversion
electrocatalysts bearing earth abundant metals include, but are not limited to, Fe Carbonyl clusters
studied by Berben et al.,
19
Ni(II) cyclam studied by Eisenberg et al.,
20,21
Fe and Co quaterpyridines
studied by Robert et al.,
22,23
and Co N4-macrocycles studied by Peters et al.
24
1.4 Co Complexes Bearing Azacalix[4](2,6)pyridine Ligands
We have previously reported extensive studies on Co(II) complex bearing azacalix[4](2,6)pyridine
ligands as efficient and selective catalysts for the electrocatalytic conversion of CO 2-to-CO. We
gained interest in this ligand due to literature precedent showing the effectiveness of pyridine and
macrocyclic complexes for CO2 conversion.
18,22–28
6
Figure 1.2. Synthesis of the six Co(II) complexes 1-6 previously studied in the Marinescu
group.
As seen in Figure 1.2, when the macrocycle is treated with [Co(H2O)6][BF44]2, the pyridine groups
bind in a square planer fashion (if solvent ligands are ignored) and the amine groups alternatingly
point above and below this square plane, resulting in a saddle-shaped geometry. This geometry
allows for the amines to be outside the primary coordination sphere in a pendant fashion, at
proximal distance to substrate bound to the metal center. This design draws inspiration from
nature. The metalloenzyme CO-dehydrogenase (CODH) is responsible for the catalytic conversion
of CO2 in biological systems at ambient conditions with unparalleled efficiency. When CO2 binds
to the CODH active site, the CO2 oxygen atoms, bearing partially negative charges, interact with
the positively charged neighboring amino acid residues. This interaction stabilizes the binding of
CO2 to the metal centers of the active sites.
29
Figure 1.3.
30
Electrochemical studies of CoL
1
. (a) CVs of 1 (0.5 mM) in 0.1 M [nBu4N][PF6] in
DMF under N2 (black and red) or CO2 (blue). (b) Linear scan voltammograms of 1 (0.5 mM) in
0.1 M [nBu4N][PF6] in DMF under CO2 and varying concentrations of methanol. Scan rates: 100
mV/s.
Our initial studies of this system focused on complexes 1 and 6, where the amines are either all
NH groups or all NCH3 groups, respectively.
30
Cyclic voltammetry (CV) studies of 1 (0.5 mM,
Figure 1.3a) in DMF (with 0.1 M [nBu4N][PF6]) with a glassy carbon working electrode, platinum
auxiliary electrode, and silver as the pseudo-reference electrode, reveal that 1 undergoes a
reversible reduction at -1.65 V vs Fc
+/0
and an irreversible reduction at -2.46 V vs Fc
+/0
. These
reductions are attributed to the Co
II/I
and Co
I/0
reductions, respectively. When the solution is
7
concentrated with CO2, a large increase in current is observed at potentials near the Co
I/0
reduction.
Additionally, the potential of the Co
I/0
reduction shifts positively in the presence of CO2,
suggesting a favorable interaction between the Co
0
complex and CO2. The addition of MeOH (and
TFE) as a Brønsted acid into the solution results in large increases in currents. Controlled potential
electrolysis (CPE) experiments of 1 (0.5 mM) performed at -2.8 V vs Fc
+/0
in the presence of TFE
(1.2 M) in DMF (with 0.1 M [nBu4N][PF6]) indeed reveal that the increase in current is due to the
catalytic conversion of CO2, confirmed by gas chromatography of the cell headspace. After 2 h of
electrolysis, 1 consumes 30.9 C of charge and produces exclusively CO with a Faradaic efficiency
(FE) of 98 % at a rate of 16,900 s
-1
.
31
Similar electrochemical studies were performed on the methylated complex 6. CV studies of 6
reveal similar reduction events to 1, with a reversible Co
II/I
reduction at -1.41 V vs Fc
+/0
and an
irreversible Co
I/0
reduction at -2.58 V. However, performing CPE experiments on 6 under similar
conditions to 1 show 6 converts CO2 to CO with a much-reduced FE of 36 % and at a rate of
merely 20 s
-1
.
31
These results highlight the importance of the pendent secondary amines in 1 for
performing CO2 conversion efficiently.
Figure 1.4.
31
Experimental catalytic rate constants, kobs (s
−1
), as a function of the number of
pendant secondary amines for complexes 1−6 measured in the presence of 1.4 M TFE and under
CO2 saturation at a scan rate of 100 mV/s. Rates are obtained from the plateau current. A linear fit
(R
2
= 0.97) is shown in gray for complexes 1 −5.
The tunability of the ligand framework allowed us to synthesize complexes 2-5, to further study
the role of these pendant amines. Complexes 1-5 all convert CO2 to CO with FE ≥ 90 %, but at
8
varying rates. As seen in Figure 1.4, there is a linear relationship between the number of secondary
amines and the rate of catalysis, with complex 1 operating at the fastest rate. It is worth noting that
the rate of catalysis for complex 5 is comparable to that of 6. This is due to the fact that the
secondary amines in this system are quite acidic, bearing pKa values between 2.48-3.10 for
complexes 1-5, allowing the complexes to self-deprotonate in solution.
Figure 1.5.
31
Proposed EECC catalytic cycle illustrated with complex CoL
1
, where E =
electrochemical, and C = chemical step.
31
Through collaborative computation studies performed by Miller et al., we were able to gain insight
on the mechanism by which 1 converts CO2 to CO. Initially, we believed the pendant secondary
amines were directly involved in stabilizing the Co-CO2 adduct via interactions between the NH
groups and the O atoms of CO2.
30
Computational studies reveal that it is actually the case that each
secondary amine group can noncooperatively bind to acid (TFE), enhancing the local
concentration of proton donors around the COOH adduct.
31
9
The azacalix[4](2,6)pyridine system, when bound to a Co(II) center, has proven to be a competent
electrocatalyst for the selective conversion of CO2 to CO. My work focuses on broadening the
scope of this work by studying more biologically relevant metals, such as Fe and Ni.
1.5 References
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10
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12
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13
CHAPTER 2
Physical and Electrochemical Characterization of Fe and Ni Analogues of an Efficient Co-based
CO2-to-CO Electrocatalyst bearing an Azacalix[4](2,6)pyridine Macrocycle
14
ABSTRACT
We report here the synthesis and characterization of Fe and Ni complexes, FeL
1
and NiL
1
,
based on the azacalix[4](2,6)- pyridine framework, and compare their reactivity with that of
CoL
1
¸an efficient CO2-to-CO reduction catalyst. FeL
1
and NiL
1
were characterized by
1
H-NMR
spectroscopy, UV-Vis spectroscopy, FTIR spectroscopy, X-ray diffraction, and electrochemically
via cyclic voltammetry and controlled potential electrolysis. FeL
1
and NiL
1
both electrochemically
reduce CO2 to produce CO with faradaic efficiencies of 19 % (and 16 % H2) and 55 % (and 45 %
H2), respectively. Unlike CoL
1
, FeL
1
and NiL
1
both deposit onto the glassy carbon electrode
surface during electrolysis. The deposited materials (ML
1
-GCE) also electrochemically reduce
CO2 to CO with faradaic efficiencies of 5 % (and 93 % H2) for FeL
1
-GCE and 14 % (and 47 %
H2) for NiL
1
-GCE. This work broadens the scope of our efficient CoL1 catalyst design by
incorporating more biologically relevant and earth abundant metal centers.
2.1 Introduction
The catalytic conversion of carbon dioxide (CO2) into valuable products is a promising
pathway for curbing the detrimental effects fossil fuels have on the environment.
1–3
CO2
conversion aims to make the future more sustainable by using CO2 as a feedstock chemical to
reduce the carbon footprint of the chemical industry.
4
In nature, carbon monoxide dehydrogenase
(CODH) catalyzes the selective reversible conversion of CO2 to carbon monoxide (CO) via a two
proton and two electron transfer. X-ray diffraction studies show that CODH converts CO2 to CO
via a bifunctional activation by two metal centers in the NiFe cluster, where CO 2 binding is
stabilized by H-bonding interactions from neighboring amino acid residues.
5
These studies suggest
that a transition-metal surrounded by pendant proton donor ligands makes for an effective design
for CO2 reduction catalysts. In this sense, scientists draw inspiration from nature in their catalyst
design in hopes of developing a CO2 reduction catalyst that is selective, efficient, and stable.
Due to the many possible products of CO2 reduction, catalyst selectivity remains a
challenge. Furthermore, an additional selectivity facet arises from competing with the hydrogen
evolution reaction if experiments are performed in the presence of proton donors. Molecular
catalysts are of interest in this regard because they allow for fine-tuning of the catalyst structure,
which may lead to alterations in the selectivity of the catalyst. Many examples exist of molecular
electrocatalysts displaying widely different activity and selectivity when the metal center is varied.
15
It has been shown that product selectivity of CO2 reduction with metal complexes bearing the same
pentadentate macrocyclic N5 ligand (2,13-dimethyl-3,6,9,12,18-pentaazabicyclo-
[12.3.1]octadeca-1(18),2,12,14,16-pentaene) can be tuned with high efficiencies by switching the
metal center between cobalt and iron.
6
Cobalt(II) complexes bearing this N5 ligand can
catalytically convert CO2 to CO with a faradaic efficiency (FE) of 97 % and a turnover number
(TON) = 270 after 22 h under photochemical conditions. Iron(III) complexes bearing the same N5
ligand selectively convert CO2 to formic acid (HCOOH) with FE = 80 % and TON = 1260 after 3
h under electrochemical conditions. This change in selectivity arises from Fe(III) being weakly π-
donating, making C-O bond cleavage in the Fe-CO2 adduct slow, leading to facile isomerization
of the adduct and formation of HCOOH. A similar example includes studies of cobalt(II) and
iron(II) complexes bearing quaterpyridine (qpy) ligands.
7,8
A Co(II)qpy complex efficiently
catalyzes the reduction of CO2 to CO with FE = 97 % at an overpotential of η = 140 mV under
electrochemical conditions, making it one of the most active molecular catalysts so far reported
for CO2 reduction. Conversely, a Fe(II) complex bearing the same ligand experience a decrease in
faradaic efficiency to FECO = 48 % under similar conditions. The drop in Faradaic efficiency for
CO arises from a closing of the catalytic cycle where the electrochemically produced Fe(I)qpyCO
is further reduced to the inactive Fe(0)qpyCO. These examples illustrate that by changing the metal
center of a catalyst, one can possibly change the reactivity of the complex.
We previously reported a series of cobalt complexes bearing varying secondary (NH) and
tertiary (NCH3) pendant amine groups incorporated in the ligand scaffold.
9,10
We found that when
all four amines are NH groups (CoL
1
, scheme 1) the complex makes for an efficient electrocatalyst
for the selective reduction of CO2 to CO with high rates.
9
A linear trend can be seen between the
number of NH groups and the activity of the catalysts; methylation of all four secondary amines
(CoL
2
, scheme 1) decreases the activity 300-fold. Electrochemical, kinetic, and density functional
theory studies suggest a mechanism in which noncooperative pendant amines facilitate a H-
bonding network that enables proton transfer from externally added acid to the activated CO2
substrate.
10
In this current work, we investigate a series of ML
1
complexes (scheme 1) with varying
metal centers to study if the reactivity of our catalyst can be modified by tuning the metal center.
This work aims to broaden the understanding of our efficient catalyst design by incorporating more
biologically relevant and earth abundant metal centers.
16
2.2 Results and Discussion
Synthesis and Characterization
The desired L
1
ligand is prepared according to reported literature procedures.
11
Addition of iron(II)
and nickel(II) precursors to L
1
in pyridine led to the formation of the corresponding metal
complexes FeL
1
and NiL
1
in moderately high yields of 70% and 86%, respectively (Scheme 1).
Both complexes are synthesized in air and are air stable. FeL
1
and NiL
1
are isolated from slow
diffusion of Et2O to yield red and brown crystals, respectively.
Scheme 1. Synthesis of the complexes mentioned in this work.
Single-crystal X-ray analysis (Figure 2.1) reveals FeL
1
and NiL
1
bear the formulae
FeL
1
∙(pyridine)2 and NiL
1
∙(pyridine)2, respectively. FeL
1
and NiL
1
adopt similar coordination
environments to CoL
1
; the pyridine nitrogen atoms bind to the metal in a square planer fashion,
with an average bond distance of 2.0 Å. The metal-NH distances average to 3.1 Å. These distances
closely match the bond distances in CoL
1
and those of typically seen for Fe
II
and Ni
II
polypyridine
complexes.
12,13
The metal centers are in an octahedral environment with axial pyridine ligands and
BF4
-
counterions that lie outside the coordination sphere. Overall, this results in a saddle-
conformation, if the pyridine ligands are excluded.
17
Figure 2.1. Side (top) and top (bottom) views of FeL
1
(a) and NiL
1
(b). Hydrogen atoms,
noncoordinating anions, solvent molecules, and axial pyridine ligands omitted for clarity.
1
H-NMR spectroscopy studies in of FeL
1
pyridine-d5 reveal a diamagnetic
1
H-NMR spectrum,
suggesting a low-spin d
6
complex, NiL
1
displays a paramagnetic
1
H-NMR spectrum with μs = 3.01
B.M., corresponding to two unpaired electrons, suggesting a high-spin Ni d
8
complex.
14
Similarly,
CoL
1
is also paramagnetic with μs = 3.63 B.M., corresponding to three unpaired electrons.
10
UV-
VIS spectroscopic studies (Figures 2.2 and S4-6) reveal FeL
1
, CoL
1
, and NiL
1
all absorbs in the
UV region at wavelengths of 344, 336, and 336 nm with molar absorptivity values of 19200,
34400, and 27500 M
-1
cm
-1
. CoL
1
has an additional absorbances at 442 nm with a molar
absorptivity of 2011 M
-1
cm
-1
. FTIR studies (s S1-S3) of the complexes reveal N-H stretching
frequencies of 3349, 3343, and 3346 cm
-1
for FeL
1
, CoL
1
, and NiL
1
, respectively. These values
suggest that the N-H bond strength in CoL
1
is stronger than those in FeL
1
and NiL
1
.
ML
1
Average M – Npy (Å) Average M – NR (Å) reference
FeL
1
2.0 3.1 This work
CoL
1
1.9 3.0
9
NiL
1
2.1 3.1 This work
Table 1. Average relevant bond distances of FeL
1
, CoL
1
, and NiL
1
.
18
Figure 2.2. UV-VIS spectra of FeL
1
(a), CoL
1
(b), and NiL
1
(c). Spectra acquired in DMF with a
1 cm quartz cuvette. a) A single absorbance is observed at 344 nm with ε = 19,200 M
-1
cm
-1
. b)
Absorbances are observed at 336 nm (ε = 34,400 M
-1
cm
-1
) and 434 nm (ε = 2,011 M
-1
cm
-1
). c) A
single absorbance is observed at 336 nm (ε = 27,500 M
-1
cm
-1
).
Electrochemistry
All complexes were studied by cyclic voltammetry at a concentration of 0.5 mM in DMF
with 0.1 M [nBu4N][PF6] as the electrolyte, at a scan rate of 100 mV/s with a glassy carbon
working electrode, platinum counter electrode, and silver wire separated from solution by a vycor
frit as the pseudo-reference electrode. All potentials are reported vs Fc
+
/Fc.
Figure 2.3. Electrochemical studies of the ML
1
complexes. (a) CVs of FeL
1
(0.5 mM) in 0.1 M
[nBu4N][PF6] in DMF under N2 (red) and CO2 (blue). (b) CVs of CoL
1
(0.5 mM) in 0.1 M
[nBu4N][PF6] in DMF under N2 (red and black) and CO2 (blue).
9
(c) CVs of NiL
1
(0.5 mM) in 0.1
M [nBu4N][PF6] in DMF under N2 (red and black) and CO2 (blue).
19
ML
1
M
2+/+
(V
vs
Fc
+
/Fc)
M
+/0
(V vs
Fc
+
/Fc)
FECO/H2
ML
1
(%)
FECO/H2
ML
1
-
GCE
icat/ip* reference
FeL
1
E = -2.35
(irrev)
N/A 19 / 16 3 / 93 This work
CoL
1
E1/2 = -
1.65
E = -2.46
(irrev)
99 / 0 N/A 208.8
9,10
NiL
1
E1/2 = -
1.51
-2.77
(irrev.)
55 / 45 47 / 14 8.185 This work
Table 2. Summary of electrochemical and catalytic properties of the ML
1
complexes. Reduction
potentials determined by cyclic voltammetry with 0.5 mM of ML
1
in 0.1 M [nBuN4][PF6] in DMF
at 100 mV/s with a glassy carbon electrode. Catalytic values determined using controlled potential
electrolysis of 0.5 mM of ML
1
in 0.1 M [nBuN4][PF6] in DMF with 1.4 M TFE using a glassy
carbon electrode. *Normalized currents (icat/ip) calculated in acid saturation conditions.
Electrochemical characterization of FeL
1
and NiL
1
(Figures 3-4 and S7-S10) reveals a single
irreversible reduction of FeL
1
at E = -2.35 V and a reversible reduction for NiL
1
with E1/2 of -1.51
V , attributed to the Fe
(II)
/Fe
(I)
and
Ni
(II)
/Ni
(I)
reductions, respectively. Scanning more negatively
reveals an irreversible NiL
1
reduction at E = -2.77 V, corresponding to the Ni
(I)
/Ni
(0)
reductions.
In this regard, NiL
1
behaves most similarly to CoL
1
, both bearing one reversible and one
irreversible reduction but at differing potentials.
When CVs are performed in a CO2 atmosphere, we see enhanced currents near the second
reductions for FeL
1
and NiL
1
, as well as a positive shift in the second reduction potentials by c.a.
100 mV (Figure 2.2a). Similar behavior is seen with CoL
1
(Figure 2.2b),
9
suggesting the doubly
reduced FeL
1
and NiL
1
complexes react favorably with CO2. Unlike the CoL
1
system, the shape
of the CV traces for FeL
1
and NiL
1
bear a peak shape, rather than a plateau. This indicates a
reaction that is limited by substrate diffusion.
15
Additionally, repetitive scans of FeL
1
and NiL
1
under CO2 (Figures S11-S12) shows irreproducible currents. Such a behavior is not observed with
20
repetitive scan under N2 (Figures S13-14), suggesting that when CO2 is present in cathodic
conditions, passivation of the electrode surface occurs.
Figure 2.4. (a) CVs of FeL
1
(0.5 mM) titrating with 2,2,2-trifluoroethanol in 0.1 M [nBu4N][PF6]
in DMF under CO2. (b) CVs of CoL
1
(0.5 mM) titrating with 2,2,2-trifluoroethanol in 0.1 M
[nBu4N][PF6] in DMF under CO2.
10
(c) CVs of NiL
1
(0.5 mM) titrating with 2,2,2-trifluoroethanol
in 0.1 M [nBu4N][PF6] in DMF under CO2.
In CO2, the peak current potentials for FeL
1
and NiL
1
are -2.78 and -2.90 V, respectively. We have
previously reported that the addition of 2,2,2-trifluoroethanol (TFE) to CoL
1
results in larger
catalytic currents, resulting from a noncooperative mechanism where the pendant amines in the
ligand system facilitate a hydrogen-bonding network which enables direct proton transfer from
acid to the activated CO2 adduct.
10
Titrating the complexes with TFE (Figures 2.4) results in
enhanced current densities, reaching -4.58 mA/cm
2
for FeL
1
at -2.68 V and -1.6 mA/cm
2
for NiL
1
at -2.90 V.
We previously reported controlled potential electrolysis (CPE) experiments of CoL
1
(0.5 mM)
performed at -2.8 V in DMF containing [nBu4N][PF6] (0.1 M) and TFE (1.4 M) with CO2 (0.2 M)
showing that CoL
1
is a competent catalyst for CO2-to-CO reduction, able to pass 30.9 C of charge
over 2 h to produce CO with a FE of 98 %.
9
CPE experiments were performed for FeL
1
and NiL
1
under similar conditions at their respective potentials of maximum current of -2.68 and -2.90 V,
respectively. CPE of FeL
1
(Figure S15) reveals a rapid decrease in current within 15 min of
electrolysis. After 1 h, no gaseous products were detected in the headspace and no formate was
observed in solution. Conversely, NiL
1
(Figure S17) passes 6.905 C of charge over 1 h, producing
19.7 μmol of CO and 16.6 μmol of H2, corresponding to faradaic efficiencies of 55 % and 45 %,
respectively. CVs of NiL
1
before and after CPE reveal a large increase in current density after the
CPE is performed (Figure S16). These results suggest decomposition of the material under these
21
conditions, confirmed visually by the formation of a brown film on the electrode surface.
Performing CPE of FeL
1
at a more positive potential of -2.40 V (Figure S16) shows behavior
similar to NiL
1
, although at much lower efficiency. After 1 h, FeL
1
passes 7.286 C of charge,
producing 7.4 μmol of CO and 6.4 μmol of H2, resulting in Faradaic efficiencies of 19 and 17 %,
respectively. Analysis of the cell solution reveals no formate is produced. Similarly, CV traces
performed before and after electrolysis reveal an increase in current and a brown film on the
electrode surface, suggesting FeL
1
similarly decomposes.
The glassy carbon electrode (GCE) surface following CPE of FeL
1
, CoL
1
, and NiL
1
were analyzed
by scanning electron microscopy (SEM) and X-ray photoelectron spectroscopy (XPS). These post-
CPE electrodes are referred to as ML
1
-GCE. Two portions of the electrode were analyzed by
SEM, one surface that was examined as-is after CPE, and a surface that was first washed with
DMF to remove any soluble species. XPS experiments were only performed on the DMF-washed
surface. SEM images of both the as-prepared and DMF-washed FeL
1
-GCE surfaces (Figure S20)
reveal a morphology of nanospheres of about 100 nm in diameter. XPS studies of these
nanospheres (Figure S23a) reveal two iron signals with binding energies of 723 and 710 eV.
Similar studies were performed for NiL
1
-GCE and a similar morphology of nanospheres, but with
smaller diameters, is observed in both the washed and as-prepared NiL
1
-GCE surfaces (Figure
S21). XPS of the NiL
1
-GCE nanospheres (Figure S23b) reveal signals with binding energies of
873, 862, and 856 eV. Similar SEM studies were performed for CoL
1
-GCE (Figures S22), and no
deposition of CoL
1
onto the GCE surface was observed for neither the DMF-washed nor the as-
prepared surfaces. SEM images of the as-prepared CoL
1
-GCE surface reveal morphologies
indicative of the electrolyte [nBu4N][PF6], whereas images of the washed surface is similar to that
of the “clean” surface of the fresh GCE electrode. XPS analysis of the DMF-washed CoL
1
-GCE
surface (Figure S23b) does not reveal any signals indicative of Co. These results suggest that
rinsing the FeL
1
-GCE and NiL
1
-GCE with DMF simply removes deposited electrolyte, but the
deposited ML
1
material is not removed. Whereas CoL
1
does not appear to deposit onto the GCE
surface.
CPE studies were also performed on FeL
1
-GCE and NiL
1
-GCE with no catalyst in solution. CPE
experiments of a FeL
1
-GCE at -2.4 V in DMF containing [nBu4N][PF6] (0.1 M) and TFE (1.4 M)
with CO2 (0.2 M) after 1 h reveal FeL
1
-GCE produces 6.3 μmol of H2 with a FE = 94 % and 0.3
22
μmol of CO with a FE = 5 %. A similar experiment was performed with NiL
1
-GCE, producing
2.3 μmol of H2 and 0.7 μmol of CO after 1h, resulting in FE values of 47 and 14 %, respectively.
Computational Studies
Figure 2.5. Alpha (left) and Beta (right) Molecular orbital diagrams of the ML
1
series, obtained
through density functional theory calculations using 6-31+G*/B3LYP.
The ML
1
series was also studied using computational methods. Using Density functional
theory (DFT), we acquired molecular orbital diagrams of the ML
1
series. Figure 2.5 shows the
Alpha (left) and Beta (right) molecular orbitals of the ML
1
series, which provide supporting
evidence of our experimental observations. Firstly, the DFT results assign the spin states of FeL
1
,
CoL
1
, and NiL
1
tb singlet, doublet, and triplet, respectively, which agree with our NMR and
Evan’s method results, mentioned above. Additionally, the LUMO and LUMO+1 give us insight
into the ML
I
and ML
0
reduced species. Both the CoL
1
and the NiL
1
complexes have Beta LUMO
and LUMO+1 orbitals which are metal-based, while for FeL
1
, the Alpha LUMO orbital is metal-
based while the Alpha LUMO+1 is ligand-based. This explains why, by CV, we see similar
reductive behavior between CoL
1
and the NiL
1
, both displaying two reduction events, while FeL
1
displays only a single reduction.
2.3 Conclusion
In conclusion, we have prepared Fe and Ni analogues of an efficient Co-based catalyst. CoL
1
, and
physically characterized them via XRD,
1
H-NMR, UV-VIS, and FTIR spectroscopy and
23
electrochemically through CV and CPE. Of the three complexes, CoL
1
remains the best catalyst
for CO2 reduction, producing the most desired product with the highest rates for the selective
conversion of CO2 to CO. The FeL
1
and NiL
1
complexes both produce mixtures of CO and H2 and
operate at sufficiently slower rates. Additionally, both FeL
1
and NiL
1
are unstable under cathodic
conditions in the presence of CO2, leading to decomposition and deposition onto the electrode
surface during electrolysis. The deposited materials, FeL
1
-GCE and NiL
1
-GCE, are also able to
perform CO2 reduction to CO, but mostly produce H2. This work broadens the scope of our
efficient CoL1 catalyst design by incorporating more biologically relevant and earth abundant
metal centers and emphasizes that the identity of the metal center of an electrocatalyst plays a
crucial role in the reactivity and efficiency of its performance.
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25
Supporting Information STUDYING FE AND NI ANALOGUES OF AN EFFICIENT CARBON
DIOXIDE REDUCTION CATALYST
General
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 δ 7.26 and pyridine-d5 δ 7.22, 7.58, 8.74) and are reported as
parts per million relative to tetramethylsilane. Elemental analyses were performed by Robertson
Microlit Laboratories, Ledgewood, New Jersey.
Electrochemical Methods
Electrochemistry experiments were carried out using a Pine potentiostat. All experiments in this
paper were referenced relative decamethylferrocene (Fc*) with the Fe
3+/2+
couple at -0.48 V as an
internal standard. All electrochemical experiments were performed with 0.1 M
tetrabutylammonium hexafluorophosphate as supporting electrolyte, which was recrystallized
before use. ML1 complex concentrations were generally at 0.5 mM and experiments with CO2
were performed at gas saturation in dimethylformamide (DMF, 0.2 M).
Cyclic voltammetry (CV) experiments were performed in a single compartment electrochemical
cell under nitrogen or CO2 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. Controlled-potential electrolysis (CPE) measurements were conducted in a two-
chambered H cell. The first chamber held the working and reference electrodes in 40 mL of 0.1 M
tetrabutylammonium hexafluorophosphate and 1.4 M TFE in DMF. The second chamber held the
auxiliary electrode in 20 mL of 0.1 M tetrabutylammonium hexafluorophosphate in DMF. The
two chambers were separated by a fine porosity glass frit. The reference electrode was separated
from solution by a Vycor tip. Glassy carbon plate electrodes (GCE) (6 cm × 1 cm × 0.3 cm; Tokai
26
Carbon USA) were used as the working and auxiliary electrodes. After CPE experiments were
performed for the ML1 complexes, the GCE electrode was saved either as is or rinsed with DMF
and dried under vacuum atmosphere until further use. These GCE electrodes were used against in
CPE experiments to test the performance of the deposited materials. Gas analysis for CPE
experiments were performed using 2 mL sample aliquots taken from the headspace of the
electrochemical cell and injected on a Shimadzu BID-2010 plus series gas chromatograph with a
2m × 1mm ID micropacked column. Faradaic efficiencies were determined by dividing he
measured gas produced by the amount of gas expected based on the charge passed during the bulk
electrolysis experiment
Scanning Electron Microscopy (SEM)
SEM images were collected using a JEOL-7001F operating at 15 kV with 5 nA of probe current.
X-Ray Photoelectron Spectroscopy (XPS)
XPS data were collected using a Kratos AXIS Ultra instrument. The monochromatic X-ray S2
source was the Al K α line at 1486.7 eV, and the hybrid lens and slot mode were used. Low
resolution survey spectra were acquired between binding energies of 1–1200 eV. Higher resolution
detailed scans, with a resolution of 0.1 eV, were collected on individual XPS regions of interest.
The sample chamber was maintained at < 9 × 10
–9
Torr. The XPS data were analyzed using the
CasaXPS software.
Fourier Transform Infrared Spectroscopy (FTIR)
FTIR spectra were acquired using a Bruker Vertex 80v spectrometer on a KBr pressed pellet.
Evan’s Method
1
Evan’s method was used to determine the total spin (S) of a metal complex by
1
H NMR
spectroscopy. Equation 1 is used to determine the MMs (Measured Molar Susceptibility). ΔHz is
the difference in hertz between the peaks of the solvent in contact with the complex and the ones
in the capillary tube, and M is the molarity of the sample (in units of mol/L), and Hz NMR is the
spectrometer frequency, in hertz (500,000,000).
MMs=
3000 x ΔHz
4π x M x Hz
NMR
(1)
27
Subsequently, eqs 2 and 3 were used to determine the number of unpaired electrons.
X
P
= MMs− X
D
, 𝑋 𝐷 =
mM
2
x 10
−6
(2)
μ
eff
= 2.84√T x X
P
(3)
XP indicates the corrected molar susceptibility, mM is the molar mass of the complex (g/mol), and
T is the temperature in Kelvin (298 K).
Rate Calculations from Cyclic Voltammetry
2
Equation 1 gives the catalytic current (icat) for an EECC process (E = electrochemical, C = chemical
step) and assumes a one-electron diffusion current and pseudo-first-order kinetics. In eq 1, F
corresponds to Faraday’s constant (96485 C/mol), S is the surface area of the glassy carbon
electrode (0.07065 cm
2
for CV experiments), C
cat
0
is the catalyst concentration ([cat] = 0.5 mM =
5 x 10
-7
mol/cm
3
), Dcat is the diffusion constant of the catalytically-active species (c.a. 5 x 10
-6
cm
2
/s), and kcat is the rate constant of the catalytic reaction.
𝑖 cat
= FSC
cat
0
√D
cat
2𝑘 cat
(4)
Equation 4 is simplified to equation 5 by standardizing with the current in the absence of substrate
(in this case CO2). In equation 5, F is Faraday’s constant (96485 C/mol), S is the surface area of
the glassy carbon electrode (0.07065 cm
2
for CV experiments), C
cat
0
is the catalyst concentration
([cat] = 0.5 mM = 5 x 10
-7
mol/cm
3
), Dcat is the diffusion constant of the catalytically-active species
(c.a. 5 x 10
-6
cm
2
/s), ν is the scan rate (0.1 V/s), R is the universal gas constant (8.31 J K
-1
mol
-1
),
and T is the temperature in Kelvin (298 K).
𝑖 p
= 0.446 x FSC
cat
0
√
D
cat
Fν
RT
(5)
Dividing equation 4 by equation 5 allows for determining icat/ip and the catalytic rate constant (kcat)
without needing to determine S, C
cat
0
, and Dcat. The ratio of equations 4 and 5 yields equation 6,
which can be rearranged to solve for kcat as seen in equation 7.
𝑖 cat
𝑖 p
=
1
0.446
x √
2k
𝑐𝑎𝑡
RT
νF
= 2.24 x√
2k
𝑐𝑎𝑡
RT
νF
(6)
28
𝑘 cat
= (
𝑖 cat
𝑖 p
)
2
νF
2RT x 2.24
2
(7)
Equation 7 simplifies into equation 8 which can be used to calculate kcat directly.
𝑘 cat
= 0.387 x (
𝑖 cat
𝑖 p
)
2
(8)
In the above calculations, the ip values used correspond to the peak current obtained from the M
I/0
reductions.
Syntheses
The ligand L1 and the complex CoL1 were prepared according to the previously reported literature
procedures.
3,4
FeL1. [Fe(H2O)6][BF4]2 (27.1 mg, 0.070 mmol) in 9:1 pyridine:DMF (1 mL) was added to a
solution of L1 (16.7 mg, 0.071 mmol) in pyridine (1 mL) giving rise to a brown solution. The
mixture was allowed to stir for 15 min and filtered through a microfiber filter. Slow diffusion with
diethyl ether produced red crystals, 70 % yield.
1
H NMR (500 Hz, pyridine-d5) δ 10.74 (s, 4H,
NH), 7.75 (t, J = 7.94 Hz, 4H, p-NC5H3), and 7.06 (d, J = 7.94 Hz, 8H, m-NC5H3). Anal. calcd for
[FeL1·2py·1DMF·1H2O]: C, 46.79; H, 4.16; N, 18.19. Found: C, 46.70; H, 4.13; N, 17.79
NiL1. [Ni(MeCN)6][BF4]2 (27.1 mg, 0.070 mmol) in 9:1 pyridine:DMF (1 mL) was added to a
solution of L1 (16.7 mg, 0.071 mmol) in pyridine (1 mL) giving rise to a brown solution. The
mixture was allowed to stir for 30 minutes. The solution was filtered through a microfiber filter.
Slow diffusion with diethyl ether produced amber crystals, 86 % yield.
1
H NMR (500 Hz, pyridine
d5) δ 49.86 (s, 4H, p-NC5H3) and 14.96 (s, 4H, NH). Anal. calcd for [NiL1·2py·1DMF·1H2O]: C,
46.63; H, 4.15; N, 18.13. Found: C, 46.71; H, 4.16; N, 17.76.
29
Figure S1. Transmittance FTIR spectrum of FeL1.
30
Figure S2. Transmittance FTIR spectrum of CoL1.
Figure S3. Transmittance FTIR spectrum of NiL1.
Figure S4. Absorbance values at varying concentrations of FeL1 in DMF. Absorbance values
measured at 344 nm.
31
Figure S5. Absorbance values at varying concentrations of CoL1 in DMF. Absorbance values
measured at 336 nm.
Figure S6. Absorbance values at varying concentrations of NiL1 in DMF. Absorbance values
measured at 336 nm.
32
Figure S7. Cyclic voltammograms of 0.5 mM NiL1 in a DMF solution containing 0.1 M
[nBu4N][PF6] under an atmosphere of N2 displaying the reversible one-electron reduction,
corresponding to the Ni
II/I
couple, (left) and the irreversible one-electron reduction (right),
corresponding to the Ni
I/0
reduction. Scan rates vary from 0.025 to 2 V/s.
Figure S8. Plot showing peak current densities for both NiL1 reductions in the CVs of 0.5 mM
NiL1 in DMF containing 0.1 M [nBu4N][PF6] under an atmosphere N2. The cathodic and anodic
peak current densities increase linearly with the square root of the scan rate. This is indicative of
a freely-diffusing species, in which the electrode reaction is controlled by mass transport.
33
Figure S9. Cyclic voltammograms of FeL1 (0.05 mM left, 0.5 mM right) in a DMF solution
containing 0.1 M [nBu4N][PF6] under an atmosphere of N2 displaying the irreversible reduction
corresponding to the Fe
II/I
reduction at -2.35 V. Scan rates vary from 0.025 to 2 V/s.
34
Figure S10. Plot showing peak current densities for the FeL1 reduction in the CVs of FeL1 (0.05
mM in DMF containing 0.1 M [nBu4N][PF6] under an atmosphere N2). The cathodic peak current
densities increase linearly with the square root of the scan rate. This is indicative of a freely-
diffusing species, in which the electrode reaction is controlled by mass transport.
35
Figure S11. Cyclic voltammogram of 0.5 mM of FeL1 in a DMF solution containing 0.1 M
[nBu4N][PF6] under an atmosphere of CO2 after repetitive scans displays irreproducible current
densities. Scan rate is 100 mV/s.
36
Figure S12. Cyclic voltammogram of 0.5 mM of NiL1 in a DMF solution containing 0.1 M
[nBu4N][PF6] under an atmosphere of CO2 after repetitive scans displays irreproducible current
densities. Scan rate is 100 mV/s.
-0.5
-0.4
-0.3
-0.2
-0.1
0
0.1
-3 -2.5 -2 -1.5 -1
Current Density (mA/cm
2
)
Potential (V) vs Fc/Fc
+
1st CO2 Scan
2nd CO2 Scan
3rd CO2 Scan
4th CO2 Scan
37
Figure S13. Cyclic voltammogram of 0.5 mM of FeL1 in a DMF solution containing 0.1 M
[nBu4N][PF6] under an atmosphere of N2 after repetitive scans displays reproducible current
densities. Scan rate is 100 mV/s.
-0.5
-0.4
-0.3
-0.2
-0.1
0
-3.3 -3 -2.7 -2.4 -2.1
Current Density (mA/cm
2
)
Potential (V) vs Fc
+
/Fc
1st N2 Scan
2nd N2 Scan
3rd N2 Scan
4th N2 Scan
38
Figure S14 Cyclic voltammogram of 0.5 mM of NiL1 in a DMF solution containing 0.1 M
[nBu4N][PF6] under an atmosphere of N2 after repetitive scans displays reproducible current
densities. Scan rate is 100 mV/s.
-0.35
-0.3
-0.25
-0.2
-0.15
-0.1
-0.05
0
0.05
0.1
-3 -2.5 -2 -1.5 -1
Current Density (mA/cm
2
)
Potential (V) vs Fc/Fc
+
1st N2 Scan
2nd N2 Scan
39
Figure S15. Controlled potential electrolysis of 0.5 mM FeL1 performed at -2.68 V, corresponding
to the potential of max current observed by cyclic voltammetry, containing 0.1 M [nBu4N][PF6]
and 1.4 M TFE.
Figure S16. Controlled potential electrolysis of 0.5 mM FeL1 performed at -2.4 V, containing 0.1
M [nBu4N][PF6] and 1.4 M TFE.
Figure S17. Controlled potential electrolysis of 0.5 mM NiL1 performed at -2.9 V, containing 0.1
M [nBu4N][PF6] and 1.4 M TFE.
40
Figure S18. Controlled potential electrolysis of a FeL
1
-GCE performed at -2.4 V, containing 0.1
M [NBu4][PF6] and 1.4 M TFE.
Figure S19. Controlled potential electrolysis of a NiL
1
-GCE performed at -2.9 V, containing 0.1
M [nBu4N][PF6] and 1.4 M TFE.
Figure S20. Scanning Electron Microscope (SEM) images (25 kV accelerating voltage) of the
DMF-washed (left), unwashed (middle), and clean (right) surfaces of a GCE used for a CPE
41
experiment of 0.5 mM FeL1 with 0.1 M [nBu4N][PF6] and 1.4 M TFE in DMF. Both the DMF-
washed and unwashed surfaces display similar morphologies of deposited FeL1 material.
Figure S21. SEM images (25 kV accelerating voltage) of the DMF-washed (left), unwashed
(middle), and clean (right) portions of a GCE used for a CPE experiment of 0.5 mM CoL1 with
0.1 M [nBu4N][PF6] and 1.2 M TFE in DMF. The DMF-washed surface shows no deposition of
material. The unwashed portion shows a morphology indicative of the electrolyte.
Figure S22. SEM images (25 kV accelerating voltage) of the DMF-washed (left), unwashed
(middle), and clean (right) portions of a GCE used for a CPE experiment of 0.5 mM NiL1 with 0.1
M [nBu4N][PF6] and 1.4 M TFE in DMF. Both the DMF-washed and unwashed surfaces display
similar morphologies of deposited NiL1 material.
42
Figure S23. XPS analysis of DMF-washed ML
1
-GCE surfaces. (a) Fe 2p core level XPS spectrum
of FeL
1
-GCE; (b) Co 2p core level XPS spectrum of CoL
1
-GCE; (c) Ni 2p core level XPS
spectrum of NiL
1
-GCE.
Table S1. Crystal data and structure refinement for FeL1.
Identification code GR0235FeNH
Chemical formula C30H26BF4FeN10
Formula weight 669.27 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal size 0.104 x 0.135 x 0.460 mm
Crystal habit clear yellow plate
Crystal system monoclinic
Space group P 1 21/m 1
Unit cell dimensions a = 10.7530(16) Å α = 90°
b = 13.695(2) Å β = 91.719(3)°
43
c = 15.596(2) Å γ = 90°
Volume 2295.7(6) Å
3
Z 2
Density (calculated) 0.968 g/cm
3
Absorption coefficient 0.371 mm
-1
F(000) 686
Diffractometer Bruker APEX II CCD Bruker APEX DUO
Radiation source fine-focus tube (MoKα , λ = 0.71073 Å)
Theta range for data
collection
1.90 to 27.48°
Index ranges -13<=h<=13, -17<=k<=17, -20<=l<=20
Reflections collected 41330
Independent
reflections
5468 [R(int) = 0.0631]
Coverage of
independent
reflections
100.0%
Absorption
correction
multi-scan
Max. and min.
transmission
0.9620 and 0.8480
44
Structure solution
technique
direct methods
Structure solution
program
SHELXT 2014/5 (Sheldrick, 2014)
Refinement method Full-matrix least-squares on F
2
Refinement program SHELXL-2017/1 (Sheldrick, 2017)
Function minimized Σ w(Fo
2
- Fc
2
)
2
Data / restraints /
parameters
5468 / 7 / 245
Goodness-of-fit on F
2
1.145
Final R indices 4353 data; I>2σ(I) R1 = 0.0549, wR2 = 0.1699
all data R1 = 0.0678, wR2 = 0.1772
Weighting scheme
w=1/[σ
2
(Fo
2
)+(0.0826P)
2
+3.0613P]
where P=(Fo
2
+2Fc
2
)/3
Largest diff. peak
and hole
0.562 and -0.449 eÅ
-3
R.M.S. deviation
from mean
0.090 eÅ
-3
Table S2. Crystal data and structure refinement for NiL1
Identification code GR02XXX_Ni_NH
Chemical formula C30H26B2F8N10Ni
45
Formula weight 758.94 g/mol
Temperature 100(2) K
Wavelength 0.71073 Å
Crystal system monoclinic
Space group P 1 21/n 1
Unit cell dimensions
a =
10.920(3) Å
α = 90°
b =
13.874(3) Å
β = 95.981(4)°
c =
23.300(6) Å
γ = 90°
Volume
3510.8(15)
Å
3
Z 4
Density (calculated) 1.436 g/cm
3
Absorption coefficient 0.633 mm
-1
F(000) 1544
Diffractometer Bruker APEX II CCD Bruker APEX DUO
Radiation source fine-focus tube (MoKα , λ = 0.71073 Å)
Theta range for data
collection
1.71 to 21.97°
Index ranges -11<=h<=11, -14<=k<=14, -24<=l<=24
46
Reflections collected 43704
Independent reflections 4301 [R(int) = 0.0685]
Coverage of independent
reflections
100.0%
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 2017/1 (Bruker AXS,
2017)
Function minimized Σ w(Fo
2
- Fc
2
)
2
Data / restraints /
parameters
4301 / 703 / 543
Goodness-of-fit on F
2
1.095
Δ/σmax 0.001
Final R indices 3636 data; I>2σ(I)
R1 = 0.0853, wR2 =
0.2019
all data
R1 = 0.0967, wR2 =
0.2085
Weighting scheme
w=1/[σ
2
(Fo
2
)+(0.0628P)
2
+42.4197P]
where P=(Fo
2
+2Fc
2
)/3
47
Largest diff. peak and hole 1.314 and -0.700 eÅ
-3
R.M.S. deviation from
mean
0.118 eÅ
-3
(1) Piguet, C. Paramagnetic Susceptibility by NMR : The “ Solvent Correction ” Removed for
Large Paramagnetic Molecules. J. Chem. Ed. 1997, 74 (7), 815–816.
(2) Costentin, C.; Drouet, S.; Robert, M.; Saveant, J. M. Turnover Numbers, Turnover
Frequencies, and Overpotential in Molecular Catalysis of Electrochemical Reactions. Cyclic
Voltammetry and Preparative-Scale Electrolysis. J. Am. Chem. Soc. 2012, 134 (27), 11235–11242.
(3) Zhang, E. X.; Wang, D. X.; Huang, Z. T.; Wang, M. X. Synthesis of (NH)m(NMe)4-m-
Bridged Calix[4]Pyridines and the Effect of NH Bridge on Structure and Properties. J. Org. Chem.
2009, 74 (22), 8595–8603.
(4) Chapovetsky, A.; Do, T. H.; Haiges, R.; Takase, M. K.; Marinescu, S. C. Proton-Assisted
Reduction of CO2 by Cobalt Aminopyridine Macrocycles. J. Am. Chem. Soc. 2016, 2016 (138),
5765–5768.
Abstract (if available)
Abstract
We report here the synthesis and characterization of Fe and Ni complexes, FelL¹ and NiL¹, based on the azacalix[4](2,6)-pyridine framework, and compare their reactivity with that of CoL¹, an efficient CO₂-to-CO reduction catalyst. FelL¹ and NiL¹ were characterized by ¹H-NMR spectroscopy, UV-Vis spectroscopy, FTIR spectroscopy, X-ray diffraction, and electrochemically via cyclic voltammetry and controlled potential electrolysis. FelL¹ and NiL¹ both electrochemically reduce CO₂ to produce CO with faradaic efficiencies of 19% (and 16% H₂) and 55% (and 45% H₂), respectively. Unlike CoL¹, FelL¹ and NiL¹ both deposit onto the glassy carbon electrode surface during electrolysis. The deposited materials (ML¹-GCE) also electrochemically reduce CO₂ to CO with faradaic efficiencies of 5% (and 93% H₂) for FelL¹-GCE and 14% (and 47% H₂) for NiL¹-GCE. This work broadens the scope of our efficient CoL¹ catalyst design by incorporating more biologically relevant and earth abundant metal centers.
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Asset Metadata
Creator
Rangel, Geovanni M.
(author)
Core Title
Studying iron and nickel analogues of an efficient carbon dioxide reduction electrocatalyst
School
College of Letters, Arts and Sciences
Degree
Master of Science
Degree Program
Chemistry
Publication Date
08/01/2019
Defense Date
05/30/2019
Publisher
University of Southern California
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Tag
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Marinescu, Smaranda C. (
committee chair
), Haiges, Ralf (
committee member
), Malmstadt, Noah (
committee member
)
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