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Design of nanomaterials for electrochemical applications in fuel cells and beyond
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Design of nanomaterials for electrochemical applications in fuel cells and beyond
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Content
i
Design of Nanomaterials
for Electrochemical Applications in Fuel Cells and Beyond
by
Juan Pablo de los Rios
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)
December 2024
Copyright 2024 Juan Pablo de los Rios
ii
“Prefiero morir de pie que vivir arrodillado”
– Emiliano Zapata
iii
To my Little Moon.
I love you.
iv
Acknowledgements
To Dr. Prakash, thank you infinitely for welcoming me into your group. You saw
my ambition to pursue this degree at the most difficult part of my graduate journey. I was
constantly uncertain, questioning it all. Thank you for extending your hand and guiding
me through this process. You gave me the opportunity to explore ideas freely and pave
my own trajectory as a scientist, always being supportive. I will never forget when you
told me something along the lines of: “The PhD is not for me to give, it is for you to attain.”
You are a true leader and teacher sir. Thank you. To my research group, both the
methanol and fluorine sides, thank you for all the science talks and shared experiences.
I also want to acknowledge all the staff at the Loker Hydrocarbon Research Institute,
which makes it absolutely the best place to work at. Thank you to Jessie, Dr. Robert, Dr.
Alan, David, and Gloria. I appreciate the immense support and guidance you all provide
to graduate students.
To Dr. Vicente Galvan, your patience, knowledge, and mentorship have brought
me to this point. Any time I have a doubt I know that I can call you. Thank you for your
friendship. Thank you for being my gym partner, too. You are a major source of inspiration
to how I want to become as both a scientist and human. I appreciate you greatly. To all
the members of the fuel cell lab. Bo, Adam, Eric, Yeshvi, and Ryuichi. I will miss you all
so much. It was both an honor and pleasure to work with you all. The environment within
the fuel cell lab was always peaceful, collaborative, and challenging. I hope that wherever
I go, I can have a similar atmosphere. Never hesitate to reach out if I can help you in any
way.
v
Drs. Alex, Cindy, and Joe, there are no words I can use to express my gratitude to
you three. I will always be here for you. Thank you for making graduate school absolutely
meaningful. I would not be here without you guys. I would like to thank Paulo Uranga. I’m
not exaggerating when I say that without him, I would have never arrived in LA. He was
an instrumental part of my graduate school journey. I want to thank Dr. Emilio Cavazos
and Dr. Claudia Bernier. I will always cherish the memories we made. Both of you have
been an inspiration for me to complete my degree. I would like to thank Enrique Licon
and Sebastian Moya for welcoming me to LA and introducing me to an amazing group of
friends, Elisa Ochoa, Jose Javier Calvo, Juan Diaz, Iñigo Sanchez, Emir Ucer, Limari
Archuleta, Regina Padilla, Ryan Kennedy, and Roberta de Leon. I will always treasure all
of our memories together. Thank you to Michael Koutombas for constantly motivating me.
Thank you for all the laughs, conversations, and cigarillos. I appreciate you deeply. I want
to thank Dr. Tai Nguyen. You were a major propelling force for never succumbing to
failure. To put endless hours of time and effort into what one is passionate about. I
appreciate you. Thank you to Matthew Coe. Your constant visits downstairs to check on
me always made my day better. Thank you for introducing me to all of your amazing
family and incorporating me as one of your own. Thank you for all of the support, the
LACC trips, Thanksgivings at VC, and Pasadena afternoons. Thank you for the laughs,
conversations, and fooling around. I appreciate you. I want to thank Alexander Knieb.
Thank you for always being your genuine self and looking out for my well-being. Your
hard work, determination, and passion for science are infectious. You inspired me greatly
throughout my academic journey.
vi
I want to especially thank my mother and father-in-law, Monika and Ingo. I would
not have finished this dissertation without your love and support. I am blessed to have
you in my life. You welcomed us and took care of Anja, Luna, and I. Thank you infinitely.
I feel incredibly fortunate that we will be a short drive away because we can’t imagine a
life far away from you. I want to especially thank my brother. You are one of a kind. I
would not be the man that I am without you in my life. Thank you for being a never-ending
source of inspiration, ideas, and philosophies. Your bravery, humility, and loyalty are to
be admired my brother. I also want to especially thank my sister. Your tenacity and
intelligence have pushed me to break all limitations I have self-imposed. Thank you for
supporting me when you moved to LA. Without you there, I would have most likely not
finished this degree. You have been a major source of inspiration for me.
I want to thank the most beautiful person I’ve ever met, Anja. You have given me
everything and then more…unconditionally. I feel eternally grateful from getting to share
this life with you. You have inspired me to take risks, to remain hopeful, and respond
courageously to the outcomes. “A man’s true character is not revealed by his failures, but
by the way he rises above them.” And with you, I can rise above anything. Thank you for
the family we created and the home we built. Thank you for taking care of Luna and I, for
being so nurturing, loving, and selfless. Thank you endlessly.
Finally, I want to thank my parents to whom this dissertation is dedicated to. The
values you instilled since childhood have propelled me to this point. Your unwavering
support and dedication to my vision makes me feel eternally blessed, grateful, and
responsible. I will not be the man I am today without your sacrifice, love, and strength.
vii
Table of Contents
Epigraph………………………………………………………………………………………....ii
Dedication………………………………………………………………………………………iii
Acknowledgements…………………………………………………………………………...iv
List of Tables………………………………………………………………………………….viii
List of Figures…………………………………………………………………………………..x
Abbreviations…………………………………………………………………………………xix
Preface………………………………………………………………………………………...xxii
Abstract………………………………………………………………………………………xxiv
Chapter 1
Synthesis of Ni-TiO2 Nanocomposites
as Enzyme-less, Amperometric Sensors
for the Electrooxidation of Glucose…………………………………………………………...1
Chapter 2
NiSn Alloys as Electrocatalysts
for the Urea Oxidation Reaction and
Integration into Direct Urea Fuel Cells……………………………………………………...20
Chapter 3
Part A: Recent Advancements in Silver-based
Catalysts for Electrochemical CO2 Reduction:
Detailed Mechanisms, Surface and Structural
Features, and Practical Applications………………………………………………………..39
Part B: Polyethyleneimine Encapsulated
Silver Nanoparticles as Cathode Electrodes
for Syngas Generation………………………………………………………………………..83
References……………………………………………………………………………………104
Appendix
Supporting Information 1…………………………………………………………….132
Supporting Information 2…………………………………………………………….153
Supporting Information 3…………………………………………………………….170
viii
List of Tables
Table 1 Selected examples of Ni–Ti electrode
systems for the electrooxidation of glucose…………………………………………………...4
Table 2 Equilibrium potentials (vs SHE) for
the electrochemical CO2 reduction to different
products in pH 7 aqueous solution…...............................................................................42
Table 3 Silver-incorporated bimetallic
catalysts for the electroreduction of CO2
into its different products. (1a) nanowires
(1b) nanospheres (1c) nanodimers
(2) normal hydrogen electrode
(3) standard calomel electrode
(4) standard hydrogen electrode……………………………………………………………...70
Table 4 EIS results derived from fitting
the experimental data to the Randles circuit
model for Pt/C and PEI-750k Ag NPs catalysts
in N2-saturated 0.5 M KHCO3 electrolyte.
The parameters include the solution resistance
(Rs), the constant phase element-T (CPE-T),
the constant phase element-P (CPE-P),
and the charge transfer resistance (Rct).
Results represent mean ± standard deviation
for both GEIS and PEIS mode………………………………………………………………...93
Table S1 Ni wt. % quantification and
comparison using (yellow) EDS, (green)
ICP–OES and (orange) XPS for the four
synthesized nanocomposites………………………………….…………………………....144
Table S2 Summary of the 60% Ni–TiO2/XC72R
nanocomposite activation in 0.1 M NaOH for
different potential windows (vs MMO) at a scan
rate of 50 mV s
–1………………………………………………………………………………148
Table S3 Elemental compositions for the
a) Ni80/C, b) Sn80/C, c) Ni60Sn20/C,
b) d) Ni40Sn40/C, e) Ni20Sn60/C materials
based on respective XRF spectra…………………………………………………………155
ix
Table S8 Ni and Sn wt. % quantification through
XRD and XRF for NiSn alloys and
monometallic analogues……………………………………………………………………..160
Table S9 Summary of electrochemical
performance for Ni-based catalysts during
activation, including maximum current density
(jmax), peak potential (E1/2), cycle number at
jmax, and diffusion coefficient (D) in alkaline media………………………………………...163
Table S10 Tabulated ECSA results for
Ni-based electrocatalysts in 1 M KOH
post-activation. The specific current density
(jsp) at 0.65 V was calculated by normalizing
to the mass of the catalyst (m = 0.05 mg)
instead of geometric surface area of RDE………………………………………………….164
Table S11 Tabulated Tafel equation values
for each electrocatalyst, the calculated exchange
current density (j0), and the correlation for
the line of best fit when normalized to the
geometric surface area of the RDE………………………………………………………….165
Table S12 Tabulated Tafel equation values
for each electrocatalyst, the calculated exchange
current density, and the correlation for the
line of best fit when normalized to their respective
ECSA…………………………………………………………………………………………..166
Table S13 Tafel slope (TS) analysis at three
different overpotential regions: η < 100 mV
(low), 100 mV< η < 250 mV (med), and
η > 250 mV (med) for Pt/C and PEI-750k Ag NPs………………………………………..172
x
List of Figures
Fig. 1 a) SEM micrograph of the
60% Ni-TiO2/XC72R nanocomposite at 50,000x
magnification and its b) EDS elemental mapping…………………………………………….7
Fig. 2 High resolution XPS spectra
for a) 20% Ni-TiO2/XC72R,
b) 40% Ni-TiO2/XC72R,
c) 60% Ni-TiO2/XC72R and d) 80% Ni/XC72R……………………………………………….9
Fig. 3 a) Activation of the
60% Ni-TiO2/XC72R catalyst in 0.1 mol L–1
NaOH from 0 to 0.9 V at a scan rate of
50 mV s–1 and at b) varying scan rates……………………………………………………….11
Fig. 4 CVs of the Ni-TiO2/XC72R
electrodes: a) 20%, b) 40%, c) 60%, and
d) 80% at a scan rate of 20 mV s–1
in
0.1 mol L–1 NaOH (solid) +
1 mmol L–1 glucose (dotted)…………………………………………………………………...13
Fig. 5 Amperometry responses at
potentials of a) 0.5 V, c) 0.6 V, and
e) 0.7 V with their corresponding calibration
curves (b, d, and f, respectively) for the
60% Ni-TiO2/XC72R electrode in
0.1 mol L–1 NaOH with 50 μmol L–1 glucose
aliquot addition every 50 s…………………………………………………………………….15
Fig. 6 a) CA responses for
60% Ni-TiO2/XC72R electrode at 0.7 V
in 0.1 M NaOH using 1 mM (black line)
and 50 μM (red line) aliquots per step
and the (b and c, respectively) calibration
curves with error bars (n = 3), d) Three
independent amperometry measurements
for 60% Ni-TiO2/XC72R electrode at
0.7 V in 0.1 M NaOH using sequential aliquot
additions of 1 mM glucose and 0.1 mmol L–1
of ascorbic acid, uric acid, and lactic acid……………………………………………………17
Fig. 7 XRD patterns of NiSn/C alloys
with varying compositions
(b; Ni20Sn60/C, c; Ni40Sn40/C, d; Ni60Sn40/C)
and the monometallic analogues
xi
(a; Sn80/C, e; Ni80/C). Peaks at 44.84°, 52.17°,
and 76.67° correspond to the (1 1 1), (2 0 0),
and (2 2 0) planes of nickel NPs, respectively……………………………………………….26
Fig. 8 a) CV for the Ni60Sn20/C electrocatalyst
at a scan rate of 50 mV s–1
in 1 mol L–1 KOH
after 15 cycles showing the 200 mV cathodic
window for ECSA; b) peak current vs. square
root of scan rate plot for N60Sn20/C (green),
N40Sn40/C (red), N20Sn60/C (blue), and
Ni80/C (black) demonstrating the slope and
correlation for each sample; c) CVs for
Ni60Sn20/C showing the last activation cycle
(solid) in 1 mol L–1 KOH and the oxidation of
0.33 mol L–1 urea (dashed) at a scan rate of
50 mV s–1
; d) oxidation of 0.33 mol L–1 urea
from initial to final (n = 25) CV…………………………………………………………………29
Fig. 9 a) Tafel plots for N40Sn40/C (red),
N20Sn60/C (blue), N60Sn20/C (green), and
Ni80/C (black) normalized to the geometric
area (0.195 cm2
) of the electrode and to
the ECSA of each electrocatalyst (inset)
with 0.33 mol L–1 urea in 1 mol L–1 KOH;
b) Constant potential urea oxidation at 0.6 V
(vs MMO) for the electrocatalysts in
alkaline media…………………………………………………………………………………..31
Fig. 10 Nyquist plot for N20Sn60/C (blue),
N40Sn40/C (red), N60Sn20/C (green),
Ni80/C (black),and Sn80/C (orange)
electrodes with 0.33 mol L–1 urea in
1 mol L–1 KOH at a) 0.45 V and b) 0.55 V.
The data was recorded over a frequency
range from 10 kHz to 100 mHz with
perturbation amplitude of 25 mA……………………………………………………………...33
Fig. 11 a) Polarization curves employing
N60Sn20/C anode and Pt/C cathode
separated by a N117 PEM, b) using urea
concentration of 1 M at increasing
temperatures, c) CA measurements at
10 mV demonstrating the stability of the
DUFC at 8 h (inset) and 20 h,
and d) increasing anolyte concentration to
7 M KOH at 25, 40, and 60 C…………………………………………………………………36
xii
Fig. 12 UV-Vis absorption spectra of
PEI-750k-Ag+ complex (blue curve) and
PEI Ag NPs (black curve) after addition of
the reducing agent; X-ray diffraction patterns
of PEI Ag NPs with varying PEI Mw of
800 (red), 25,000 (yellow), and
750,000 (green)………………………………………………………………………………...88
Fig. 13 SEM micrographs PEI-750k
Ag NPs at magnifications of a) 65,000x
and b) 150,000x with accelerating voltage
of 10 kV………………………………………………………………………………………….90
Fig. 14 a) TEM image of PEI-750k Ag NPs
and b) high-resolution TEM image of an
Ag NP representative………………………………………………………………………….91
Fig. 15 Nyquist plots for 5% Pt/C (blue)
and PEI-750k Ag NPs (green) electrocatalysts
in N2-saturated 0.5 M KHCO3 electrolyte.
The data was recorded over a frequency range
from 1 kHz to 1 Hz at OCV. The left panel
presents the galvanostatic mode with a
perturbation amplitude of 50 mA, while the
right panel shows the potentiostatic mode with
a perturbation amplitude of 50 mV. The
experimental data are fitted to a Randles circuit
model (inset), as shown by the black curves………………………………………………...93
Fig. 16 a) LSVs after iR correction of
PEI-750k Ag NPs (green) and Pt/C (blue) in
N2-saturated 0.5 M KHCO3 at a scan rate
of 10 mV s–1
, Tafel plots of c) PEI-750k Ag NPs
and d) Pt/C at three overpotential regions, and
d) Tafel analysis at higher overpotentials
demonstrating the more sluggish kinetics (inset)
of PEI-750k Ag NPs compared to Pt/C……………………………………………………….95
Fig. 17 a) LSV curves of pristine TCP
(dashed), polycrystalline Ag foil (blue) and
PEI-750k Ag NPs (green) in CO2-saturated
0.5 M KHCO3 at a scan rate of 50 mV s
–1
b) LSVs of PEI-750k Ag NPs at different
scan rates, c) plot of the square root of
scan rate vs. peak cathodic current density
xiii
and d) plot of log scan rate vs log peak
cathodic current density for PEI-750k Ag NPs
at three different potential regions……………………………………………………………97
Fig. 18 LSVs of Ag NPs in PEI with varying
molecular weights (750k, green; 25k, yellow;
800, red) in CO2–saturated 0.5 M KHCO3 at
a) 10 mV s–1 and b) 50 mV s–1…………………………………………………………………99
Fig. S1 SEM image and EDS elemental
mapping of the a) 20%, b) 40%, and c) 80%
Ni–TiO2/XC72R nanocomposites…………………………………………………………...134
Fig. S2 EDS spectrum and quantification
of the a) 20%, b) 40%, c) 60%, and d) 80%
Ni–TiO2/XC72R nanocomposites…………………………………………………………...137
Fig. S3 ICP–OES measurements for
20% Ni–TiO2/XC72R, 40% Ni–TiO2/XC72R,
60% Ni–TiO2/XC72R, and 80% Ni/XC72R
at a) 221.647 nm and b) 231.604 nm……………………………………………………….141
Fig. S4 XRD patterns for the four Ni–TiO2
nanocomposites: 20% Ni–TiO2/XC72R (green),
40% Ni–TiO2/XC72R (blue),
60% Ni–TiO2/XC72R (red), and
80% Ni/XC72R (gray). Double dagger (‡)
signals correspond to TiO2 anatase
and asterisk (*) signals correspond to Ni NPs……………………………………………...142
Fig. S5 XPS survey spectra four Ni–TiO2
nanocomposites: 20% Ni–TiO2/XC72R (green),
40% Ni–TiO2/XC72R (blue),
60% Ni–TiO2/XC72R (red), and
80% Ni/XC72R (gray)………………………………………………………………………...143
Fig. S6 Activation of a) 20%, b) 40%,
c) 60% Ni–TiO2/XC72R electrodes in
0.1 M NaOH at a scan rate
of 50 mV s–1
from 0 to 0.6 V………………………………………………………………….145
Fig. S7 Activation after 8 cycles of a) 20%,
b) 40%, and c) 80% Ni–TiO2/XC72R electrodes
in 0.1 M NaOH at a scan rate of 50 mV s–1
from 0 to 0.9 V……………………………………………………………………………......149
xiv
Fig. S8 Potentiostatic measurements for the
Ni–TiO2/XC72R nanocomposites in
0.1 M NaOH at 0.7 V (vs. MMO) for 350 s………………………………………………..152
Fig. S9 Amperometric studies for
60% Ni–TiO2/XC72R electrode at 0.7 V
in 0.1 M NaOH (black line) and
0.1 M NaOH + 0.1 M NaCl (green line) and
its calibration curve with error bars (n = 3)………………………………………………..153
Fig. S10 XRF spectra for a) Ni80/C, b) Sn80/C,
c) Ni60Sn20/C, d) Ni40Sn40/C, and e) Ni20Sn60/C
powders……………………………………………………………………………………….155
Fig. S11 CVs at a scan rate of 50 mV/s for
a) Ni20Sn60/C, b) Ni40Sn40/C, c) Ni80/C, and
d) Sn80/C in 1 M KOH demonstrating the
Ni(OH)2/NiOOH redox couple………………………………………………………………161
Fig. S12 CVs of the Ni20Sn60/C, Ni40Sn40/C,
Ni80/C, and Sn80/C electrodes showing the
last activation cycle (solid) in 1 mol L–1 KOH
and the oxidation of 0.33 mol L–1 urea
(dashed) after 25 cycles at
a scan rate of 50 mV s–1…………………………………………………………………….162
Fig. S13 EIS for N20Sn60/C (blue), N40Sn40/C (red),
N60Sn20/C (green), Ni80/C (black), and
Sn80/C (orange) electrodes with 0.33 mol L–1
urea in 1 mol L–1 KOH at 0.5 V.
The data was recorded over a frequency range
from 10 kHz to 100 mHz with perturbation
amplitude of 25 mA…………………………………………………………………………..167
Fig. S14 Polarization and power density curves
for the DUFC at a) increasing anolyte and
catholyte flow rates under r.t. conditions
and b) at different temperatures with the
optimized flow rate of 580 mL min–1……………………………………………………….168
Fig. S15 Chronoamperometric response
showing the current density versus time
at a constant applied potential of 0.2 V,
illustrating the decay in current over time
of the DUFC at r.t…………………………………………………………………………….169
xv
Fig. S16 Polarization and power density
curves of the DUFC using
mTPN1-TMA anion exchange membrane
at varying temperatures……………………………………………………………………...170
Fig. S17 Nyquist plots under
galvanostatic mode at OCV from
10 kHz to 1 Hz for a) PEI-750k Ag NPs
and b) Pt/C in N2-saturated 0.5 M KHCO3
at sinusoidal current perturbations of
1 mA (yellow), 10 mA (red), 25 mA (black),
and 50 mA (blue). K-K relations for
c) PEI-750k Ag NPs and d) Pt/C comparing
the real and imaginary components
of impedance and their
frequency dependance………………………………………………………………………173
Fig. S18 Nyquist plots under
potentiostatic mode at OCV from
10 kHz to 1 Hz for a) PEI-750k Ag NPs
and b) Pt/C in N2-saturated 0.5 M KHCO3
at sinusoidal current perturbations of
1 mV (yellow), 10 mV (red), 25 mV (black),
and 50 mV (blue). K-K relations for
c) PEI-750k Ag NPs and d) Pt/C
comparing the real and imaginary
components of impedance and their
frequency dependance………………………………………………………………………174
Fig. S19 Maximum deviation |Δ|max
in the real and imaginary components
of impedance as a percentage for
different perturbation amplitudes.
Panels (a) and (b) show GEIS |Δ|max
for current perturbations (in mA) for
PEI-750k Ag NPs and Pt/C, respectively.
Panels (c) and (d) show PEIS |Δ|max
for voltage perturbations (in mV) for
PEI-750k Ag NPs and Pt/C,
respectively. Blue bars represent the
real part of the impedance, while
orange bars represent the imaginary part.
These plots indicate the perturbation
amplitudes at which the impedance data
show minimal deviation, thereby
validating the data's consistency and
xvi
suitability for further EIS analysis……………………………………………………………175
Fig. S20 Chromatograms after
constant potential electrolysis of
15 min at six different potentials for
PEI-750k Ag NPs in
N2-saturated 0.5 M KHCO3.
Hydrogen, nitrogen, and carbon
dioxide signals with corresponding
retention times of 2.21, 2.51, and
5.42 min………………………………………………………………………………………..176
Fig. S21 LSVs of pristine glassy carbon
electrode (GCE; dashed) and PEI-750k
Ag NPs on GCE (green) at a potential
window from 0 to –0.8 V (vs RHE) in
CO2-saturated 0.5 mol L–1 KHCO3 under
varying scan rates…………………………………………………………………………….177
Fig. S22 LSVs of pristine microporous
layer carbon paper (MPL; dashed) and
PEI-750k Ag NPs on MPL (green) under
a potential window from 0 to –0.8 V in
CO2-saturated 0.5 mol L–1 KHCO3 under
varying scan rates…………………………………………………………………………….178
Fig. S23 LSVs of pristine Toray carbon
paper (TCP; dashed) and PEI-750k Ag
NPs on TCP (green) at a potential window
from 0 to –0.8 V (vs RHE) in
CO2-saturated 0.5 mol L–1 KHCO3 under
varying scan rates…………………………………………………………………………….179
Fig. S24 LSVs of pristine glassy carbon
electrode (GCE; dashed) and PEI-750k
Ag NPs on GCE (green) at a potential
window from 0 to –1.2 V (vs RHE) in
CO2-saturated 0.5 mol L–1 KHCO3 under
varying scan rates…………………………………………………………………………….180
Fig. S25 LSVs of pristine microporous
layer carbon paper (MPL; dashed) and
PEI-750k Ag NPs on MPL (green) at
a potential window from 0 to –1.2 V in
CO2-saturated 0.5 mol L–1 KHCO3 at
different scan rates…………………………………………………………………………...181
xvii
Fig. S26 LSVs of pristine Toray carbon
paper (TCP; dashed) and PEI-750k Ag
NPs on TCP (green) at
a potential window from 0 to –1.2 V in
CO2-saturated 0.5 mol L–1 KHCO3 under
different scan rates…………………………………………………………………………...182
Fig. S27 Chromatograms after constant
potential electrolysis of 15 min at
six potentials for a) pristine MPL and
b) PEI-750k Ag NPs in CO2-saturated
0.5 mol L–1 KHCO3. Hydrogen, nitrogen,
carbon monoxide, and carbon dioxide signals
with corresponding retention times
of 2.19, 2.53, 2.70 and 5.33 min…………………………………………………………….183
Fig. S28 Chromatograms after constant
potential electrolysis of 15 min at six potentials
for a) pristine MPL and b) PEI-750k Ag
NPs in CO2-saturated 0.5 mol L–1 KHCO3.
Hydrogen, nitrogen, carbon monoxide,
and carbon dioxide signals with
corresponding retention times of
2.19, 2.53, 2.70 and 5.33 min…………………………………………………………….….184
Fig. S29 Chromatograms after constant
potential electrolysis of 15 min at six potentials
for (a) pristine TCP and (b) PEI-750k Ag
NPs in CO2-saturated 0.5 mol L–1 KHCO3.
Hydrogen, nitrogen, carbon monoxide, and
carbon dioxide signals with corresponding
retention times of 2.19, 2.53, 2.70 and 5.33 min…………………………………………...185
Fig. S30 Chromatograms after constant
potential electrolysis of 15 min at three potentials
for PEI-750k (green), PEI-25k (yellow), and
PEI-800 (red) Ag NPs in CO2-saturated
0.5 mol L–1 KHCO3. Hydrogen, nitrogen,
carbon monoxide, and carbon dioxide signals
with corresponding retention times
of 2.19, 2.53, 2.70 and 5.33 min…………………………………………………………….186
Fig. S31 TGA isotherm of PEI-Ag NPs
with PEI Mw of a) 750k, b) 25k, and
c) 800, Ag NPs at cycles with
xviii
varying temperatures of 25, 55,
and 85°C………………………………………………………………………………………187
xix
Abbreviations
Anion Exchange membrane AEM
Carbon capture and sequestration CCS
Carbon capture and utilization CCU
Charge transfer resistance Rct
Chronoamperometry CA
Constant phase element CPE
Covalent organic frameworks COFs
Current i
Current density j
Cyclic voltammetry CV
Density functional theory DFT
Diffusion coefficient D
Direct urea fuel cell DUFC
Electrochemical CO2 Reduction Reaction ECO2RR
Electrochemical impedance spectroscopy EIS
Electrochemical surface area ECSA
Energy dispersive x-ray spectroscopy EDS
Exchange current density j0
Faradaic efficiency FE
Galvanostatic electrochemical impedance spectroscopy GEIS
Gas chromatography GC
Gas diffusion electrode GDE
Glassy carbon electrode GCE
Glassy carbon rotating ring disk electrode GC RDE
xx
Glucose oxidation reaction GOR
High resolution transmission electron microscopy HRTEM
Highest occupied molecular orbital HOMO
Hydrogen evolution reaction HER
Inductively coupled plasma ICP
Kromers-Kronig K-K
Linear Sweep Voltammetry LSV
Lowest unoccupied molecular orbital LUMO
Mass current density jsp
Maximum current density jmax
Membrane-electrode assembly MEA
Mercury/mercury oxide MMO
Metal-organic frameworks MOFs
Microporous layered carbon paper MPL
Nanoparticles NPs
Open circuit voltage OCV
Optical emission spectroscopy OES
Overpotential η
Oxygen evolution reaction OER
Oxygen reduction reaction ORR
Peak current ip
Peak potential E1/2
Photoelectron spectroscopy XPS
Polyethyleneimine PEI
Potential E
xxi
Potentiostatic electrochemical impedance spectroscopy PEIS
Proton exchange membrane PEM
Proton-coupled electron transfer PCET
Rate determining step RDS
Reversible hydrogen electrode RHE
Rotating ring disk electrode RDE
Saturated calomel electrode SCE
Scan rate v
Scanning electron microscopy SEM
Scanning tunneling microscopy STM/STEM
Silver/silver chloride Ag/AgCl
Single atom catalysts SACs
Solution resistance Rs
Standard hydrogen electrode SHE
Standard reduction potential E0
Tafel slope TS
Thermal conductivity detector TCD
Thermogravimetric analysis TGA
Toray carbon paper TCP
Urea oxidation reaction UOR
Vulcan carbon XC72R
X-ray diffraction XRD
X-ray fluorescence spectrometry XRF
xxii
Preface
For my graduate work, I chose to explore various research ideas rather than
focusing on a single principal proposal. Although each chapter represents an independent
project, they are all connected through the fields of materials science, electrochemistry,
and engineering. I found this approach to be the most effective in conveying the
overarching theme of this dissertation: the design of nanomaterials for tailored
electrochemical transformations.
I sought to fulfill this theme through three guiding principles. Firstly, it was
particularly important that the synthetic methodologies be pragmatic. The one-pot
approach and relatively mild conditions highlighted in this work demonstrate the
commercial viability of these processes. The compositions of the targeted materials were
further confirmed through various characterization techniques, validating both the
synthesis method and the materials’ intended structure. Secondly, the characterized
materials were tested as electrocatalysts for specific redox transformations of small
molecules. These electrodes were subsequently optimized through a range of
electrochemical methods and techniques. Lastly, different experimental conditions were
tested to further enhance the performance of the electrochemical devices, with a focus
on maximizing current output, power generation, and both short- and long-term stability.
The nature of this work embodies my broader scientific philosophy: curiositydriven, but solution-oriented. My decision to pursue these diverse research directions
reflects my desire to address real-world challenges, particularly in the development of
sustainable energy solutions. By bridging fundamental research with practical
xxiii
applications, I hope this dissertation contributes to the scientific community. Ultimately, I
aim to support the ongoing global pursuit of sustainable, renewable energy technologies.
xxiv
Abstract
The interchangeable conversion between chemical energy and electrical energy is
crucial for achieving a sustainable infrastructure. Fuel cells and electrolytic devices will
serve as essential technologies for these energy conversions, enabling both efficient
generation and storage of renewable power. While significant progress has been made,
further innovation at the materials level, ranging from the micro to the macro scale,
remains necessary. Specifically, catalytic materials used in anode and cathode electrodes
continue to be a focus of intensive research. Ideally, these electrocatalysts should be
non-precious metal-based (i.e., cheap and abundant), highly active, and stable under
operational conditions. This dissertation presents practical synthetic methodologies for
the development of innovative nanomaterials designed for electrochemical devices,
which include glucose biosensors, AEM/PEM fuel cells, and CO2 electrolyzers. Different
compositions of a targeted catalyst were synthesized, characterized using advanced
analytical techniques (e.g., microscopy, spectroscopy, diffraction) and studied for their
tailored redox transformations. The material’s structure-activity relationships were
analyzed via a variety of methods and techniques, such as electrochemical kinetics (RDE
and Tafel analysis), polarization (CSV/LSV), constant bias (CA/CP), and impedance
spectroscopy (EIS), all aimed at enhancing reactivity. The optimized electrocatalyst was
then integrated into a device, engineered to maximize energy output and operational
stability. These findings demonstrate the feasibility of employing cost-effective, highperformance nanomaterials in market-ready, next-generation electrochemical
technologies.
1
Chapter 1
Synthesis of Ni-TiO2 Nanocomposites as Enzyme-less, Amperometric
Sensors for the Electrooxidation of Glucose
Abstract
The simple synthesis of a Ni-TiO2 nanocomposite supported on Vulcan carbon (XC-72R)
for the electrooxidation reaction of glucose is reported. Four transition metal weight ratios
were synthesized and characterized. Cyclic voltammetry (CV) studies in 0.1 M NaOH
demonstrate that the four metal catalysts can effectively oxidize 1 mM glucose, with the
3:1 (60%) Ni to Ti nanocomposite yielding the highest current. The 60% Ni-TiO2/XC72R
catalyst was used to construct an enzyme-less, chronoamperometric sensor for glucose
detection in alkaline medium. Using 50 µM aliquots of glucose at a potential of 0.7 V (vs
MMO), the sensor responded rapidly (< 3 s), provided a sensitivity of 3300 µA mM–1
cm–2
, detection limits of 144 nM (Signal/Noise = 3), and excellent selectivity and
reproducibility. The glucose aliquot concentrations were then increased to 1 mM to mimic
physiological blood conditions of 1-20 mM. At the same potential, the sensor continued
to respond rapidly (< 1 s), showed a sensitivity of 273.7 µA mM–1 cm–2
, detection limits of
3.13 µM (S/N = 3), and excellent selectivity and reproducibility. The catalyst also exhibited
an ideal anti–poisoning capability to free chloride ions and negligible signals towards
other interfering species.
2
Introduction
In 1994, The Centers for Disease Control and Prevention declared that diabetes had
reached epidemic proportions. Since then, however, little has been done to suppress the
yearly increasing statistics.1,2 As of 2019, global diabetes prevalence reached 463 million
people (9.3%) and it is expected that by 2045, an astounding 700 million people (10.9%)
worldwide will be diagnosed with this condition.3 Studies throughout the 2020 coronavirus
(COVID-19) pandemic demonstrated that older age and presence of diabetes mellitus,
hypertension, and obesity significantly increased the risk for hospitalization and death in
COVID-19 patients.4 Presently, a positive diabetes diagnosis undoubtedly diminishes an
individual’s health and quality and duration of life. Efforts should focus on both raising
awareness at a global scale and developing preventive measures and effective
treatments as part of fundamental care and health.5 Reliable sensors for monitoring
glucose levels continue to be of major importance for diabetes control and treatment.
Enzymatic biosensors with incorporated transductors, specifically glucose oxidase
biosensors, have been studied and developed extensively for this mission.6–8 Despite
their low detection limit and high selectivity, enzymatic biosensors often require complex
immobilization techniques of the enzyme onto a substrate electrode and often suffer from
leakage and poor stability. Glucose oxidase biosensors can also only work properly under
specific temperatures, pHs, and chemical environments.9–11 Alternatively, nonenzymatic
amperometric biosensors continue to be investigated as potential substitutes to
enzymatic biosensors. Previous literature has focused on the development of
nanostructured metals12–14
, metal alloys9,15–18
, and metal oxides19–23 as electrocatalytic
3
materials for glucose oxidation. Many of these materials display faster response times,
lower detection limits, and better stability compared to enzymatic biosensors. However,
their selectivity towards other carbohydrates besides glucose are often discernable and
they can suffer from poisoning by intermediates and other ions.23,24
Electrochemical oxidation of glucose has been demonstrated to be enhanced by
nickel catalysts given the existence of the Ni(OH)2/NiOOH (Ni2+/Ni3+) redox couple in an
alkaline medium.13 Nickel’s popularity stems from its spontaneous transformation into
NiO, which consequently forms Ni(OH)2 due to adsorption of hydroxide ions.12 At specific
potentials, the Ni(OH)2 is converted to NiOOH, which is the active form for the
electrooxidation of glucose.25 Nickel is also an abundant, nontoxic, inexpensive metal
demonstrating commercial and industrial applicability. Continually, the incorporation of
transition metal oxides as substrates for catalysis is common in the field of chemical
sensors. Specifically, TiO2 has been widely studied as a substrate for chemical sensors
due to its multiple morphologies, biocompatibility, nontoxicity, and chemical and thermal
stability. It’s classification as an n-type semiconductor improves the transport of surface
reaction electrons to the metal substrate.26 This ultimately enhances the performance of
the sensor.
Previous studies have examined the synergistic effects of nickel and titanium for
the oxidation of glucose, as shown in Table 1. However, these investigations often
employ complex morphologies, such as Ti nanowire arrays or sheets, which require pre–
treatment protocols and more elaborate synthetic methods.27–31 Herein we expand on
previous examples by presenting a simple chemical reduction of Ni nanoparticles unto a
TiO2 anatase substrate with Vulcan carbon for the development of a nonenzymatic sensor
4
for glucose electrooxidation. Multiple wt. % of Ni and TiO2 were investigated to optimize
the ratio of the chemical species that would be the most active towards the oxidation
reaction. The two linear ranges of 50 µM-1 mM and 1-20 mM glucose demonstrated a
versatile sensor with exceptional response times and selectivity. Our studies using a
nanocomposite of 60% Ni on 20% TiO2 as an anode electrode reports the highest
sensitivity of 3300 µA mM–1 cm–2 with the lowest detection limit of 144 nM.
Materials and Methods
Chemicals.– Ni(NO3)2 · 6H2O (99.9% purity) was purchased from Strem Chemicals,
anatase TiO2 catalyst support (1/8” pellets) was purchased from Alfa Aesar, D-Glucose
Anhydrous (granular) purchased from Macron Chemicals, Vulcan carbon (XC-72R) from
Table 1 Selected examples of Ni–Ti electrode systems for the electrooxidation of glucose.
Electrode Linear Range (mM)
Sensitivity
(µA cm-2 mM- 1)
LOD (µM) Ref.
Ni nanoflakes/Ti sheets 0.05 - 0.6 7.32 x 10- 3 1.2 27
Ni-Ti/TiO2 nanotube arrays 0.1 - 1.7 200 4 28
Ni-TiO2 nanowire arrays 0.2 - 2 1472 10 29
Ni-Ti-O nanotubes 0.002 - 0.2 83 0.13 30
NiTi sheets 0.03 - 14 192 8 31
Ni/TiO2 0.05 - 1 3300 0.144 this work
1 - 20 273.7 3.13 this work
5
FuelCell Store, and Liquion™ LQ-1105 5% Nafion binder solution from Ion Power.
Analytes were used as received without further purification: L-ascorbic acid sodium salt
(99%) purchased from Alfa Aesar, DL-lactic acid purchased from Sigma Aldrich, and uric
acid purchased from Strem Chemicals. Millipore water (18.2MΩ cm) was used for all
solution preparations. All other reagents were of analytical grade and used without further
purification. The electrochemical measurements were performed in 0.10 M NaOH
solution.
Preparation of Ni-TiO2/XC72R catalyst.– To synthesize the 60% Ni-TiO2/XC72R
catalyst, 297.3 mg Ni(NO3)2 · 6H2O and 33.4 mg anatase TiO2 were dispersed in 20 mL
ethylene glycol. Then 20.0 mg of Vulcan carbon XC-72R was added into the solution and
stirred for 30 min followed by slow addition of 936 µL of 35% hydrazine hydrate. The
solution was stirred for an additional 30 min. The mixture was transferred to a Teflon lined
vessel and heated to 150 °C for 6 h in an oven. After the 6 h had elapsed, the solution
was transferred to a centrifuge tube and 20 mL of deionized (DI) H2O were added. The
solution was centrifuged at 6000 rpm for 5 min and the supernatant was discarded. The
washing process was repeated two additional times with water and finally dried in oven
for 24 h at 60 °C resulting in a black powder. Similar syntheses were completed but
varying the weight ratios 1:3 (20%), 1:1 (40%), and 4:0 (80%) of Ni:Ti.
Electrochemical testing.– The electrochemical circuit consisted of a glassy carbon
roating ring disk electrode (GC RDE) with a geometric surface area of 0.195 cm2 attached
to a Pine Research MSR Rotator, 4.24 M KOH MMO (mercury/mercury oxide) reference
electrode, and Pt wire counter electrode. Prior to any measurements the RDE was
polished with a microcloth and 0.05 µm alumina slurry bought from Electron Microscopy
6
Sciences. All voltammetric determinations were carried out with a Solartron Analytical
1287A Electrochemical Interface potentiostat/galvanostat instrument. The
electrochemical measurements were performed in 0.10 M NaOH solution and 1 mM
glucose. Nitrogen gas was flushed for 30 min prior to experiments and steadily flowed
during all voltammetric measurements. Chronoamperometry (CA) measurements were
performed in triplicate. Catalyst inks were obtained by dissolving 2.5 mg of dried catalyst
and 7.5 mg (~10 µL) of 5% Nafion binder into 100 µL isopropanol/900 µL DI H2O mixture.
The solution was sonicated for 24 min. Then 20 µL of the catalyst ink was dropcast unto
the GC RDE and dried in an oven at 60 °C for 30 min. All electrochemical experiments
were normalized to the geometric surface area of the GC RDE.
Physical characterization.– Powder X-ray diffraction (XRD) measurements were
performed on a Rigaku Ultima IV X-ray diffractometer with Cu Kα (0.154056 nm) radiation
source and a scan rate of 6° min-1
from a 2Ɵ value of 10° to 90°. Scanning electron
microscopy (SEM) images were obtained from a Nova NanoSEM 450 field emission
scanning electron microscope with an acceleration voltage of 18.0 keV. Energy dispersive
X-ray spectroscopy (EDS) measurements were conducted in a JSM 7001F Field
Emission Scanning Electron Microscope. Inductively Coupled Plasma Optical Emission
Spectroscopy (ICP-OES) measurements were conducted in a Thermo Scientific iCAP
7400 ICP-OES Analyzer monitoring both Ni spectral lines of 221.647 nm and 231.604
nm. X-ray Photoelectron (XPS) spectra were collected on a Kratos Axis Ultra DLD using
a mono Al anode with a pass energy of 160 keV for the survey scan and 20 keV for the
high-resolution scans.
7
Results & Discussion
Physical characterization.– The four nanocomposites, with varying nickel to titania wt.
%, were first synthesized and characterized by SEM. Figure 1a demonstrates the SEM
image for the 60% Ni-TiO2/XC72R catalyst at 50,000x magnification. Figures 1b shows
an SEM image and the EDS mapping for this nanocomposite. It is observable from the
mapping images that the material is composed of Ti, Ni, O, and C only. The SEM images
and corresponding EDS mapping images are demonstrated in Fig. S1 for the 20% NiTiO2/XC72R, 40% Ni-TiO2/XC72R, and 80% Ni/XC72R, respectively. The EDS spectra in
Fig. S2 shows the signal for all constituent elements and their corresponding wt. % in the
Ni-TiO2/XC72R nanocomposites without any other impurities present. According to EDS,
the nickel content in the 20% Ni-TiO2/XC72R, 40% Ni-TiO2/XC72R, 60% Ni-TiO2/XC72R,
and 80% Ni/XC72R nanocomposites were 24.02%, 45.30%, 58.48%, and 80.60%,
respectively. ICP-OES measurements were conducted to further quantify the Ni content
C K
Fig. 1 a) SEM micrograph for the 60% Ni-TiO2/XC72R nanocomposite at 50,000x magnification
and its b) EDS elemental mapping.
Ti K O K Ni K C K
a b
8
in the synthesized nanocomposites. Solutions of 55 ppm in 2% HNO3 were prepared for
each of the four nanocomposites and the Ni spectral lines of 221.647 nm and 231.604
nm were selected. Figure S3 demonstrates that, as expected, the nickel loading
increased analogously to the incrementing nickel-to-titania ratio for both spectral lines.
The nickel loading in the nanocomposites referencing the 221.647 nm and 231.604 nm
spectral lines respectively were determined to be 16.89% and 19.06% for the the 20% NiTiO2/XC72R, 39.03% and 41.11% for the 40% Ni-TiO2/XC72R, 58.99% and 60.86% for
the 60% Ni-TiO2/XC72R, and 77.63% and 79.34% for the 80% Ni/XC72R.
XRD measurements were conducted to obtain the catalysts’ crystallographic data
and confirm the formation of a Ni-TiO2 nanocomposite with varying metal ratios. The XRD
patterns in Figure S4 shows reflections for Ni nanoparticles (*) on an anatase TiO2
substrate (‡) . The peaks with 2θ values of 44.59°, 51.91°, and 76.49° correspond to the
(1 1 1), (2 0 0), and (2 2 0) planes in nickel nanoparticles, respectively.32 Whereas, the
peaks corresponding to TiO2 anatase are given at 2Ɵ of 25.44°, 37.10°, 44.59°, 54.00°,
55.18°, 62.81°, 68.91°, 70.40°, 75.17°, 76.22°, and 82.82°.33 As the titania wt. %
decreases and more nickel is incorporated into the catalysts, the intensity from the titania
reflections diminish. Eventually, those reflections disappear at the 80% Ni/XC72R catalyst
and those from the nickel nanoparticles become more apparent. As expected, the 20%
Ni-TiO2/XC72R nanocomposite displays the lowest intensity for the Ni nanoparticle
reflections and the highest for the TiO2 substrate. Furthemore, the nickel crystallite size
increases as the wt. % of TiO2 decreases in the catalysts. The average crystallize size of
the catalysts was calculated using the Debye-Scherrer equation and the broadening of
the (1 1 1) peak. The average crystallite size for the 20% Ni-TiO2/XC72R, 40% Ni-
9
TiO2/XC72R, 60% Ni-TiO2/XC72R, and 80% Ni/XC72R were determined to be 9, 11, 12,
and 13 nm, respectively. The nanocomposites were then characterized through XPS in
order to understand the surface compositions.
XPS wide scans are shown in Fig. S5. The four different nanocomposites show the
presence of only the expected elements. Carbon, oxygen, and nickel signals are revealed
in all the samples while titanium cannot be observed in the 80% Ni/XC72R
nanocomposite, as expected. These XPS survey scans help validate the presence of
titania at the surface of the nanocomposites except for the 80% Ni/XC72R catalyst.
Furthermore, the nickel wt. % in each nanocomposite surface were determined from the
survey scans, shown in Table S1. The 20% Ni-TiO2, 40% Ni-TiO2, 60% Ni-TiO2, and 80%
Fig. 2 High resolution XPS spectra for a) 20% Ni-TiO2/XC72R, b) 40% Ni-TiO2/XC72R, c) 60%
Ni-TiO2/XC72R and d) 80% Ni/XC72R.
a b
c d
10
Ni/XC72R nanocomposites had wt. % of 11.93%, 16.58%, 21.13%, and 34.63%,
respectively. These values are lower compared to the EDS data given that the penetration
depth of the XPS electron beam occurs strictly at the surface of the material. Nonetheless,
the nickel wt. % at the surface are in accordance with the expected trend as nickel content
is increased in the catalysts. Figure 2 demonstrates the high-resolution scans for the four
nanocomposites monitoring the Ni 2p core level. The binding energies of the Ni 2p can
be resolved into a doublet of Ni 2p1/2 and Ni 2p3/2, as a result of spin-orbital coupling. The
Ni 2p3/2 curves are deconvoluted into two regions corresponding to the oxidation states of
Ni(II) at 855.5 eV and Ni(0) at 852 eV. Therefore, the high resolution XPS scans can
validate the coexistence of Ni(II) and Ni(0) on the surface of the material. The presence
of metallic nickel in the 20% Ni-TiO2, 40% Ni-TiO2, 60% Ni-TiO2, and 80% Ni/XC72R are
2.48%, 3.55%, 7.82%, and 1.80%, respectively. This quantification demonstrates that the
80% Ni/XC72R is mostly in the Ni(II) state. It also shows that an increase in titania content
in the nanocomposites leads to more nickel nanoparticle formation, with the 60% Ni-TiO2
having the highest presence at 7.82% Ni(0).
Catalyst activation & stability.– Prior to glucose oxidation measurements, the activation
and stability of the synthesized catalysts were investigated. The 60% Ni-TiO2/XC72R
nanocomposite was subjected to activation treatment in 0.1 M NaOH at various potentials
by applying a sweep rate of 50 mV/s until the maximum currents were obtained. Figure
S6 demonstrate the activation of the 60% Ni-TiO2 nanocomposite at different potential
windows. Table S2 summarizes these results and shows that widening the potential
window leads to more rapid stabilization of the active nickel species by decreasing the
number of CV scans. It also demonstrates that higher currents are generated at larger
11
potential windows. Figure 3a displays 8 consecutive CV scans of the 60% Ni-TiO2/XC72R
electrode at 0.9 V. The second scan demonstrates the appearance of the cathodic and
anodic peaks at 398 mV and 598 mV, respectively. This redox couple can be attributed
to the nickel oxidation from Ni2+/Ni3+. The cathodic peak shifts to more negative potentials
and the anodic peak to more positive ones as the cycle numbers are increased. These
translations in peak position are due to the changes in the crystal structures of the Ni(OH)2
and NiOOH constituents of the surface film.13 The gradual increase in maximum peak
height demonstrates the progressive nucleation of NiOOH species from Ni(OH)2 to form
an active layer that stabilizes at 305 mV and 705 mV after just 8 cycles. The remaining
20% Ni-TiO2, 40% Ni-TiO2, and 80% Ni nanocomposites were activated in 0.1 M NaOH
for 8 cycles as shown in Fig. S7. From the catalyst activation experiments, it is noticeable
that the 60% Ni-TiO2 yields the highest current density at 8 cycles. This demonstrates that
the wt. % ratio of 3:1 nickel-to-titania is the most electrochemically active. To gain further
insight into the difference in the observed current, the electrochemical surface area
(ECSA) for each catalyst was calculated using the “‘Beta’ method” by integrating the
stabilized transition of Ni(II) to Ni(III).34 For the 20% Ni-TiO2 and the 40% Ni-TiO2 catalysts,
Fig. 3 a) Activation of the 60% Ni-TiO2/XC72R catalysts in 0.1 mol L–1 NaOH from 0 to 0.9 V at a
scan rate of 50 mV s
–1 and at b) varying scan rates.
a b
12
the integration was calculated in the potential window of 0.5 to 0.7 V and for the 60% NiTiO2 and the 80% Ni catalysts, the integration was calculated in the potential window of
0.5 to 0.75 V. The calculated ECSA for the 20% Ni-TiO2/XC72R, 40% Ni-TiO2/XC72R,
60% Ni-TiO2/XC72R, and 80% Ni/XC72R are 97.4, 139.04,161.12, and 89.44 m2 g
–1
,
respectively. These values further confirm that the higher current density of the 60% NiTiO2 is a result of increased ECSA.
Furthermore, Fig. 3b demonstrates the electrode activation in the potential window
from zero to 0.9 V at various scan rates. At a potential sweep rate of 10 mV s
–1
two anodic
peaks can be visualized, which become less apparent as the scan rate increases. These
peaks demonstrate the distinct nickel phases.35 Bode et al. proposed a scheme that
involved two phases of nickel hydroxide, α– and β–Ni(OH)2, and the two phases of the
nickel oxyhydroxide species, γ– and β–NiOOH.[36] The β–Ni(OH)2 phase has a brucite
structure with short interlayer distances of 4.605 Å while the α–Ni(OH)2 phase is less
defined, highly hydrated, and has large interlayer distances (> 8 Å).37 Upon
electrochemical exposure in base, the β–Ni(OH)2 phase is oxidized to β–NiOOH while
retaining its highly dense, packed structure and the α–Ni(OH)2 is converted to the γ–
NiOOH phase while preserving its disordered and poorly defined configuration. Recent
studies, however, have demonstrated that a mixing of the two phases is possible with
both long α– and short β– interlayer distances coexisting in the material. These materials
were first proposed as “badly crystallized β” or βbc.
38 More extensive studies are
necessary to decouple the αII/γIII and βII/βIII phases given that their presence is
nonstoichiometric. Nonetheless, it can be seen from the inset in Fig. 3b, plotting the
maximum peak current (ipa) vs. square root of the voltage scan rate (v
1/2) gives a linear
13
relationship. This suggests that the oxidation of Ni(OH)2 to the NiOOH is a diffusion limited
process, which agrees with previous reports.39 Furthermore, the short-term stability of the
Ni-TiO2/XC72R catalysts at a potential of 0.7 V in 0.1 M NaOH is shown in Fig. S8. All
catalysts demonstrate no loss in current in the tested time period, indicating the overall
catalyst stability under the alkaline conditions.
Glucose oxidation.– Figure 4 demonstrates the last activation scan for each electrode
in 0.1 M NaOH at a scan rate of 20 mV s
–1 and the subsequent addition of 1 mM glucose.
The anodic peaks shift to positive values and are enhanced upon the addition of glucose,
indicating an interaction between glucose and the catalyst. Even though all catalysts
exhibited electrooxidation of glucose from 0.50 to 0.75 V, the 60% Ni-TiO2/XC72R
Fig. 4 CVs of the Ni-TiO2/XC72R electrodes a) 20%, b) 40%, c) 60%, and d) 80% at a scan rate
of 20 mV s
–1
in 0.1 mol L–1 NaOH (solid) + 1 mmol L–1 glucose (dotted).
a b
c d
14
catalyst produced the highest activity with a maximum current density of 9.80 mA cm–2
.
This current density was higher due to the greater wt. % of nickel when compared to the
20% and 40% Ni-TiO2/XC72R nanocomposites. Additionally, when compared to the 80%
Ni/XC72R, the incorporation of titania anatase in the 60% Ni-TiO2/XC72R plays a
significant role. First, it demonstrated that more Ni(0) species are present at the electrode
surface, which can spontaneously transform to Ni(III) in alkaline media. Second, it yielded
the largest ECSA at 161.12 m2 g
–1
, which would allow for more reactivity towards the
analyte of interest.
Chronoamperometric glucose sensing.– Given the improved catalytic activity of the
60% Ni-TiO2/XC72R catalyst, CA experiments in glucose were conducted at different
potentials; shown in Fig. 5 with their corresponding calibration curves. Aliquots of 50 μM
glucose were added into the 0.1 M NaOH electrolyte every 50 s while the RDE was
rotating at 4200 rpm. These aliquot concentrations were selected because the
nanocomposite would be applied as the noninvasive alternative for glucose monitoring.
This is because statistically significant correlation can be found between blood glucose
levels and salivary glucose levels. In a previous study, it was found that patients with
blood glucose levels between 100 and 280 mg dL–1
(5.6 mM and 15.5 mM) had a mean
salivary glucose level of 1.002 mg dL–1
(0.056 mM) and those with blood glucose levels
between 180 and 440 mg dL–1
(10 mM and 24 mM) had a mean salivary glucose level of
2.31 mg dL–1
(0.13 mM).40 Therefore, the selected aliquots of 50 μM adequately covered
the salivary glucose level range for a noninvasive approach of the sensor. It is evident
from Fig. S5 that increasing the potential leads to an increase in current. This enforces
the idea that the applied potential directly affects the amperometry response of the
15
enzyme-less biosensor. At a potential of 0.5 V, a linear range exists from 50-250 μM (R2
= 0.98) with a sensitivity of 900 μA mM–1 cm–2
. When increasing the potential to +0.6 V
as seen in Figure 5c, the linear range expanded significantly to 750 µM (R2 = 0.99) and
the sensitivity of the electrode increased to 1900 μA mM–1 cm–2
.
Fig. 5 Amperometry responses at potentials of a) 0.5 V, c) 0.6 V, and e) 0.7 V with their
corresponding calibration curves (b, d, and f, respectively) for the 60% Ni-TiO2/XC72R electrode
in 0.1 mol L–1 NaOH with 50 μmol L–1 glucose aliquot addition every 50 s.
a b
c d
e f
+0.5V
+0.6V
+0.7V
16
At the highest measured potential of 0.7 V, the sensitivity of the electrode plot was
enhanced to 3600 μA mM–1 cm–2
, and the concentration range became linear from 0.05-
1 mM (R2 = 0.99). The noise level at 0.7 V also increases at higher glucose concentrations
compared to 0.6 V, which may be associated with more intermediate species adsorbed
onto the catalyst surface.23 Given that the amperometry responses showed phenomenal
sensitivity and linearity at low glucose concentrations, higher concentrations of 1 mM
were used to replicate a concentration range analogous to human physiology. The Center
for Disease Control and Prevention states that fasting blood sugar levels for a healthy
individual are below 140 mg dL–1
(< 8 mM), prediabetics are between 140-199 mg/dL (8-
11 mM), and diabetics are anywhere above 200 mg dL–1
(> 11.1 mM). Figure 6 displays
the electrooxidation of glucose employing 1 mM (black line) and 50 μM (red line) aliquots
per step. As expected, higher currents are produced when increasing the aliquot
concentrations from 50 μM to 1 mM. The calibration curve for the 1 mM aliquots, shown
in Fig. 6b, generated a correlation of R2 = 0.99, sensitivity of 273.7 μA mM–1 cm–2
, and
detection limit of 3.13 µM (S/N = 3) while that of the 50 μM aliquot additions, shown in
Figure 6c, produced a regression equation with correlation of R2 = 0.99, sensitivity of 3300
μA mM–1 cm–2 and detection limit of 144 nM (S/N = 3) at an applied potential of 0.7 V.
This demonstrates that the 60% Ni-TiO2/XC72R electrode can effectively cover the
glucose range found in human saliva (50-1000 μM) and those in human blood (1-20 mM).
Interference studies.– Prior studies have demonstrated that the presence of chloride
ions can interfere with the performance of non-enzymatic sensors.19,20,41 Thus,
amperometry studies were conducted using 0.1 M NaOH as the electrolyte while
incorporating 0.1 M NaCl. Figure S9 displays the amperometric responses for the 60%
17
Ni-TiO2/XC72R electrode at 0.7 V using glucose aliquots of 1 mM per step. Contrary to
previous studies where a larger ionic activity would lead to the generation of higher
currents, a minimal difference in current response was observed. Even though the
sensitivity of the sensor increased to 287.7 μA mM–1 cm–2
, there is not a significant
difference between the electrolyte containing 0.1 M NaCl and that just containing 0.1 M
NaOH. From these almost constant responses it can be concluded that the electrode is
still active for glucose oxidation regardless of chloride ion presence. The presence of
other analytes including uric acid (UA), lactic acid (LA), and ascorbic acid (AA) have been
shown to directly interfere with the electrochemical oxidation of glucose, especially with
Fig. 6 a) CA responses for 60% Ni-TiO2/XC72R electrode at 0.7 V in 0.1 M NaOH using 1 mM
(black line) and 50 μM (red line) aliquots per step and the (b and c, respectively) calibration curves
with error bars (n = 3), d) Three independent amperometry measurements for 60% NiTiO2/XC72R electrode at 0.7 V in 0.1 M NaOH using sequential aliquot additions of 1 mM glucose
and 0.1 mmol L–1 of ascorbic acid, uric acid, and lactic acid.
a b
c d
18
nonenzymatic sensors.19 Given that these compounds coexist with glucose in real blood
samples, studies were performed to understand the selectivity of the new sensor towards
glucose in the presence of these analytes. Figure 6 shows the amperometry responses
of the 60% Ni-TiO2/XC72R electrode at 0.7 V in 0.1 M NaOH using sequential aliquots of
1 mM glucose, 0.1 mM AA, 0.1 mM UA, and 0.1 mM LA. It can be observed that these
compounds do not produce a significant current response at this potential and thus, do
not interfere with the electrochemical oxidation of glucose. The selectivity of the 60% NiTiO2/XC72R in 0.1 M NaOH at 0.7 V is exclusive to glucose.
Conclusion
This presented report showcased a convenient synthesis of a nickel-titania on
carbon nanocomposite system for the amperometry electrooxidation of glucose. Varying
wt. % of nickel precursor were reduced to Ni nanoparticles on a TiO2 anatase/Vulcan
carbon substrate by reduction with 35% hydrazine. The 60% (3:1) Ni-TiO2/XC72R
nanocomposite was the most active towards glucose oxidation and thus studied as an
enzyme-less, chronoamperometric glucose sensor. Using 50 µM aliquots of glucose in
0.1 M NaOH alkaline medium at a potential of 0.7 V (vs MMO), the sensor responded
rapidly (< 3 s) providing a sensitivity of 3300 µA mM–1 cm–2
(n = 3) and detection limits of
144 nM (S/N = 3). Using aliquot concentrations of 1 mM glucose to mimic physiological
blood conditions, the sensor continued to respond rapidly (< 1 s) showing a sensitivity of
273.7 µA mM–1 cm–2
(n = 3) and detection limits of 3.13 µM (S/N = 3). The nanocomposite
is stable, has an anti-poisoning capacity by chloride ions, and is selective towards glucose
amongst under analytes including 0.1 mM ascorbic acid, 0.1 mM uric acid, and 0.1 mM
lactic acid. Results demonstrate that the 60% Ni-TiO2/XC72R nanocomposite is an
19
incredibly effective catalyst for glucose electrooxidation and a promising candidate for an
enzyme-less sensor for glucose detection and monitoring.
20
Chapter 2
NiSn Alloys as Electrocatalysts for the Urea Oxidation Reaction and
Integration into Direct Urea Fuel Cells
Abstract
The use of the Ni electrode for the urea oxidation reaction (UOR) in alkaline media has
been reported as the most commercially viable option for wastewater treatment to
electrical energy. However, due to the sluggish kinetics involved in the 6-electron
oxidation process and the passivation of the electrode caused by reaction intermediates,
ongoing research has focused on the development of different Ni electrocatalysts to
address these specific challenges. Herein, we present a chemical reduction method for
producing Ni nanoparticles, leading to the synthesis of NixSny/C alloys with varying
metallic weight percentages. The compositions of the powders were analyzed through
XRD and XRF. RDE electrochemical studies in 1 M KOH revealed an enhanced activation
to the Ni(III) active species for the Ni60Sn20/C electrocatalyst, with calculated ECSA of
18.68 cm2 mg–1 and a proton diffusion coefficient of 2.27 10–13 cm2 s
–1
. In 0.33 M urea,
the electrocatalyst generated the highest current density of 60.35 mA cm–2 at 0.65 V (vs
MMO). Tafel analysis, stability measurements, and EIS results further elucidated the
optimized electrocatalytic activity of Ni60Sn20/C compared to the other alloys and the
monometallic materials. The Ni60Sn20/C catalyst was then tested in a membraneelectrode assembly (MEA) configuration for the direct urea fuel cell (DUFC). At room
temperature, the DUFC generated current and power densities of 10 mA cm–2 and 1.25
mW cm–2 near 0.4 V, respectively. The DUFC demonstrated enhanced electrochemical
rates with increasing temperatures and stability for up to 24 h.
21
Introduction
Urea, a final biological product excreted at 2-2.5 wt % in 1.5 L urine per day in
adults (11 kg urea per year), is energetically equivalent to 18 kg of liquid 42,43. The higher
theoretical energy density of urea (16.9 MJ L
–1
) compared to compressed (5.6 MJ L
–1
)
and liquid (10.1 MJ L
–1
) hydrogen, high solubility and low volatility, and easy transport
and storage, makes it an attractive candidate for a variety of applications.44 The
electrochemical oxidation of urea can obtain a significantly lower theoretical potential than
that of the hydrogen-oxygen electrolytic cell, 0.072 V compared to 1.229 V, respectively.
This can reduce the price of hydrogen energy by more than the 36% (from 4.13 to 2.63
USD per kg) previously reported.
42 The urea oxidation reaction (UOR) may reduce the
energy invested in water electrolysis by up to 93 percent.45 Laterally, fuel cell technology
continues harnessing major research attention in the quest for renewable energy
sources.46 The conversion of chemical energy to electrical energy in the direct urea fuel
cell (DUFC) has the potential to supplement traditional fossil fuel-based power sources,
turning waste into power.47 Urea offers the dual benefit of wastewater treatment and
electricity generation, presenting economic advantages to both the environmental and
energy sectors.48–50
Non-precious Ni-based materials have established their superiority for the urea
oxidation reaction (UOR), particularly in alkaline media.51–53 The UOR mechanism in
alkaline media can be described by the following reactions:
Anode
CO(NH2)2 + 6OH–
→ CO2 + N2 + 5H2O + 6e–
E0 = –0.756 V (vs SHE)
54
22
Cathode
3
2
O2 + 3H2O + 6e–
→ 6OH–
E0 = 0.4 V (vs SHE)
Cell
CO(NH2)2 +
3
2
O2 → CO2 + N2 + 5H2O E0 = 1.156 V (vs SHE)
Hydrogen peroxide can replace humidified oxygen as oxidant in both acidic and basic
conditions.55 The oxidation of H2O2 has a higher standard reduction potential compared
to that oxygen under both pH extremes. In basic pH, it is 0.878 vs. 0.401 V, respectively,
and 1.776 vs 1.229 V in acidic medium. This gives a theoretical cell potential of 2.51 V
and 1.15 V, respectively.
42 A major limitation in both urea electrolysis and the DUFC under
different conditions, however, lies in the sluggish kinetics of the six-electron transfer
process required for UOR at the anode.43 As a result, research has largely focused on
developing high-performance Ni electrocatalysts to reduce the high overpotentials and
achieve practical efficiency.
Another challenge in the electrooxidation of urea is the degradation of the catalyst
by carbon monoxide, which has led to the incorporation of other CO-tolerant metallic
species in catalyst design.56–58 Specifically, Ni–Sn bimetallic electrocatalysts continue
harnessing attention due to the cost-effectiveness of tin and its reported activity for CO
oxidation and desorption.
59–61 Most reports, however, present elaborate synthetic
methodologies with energy-intensive conditions and are often limited to half-cell
studies.62,63 Herein, we report on the one-pot synthesis of NiSn alloy powders for UOR in
alkaline media. Electrocatalysts with varying wt. % of Ni and Sn were characterized
through XRD and XRF. These were then studied using CV and electrochemical
23
impedance spectroscopy (EIS) for enhanced urea oxidation. The optimized catalyst was
then integrated as an anode electrode in a DUFC.
Materials and Methods
Chemicals.– Ni(NO3)2 • 6H2O (99.9%) was purchased from Strem Chemicals, tin powder
(~100 mesh, 99.5%) from Alfa Aesar, Vulcan carbon XC-72R and 5% Pt/ C from Fuel Cell
Store, and Liquion™ LQ-1105 5% Nafion binder solution from Ion Power. All other
reagents were of analytical grade and used without further purification. Millipore high
purity water (18.2MΩ cm) was used for all preparations.
Synthesis of NiSn/C catalysts and ink preparation.– For the synthesis of the
Ni60Sn20/C catalyst, 297 mg of Ni(NO3)2 • 6H2O, 20 mg of tin powder, and 20 mg of Vulcan
carbon were stirred in 20 mL ethylene glycol until dissolution. Then, 936 µL of hydrazine
solution (35% wt. % in H2O) were slowly added into the mixture and stirred for 15 minutes.
The solution was transferred to a Teflon lined vessel, sealed, and subjected to 150°C in
oven for 6 h. After heating, the catalyst solution was washed with 30 mL H2O and
centrifuged at 4200 rpm for 5 min. The supernatant was discarded, and the washing
process was repeated two additional times to remove unreacted materials. The catalyst
was dried at 100°C for 24 h, resulting in a blue-gray powder. The same synthetic
methodology was followed for isolation of the Ni40Sn40/C, Ni20Sn60/C, Ni80/C, and Sn80/C
materials but varying the wt. % of metal precursors. Catalyst inks were prepared by
dissolving 2.5 mg of powder in 100 µL isopropanol/900 µL H2O mixture followed by
addition of 7.5 mg (~10 µL) 5% wt. Nafion binder. The catalyst inks were then sonicated
by probe for 3 min. Twenty µL of the catalyst ink were dropcast on the GC RDE and dried
at 60°C for 30 min.
24
Membrane-electrode assembly (MEA) fabrication.– Cathode and anode inks were
prepared by dissolving 50 mg of Pt/C and Ni60Sn20/C catalyst powder in 200 µL
isopropanol/500 µL H2O. To each mixture, 300 µL of 5% Nafion binder dispersion was
added. The cathode and anode inks were then sonicated for three cycles of 480 s each
using a benchtop sonicator, followed by three cycles of 60 s using a probe sonicator. The
NiSn/C anode and Pt/C cathode inks were painted onto a Toray Carbon paper electrode
(E-TEK, TGH-060, 10% wet proofing, Fuel Cell Store) and Toray Carbon with a
microporous layer (Sigracet 22BB, Fuel Cell Store), respectively, with a surface area of 4
cm2
. The electrodes were dried in an oven at 60°C overnight. Both Nafion 117 PEM and
meta-terphenyl-fluoro-alkylene trimethylammonium (mTPN1-TMA) anion exchange
membranes (AEM) were soaked in 1 M KOH for 4 h to exchange the Br–
into the hydroxide
ions and then rinsed with DI H2O. The MEA was prepared by sandwiching the mTPN1-
TMA membrane between the cathode and anode, followed by hot pressing at 140°C for
5 min at 500 lbs.
Half-cell electrochemical testing.– Voltammetry, constant potential, and impedance
measurements were conducted using a Solartron Analytical 1287A Electrochemical
Interface potentiostat/galvanostat instrument. Studies were performed in a threeelectrode configuration half-cell using a glassy carbon (GC) rotating ring disk electrode
(RDE; Pine Research MSR Rotator) with geometric surface area of 0.195 cm2 as the
working electrode. GC RDE was polished with 0.05 µm alumina slurry (Electron
Microscopy Sciences) and microcloth prior to its use. A mercury/mercury oxide (MMO) in
4.24 M KOH and a Pt wire were used as the reference and counter electrode,
respectively. The 1 M KOH electrolyte was flushed with N2 gas for 30 min prior to testing
25
and steadily flowed during all electrochemical measurements. Potential conversion from
the MMO reference electrode to the reversible hydrogen electrode (RHE) followed
equation:
E (vs RHE) = E (vs Hg/HgO) + 0.098 V + 0.0591 V × pH
Direct urea fuel cell measurements.– The fuel cell polarization measurements were
performed using a Fuel Cell Test System 890B (Scribner Associated). Anolyte and
catholyte solutions constituting of 0.33 M urea in 1 M KOH and 2 M H2O2 in 2 M H2SO4,
respectively, were steadily flowed into the fuel cell at an optimized flow rate of 580 mL
min–1
.
Physical characterization.– A benchtop Rigaku Miniflex 600 with Cu Kα radiation source
was used for powder X-ray diffraction (XRD) measurements at a scan rate of 4° min–1
from 2θ of 10° to 90°. The wt. % relation in the electrocatalysts was further analyzed using
a Bruker S8 Tiger Wavelength Dispersive X-Ray Fluorescence (XRF) Spectrometer.
Results and Discussion
Physical characterization.– XRD measurements were conducted to identify the
composition of the isolated NiSn/C powders, as illustrated in Fig. 7. The black curve
represents the diffraction pattern of Ni80/C, showing peaks at 2θ values of 44.84°, 52.17°,
and 76.67°, which correspond to the (1 1 1), (2 0 0), and (2 2 0) planes of nickel
nanoparticles, respectively. As anticipated, the Sn80/C powder did not exhibit the same
reflections as the Ni NPs material. Instead, Sn80/C displayed reflections at 30.65°, 32.03°,
43.89°, 44.92°, 55.36°, 62.55°, 63.84°, 64.61°, 66.97°, 72.44°, 73.19°, and 79.53°. As
shown in Fig. 7, the intensity of these reflections diminished as the wt % of Sn decreased
in the samples. Due to the overlap of the Ni (1 1 1) phase to the Sn (2 2 0) and (2 1 1)
26
phases at 43.89° and 44.92°, respectively, the crystallite size in the alloys was calculated
using Scherrer’s equation based on the Ni (2 0 0) plane. The crystallite sizes for
Ni20Sn60/C, Ni40Sn40/C, Ni60Sn20/C, and Ni80/C were 11.8, 9.5, 11.8, and 12.2 nm,
respectively. The metallic content (%) was calculated through the reference intensity ratio
(RIR) method using the (2 0 0) planes for both Ni and Sn. Table S8 shows the calculated
content of tin and nickel in the Ni20Sn60/C, Ni40Sn40/C, and Ni60Sn20/C alloys, and the trend
aligns with the inteded synthesis goals. X-ray fluorescence (XRF) was then employed for
the characterization and composition of the metalllic constituents in the isolated powders.
Figure S10-S14 shows the spectra for the Ni20Sn60/C, Ni40Sn40/C, Ni60Sn20/C, Ni80/C, and
Sn80/C materials, along with the respective tabulated elemental compositions (Tables S2-
S7). Given the low energy of the characteristic X-rays emitted by carbon, its detection via
Fig. 7 XRD patterns of NiSn/C alloys with varying compositions (b;
Ni20Sn60/C, c; Ni40Sn40/C, d; Ni60Sn40/C) and the monometallic
analogues (a; Sn80/C, e; Ni80/C). Peaks at 44.84°, 52.17°, and 76.67°
correspond to the (1 1 1), (2 0 0), and (2 2 0) planes of nickel
nanoparticles, respectively.
a
b
c
d
e
27
XRF was paticularly challenging, resulting in an apparent increase in the concentration
percentages of the detecable species Ni and Sn. Nonetheless, contamination by other
species was neglible, which helped demonstrate synthetic reliability and practicallity. This
was also observed for the isolated NiSn alloys, showing purities of over 95%. Additionally,
the wt. % trend via XRF also aligned with the synthesis for these materials, as presented
in Table S8.
Catalyst activation.– We employed CV to study the Ni(OH)2/Ni(OOH) redox couple in
an alkaline electrolyte prior to the electrooxidation of urea. In basic media, Ni NPs
spontaneously convert to Ni(OH)2 and upon the application of a positive potential bias,
the electrocatalyst transitions to the NiOOH active species.
64 Figure 8a shows the
transition from Ni(II) to Ni(III) in 1 M KOH at a scan rate of 50 mV s
–1
for the Ni60Sn20/C
electrocatalyst. The physiochemical convertion to the oxyhydroxide phase begins at a
potential of 0.46 V (vs MMO), quickly reaching a maximum current density (jmax) at 0.53
V over 7 cycles of approximately 23 mA cm–2
. At E ≈ 0.7 V, the oxygen evolution reaction
(OER) starts. The electrochemical surface area (ECSA) of 9.97 cm2 mg–1 was calculated
by integrating the Ni transition in the 200 mV cathodic window (from 0.3 to 0.5 V) to obtain
the coulombic charge (q), dividing by the theoretical charge density of 420 C cm–2
, and
normalizing to electrocatalyst mass dropcast on the RDE.65 ECSA Values of 1.31, 6.16,
and 5.28 cm2 mg–1 were calculated for Ni20Sn60/C, Ni40Sn40/C, and Ni80/C, respectively.
As seen in Fig. S15 the CVs of all electrocatalysts in 1 M KOH demonstrated a
similar quasi-reversible behavior for the Ni(II)/Ni(III) redox couple. A major difference can
be appreciated for the Ni80/C material in which a second peak was observed at higher
oxidative potentials. According to Bode’s diagram, this reflects the different phases of
28
nickel due to the oxidative state change from Ni(II) to Ni(III). The ip at different scan rates
was plotted to understand the kinetics of the Ni(II)/Ni(III) redox couple in 1 M KOH, as
shown in Fig. 8b. The conversion from Ni(OH)2 to NiOOH was determined to be diffusionlimited given the linearity of the line of best fit for the Ni electrocatalysts (R2 > 0.99).
Additionally, the largest slope of 4.12 mA mV–1/2 s
1/2 for Ni60Sn20/C show faster charge
transfer kinetics at the electrode surface. Studies have indicated that proton diffusion is a
rate-determing step (RDS) in the overall process for Ni(OH)2 to NiOOH conversion in
alkaline media.66 Thus, we calculated the diffusion coefficients (D) at 50 mV s–1
through
the Randles-Sevcik equation:
𝑖𝑝 = 2.7 × 105𝑛
3/2𝐴𝐷
1/2𝐶𝑣
1/2
where 𝑖𝑝 is the maximum current density, 𝑛 is equal to the 1e–
for the Ni(II)/Ni(III)
transition, 𝐴 is the geometric surface area of the electrode, 𝐶 is the concentration of
limiting species, and 𝑣 is the scan rate.59 Proton concentration can be stoichiometrically
quantified by the molecular weight of Ni(OH)2 and its density of 3.97 g cm–3
. As reported
in Table S9, the larger DH+ for the Ni60Sn20/C electrocatalyst suggests a faster RDS
compared to various other electrodes.67
Urea oxidation reaction.– Given the faster kinetics for the formation of NiOOH and its
larger ECSA, we expected a larger current density generation by the Ni60Sn20/C
electrocatalyst compared to the other materials. Fig. 8c demonstrates the last activation
CV for Ni60Sn20/C in 1 M KOH and the oxidation of 0.33 M urea after several cycles (n =
25). With this electroactalyst, urea oxidation has an onset potential of 468 mV and rapidly
reaches a current density of 51.16 mA cm–2 at 0.65 V. This improvement in current density
was specifically attributed to the urea oxidation given that the onset of OER was observed
29
at potentials larger than 0.65 V. In comparison, Ni20Sn60/C, Ni40Sn40/C, and Ni80/C
generated current densities of 23.97, 38.92, and 39.94 mA cm–2
, respectively, for the
same scan rate. It must be noted that both jmax and jsp at 0.65 V were calculated for UOR
at n = 25. As seen in Fig. 8d, the current density gradually decreased from 65.03 mA
cm–2 at the first cycle to 51.16 mA cm–2 at the last. This gradual current drop was observed
for the Ni-based electrocatalysts, which indicate surface passivation right after urea
oxidation. There are losses at the beginning of urea exposure that eventually stabilize
after continuous oxidation. The maximum current (ip) at 0.65 V for each electrocatalyst
Ni60Sn20/C
Fig. 8 a) CV for the Ni60Sn20/C electrocatalyst at a scan rate of 50 mV s
–1
in 1 mol L–1 KOH after
15 cycles showing the 200 mV cathodic window for ECSA; b) peak current vs. square root of scan
rate plot for N60Sn20/C (green), N40Sn40/C (red), N20Sn60/C (blue), and Ni80/C (black) demonstrating
the slope and correlation for each sample; c) CVs for Ni60Sn20/C showing the last activation cycle
(solid) in 1 mol L–1 KOH and the oxidation of 0.33 mol L–1 urea (dashed) at a scan rate of 50 mV
s
–1
; d) oxidation of 0.33 mol L–1 urea from initial to final (n = 25) CV.
a
c d
b
30
was divided by the mass of Ni dropcast (0.05 mg * wt. % Ni) on the RDE to compare
specific mass activity. The specific mass activity (jsp) for Ni20Sn60/C, Ni40Sn40/C,
Ni60Sn20/C, and Ni80/C was calculated to be 479, 389, 341, and 215 mA mg–1
. The specific
mass activity as Sn content increases suggests that it may enhance the efficiency of the
catalyst in the reaction mechanism. Table S10 summarizes electrochemical performance
parameters for the different electrodes including the columbic charge used to calculate
ECSA and jsp.
As seen in the inset of Fig. 8c, upon the electrode’s exposure to urea, there was
a major decrease in jmax for the reduction of NiOOH to Ni(OH)2. This suggests that NiSn
alloys follow an indirect mechanism for the urea oxidation where catalytic active site
regeneration is necessary. In this process, urea molecules react with electrogenerated
NiOOH active sites by an EC mechanism, releasing N2 and CO2 as products, and
generate the Ni(OH)2 catalyst.42,68 The Ni(OH)2 species leads to irreversibility
pronounciation of the NiOOH/Ni(OH)2 wave. Nonetheless, sequenital oxidation by
hydroxide ions to NiOOH is possible given the positive potential bias. Thus, the calculated
high D for the Ni60Sn20/C electrocatalyst demonstrates the enhanced propensity for
regeneration of the active sites for further UOR.
69 It is noteworthy to mention that the
oxidation of Ni(II) to Ni(III) is the competing reaction to UOR, also contributing to current
generation.70 Interestingly, Ni80/C displayed distinct electrochemical behavior. Figure
S16 demonstrates the last activation cycle (1 M KOH) of the electrodes and their 25th
cycle after 0.33 M urea oxidation. The reversibility of the NiOOH/N(OH)2 wave for Ni80/C
suggests the more complex, direct mechanism for urea oxidation. Additionally, the
sudden current density jumps on the reverse scan (0.9 V to zero) near 0.81 V parallels
31
that of Pt/C for methanol oxidation in alkaline media. Theoretical studies have proposed
passivation of the NiOOH surface due to surface blockage caused by CO intermediates
generated during urea oxidation, suggesting that desorption of (
*COO) intermediates
might be distincty occuring at the surface of NiSn alloys.71,72
Tafel, Stability and EIS.– We then studied the electrochemical kinetics for the catalysts
for UOR via Tafel analysis to further understand the role of tin. Figure 9a demonstrates
the Tafel plots of the Ni-based electrocatalytic materials in 1 M KOH at a scan rate of 0.5
mV s–1
. As demonstrated, the Ni80/C electrocatalyst generated the largest tafel slope (b)
and smaller exchange current density (j0) compared to the NiSn alloys. We observed that
UOR kinetics by bimetallic electrodes became enhanced with increasing wt. % of Sn.
Tafel slopes (TS) of 67.3, 70.9, 77.8, and 90.02 mV dec–1 were calculated for Ni20Sn60/C,
Ni40Sn40/C, Ni60Sn20/C, and Ni80/C, respectively. These findings help demonstrate that
Sn-incorporating catalysts likely facilitate the desorption of byproducts and/or
intermediates given the lower TS. After normalization of the current to the respective
a b E = 0.6 V (vs MMO)
Fig. 9 a) Tafel plots for N40Sn40/C (red), N20Sn60/C (blue), N60Sn20/C (green), and Ni80/C (black)
normalized to the geometric area (0.195 cm2
) of the electrode and to the ECSA of each
electrocatalyst (inset) with 0.33 mol L–1 urea in 1 mol L–1KOH; b) Constant potential urea oxidation
at 0.6 V (vs MMO) for the electrocatalysts in alkaline media.
a
32
calculated ECSAs (as observed in the inset), the Ni20Sn60/C, Ni40Sn40/C, Ni60Sn20/C, and
Ni80/C yielded Tafel slopes of 69.5, 71.8, 81.3 and 94.7 mV dec–1
, respectively. The j0
values demonstrated more variability than those normalized to the geometric surface area
of the electrode, as comparable in Tables S11-S12. The highest j0 of 34.47 mA cm–2
ECSA
was calculated for the Ni60Sn20/C alloy. We thus determined that the Ni60Sn20/C alloy
offers an optimized balance between catalytic activity and the ability to manage
byproducts and intermediates at the surface.
Continually, we observed gas evolution management by the Sn-incorporating
electrocatalysts at the RDE through chronoamperometry. Figure 9b shows the constant
potential eletrolysis experiments at 0.6 V (vs MMO) for all electrodes in 0.33 M urea. As
expected, Sn80/C produced a neglible current density throughout the studied time window
while the Ni20Sn60/C, Ni40Sn40/C, Ni60Sn20/C, and Ni80/C electrodes generated current
densities of 6.22, 20.89, 27.57, and 19.57 mA cm–2
, respectively. The alloys
demonstrated a stable current response while that of the Ni80/C material had a different
output. The current reponse gradually stabilized to around 20 mA cm–2 but then
progressively fluctuated over an extended time frame. Suddenly, the signal restabilized
and the evolving noise cycle started once again. This was physically manifested by
bubble formation and adhesion on the RDE. Bubble formation began in less than five
minutes, continued expanding for about 700 s, and finally bursted to reveal a short-lived,
stable current response. Interestignly, this behavior was exclusive to the nonincorporating tin electrode.
We performed EIS measurements to visualize both the catalyst activation and
UOR processes by the five electrocatalysts from 450 to 550 mV in 0.33 M urea and 1 M
33
KOH electrolyte, as seen in Fig. 10. At 450 mV, two depressed semicircles can be
appreciated for the Ni60S260/C electrode, which demonstrates two processes occuring at
different time domains.
63 The semicircle in the high frequency region represents the
Ni(II)/Ni(III) redox transformation while that in the low frequency region was attributed to
the UOR.73 The diameter of the first semicircle (i.e., high frequency) closely represents
the charge transfer resistance (Rct) for the reaction of Ni(OH)2 to NiOOH, which was about
24 Ω. The semicircle for urea oxidation had Rct of 20 Ω. When increasing the potential to
0.50 V, as seen in Fig. S17, the conversion to the NiOOH active species can be observed
for all the electrocatalytic electrodes. At 550 mV, in the case of the Ni60Sn20/C electrode,
the Rct for both processes decreased equally to 6 Ω. These results indicate that the
current density output at this potential is evenly distributed between Ni(OH)2/NiOOH
conversion and UOR through Ohm’s Law. Additionally, the Ni80/C electrode had a more
Fig. 10 Nyquist plot for N20Sn60/C (blue), N40Sn40/C (red), N60Sn20/C (green), Ni80/C (black),and
Sn80/C (orange) electrodes with 0.33 mol L–1 urea in 1 mol L–1 KOH at a) 0.45 V and b) 0.55 V.
The data was recorded over a frequency range from 10 kHz to 100 mHz with perturbation
amplitude of …
a b
34
than two-fold Rct increase for NiOOH formation when compred to Ni60Sn20/C. The electrooxidation of urea at the Ni80/C electrode can also be seen at this potential bias.
Direct Urea Fuel Cell.– We proceeded by painting the Ni60Sn20/C ink onto carbon paper
(4 cm2
) to serve as the anode electrode for the DUFC. The cathode electrode consisted
of Pt/C and parted from the anode by a Nafion proton-exchange membrane (PEM) in a
zero-gap MEA. Fuel input was firstly optimized by varying the flow rate of both anolyte
and catholyte into the cell. As shown by the polarization curve in Fig. S18, increasing the
flow rate from 120 to 260 to 580 mL min–1
leads to enhanced electrochemical rates. The
DUFC was then tested at different temperatures at the optimized flow rate. Figure 11a
demonstrates the temperature-dependent polarization curves using 0.33 M urea in 1 M
KOH urea as the anolyte. At room temperature, a current and power density of 10 mA
cm–2 and 1.25 mW cm–2 were produced near 0.4 V, respectively. The elevated
temperatures led to improved electrochemical rates, thus increasing both current and
power outputs. At the same potential bias of 0.4 V and 80 C, the DUFC generated current
and power density of 26 mA cm–2 and 3.25 mW cm–2
, respectively, with a jmax close to 50
mA cm–2
. In an attempt to enhance the power density output, we sought to increase the
urea concentration in the anolyte to 1 M urea. However, as shown in Fig. 11b, there was
no major enhancement in electrochemical performance when the urea concentration was
increased. This elucidates that Ni electrocatalytic sites are saturated at 0.33 M urea.
Theoretically, a larger OH– concentration should increase the availability of NiOOH
species for additional UOR. Our hypothesis demonstrated to be true when increasing
electrolyte concentration from 1 M to 7 M KOH. As seen in Fig. 11d, the power density
incremented 3x-fold using 7 M KOH at all applied temperatures. Interestingly, even
35
though the current density output was larger at the higher OH– concentration, j becomes
limited by mass-transport losses.
The stability of the DUFC was then tested under a constant potential bias at
different time ranges. Under a potentiostatic hold of 0.2 V, the current density gradually
decreases for the first 4 h and eventually stabilizes at 1 mA cm–2
for the next 20 hours ,
as shown in Fig. S19. Continually, we tested the stability of the DUFC at a higher current
density output. The inset on Fig. 11c shows that the MEA can sustain a j of about 15 mA
cm–2 for 8 h when applying a constant potential bias of 10 mV at room temperature. After
8 h, the current density begins to gradually decrease, eventually reaching 12.5 mA cm–2
after 20 h. These studies demonstrate the robustness of the Ni60Sn20/C electrode for
integration in the DUFC. We subsequently tested the same electrode system using an
mTPN1-TMA AEM to compare its performance with the MEA utilizing Nafion 117.
74
Figure S20 demonstrates the polarization and power density curves for mTPN1-TMA at
varying temperatures of 25, 40, and 80 C. At room temperature and 10 mA cm–2 , the
power density output using AEM is higher (2 mW cm–2
) compared to Nafion PEM (1.5
mW cm–2
). This power density output with AEM increases at higher temperatures,
doubling at 40 C. Interestingly, current density differences are negligible despite there
being a significant decrease in activation losses (500 mV drop for PEM compared to 200
mV for AEM). This was attributed to faster charge transport by hydroxide ions, suggesting
that the system’s overall performance could be further optimized by leveraging the
enhanced ionic conductivity in alkaline environments for improved long-term stability and
efficiency.
36
Conclusion
Generating electrical energy from renewable sources is crucial for achieving a sustainable
infrastructure. The electrochemical oxidation of urea offers a promising avenue for this
mission by addressing both wastewater treatment and production of clean energy. The
non-precious, state-of-the-art electrocatalysts for the UOR rely on Ni in alkaline media.
a b
Fig. 11 a) Polarization curves employing N60Sn20/C anode and Pt/C cathode separated by a N117
PEM, b) using urea concentration of 1 M at increasing temperatures, c) CA measurements at 10
mV demonstrating the stability of the DUFC at 8 h (inset) and 20 h, and d) increasing anolyte
concentration to 7 M KOH at 25, 40, and 60 C.
c d
37
However, significant limitations remain due to slow reaction kinetics and electrode
poisoning by intermediates. Tin has been reported to mitigate these issues for both
alcohol and urea oxidation. In this study, we reported the chemical reduction of Ni NPs
for the facile synthesis of NiSn/C alloys with varying wt. % of metallic constituents. The
isolated powders were characterized through XRD, XRF, and ICP-OES to confirm their
composition. Among the various Ni-Sn compositions, the Ni60Sn20/C alloy demonstrated
improved activation for the NiOOH active layer, achieving a high ECSA of
18.68 cm2 mg–1 and a diffusion coeffcient of 4. 99 10–9 cm2 s
–1 in 1 M KOH. In the
presence of 0.33 M urea, the catalyst produced a maximum current density of 60.35 mA
cm–2 at 0.65 V (vs MMO). The enhanced electrochemical performance of Ni60Sn20/C is
attributed to the incorporation of tin, which facilitates the adsoprtion of hydroxyl ions and
desorption of byproducts and/or intermediates at the electrode surface. Enhanced
electrochemical kinetics, as determined by Tafel analysis and EIS studies, indicate that
Ni60Sn20/C represents the optimized composition for the UOR. The DUFC utilizing the
Ni60Sn20/C electrocatalyst achieves competitive electrochemical performance at 25 C,
which is further boosted at elevated temperatures.
38
39
Chapter 3
Part A: Recent Advancements in Silver-based Catalysts for
Electrochemical CO2 Reduction: Detailed Mechanisms, Surface and
Structural Features, and Practical Applications
Abstract
The electrochemical CO2 reduction reaction (ECO2RR) into its value-added products
represents a promising strategy to mitigate anthropogenic CO2 emissions. Research has
demonstrated that Ag-based materials are the leading candidates for carbon monoxide
formation due to their high faradic efficiency and current density at lower overpotentials.
Herein, we present an overview of the mechanistic insights in diverse molecular systems
but primarily emphasize those of Ag-based heterogeneous electrocatalysts. Multiple
examples of different nanostructures highlighting size, surface, composition, and other
morphological characteristics are discussed and their relation to electrocatalytic
performance (efficiency, selectivity, stability, etc.) evaluated and compared. Additionally,
we demonstrate recent cell design advances with the overall goal of showcasing the nearfuture, practical applications of this technology. The review acknowledges that there are
still current challenges and valuable opportunities in the field of ECO2RR including: (1)
developing industrially viable methodologies for the synthesis of electrocatalysts, (2)
understanding complex mechanistic pathways at the electrode-electrolyte interface via in
situ techniques, and (3) optimizing membrane-electrode assembly fabrication, electrolytic
conditions, and cell stack design.
40
Introduction
The greenhouse gas effect refers to the entrapment of solar energy due to carbon
emissions in the atmosphere. Scientists agree that these carbon-based emissions
correlate to a prominent increase in global temperatures, a phenomenon termed as global
warming. Global warming has led to critical environmental issues including the melting of
ice caps and rise in sea levels, amongst other severe weather patterns. Specifically, it
has been shown that global CO2 concentrations have steadily increased since the
Industrial Revolution.75,76 These anthropogenic CO2 emissions are primarily attributed to
the combustion of fossil fuels as a source for energy. Moreover, energy demands
continue to rise and the greenhouse effect, global warming, and concerning
environmental issues have become more apparent.77 The need for cost and energy
efficient technologies that can capture and utilize CO2 to reduce its global, atmospheric
concentration is therefore of utmost importance.
Multiple approaches have been studied to mitigate the continuous rise of CO2
levels in the atmosphere. One of the long-considered strategies to reduce atmospheric
CO2 is carbon capture and sequestration (CCS). Presently, most of the industrially-viable
CCS technologies are employed at the point source (i.e., power plants and oil refineries)
using chemical absorption and desorption in amine-based solutions through temperature
and/or pressure manipulations.78 These amine-based technologies, however, have
presented disadvantages for their application at diluted point sources; the lower CO2
concentration decreases the efficacy of the systems and thus, the search for improved
CCS technologies continues. This includes carbon dioxide capture through solid amine
sorbents which have the potential to keep some of the advantages shown by aqueous
41
amines while minimizing the disadvantages. However, this area of CCS through solid
amine sorbents is still in its infancy and substantial progress must be demonstrated before
its mainstream industrialization.79 A second strategy for reduction of CO2 levels is to
capture and utilize it as a precursor for value-added products in a process called carbon
capture and utilization (CCU). There are several approaches for carbon utilization
including electrochemical, photochemical, and thermochemical strategies.77
Photochemical and thermochemical methodologies, however, have shown intrinsic
limitations that are difficult to evade. Photochemical systems show little selectivity and
low production rates while thermochemical systems require energy intensive conditions
(i.e., high temperatures and pressures) and the use of H2 as a reducing agent.80–83
Electrochemical approaches have been preferred given that reactions are typically
conducted at ambient conditions, can be highly selective towards a particular product,
and can be easily controlled by tuning the overpotential. Additionally, the post-separation
costs in electrochemical devices are greatly minimized given the natural separation of the
anodic and cathodic chambers. The electrochemical CO2 reduction reaction (ECO2RR)
has the potential to complete the anthropogenic carbon cycle by recycling CO2 from
various sources into feedstock materials for fuels and chemical manufacturing.84
The Electrochemical CO2 Reduction Reaction
The ECO2RR continues being a promising candidate for remission of
anthropogenic CO2. Unfortunately, linear CO2 is stable and chemically inert with a low
electron affinity and a large energy gap (13.7 eV) between its lowest unoccupied
molecular orbital (LUMO) and highest occupied molecular orbital (HOMO). Carbon
dioxide can be subject to chemical transformation at the electrophilic carbon but a
42
substantial input of energy is required; the bond energy for C=O dissociation is ~750
kJ/mol.85 For its electrochemical reduction into its value-added products, potentials more
negative than the respective thermodynamic values are required.84 Table 2 demonstrates
the equilibrium electrode potentials (vs SHE) for the half-reactions in aqueous solution
(pH = 7) for CO2 reduction products.86 As shown, the multi proton-coupled electron
reactions require a less negative potential compared to the other reduction processes,
indicating more thermodynamic favorability. Deviation occurs at the one electron where a
high thermodynamic energy of 𝐸
0 = −1.90 𝑉 reduction is required for CO2 to CO2
●–
conversion. This high thermodynamic potential is attributed to the structural differences
between linear CO2 and bent CO2
●–
. This step makes the ECO2RR an energy intense
process with sluggish kinetics in which high overpotentials are often necessary and thus,
electrocatalysts are employed. An ideal catalyst will (amongst other features) enable
electrochemical reductions to proceed with high activity and selectivity by minimizing the
overpotential.
In the context of the ECO2RR, electrocatalysts can be classified as heterogeneous
and homogeneous catalysts depending on their physical state and distribution in the
reaction mixture. Heterogeneous electrocatalysts exist in a different physical state from
the reactants. Firstly, tethering electrocatalytically active molecules to conductive
surfaces combines the advantages of highly efficient conversion of electrical energy to
chemical energy. This leads to enhanced electron mobility and conductivity. Secondly,
mass transport limitations become negligible given the direct contact and interaction of
the catalyst with the support materials and therefore, no bulk solution of the catalyst in
the electrolyte exists. Thirdly, recycling processes become more efficient and cost
43
effective because of the catalyst’s direct adherence to the electrode.87 The literature on
heterogeneous catalytic materials for ECO2RR is extensive: bulk metals88, nanoparticles
(NPs)89–91, metal-alloys92–94, metal oxides95–97, metal-organic frameworks (MOFs)98–100
,
and single-atom catalysts (SACs)101–103 have all been reported as electrocatalysts for the
reduction of CO2. The reader is encouraged to explore some of the most recent reviews
for heterogeneous electrocatalysts.87,104,105
Conversely, homogeneous electrocatalysts exist in the same physical state as the
reactants. These often include molecular catalysts such as transition metal complexes
and other organometallic compounds. Briefly, in the homogeneous electrochemical
catalytic system, the molecular catalyst acts as an electron shuttle for indirect electrolysis
of the analyte.106 First, a more highly reduced state of the complex is often produced after
accepting electrons from an electrode. Second, this reduced state of the molecular
catalyst donates electrons to the targeted analyte in solution. This returns the catalyst to
its initial state and allows the electrochemical transformation of the analyte. Under these
Half-cell Reaction E
o
(V)
CO2 + e– ⟶ CO2
•– –1.90
CO2 + 2H+ + 2e– ⟶ HCOOH –0.61
CO2 + 2H+ + 2e– ⟶ CO + H2O –0.53
CO2 + 4H+ + 4e– ⟶ C + H2O –0.20
CO2 + 4H+ + 4e– ⟶ HCHO + H2O –0.48
CO2 + 6H+ + 6e– ⟶ CH3OH + H2O –0.38
CO2 + 8H+ + 8e– ⟶ CH4 + 2H2O –0.24
Table 2 Equilibrium potentials (vs SHE) for the electrochemical CO2
reduction to different products in pH 7 aqueous solution.
44
conditions, detailed kinetic and mechanistic insights of the electrochemical reactions can
be studied.106–109 Given the synthetic control over electronic and steric properties in the
vicinity of the catalytic center, structure-function relations can be understood. This
approach permits the rational design of new catalyst structures with improved
performances.110 Therefore, a diverse library of organic frameworks, metal centers, and
experimental conditions have been developed and their respective faradaic efficiency and
selectivity towards different reduction products have been reported. The most promising
metal complex catalysts for ECO2RR are those based on earth-abundant metals. These
transition metal complexes have been carefully reviewed with descriptions of their
principal frameworks.111–113 The application of molecular catalysts under homogeneous
electrocatalytic conditions, however, still present several limitations. First, molecular
complexes have poor electrical conductivity and charge mobility. Their interactions with
electrodes are weak, which lead to mediocre electron transfer efficiencies. Second, mass
transfer resistance is high given that only molecules at the electrode surface can be
oxidized and reduced for successive catalytic processes. Most molecules are present in
solution and do not participate in the electrocatalytic process; a large quantity of molecular
electrocatalyst must often be used. Third, at industrial scales, major difficulties are
confronted for the recovery, separation, and reusability of the catalysts at the end of their
respective cycles.114 Finally, advanced synthetic techniques are required for the
construction of the frameworks and isolation of the overall transition metal complexes.
The differences between homogeneous and heterogeneous electrocatalytic
systems are numerous, including experimental conditions and cell design. Nonetheless,
the homogeneous approach towards ECO2RR was previously presented to showcase the
45
fundamental necessity to study and understand these electron vehicles. Over the last
decade, research endeavors for homogeneous electroreduction of CO2 have primarily
focused on the synthetic design and the structure-activity relationships of transition metal
complexes to enhance selectivity, efficiency, and performance. Presently, the field of
ECO2RR has moved towards bridging the intrinsically distinct homogeneous and
heterogeneous approaches. Through immobilization or adherence of molecular catalysts
onto conductive substrates and/or electrodes, the advantages from both approaches can
be exploited. However, this is not a trivial process. First, molecular catalysts must be
uniformly dispersed onto the substrate to prevent deactivation through aggregation or
dimerization from neighboring molecules.115 Second, multiple heterogenization
methodologies must be considered. These include the three principal reported
immobilization methods of non-covalent (i.e., electrostatic and π–π interactions), covalent
(i.e., chemical and electrochemical reactions), and periodic interactions (i.e., MOFs and
porous organic polymers), which all have their respective constraints. Third, the
conductive support or substrate can influence the rate and mechanism of CO2 conversion.
Finally, the electrolyte selection after heterogenization (going from organic to aqueous
phase) and long-term stability of the molecular catalyst are other limitations that must be
carefully addressed.110,111,115,116
Mechanistic Insights of the ECO2RR
The solid-state catalysts play a role in the electrochemical reduction of CO2, where
interactions occur among the adsorbed CO2 molecules, electrons, and protons. Various
experimental parameters, including cathodic potential, electrolyte salts, CO2 pressure,
and catalyst nature (i.e., surface morphology), can influence the reaction pathways and
46
product formation.117 Different product distributions result from the distinct reaction
pathways or combinations thereof.118 As mentioned earlier, electrocatalytic materials are
necessary to bind CO2 selectively and thus reduce the high overpotentials that the
reduction entails. Efforts have principally focused on addressing this issue by studying
different metal catalysts and the various products that they respectively form. Three major
categorizations of electrocatalysts for CO2 reduction are generally established in the
literature: metals that mainly form (1) formate, HCOO–
(i.e., Pb, Hg, Bi, In, Cd, Tl, and
Sn), (2) carbon monoxide, CO (i.e., Au, Ag, Zn, Pd, Ga), and (3) hydrocarbons and
alcohols (i.e., Cu). This review will mainly discuss electrocatalytic materials and
mechanisms that selectively yield 2e– products. Ample literature precedence exists for
copper-based electrocatalytic materials, which is encouraged for further understanding of
C2 and C2+ products.119–121
As explained earlier, compared to proton-coupled electron transfer (PCET) steps
in ECO2RR, the direct electron transfer step to generate CO2
•−
is much more
thermodynamically demanding. This is due to the energy required for reorganization of a
stable, linear molecule to its radical, bent configuration.122 However, even the earliest
reported mechanism for CO2 reduction over metal electrodes proposed the formation of
the CO2
•−
intermediate at the catalyst surface. Hori et al. proposed that the coordination
of CO2
•−
to transition metal atoms occurs through electron transfer of CO2 to the
unoccupied orbitals of the metal. Molecular orbital studies demonstrated that the CO2
bond becomes mostly stabilized by backbonding donation from the transition metal’s
highest occupied molecular (d) orbitals to the antibonding (
*
) lowest molecular orbital of
CO2.
88 This bonding interaction between metal center and CO2 can occur in three modes:
47
1) side-on coordination, 2) carbon coordination, or 3) end-on coordination. Pacansky et
al. studied the molecular orbital energies and atomic population analysis of the CO2
•−
intermediate and concluded that the highest occupied orbital is localized at the C atom at
84%. These results suggested the nucleophilic nature of CO2
•−
to bond with the metal
atom through carbon coordination.123 The conclusions by Hori et al. and Pacansky et al.
were experimentally supported by other reports. Sakaki used a molecular orbital
approach to describe NiIF(NH3)4(CO2), [NiI
(NH3)4(CO2)]+
, and [NiIIF(NH3)4(CO2)]+ as
model complexes of intermediate species in ECO2RR mediated by Ni(cyclam)Cl2. He
described favorable bonding to the transition metal through carbon coordination and
showed that the complex was stabilized by strong charge transfer through backdonation
from Ni to carbon dioxide.124 These results also supported the reaction mechanism
described by Sauvage et al. for CO production by a Ni cyclam electrocatalyst.125
In the electrochemical approach, CO2 is adsorbed at an active site (*) in the catalyst
surface to generate the *CO2
•−
reduced species. A proton transfer (most likely from the
bicarbonate electrolyte) then forms either the carboxyl intermediate (*COOH) or formate
intermediate (*OCHO) adsorbed at the active site.126,127 Even though the first electron
step is thermodynamically unfavorable, the formation of subsequent intermediates is
more kinetically challenging, indicating a shift in the rate-determining step from electron
transfer to protonation.128 In a 2008 report, bonding of CO2
•− at the Ni(110) electrode
surface occurred through the carbon at 90 K. However, when the temperature of the
reaction was elevated and H was in the vicinity, the *COOH intermediate flipped and
bound through the two oxygen atoms. Therefore, the *OCHO intermediate could be
observed, yielding formate as the main product.129 Conversely, in the previously
48
mentioned report by Sakaki, an increase in negative charge on oxygen (from –0.33 to –
0.58 eV) facilitated its protonation to the coordinated CO2, which yielded CO formation.
In this comparison, it was the hydrogenation of carbon versus oxygen which determined
the formation of formate versus carbon monoxide.121 However, the *COOH is
energetically favored because the bent conformation of CO2
•−
lowers the necessary
energy for electron transfer and thus, promotes hydrogenation of the O atom given its
extra negative charge. Depending on the electrocatalysts interaction with the *COOHads
intermediate, either HCOO–
/HCOOH or CO will be produced.130 In both Group 1 and 2
electrocatalytic systems, a subsequent PCET step results in the main products being
formate and carbon monoxide, respectively. The selective generation of CO, however,
relies on the electrocatalytic material tightly binding the *COOHads intermediate while
exhibiting a low binding affinity for *CO species. In this pathway, the *COOHads undergoes
a subsequent PCET, simultaneously liberates an equivalent water molecule (i.e.,
dehydrogenation) and delivers the *CO intermediate, finally releasing CO as a product.
In the other pathway, the *COOHads intermediate undergoes hydrogenation to form the
*HCOOH intermediate and hence, the HCOO–
/HCOOH product.126 Adjusting the binding
strength of intermediates on catalysts holds significant importance in crafting ECO2RR
catalysts for converting CO2 to CO with enhanced activity and selectivity.131
Silver-Based Heterogeneous Catalysts for the Reduction of CO2 to CO
Electrocatalysts like Au, Ag, Zn, Pd, and Ga have been identified for producing
mixtures of CO and H2 with different ratios based on the applied voltage.132 Among these
materials, Au and Ag demonstrate the highest activity and selectivity for CO over H2.
Given that Ag is more abundant and cost-effective than Au, it emerges as the most
49
favorable electrocatalyst for the large-scale production of CO. Nonetheless, often labeled
as the competitive reaction to the ECO2RR in aqueous electrolytes, the hydrogen
evolution reaction (HER) is still unavoidable. Under analogous conditions, *H may also
absorb on the catalysts and undergo the same proton-electron step to generate the
undesired HER.133 Moreover, when applying a high overpotential, the rate of the HER is
also enhanced. Thus, silver materials are often classified as promising candidates if they
can either produce high current densities (mA cm–2
) at low overpotentials (η) and high
faradaic efficiency (FE) and selectivity towards CO. The ECO2RR presents a multifaceted
challenge characterized by both energetic and atomic inefficiencies, primarily attributed
to substantial kinetic overpotentials and the intricate control of product selectivity. This
becomes particularly evident in aqueous media where the competitive HER becomes a
parasitic reaction. Addressing these obstacles requires the development of catalysts that
exhibit heightened efficiency, selectivity, and durability. In pursuit of addressing these
challenges, research efforts for electrochemical reduction of CO2 using Ag-based
electrocatalytic materials have primarily focused on morphological and facet modulation,
binary alloy systems, mechanistic surface interactions, and conductive substrate
supports. The present report details the progress for Ag-based electrocatalysts as
heterogeneous materials for the reduction of CO2 into its value-added products.
1. Surface Compositions and their Mechanistic Interactions
Ag-based materials have been elucidated as the state-of-the-art catalysts for CO2
reduction into CO and their application as cathode silver electrodes for ECO2RR have
been extensively reported.134 Particularly, silver-oxides have demonstrated that the
interface confinement effect in supported nanostructures provide and stabilize highly
50
active sites for molecular activation.135–137 Ma et al. electrochemically reduced Ag2O to
derive a silver-oxide (OD–Ag) electrocatalyst, which was studied for reduction of CO2 to
CO. Effectively, Ag oxide electrochemically reduced CO2 with a FECO of ~80% and shifted
the overpotential by more than 400 mV compared to untreated polycrystalline Ag.
Moreover, Tafel plots of the CO partial current density (jCO) were reported to be 133 mV
dec-1 and 77 mV dec-1
for polycrystalline Ag and OD–Ag, respectively. They proposed that
the enhanced electrochemical activity was due to 1) stabilization of the *COOH
intermediate by OD–Ag and a 2) local alkaline pH which led to HER suppression and
improved ECO2RR.138 These electrocatalysts were then further investigated by Firet et
al. in which, through operando X-ray fine structure studies, verified the propositions by
Ma et al. Enhancing the oxygen presence and nanostructuring of the catalyst (aimed at
boosting local pH) play crucial roles in improving the activity and selectivity for CO
production on silver electrocatalysts. These factors are essential components for an
active and selective material for ECO2RR.139
In another study to experimentally verify the importance of atomic oxygen at the
surface, Jiang et al. prepared Ag NPs supported on carbon fiber papers via a dip-coating
and thermal annealing method. Air-annealed Ag catalysts were compared to H2-annealed
Ag. Despite parallel morphologies, structures, and phases in both materials, a dramatic
difference in CO2RR was observed. The FECO of H2–Ag was at approximately 30% while
that of Air–Ag was >90% at a current density of ~21 mA/cm2
. Interestingly, nearly identical
Tafel slopes were reported for Air–Ag and H2–Ag, demonstrating similar reaction kinetics.
Nonetheless, the improved selectivity of CO over H2 generation was attributed to surface
adsorbed oxygen species. In H2–Ag surface bonded O–Agδ+ interactions are eliminated,
51
which when compared to Air–Ag, lead to sluggish CO2 conversion and poor CO
selectivity.136 In 2023, Mattarozzi et al. synthesized pristine silver nanowires via the polyol
method and by applying an oxidative potential bias in 0.2 M NaOH, they generated
different wt. % of AgxO.140 After oxidation at 0.5 V (vs. Ag/AgCl), about 18% of Ag was in
the form of Ag2O and when increasing the oxidation potential to 0.7 V, Ag2O percentage
increased to 65. At the highest potential of 1.0 V, the nanowires consisted of 66% AgO
and 34% Ag2O, leaving no trace of metallic silver. Interestingly, at higher oxidation
potentials, the morphological characteristics of the nanowires became less defined, and
the surface roughness of the catalyst increased. The most oxidized sample produced a
3.3-fold larger CO partial current density (–294.2 mA m–2
Ag) compared to the pristine
nanowires (–89.4 mA m–2
Ag).141 Like Ma et al. and Mattarozzi et al.’s approaches, Yang
et al. prepared an oxide-derived silver (OD–Ag) electrocatalyst from polycrystalline Ag foil
through oxidation-reduction treatment. They reported a FECO of 87% for the OD–Ag
material at an overpotential of ca. 0.7 V, which was significantly higher than the untreated
Ag foil.142
The surface composition in catalytic materials for ECO2RR in silver oxides is
significant for understanding product selectivity and performance. The presented
experimental studies have demonstrated fundamental insights into the mechanistic
effectiveness of atomic oxygen at the surface in silver-based electrocatalytic materials.
Nonetheless, the stabilization effects of reaction intermediates in ECO2RR have also
been investigated with organic functional groups in silver electrocatalysts. In a recent
study, Chen, J et al. modified polycrystalline Ag foil with p–mercaptobenzoic acid (MPA)
followed by electroreduction to produce a thiol ligand modified Ag nanostructure, re–AgS4.
52
The performance of re–AgS4 for electrochemical CO2 reduction was then evaluated and
compared to synthesized Ag NPs. The re–AgS4 electrocatalytic material produced an
FECO of 98.13% compared to their Ag NPs, which had FECO = 89.23% at relatively the
same potential of approximately –1.1 V (vs RHE). The Tafel analysis observed slopes of
111.7 and 163.8 mv dec–1
for re–AgS4 and Ag NPs, respectively. This indicated faster
kinetics by re–AgS4 with the first, direct electron transfer mediated CO2 absorption being
the rate-determining step (RDS) followed by fast proton transfer to form the surface
intermediate *COOH.143
To further expand on thiol-capped silver motifs, Chen, Y et al. isolated cysteine–
Ag NPs to sequentially produce carbon supported thiol-capped Ag (Ag–TC) after
electrochemical reduction. This material was used as the working electrode to study the
ECO2RR in 0.1 M KHCO3. At a potential of –0.7 V (vs RHE), Ag–TC had an FECO of ~10%,
which then reached a maximum of 86.7% at –1.0 V. Unlike Chen’s re–AgS4 material, HER
suppression for Ag–TC was less likely at potential biases more negative than –0.9 V and
thus led to a decrease in CO selectivity. Moreover, at overpotentials larger than 600 mV,
Tafel analysis indicated slopes of 220, 316, and 196 mv dec–1 for polycrystalline Ag foil,
cysteine–Ag NPs, and Ag–TC, respectively.144 This indicated accelerated reaction
kinetics for the thiol–capped Ag–TC with direct electron transfer being the RDS.
Furthermore, they performed electrochemical reduction studies of cysteamine-modified
Ag foil (and numerous other functional groups) and verified the necessity of the Ag–S
bond for larger FECO.
In 2017, Kim et al. synthesized three types of Ag NPs with different capping agents:
dodecanethiol (DDT Ag/C) with a thiol functional group, olelyamine (OLA Ag/C) having an
53
amine functional group, and oleic acid (OA Ag/C) having a carboxyl functional group. The
three types of Ag NPs averaged a uniform size of 5 nm and functional group molar
concentration of 0.3 mM per mg of Ag. Their respective electrochemical activity for CO2
reduction was then assessed in 0.5 M KHCO3 electrolyte. At a potential of –0.7 vs RHE,
the OLA Ag/C had the highest FECO (94.2 ± 1.5%) compared to OA Ag/C (89.1 ± 1.1%)
and DDT Ag/C (65.5 ± 6.7%). The better selectivity of OLA Ag/C was attributed to
exceptional suppression of HER compared to DDT Ag/C, which exhibited an
indiscriminate increase for both HER and ECO2RR. Additionally, they isolated
cysteamine-anchored Ag NPs and post-treated them with hexylamine and hexanethiol to
elucidate the functional group participation in ECO2RR.90 Effectively, the amine
functionality demonstrated a FECO of 78% with overpotential of 540 mV compared to the
thiol functional group which had FECO of 77% at 620 mV. This experiment supported that
the amine functional group enhances ECO2RR by improving efficiency towards CO over
HER. In addition, density functional theory (DFT) calculations consistently suggested
stabilization of the *COOH intermediate with amine-capped Ag NPs.
The DDT Ag/C and re–AgS4 materials yielded CO faradic efficiencies of ~60% at
–0.7 V (vs RHE), significantly higher than those of Ag–TC (FECO = 10%) . However, unlike
the results reported for the re–AgS4 electrocatalyst, DDT Ag/C led to lower FECO and
higher H2 generation at more-negative potential biases. This was similar to what was
observed for Ag–TC. Additionally, DFT studies by Kim et al. were comparable to those by
Chen, J et al., indicating the stabilization of the surface binding energy of *COOH
intermediate by the sulfur-based ligand in Ag NPs. A conclusion also derived by Chen, Y
et al., too. Kim et al. also demonstrated, both theoretically and experimentally, that amine-
54
capped Ag NPs led to better electrocatalytic performance for CO2 conversion. These
results were shortly validated by Wang et al. where they showed that the –NH2
functionalities in cysteamine ligands decreased overpotentials by 300 mV at 1 mA cm–2
.
They reported FECO of 93% at –0.6 V (vs RHE) for their Ag–S(CH2)2NH2 electrocatalyst.145
Similarly, Abbas et al. indicated that functionalization of Ag NPs with amine-based,
oleylamine ligands decreased the contributions from HER and consequently increased
CO generation.146
The importance of silver-based nanostructures in catalyzing the electrochemical
reduction of carbon dioxide into carbon monoxide is evident. Specifically, silver oxides
play a crucial role in creating active sites for molecular activation. These improvements
were attributed to the stabilization of reaction intermediates and the creation of a local
high pH environment at the surface, suppressing competing HER and improving FECO.
Additionally, other surface modifications, such as thiol- and amine-capped silver motifs,
show promise in accelerating reaction kinetics and achieving high faradic efficiencies.
Future research endeavors should continue exploring the local electronic effects of
amine-capped silver nanostructures and elaborating on respective theoretical studies to
effectively model the mechanism for ECO2RR. Finally, even though DFT simulations have
provided validations for the specific experimental mechanisms, it is essential to simulate
all aspects of the electrochemical system and not only the surface chemistry at the
cathode electrode.147
2. Size, Morphology, and Facet Design in Electrocatalysts
There are a variety of other factors influencing the electrochemical activity of a
nanomaterial towards a specific redox transformation besides surface composition.
55
These include morphological characteristics such as atomic ratios at the corner, edge,
and surface sites, which are strictly dependent on the size of the nanomaterial. In
principle, the most promising strategy for developing an effective ECO2RR catalyst is to
synthetically control its surface structure and size. Accordingly, nanoparticles with
different sizes have exhibited different electrochemical properties towards CO2 reduction.
For example, Gao et al. reported the size-dependency of Pd NPs for the ECO2RR. The
NPs, ranging from ca. 2–10 nm, demonstrated significantly different FECO at –0.87 V (vs
RHE), which varied from 5.8% over 10.3 nm NPs to 91.2% over 3.7 nm NPs. DFT
calculations and free Gibbs energy diagrams determined that the mechanistic rates varied
by Pd NP size due to the atomic ratio tunability of the corner, edge, and surface sites.89
In a more related 2015 example, Kim et al. reported the novel synthesis of cysteamineanchored silver NPs with three varying sizes (3, 5, and 10 nm) supported on carbon
Ketjen black. After characterization of the nanosized Ag/C materials, the electrochemical
activity for CO2 to CO conversion was explored under aqueous media. The 5 nm Ag/C
demonstrated a higher CO activity over the 3 nm and 10 nm Ag/C samples, respectively,
reaching a maximum at approximately –0.9 V (FECO = 84.4%). Specifically, the 5 nm Ag/C
demonstrated a positive shift of 300 mV at 1 mA/cm2 compared to the polycrystalline Ag
foil. The 3 and 10 nm Ag/C electrodes both displayed an anodic shift of 200 mV at the
same current density. DFT calculations revealed that the Ag–S bond promotes surface
localization of the unpaired electrons, which induces stability of the intermediate and
consequently leads to better catalytic activity. Nonetheless, this report presents how
differences in Ag NP size affects current density, overpotentials, and product selectivity
for electrochemical CO2 reduction.148
56
Given the complementary results by Salehi–Khojin et al. and Kim et al. indicating
enhanced performance of 5 nm Ag NPs towards electrochemical reduction of CO2 to CO,
Deng et al. studied the size dependency of sub-5 nm Ag NPs for both ECO2RR and
HER.149 These Ag NPs were synthesized through ultrahigh vacuum onto highly oriented
pyrolytic graphite via metal evaporation, which allowed sub-nm precision without capping
or other passivating ligands. Additionally, scanning tunneling microscopy (STM) images
were used to estimate the average particle size. At –1.2 V (vs RHE), ca. 1.9 ± 0.7 nm–
diameter NPs produced an FECO of ca. 20%, which increased to ca. 90–100% for particle
diameters > 3 nm. Consequently, FEH2 decreased from 60% to < 5% with increasing NP
diameter. The highest activity and FECO was observed by NPs with average diameter of
3.7 ± 0.7 nm.
Deng et al. then sought to gain atomic-level insight into HER and ECO2RR rates
through computational modeling (applying the hydrogen electrode model) of Ag NPs with
diameters between 1 and 10 nm. Calculations found that the concentration of inaccessible
bulk atoms increased with NP diameter. Moreover, the accessible population of corner
and edge sites decreased with particle diameter while Ag(100) and Ag(111) surface sites
increased. Among the sites considered, surface Ag(100) sites and edge sites had the
smallest reaction barrier and largest rate constant contribution for ECO2RR and HER,
respectively. At smaller NP diameters the population of edge sites dominated, which
promoted HER rates. Corner sites did not contribute significantly to HER activity.
Analogously, a larger Ag particle diameter population led to an increase in Ag(100) sites
and a decrease in population at the edge site, which favored ECO2RR. At diameters > 4
nm, the Ag(100) sites and inaccessible bulk atom population increased, which showed a
57
decrease in catalyst utilization based on metal content.150 The presented results showed
the size dependence of Ag NPs for electrochemical reduction of CO2 to CO. In theory,
smaller NPs should yield better activities for ECO2RR given their higher surface area per
unit volume. Reports have indicated the optimal Ag NP size to be between 3–5 nm.
Despite the Ag NPs by Deng et al. demonstrating higher FECO over those by Kim et al.,
their synthetic approach is much more elaborate and energy intensive. Ideally, synthetical
methodologies of electrocatalytic nanomaterials should be simple, precise, and costeffective for their proper implementation into large-scale applications. Much interest has
been focused on synthesizing uniform NPs and understanding their respective sizedependency for the ECO2RR. Other electrocatalytic materials with distinct morphological
characteristics have also been explored. In 2017, Liu et al. reported the shapedependency of electrochemical reduction of CO2 to CO on triangular silver nanoplates
(Tri–Ag–NPs).151 The formation of the Tri–Ag–NPs were visualized through a UV-Vis shift
and further characterized through TEM images. When studied for the ECO2RR, Tri–Ag–
NPs demonstrated enhanced current densities and FECO at lower overpotentials
compared to similarly sized Ag NPs (SS–Ag–NPs) and bulk Ag. Specifically, at an
overpotential of just 96 mV, Tri–Ag–NPs generated a FECO of 8.1%. To match the FECO
of the triangular morphology, it required 346 and 546 mV of overpotential for SS–Ag–NPs
(8.0%) and bulk Ag (8.6%), respectively. Additionally, Tri–Ag–NPs reached a maximum
FECO of 96.8% at ca. –0.86 V (vs RHE) and stability over 7 days. DFT calculations
demonstrated the lowest *COOH Gibbs free energy for the Ag(100) facet compared to
Ag(111) and Ag(110), suggesting higher catalytic activity at the Ag(100) sites. Given that
Tri–Ag–NPs are only enclosed by Ag(100) and Ag(111) facets, the electrocatalyst
58
demonstrated enhanced performance compared to SS–Ag–NPs. Finally, contrary to
Deng et al.’s theoretical studies, Lu et al. attributed the superior CO selectivity to the
larger population of edge sites (optimized edge-to-corner ratio) at Tri–Ag–NPs.
Rosen et al. revealed the fundamental mechanism of ECO2RR to CO on highly
nanostructured silver surfaces. Changes in relative free energy of adsorbed intermediated
on model surfaces, Ag(100), Ag(111), Ag(110), and Ag(211), were computed through DFT
with the Vienna ab initio simulation package. The smaller Tafel slopes for Ag NPs and
nanoporous Ag (64 mV dec–1 and 59 mV dec–1
, respectively) compared to its
polycrystalline form (134 mV dec–1
) indicated a better electrocatalytic activity for the
nanostructures.
152 In another study, researchers sought to gain insight into the facetdependence of CO formation over different Ag thin films. With a novel methodology, Clark
et al. grew silver thins with (111), (100), and (110) orientations on single-crystal silicon
wafers. They concluded that Ag(110) exhibited superior electrochemical reduction activity
for the conversion of CO2 to CO compared to both Ag(111) and Ag(100). Thus, the higher
activity over the Ag(110) thin film was related to a larger concentration of under
coordinated sites than those on Ag(111) and Ag(100) thin films.153 Theoretical studies also
found that defects along edge sites, Ag(211), showed greater activity compared to Ag(111)
or Ag(100) and slightly higher activity compared to Ag(110) surface sites. Finally, they
reported that stabilization of the key *COOH intermediate was instrumental for efficient
electroreduction of CO2 to CO.
A major limitation of highly nanostructured electrocatalysts is their instability over
extended working conditions due to support degradation and nanoparticle aggregation.154
Research efforts have sought to study monolithic, nanoporous structures for different
59
electrochemical transformations to circumvent the stability issues. In general, nanoporous
metals consist of 3D interconnected backbones (ligaments) and pores (channels/voids) ,
which allow for easy transport of electrons and provide targeted, catalytic sites with large
surface areas.155 These materials contain more exposed facets, coordinated sites, and
defect-rich surfaces, which have been attributed to enhanced electrochemical
properties.155–159 For example, nanoporous silver (np–Ag) has been reported to display
equivalent electrocatalytic activity for the oxygen reduction reaction (ORR) against the
state-of-the-art Pt catalyst.158 This same type of material was also used as an anode
electrode for the oxidation of ammonia-borane in a direct alkaline fuel cell.160 However,
given silver’s status as the state-of-the-art metal for conversion of CO2 to CO, np–Ag
continues harnessing attention for ECO2RR. Even though multiple synthetic approaches
have been reported for the formation of nanoporous electrocatalytic materials, this review
will focus on dealloying, and electrochemical methods given their intrinsic benefits.157,159
In this notable 2014 report by Lu et al., a nanoporous Ag catalyst was obtained by a twostep dealloying/etching process of an Ag–Al precursor. This electrocatalyst was able to
reduce CO2 to CO with ~92% selectivity at a significantly higher rate compared to its
polycrystalline counterpart at moderate overpotentials < 500 mV. The Tafel slope for np–
Ag of 58 mV dec–1 compared to 132 mV dec–1 for polycrystalline Ag demonstrated the
enhanced capacity of the nanoporous frameworks to stabilize reaction intermediates over
flat surfaces.127 Yang and coworkers reported a novel np–Ag architecture prepared by
dealloying a Mg–Ag precursor. They prepared two np–Ag samples with average
ligaments of 21 and 87 nm and studied them as electrocatalysts for ECO2RR.
Interestingly, the 21 nm np–Ag electrode outperformed the 87 nm analogue, achieving an
60
FECO of 85.0% and 41.2% at –0.8 V (vs RHE), respectively. These studies demonstrated
the ligament dependency of np–Ag on ECO2RR.159
In a 2015 example, Hsieh et al. reported an Ag nanocoral electrocatalyst (ca. 150
nm) for the reduction of CO2 to CO. The coral like np–Ag was prepared from
polycrystalline Ag foils via an oxidation-reduction process in the presence of chloride
anions (0.1 M KCl). The electrocatalyst displayed a FE of 95%, overpotential of 370 mV,
and current density of 2 mA cm–2
. Tafel analysis indicated a slope of 58.5, 87.9, and 137
mV dec–1
for Ag nanocoral, Ag NPs (50 ± 24 nm), and bulk Ag foil, respectively. CO partial
current densities for Ag nanocorals and Ag NPs were 660 and 40 times higher compared
to the polycrystalline foil. Additionally, the nanocorals had a specific activity of 162 µA
cm–2 compared to the ca. 22 µA cm–2 of Ag NPs despite the larger particle size, which was
attributed to surface-bound chloride ions.156 Inspired by these findings, Yuan and
coworkers prepared a series of np–Ag electrocatalysts by oxidation of a silver sheet to
form AgCl (in 0.1 M NaOH and 0.1 M NaCl), followed by reduction of AgCl to form
nanoporous Ag. These electrocatalysts were then studied for ORR. Briefly, the enhanced
activity of np–Ag was attributed to an increase in ECSA and facilitated ion diffusion within
the interconnected pores.157
In 2017, Lee et al. produced AgCl nanosheets through oxidation of an Ag foil in
silver chloride. Sequential electroreduction in 0.5 M NaHCO3 yielded hierarchical Ag
nanosheets. High resolution STEM images revealed that the nanosheets were composed
of interconnected Ag NPs (d = 45 ± 10 nm). At a potential of –0.6 V (vs RHE), the Ag
nanosheet electrode produced a current density 17 times that of the polycrystalline Ag
foil. Additionally, the calculated FECO were of ca. 95% and 5% for Ag nanosheet and foil,
61
respectively, at an overpotential of 290 mV.161 Similarly, Dutta et al. prepared an Ag foam
via electrodeposition (j = –3.0 A cm–2
) of a plating solution (1.5 M H2SO4, 0.02 M Ag2SO4,
and 0.1 M Na3C6H5O7•2 H2O) onto an Ag foil. SEM images demonstrated a clear
distinction between the Ag foil substrate and the Ag foam. Moreover, the foam thickness,
pore diameter, and mechanism of formation are accurately detailed.134 At potentials
greater than –1.1 V (vs RHE), the dominant product of the Ag nanofoam was CO.
Interestingly, FECO dropped at more negative potential biases (E < –1.2 V) and
hydrocarbon formation became notable. At –1.5 V (vs RHE), FE for CH4 and C2H4
reached 51% and ca. 9%, respectively. This study demonstrated how a synthetic
approach can customize the morphological/structural characteristics of an Ag-based
catalyst that primarily produces CO into a catalyst resembling Cu and capable of
producing hydrocarbons. Qian et al. then studied the formation mechanism in the
oxidative-reductive synthesis approach for np–Ag.162 They employed SEM of the surface
and cross-section during the anodic formation and reduction of the AgCl layer. However,
they used a 0.1 M HCl electrolyte solution, varying from what was demonstrated by
previous authors. Analogous to Lee et al.’s interpretation, Qian and coworkers described
the np–Ag electrocatalyst as filled with nanoparticles with an average size of 80 nm.
These interconnected NPs formed secondary pores with an average diameter of ~ 93 nm.
At –0.6 V (vs RHE), the np–Ag displayed jCO of ca. 5 mA cm–2
, FECO < 90%, and Tafel
slope of 75 mV dec–1
, which like Hsieh et al.’s and Lee et al.’s reports, indicated a rapid
initial electron transfer step.
Recently, Wu et al. electrochemically generated a porous Ag–D catalyst from an
Ag2O precursor. They reported breakage of the Ag–O bond by electrons generated at the
62
cathode; the Ag–D structure did not demonstrate XRD peaks corresponding to Ag2O.
Additionally, dissolved CO2 in the electrolyte prevented agglomeration of active sites and
were helpful in the formation process to form Ag nanocrystals with defect structures. The
Ag–D electrode demonstrated enhanced electrocatalytic performance observed by a
larger total current density and CO partial current density compared to Ag foil. At a
potential of –0.7 V (vs RHE) using an H-cell, the Ag–D electrocatalyst generated FECO of
100%. Moreover, when tested in a flow cell using 1 M KOH electrolyte, the Ag–D gas
diffusion electrode (GDE) demonstrated FECO of almost 100% and current density of 180
mA cm–2 at –1.0 V (vs RHE).163 Despite the experimental comparison to a silver oxide
analogue, theoretical results attributed the enhanced performance to the abundant defect
sites in the Ag–D electrocatalyst. In a more recent report, Ag2S hollow cubes were
prepared by the sulfidation and etching of Ag2O cubes. The catalyst was then
reconstructed in situ to produce a Ag2S/Ag interface using 20% H2/Ar (volume ratio, 1 atm
at 350°C). XRD, SEM, and EDS demonstrated that the sulfide was effectively
removed/etched from the Ag2S hollow cube matrix to yield the Ag adduct. However,
instead of employing the traditional H-cell electrolytic system, Ye et al. integrated a flow
cell device. Herein, Ag2S/Ag were painted on carbon paper with a microporous layer as
their cathodic GDE. At a potential of –0.7 V (vs RHE), the electrocatalysts attained a
current density of 421.7 ± 14.4 mA cm–2 and FECO of ca. 99% over 50-hour operation.
DFT calculations demonstrated enhanced stability of *COOH intermediate by the
interfacial interactions between Ag2S and Ag structure.164 Continually, Liu et al.
synthesized hollow porous Ag NPs (using a Cu2O template method) and studied their
application as electrocatalytic materials for ECO2RR. At a potential of ca. –0.85 V (vs
63
RHE), the porous materials generated a FECO of 94% and jCO of 3 mA cm–2 compared to
the 60% and 0.5 mA cm–2 of solid Ag microspheres.165
Nanoparticles can exist in a plethora of shapes and morphologies dependent on
the applied synthetic methodologies. From a geometric perspective, particles are
surrounded by multiple facets, where the meeting point of these facets forms an edge
and the convergence of two edges creates a corner. Structural modifications essentially
involve adjusting these facets, edges, and corners. Step edge defects and the
undercoordination of atoms at the surface have been demonstrated to increase the
electrocatalytic activity of the material for conversion of CO2 to CO. The stability of key
intermediates such as *COOH has been confirmed fundamental to enhance selectivity
towards the ECO2RR pathway. Nonetheless, these electrocatalysts are often deposited
on conductive substrates to maximize catalyst dispersion and utilization. Revealing the
convolution behind selective catalytic activity is a difficult yet valuable endeavor.
Research efforts are needed to further understand the intrinsic relationship between a
nanomaterial’s size and morphological characteristics as well as elucidating the true
origin of catalytic activity. Alternatively, nanoporous metals have rapidly gained ground as
catalysts for ECO2RR given their distinct, highly desirable features. The tunability of
ligaments and pores, increased surface area, and targeted catalytic sites open avenues
for truly understanding catalytic activity and enhancing overall performance and stability.
However, these materials present limitations that should be explored further. First, the
dealloying process, which is the most common to introduce nanoporosity, causes residual
metals to leach into the ligament networks of the material. This may affect catalytic
properties and favor unwanted side reactions. Second, given the large void volumes,
64
nanoporous metals are not mechanically strong. Thus, incorporating them as electrodes
in commercial applications (e.g., fuel cells and electrolyzers) remains challenging.
Ultimately, the synthetic preparation of promising electrocatalytic candidates must ideally
be simple, consistent, and cost-effective for their proper implementation.
3. Binary Silver–Metal Alloys
Alloying metals have been reported to play a pivotal role in enhancing the
electrochemical properties of catalysts. The process of alloying involves combining two
or more metallic elements to create a new material with unique properties that surpass
those of its individual components. In electrochemistry, alloying materials exhibit
enhanced characteristics including improved conductivity, enhanced stability, and tailored
reactivity. Amongst the various strategies presently under consideration, such as size,
morphology, and substrate support, alloy catalysts have proven effective at lowering
kinetic overpotentials and controlling selectivity. This is due to electronic and geometrical
changes of the material which modify the chemical binding of intermediates. Silver
electrocatalytic materials have been alloyed with different elements, which have produced
materials with distinct morphologies and electrochemical properties, as shown in Table
3. Specifically, significant emphasis has been placed on incorporating copper into silverbased nanostructures given its ability to produce hydrocarbons. In 2016, Choi et al.
prepared Ag–Cu dendrites on a Cu foil through electrodeposition and by varying the
concentration of Ag precursor, they modified the composition ratios of Ag and Cu. Three
dendritic Ag–Cu materials, with varying atomic ratios, were compared to Ag100 dendrites
for electrochemical conversion of CO2. Interestingly, the monometallic Ag100 dendrites
yielded the highest FECO of 64.6% at –1.7 V (vs SCE) against the three bimetallic
65
analogues. However, all the Ag–Cu catalysts demonstrated a higher mass activity for CO
than the Ag100 material. Specifically, Ag57Cu43 had a mass activity of 30.5% mg-1 of Ag at
–1.5 V (vs SCE) while Ag100 had a mass activity of 16.3% mg-1 of Ag at a slightly more
negative potential bias. Thus, by controlling the Ag : Cu composition ratio, direct syngas
formation and enhancement in CO mass activity was obtainable.166
The incorporation of metallic constituents that favor the formate producing
mechanism, such as indium and tin, have also been reported. In addition, these metals
have been demonstrated to largely suppress HER activity.167,168 Luc et al. synthesized a
core-shell nanostructure containing an Ag–Sn bimetallic center enclosed by a partially
oxidized SnOx shell. The molar ratio of Sn and Ag precursors were tuned to produce
AgSn/SnOx materials with different compositions to then study their respective
electrochemical activity towards CO2 reduction. Interestingly, a volcano-type correlation
was observed between Sn content and FEHCOO-, demonstrating a maximum FE at 24
atomic percent. Through STEM and electron energy loss analysis, the Ag76Sn24 bimetallic
core displayed an SnOx shell with thickness of ca. 1.7 nm.169 DFT calculations
demonstrated the favorable formation of the *OCHO intermediate, stabilized by oxygen
vacancies at the surface. However, electrical conductivity loss occurred with increasing
thickness of the SnOx layer, which explained the volcano-shaped correlation of the
FEHCOO-. In a following report by Cai et al., a SnOx/Ag heterostructure nanomaterial was
reported for the electrochemical reduction of CO2. The SnOx/Ag electrocatalyst was
synthesized via hydrolysis of Na2SnO3 to form amorphous SnOx NPs (~5 nm) on Ag NPs
(~30 nm). Interestingly, the metal and oxide components, were reported to switch roles
as the major catalytic component by varying the working electrode potential. In the
66
potential range from –0.6 to –0.8 V (vs RHE) the oxide promoted CO formation while from
–0.8 to –1.1 V (vs RHE) the metal facilitated production of HCOOH. The binary composite
material of SnOx/Ag featured Ag as an active component for CO2 to CO conversion and
SnOx for CO2 to HCOOH conversion and exhibited combined electrocatalytic capabilities
surpassing those of its individual constituents at their optimum performance levels.170
In a more recent example, Yun et al. explored a membrane-electrode assembly
(MEA) for ECO2RR using AuAg NPs catalyst on a carbon support as cathode electrode.
The AuAg bimetallic electrocatalyst was prepared by galvanic replacement with
chloroauric acid in which AuCl4
– anions led to the spontaneous oxidation of Ag to Ag+ and
reduction of Au3+ to Au (thus generating AgCl as byproduct). Dependent on the applied
synthetic temperature, either room (AuAg–R) or boiling temperature (AuAg–B), different
morphologies of AuAg NPs were observed. In an H-cell setup at –0.9 V (vs RHE), the
AuAg–R generated an FECO over 80% with 10 mA cm–2
. Yun et al. then optimized the
catalyst by incorporating two different Ketjen black carbon supports (300J and 600JD)
and used them as cathode electrodes in a zero-gap MEA electrolyzer to circumvent the
mass transport limitations from the H-cell. They reported an FECO of 83% at 3.5 V and
618 mA cm–2
for the AuAg–R/CKB600. Additionally, in situ/operando X-ray absorption
spectroscopy confirmed that Ag+ ions preserved from the catalyst synthesis formed the
active sites with more oxidative and undercoordinated Ag surfaces.171
The synergistic effects of other metallic elements into silver–based materials for
ECO2RR are plenty. Monometallic Ag electrocatalysts largely favor CO production but
when unified with other metals to generate bimetallic systems, the product distribution is
widespread. As discussed, variations in atomic composition ratios in electrocatalysts can
67
yield a tailored distribution of products. Silver-copper alloys continue harnessing major
research interest due to reported generation of hydrocarbons. However, investigations
should aim to optimize conditions in the H-cell to then transcribe the respective systems
into an MEA. Zero-gap MEAs can effectively resolve the mass transfer limitations
introduced by the low solubility of CO2 in aqueous electrolyte.
4. Carbonaceous Supports and Other Substrate Materials
Carbon-supported electrocatalysts are by far the most studied structures for
ECO2RR. Carbon is widely available and economically advantageous, exhibiting
inertness and electrochemical stability. It is obtainable in bulk forms such as activated
carbons and graphite, as wells as in nanoscale structures like fullerenes, nanotubes, and
graphenes. In the most practical synthetic approach, the metal precursor salt is converted
into metal NPs in the presence of the carbon support. It is well-accepted that the
supporting material acts as a nucleation point for uniform dispersion of nanoparticles and
serves to prevent their agglomeration.172 Undoubtedly, the metal-support interaction
should always be considered and ideally, the experimental results normalized to the
electrochemical surface area for proper catalytic comparison. In recent years, the
introduction of heteroatoms (N, P, B, S, Se) in the carbon network have received
significant attention.173 Particularly, there has been notable advancement in N-doped
carbon-supported metal catalysts, drawing considerable interest due to nitrogen doping’s
ability to modify carbon properties for various applications. These catalysts have
demonstrated enhanced catalytic performance in different transformations compared to
their pristine carbon supported counterparts.174 In the case for ECO2RR, these
improvements have been attributed to the basicity of the N functional groups which result
68
in electron abundancy or deficiency for the stabilization of intermediate species.175–177
Additionally, CxNy structures provide channels for electrolyte diffusion and improve
surface wettability and electrical conductivity.178
Carbon nitride (C3N4) has shown promise for stabilization of metallic active sites
due to its high content of lone electron pairs and coordination ability. For example, Sastre
et al. incorporated Ag NPs (with different Ag wt. % loading) on graphitic carbon nitride (g–
C3N4) for the generation of synthesis gas. They reported that the H2 : CO product ratio
could be controlled by manipulating Ag loading and the applied potential. However, their
electrolytic system required a minimum operation time of 5 h to reach steady state current.
Nonetheless, for potentials between –1.05 V and –1.15 V (vs RHE), an ideal syngas ratio
of 2 was demonstrated under phosphate electrolytes.179 In a 2022 report, Lu et al.
anchored Ag NPs on boron-doped g–C3N4 (based on theoretical simulations) and studied
the material’s electrocatalytic properties for the reduction of CO2 to CO. At a potential of
–0.8 V (vs RHE), the Ag–B–g–C3N4 nanocomposite generated a current density of ca. 2
mA/cm2
, FECO of 93.2%, and 12-h operative electrocatalysis.180 Zhang et al. reported on
the rich oxygen sites of Ag NPs on g–C3N4 for the ECO2RR. They synthesized Ag NPs
supported on carbon nitride through the hydrothermal decomposition of Ag2O precursor.
However, due to the incomplete decomposition of Ag2O, this species coexisted with the
nanoparticles in the final composite material. At a potential of –0.8 V (vs RHE), the Ag/g–
C3N4 composite yielded a mass activity of ca. 80 mA mg–1 compared to the ca. 30 mA
mg–1 of Ag NPs and physically mixed (P–Ag/g–C3N4). This electrocatalytic enhancement
was attributed to the stabilization of reaction intermediates via the ultra-stable oxygen
species and synergistic effect between Ag NPs and g–C3N4.
181
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Even though N-doped carbon supports have received most research interest,
studies have also aimed to understand the synergistic effects from other substrate
materials. Specifically, metal oxides have been studied as substrate supports for Ag
nanomaterials in ECO2RR.182–187 In this 2014 report, Ma et al. incorporated sub-10 nm Ag
NPs of varying weight percent on TiO2 and compared their electrochemical activity to 40
wt. % Ag on carbon black (Vulcan XC-72R) using 1 M KOH electrolyte. At a potential of
–1.7 V (vs Ag/AgCl), 40 wt. % Ag/TiO2 and 40 wt. % Ag/C generated jCO of 60 and 28 mA
cm–2
, respectively. This demonstrated the enhanced activity of TiO2 for generation of CO.
The improved electrocatalysis was attributed to the mechanistic participation of the
TiIV/TiIII redox couple, which stabilized reaction intermediates.182 In another study by
Larrazabal et al., the synergistic effects of silver-indium electrocatalysts with different
architectures were assessed for carbon dioxide reduction. For benchmark supports, they
stabilized Ag NPs on carbon black and P25 TiO2 (T). Two electrocatalysts were then
synthesized on In-based supports, Ag NPs on In2O3 (IO) and on In(OH)3 (IH). Finally, a
composite support of Ag NPs on P25 TiO2 + In(OH)3 (T+IH) was also studied. Silver
electrodes decorated with oxidized indium deposits exhibited an improved current
efficiency for CO production at moderate overpotential. To explore the specific
interactions between silver and the oxidic indium phases (IH and IO), catalysts were
tailored with varying silver loadings (ranging from 5 to 40 by wt. %) supported on each
material. Their respective performances for ECO2RR were assessed at a potential of –
0.6 V (vs RHE). Compared to the pristine supports and catalysts featuring similar silver
loadings on carbon black, the Ag/IO and Ag/IH electrocatalysts demonstrated enhanced
jCO at moderate overpotentials.
70
Table 3 Silver-incorporated bimetallic catalysts for the electroreduction of CO2 into its different products. (1a) nanowires (1b)
nanospheres (1c) nanodimers (2) normal hydrogen electrode (3) standard calomel electrode (4) standard hydrogen electrode.
Entry catalyst electrolyte product(s) E (V vs RHE) FE (%) j (mA cm–2
} ref
1 Pd4Ag NWs1a 0.1 M KHCO3 HCOO– –0.08 to –0.24 > 95 3 at –0.25 V 188
2 PdAg NSs1b 0.1 M
NaHCO3
HCOO– –0.1 to –0.3 > 90 5.3 at –0.35 V 189
3 AgZn foil 0.1 M KHCO3 CO –1.2 67 - 93
4 Ag–Cu NDs1c 0.1 M KHCO3 C2H4 , H2 –1.1 40, 30 1 190
5 CuAg 0.1 M KHCO3
C2H4 ,
C2H OH 5
–0.7 60, 25 300 191
6 AgIn 0.5 M KHCO3 CO –1.3 (vs NHE)2 60 5 192
7 Ag@Cu 0.1 M KHCO3 H2 , CH4 –1.16 ca. 40 , 20 3.75 193
8 Ag–Zn dendrites 0.1 M
CsHCO3
CO –0.9 97 22.5 194
9 AgSn/SnOx
0.5 M
NaHCO3
HCOO–
, CO ,
H2
–0.85 ca. 80 , 10 ,
10 17.5 169
10 Ag57Cu43
dendrites 0.5 M KHCO3 H2 , CO –1.7 (vs SCE)3 ca. 70, 15 12.3 166
11 Ag–Co 0.5 M KHCO3 H2 , CH4 , CO –2.0 (vs SHE)4 ca. 50, 20, 8 75 195
12 AuAg NPs 0.1 M KHCO3 CO –0.9 > 80 10 171
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This phenomenon implied a mutually beneficial interaction between silver and the oxidic
indium supports, markedly evident in the case of Ag/In(OH)3 at higher silver loadings.183
Different silver architectures have also been used as substrate materials for ECO2RR.
In a 2017 report, Lee et al. produced AgCl nanosheets through oxidation of an Ag foil in
silver chloride. Sequential electroreduction in 0.5 M NaHCO3 yielded hierarchical Ag
nanosheets. High resolution STEM images revealed that the nanosheets were composed
of interconnected Ag NPs (d = 45 ± 10 nm). At a potential of –0.6 V (vs RHE), the Ag
nanosheet electrode produced a current density 17 times that of the polycrystalline Ag
foil. Additionally, the calculated FECO were ca. 95 and 5 percents for Ag nanosheet and
foil, respectively.161 Similarly, Dutta et al. prepared an Ag foam via electrodeposition
(j = –3.0 A cm–2
) of a plating solution (1.5 M H2SO4, 0.02 M Ag2SO4, and 0.1 M Na3C6H5O7
• 2 H2O) onto an Ag foil. SEM images demonstrated a clear distinction between the Ag
foil substrate and the Ag foam. Moreover, the foam thickness, pore diameter, and
mechanism of formation are accurately detailed.134 At potentials greater than –1.1 V (vs
RHE), the dominant product of the Ag nanofoam was CO. Interestingly, FECO dropped at
more negative potential biases (E < –1.2 V) and hydrocarbon formation became notable.
At –1.5 V (vs RHE), FE for CH4 and C2H4 reached 51 and ca. 9 percents, respectively.
This study demonstrated how a novel synthetic approach can customize the
morphological/structural characteristics of an Ag-based catalyst that primarily produces
CO into a catalyst resembling Cu and capable of producing hydrocarbons. In a more
recent report, Wang et al. synthesized an Au film on an Ag nanowire core electrocatalyst
for CO2-to-CO conversion. As seen through scanning transmission electron microscopy,
the Au shell had a thickness of 20-30 nm and was uniformly distributed on the Ag nanowire
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structure with an average diameter of 130 nm. At a potential of –1.2 V (vs RHE), the
Ag@Au core-shell catalyst achieved nearly 100% FECO in CO2-saturated 0.1 M KCl.196
5. Single–Atom Catalysts
As previously discussed, NP diameters larger than 4 nm increase the inaccessible
bulk atom population, which demonstrated a decrease in catalyst utilization based on
metal content. By decreasing NP size in an attempt to increase atomic surface population
and enhance catalyst utilization, other morphological changes become evident; thus,
changing the rates and selectivity for ECO2RR.150 Furthermore, metal atoms at the
surface can be exposed to different chemical environments, which bring limitations into
understanding structure-activity relationships. In this regard, single atom catalysts (SACs)
have rapidly emerged at the forefront for various electrochemical transformations. SACs
have introduced discrete energy-level distribution and a distinctive LUMO-HOMO gap due
to the unique quantum size effect and have exhibited enhanced catalytic activity and
selectivity due to their unsaturated coordination sites and distinctive electronic structures.
These features not only augmented electrochemical performance but have also led to
substantial reductions in the usage of catalytic metals by maximizing the atom
efficiency.197–199 Qiao et al. first time proposed and reported the synthesis of isolated Pt
SACs anchored to the surfaces of FeOx nanocrystallites. The Pt SACs exhibited
significantly high activity and stability for CO oxidation (double the Au/Fe2O3 standard)
and preferential oxidation of CO in hydrogen gas.200 Motivated by these findings, Back et
al. proposed the potential of different SACs and their applications in CO2 reduction based
on DFT computational studies.
133 Since then, a plethora of review reports have been
published on SACs including theoretical insights, basic structural characteristics,
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respective synthetic strategies, and electrochemical applications.198,201–204 For the
purpose of this review, the following discussion will principally expand on silver-based
SACs and their recent trends for ECO2RR. Nonetheless, readers should explore the
ample literature on SACs and their promising applications in various other
electrochemical transformations.
Preparation methods of SACs are challenging due to low catalyst loading and
particle agglomeration during synthesis.199 Recently, adopting high surface area supports
such as carbon-based supports, metal oxides, and other porous structures have gathered
interest to finely disperse and stabilize SACs. Three common strategies for SAC
stabilization exist in the reported literature, which include 1) spatial confinement, 2)
coordination frameworks (i.e., MOFs and COFs), and 3) defect/vacancy design.198
However, this review will not expand on coordination frameworks given the extensive
library on these types of systems.205–208 The term ‘confinement’ in this context implies that
individual atoms can be stabilized within the lattice or coordination environment of 2D
materials through robust covalent bonds. This way, the intrinsic nature of single particles
to aggregate into larger-sized particles is prevented. It is important to note that the
confined single atoms maintain a coordinatively unsaturated state, thereby opening the
active sites.209 The defect/vacancy design strategy, as discussed earlier, refers to the
opening of unsaturated coordination sites on substrates (i.e., carbon materials) where
metal atoms can be effectively anchored. Even though the synthesis of different Ag SACs
has been reported, their electrochemical applications are still limited.210–212
In the beforementioned report by Li et al., graphene was used as the carbon-based
host support for their Ag complex ({[Ag(NO3–O)(phtz–N)]2(μ–phtz–N,N′)2}) via strong π-π
74
interactions. This stabilization technique is intrinsically different than those reported by
Siu et al. and Mu et al., in which the respective carbon substrates were doped with
nitrogen (N–Cx) and used to anchor Ag SACs. Given the weak interactions between the
single metal atom and carbon substrate, the stabilization can only be achieved through
processes such as doping or anchoring. Constructing supports with atoms like N, O, and
S, which possess lone pairs of electrons, have facilitated synthesis of SACs due to the
robust interaction between metal species and these coordinating atoms acting as
anchoring sites.209,213,214 Nitrogen-doped materials were found to be one of the most
widely studied strategies for stabilization of SACs.215 Continually, graphitic carbon nitride
has also harnessed significant attention for the stabilization of numerous SACs.216
Nevertheless, reports of Ag SACs on g–C3N4 for catalytic reduction of CO2 are strictly
limited. Recently, Hu et al. used similar impregnation techniques to those reported for Ag
NPs to generate Ag SACs anchored on hollow porous polygonal C3N4 nanotubes (PCN).
After isolation of their PCN material, silver atoms were incorporated by a facile
“impregnation + pyrolysis” strategy. Following impregnation, the solid samples were
pyrolyzed at 150C under inert atmosphere for 1 h to then isolate noble Ag1(x)@PCN
photocatalytic materials (x: Ag wt. %). Experimental and theoretical studies demonstrated
that the Ag–N3 coordination allowed for efficient CO2 photoreduction (FECO > 94%) using
water as the reductant.217
In a 2020 report, Li et al. synthesized a binuclear Ag complex (with weak Ag–Ag
interaction) coordinated with three N atoms from phthalazines, which was then adsorbed
to graphene through strong π-π interactions. The material then underwent pyrolysis with
simultaneous annealing to protect the N coordination and prevent particle aggregation.
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The same experimental protocol was used to synthesize the mononuclear analogue. The
Ag1–G and Ag2–G materials were isolated and characterized. Interestingly,
characterization showed an Ag–N coordination of 4 and 3 for the Ag1–G and Ag2–G,
respectively, which were consistent with the precursor complex structure. When tested
for ECO2RR in 0.5 M KHCO3, the Ag2–G electrocatalyst produced FECO = 93.4% with a
current density of ca. 12 mA cm–2 at –0.7 V (vs RHE). The Ag1–G electrocatalyst had a
FECO of 79% at the same potential. The enhanced activity was experimentally and
theoretically attributed to the formation of AgN3–AgN3 sites in Ag2–G which was found to
promote CO2 adsorption and stabilization of intermediates.218 The following year, Siu et
al. again explored the Ag single atom three nitrogen coordinated motif on porous concave
N-doped carbon (Ag1–N3/PCNC) for the ECO2RR. At a potential of just –0.37 V (vs RHE),
Ag1–N3/PCNC generated FECO = 95%. Differently than Li et al.’s approach to compare
their bimetallic system to the monometallic analogue, Sui et al. compared their Ag1–N3
material to the Ag1–N2 coordinated moiety. Interestingly, the CO faradic efficiency for the
Ag1–N3/PCNC electrocatalyst was higher at all measured potentials compared to the Ag1–
N2/PCNC material. This was attributed to enhanced desorption of CO on Ag1–N3/PCNC,
which was confirmed by ATR-SEIRAS measurements.219
Recently, the nitrogen coordination environment was explored by Mu et al. by
preparing a series of Ag SACs anchored on nitrogen-doped carbon matrixes (i.e., glucose
or reduced graphene oxide). To investigate different degrees of N coordination, they
prepared Ag–N–Cx and Ag–N–rGOx materials with x = 0.0125 M, 0.025 M, and 0.05 M.
Independent of carbon source concentration, N–Cx glucose support demonstrated ca.
70% pyridine nitrogen content while N–rGOx demonstrated ca. 40% pyrrolic N. The six
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materials were then studied for electrochemical reduction of CO2. At a potential of –0.97
V (vs RHE), Ag–N–C0.05M showed an ethanol FE of 42% while the Ag–N–rGO0.05M had
FEC2H5OH of ca. 5%. Instead, the major product using Ag–N–rGOx electrocatalyst was CO
with a maximum FE occurring at –0.57 V (vs RHE) for x = 0.025 M. At this lower potential
bias, Ag–N–Cx materials favored H2 evolution with no C2H5OH production. Based on
comprehensive experimental and theoretical studies, these results showed the product
pathway conversion from CO2-to-CO to the CO2-to-C2H5OH based on catalyst
coordination design.220 Zhang et al. synthesized Ag1 SACs by the thermal transformation
of Ag NPs on nanowire-like MnO2. During the thermal approach, surface reconstruction
of the MnO2 substrate occurred due to strong collisions from the NPs, which was
attributed to be the driving force for SAC formation. When studied for electrochemical
conversion of CO2, they generated a CO faradic efficiency of 95.7% at –0.87 V (vs RHE).
At the same potential, Ag NPs/MnO2 and MnO2 produced FECO less than 65 and 5%,
respectively.221
Single-atom catalysts incorporate many of the advantages from homogeneous and
heterogeneous catalysts. In summary, the SAC motif facilitates 100% atom utilization and
exhibits distinctive geometric and electronic properties due to the absence of metal-metal
bonds and the cationic (occasionally anionic) nature of isolated catalytic sites. However,
the now mainstream field of SACs still present several limitations. Most fundamentally, it
is particularly challenging to develop synthetic methodologies for high uniformity and high
mass loadings on an industrial scale. Moreover, little theoretical and experimental
mechanistic investigations have been explored and the real active sites of SACs have not
yet been recognized. To conclude, even though significantly promising, the field of SACs
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is still in its infancy and several challenges need to be addressed before their proper
implementation is widespread.
A Practical CO2 Electrolytic System
There are five figures of merit when discussing the ideal, CO2 electrolytic device:
current density, faradic efficiency, energy efficiency/applied potential, durability of
equipment, and size of the electrolyzer.222 An industrially applicable CO2 electrolytic
device design should reach high faradic efficiencies (above 90%) at high current densities
(greater than 150 mA cm–2
), using low-cost materials with low overpotential (ηcell < 1 V)
and high energy efficiency (> 50%).223,224 These performance parameters are key to the
commercial success of CO2 electroreduction technologies. An exceptionally efficient
electrocatalyst for CO2 reduction is crucial for improving operating current density and
enhancing faradic energetic efficiencies. Thus, most research in CO2 electroreduction has
focused on the search for superior electrocatalysts.120 However, the cell design has also
harnessed significant attention for improving CO2 reduction rates. Traditionally, CO2
electrolysis is conducted in an H-cell: the anode and cathode electrodes are submerged
in an electrolyte (usually a bicarbonate salt), in which CO2 gas is bubbled into, and where
the two half-cells are separated by a membrane. In 2017, the three-compartment
microfluidic cell demonstrated a total current density of ca. 1 A cm–2 and surpassed the
performance in the H-cell (ca. 1 mA cm–2
).225 The microfluidic cell has been recently
implemented for the evaluation of CO2 reduction electrocatalysts.226–228
Nonetheless, the H-cell and microfluidic cells share intrinsic limitations in their
design: high internal resistance and the low solubility of CO2 on aqueous electrolyte
solutions (~30 mM in H2O at atmospheric pressure).229 This opened a new road for MEA-
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based, zero-gap configuration, two-compartment flow reactors to take the lead. MEAs for
zero-gap configuration using different types of ion-exchanging polymer membranes,
especially anion exchange membranes have been reported.171,230–233 Most notable was
Liu et al.’s report in 2018 were, using Sustanion membranes and ionomers, optimized a
silver cathode electrode (6 mg cm–2
) and achieved 600 mA cm–2 at ca. 3.2 V with FECO >
95%. Through circulation of humidified CO2 at the cathode and 10 mM KHCO3 at the
anode for water management of the membrane, the cell sustained an operation time of
up to 3800 hours at 200 mA cm–2 and 3 V.232 As discussed earlier, Yun et al. reported
competitive results using their AuAg–R/CKB600 electrocatalyst, which displayed a current
density of 618 mA cm–2 at 3.5 V and FECO of 83%. Even though the mass activity of 0.824
A mg–1
for AuAg–R/CKB600 electrocatalyst far surpassed Liu et al.’s Ag NPs (0.1 A mg–1
),
there are other factors worth considering. From a cost-analysis perspective, Yun et al.’s
MEA would be significantly more difficult to commercialize given the incorporation of Au
metal, a more elaborate synthetic methodology, and 10x larger electrolyte concentration.
Reports in the literature have also investigated the cell design, cathode fabrication, and
electrolyte effects in MEA systems for ECO2RR.233–236 However, engineering the
microenvironment at the catalyst-electrolyte boundary (i.e., catalyst, ionomer, electrolyte,
and gases) as a means to obtain the desired ECO2RR product distribution is vital.
Recently for example, the catalyst-ionomer ratio using Ag NPs supported on Vulcan XC72 was studied by Romiluyi et al. They demonstrated a maximized FECO ratio using an
ionomer:catalyst ratio of 3. Additionally, they experimentally demonstrated that reducing
the catalyst loading from 0.1 to 0.01 mgAg cm–2 significantly increased FECO. At a potential
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of –3.2 V, catalyst loadings of 0.1 and 0.01 mgAg cm–2 generated FECO of 78% and 91%,
respectively.
The future of MEA-based, CO2 electrolyzers seems bright. However, research
areas worth investigating remain open. First, given its status as the state-of-the-art
catalyst, most examples in the literature report the use of IrO2 as the anode electrode for
the OER. Research endeavors should aim on incorporating other high-performance OER
electrocatalysts into CO2 electrolyzers to decrease device production costs. Second,
integrating in situ/operando techniques is paramount for understanding mechanistic
pathways (especially when integrating novel cathode and/or anode electrodes) and
optimizing cell design. The complex microenvironment at the electrode surface should be
studied systematically to optimize cathode electrode fabrication and utilization,
understand cation and anion electrolyte effects, and manipulate product selectivity.
Finally, to the best of our knowledge, reports demonstrating the scalability of zero gap
MEA-based, CO2 electrolyzers are strictly limited. Assembly and operation of multilayer
electrolyzer stacks should be further explored to accelerate the industrial implementation
of this technology.
Conclusion
Silver-based catalysts have become incredibly promising for the reduction of CO2 due to
their high performance and tailored selectivity. Additionally, their lower price compared to
other noble metals has propagated their necessity as cathode electrodes in electrolytic
devices. This review discussed several key advancements:
(1) Mechanistic Insights: Studying mechanistic pathways over silver electrocatalysts has
provided fundamental understanding of intermediate species and the role each play
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for product selectivity. The discussion of these mechanisms involving CO2 binding and
the stabilization of reaction intermediates at the catalyst surface have proven
instrumental for the design of efficient electrocatalysts. In particular, the stabilization
of the COOH intermediate is responsible for achieving high selectivity towards the
carbon monoxide product. From both computational and experimental analysis, it is
evident that the electronic structure and surface morphology of silver exert an
unquestionable and complex effect on the catalytic pathway and electrocatalytic
performance.
(2) Design of Silver Electrocatalysts: Nanostructured silver has extensively demonstrated
superb electrocatalytic performance in CO2 reduction due to its large surface area and
variety of morphological characteristics. Specifically, nanoporous structures
(consisting of a 3D interconnected backbone with pores) provide higher surface area,
allow easy transport of electrons, and offer targeted catalytic sites. Surface species
such as oxygen or functional organic groups have also demonstrated improvements
in catalytic activity and selectivity by stabilizing reaction intermediates and
suppressing the undesirable HER.
(3) Alloying with Other Metals: Tuning the binding energy if intermediates and enhancing
reaction kinetics has been presented by alloying silver with other metals, like copper,
indium, and tin. These bimetallic systems show a synergistic effect which enhance
mass activity and selectivity for targeted products. For example, Ag-Cu alloy facilitates
the formation of hydrocarbons and alcohols. Studies reveal that In and Sn improve
selectivity toward CO. Control of composition and distribution of these alloying metals
is essential for tuning catalytic performance toward specific products.
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(4) Support Materials: Given their high surface area, enhanced electrical stability, and
chemical stability, carbon-based supports are the most reported substrates for metallic
nanostructures. Moreover, these types of supports uniformly disperse nanoparticles
and prevent their agglomeration, which leads to high activity over prolonged times.
Doping carbon substrates with heteroatoms, such as N, B, and S has also been
demonstrated to introduce additional catalytic sites and modify the electronic
environment of silver catalysts, resulting in higher reactivity and selectivity. Finally,
metal oxide promoters, like TiO2 and In2O3, bring stabilization of reaction intermediates
and further improve the performance of Ag-based electrocatalysts.
(5) Recent Advances in Cell Design: Zero-gap MEA flow systems are the most promising
configurations in CO2 electrolytic devices for their practical implementation. Some of
the most important features of these electrolytic systems for CO2 conversion include:
(a) reduction of ohmic resistance by minimizing the distance between the anode and
cathode electrodes, (b) enhanced mass transport by allowing direct CO2 feeding into
the catalyst layer, which is integrated directly into the membrane, (c) structural
simplicity and stability that simplify the scale-up process for industrial use, and (d)
wide range of operative conditions including temperature, flow rates, and electrolyte
compositions.
Nevertheless, research areas worthy of investigation remain. First, simple synthetic
methodologies of promising Ag-based electrocatalytic materials should be presented to
enhance their appeal for industrial use and commercialization. Ideal synthesis conditions
will involve optimizing the nanostructure, employing suitable supports, and introducing
dopants and/or mediators to enhance activity, selectivity, and stability. Second, the
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processing, fabrication, and operation of the MEA demands further consideration and
optimization. The multiple components in the sandwich assembly such as ionomer nature,
catalyst loading, and membrane type require further understanding to create an efficient
electrolytic device. Furthermore, in situ techniques should be applied during operation of
the MEA to garner further insight into the mechanisms and their relation to efficiency,
selectivity, and stability. Third, designing cell stacks in CO2 electrolytic systems is
necessary for achieving efficient and scalable CO2 conversion. Cell stacks provide longterm stability and durability ensuring consistent performance and reducing operational
costs over time.
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Chapter 3
Part B: Polyethyleneimine Encapsulated Silver Nanoparticles as
Cathode Electrodes for Syngas Generation
Abstract
This report presents the simple synthesis of Ag NPs embedded in a PEI matrix via green
conditions. The formation of Ag NPs was observed through UV-Vis and the isolated PEIAg nanocomposites were characterized by XRD, TGA, SEM, and high-resolution TEM.
The PEI-750k Ag NPs were studied for the hydrogen evolution reaction in N2-saturated
bicarbonate electrolyte and compared to the state-of-the-art Pt/C electrocatalyst. EIS
modeling results indicated a higher degree of porosity, enhanced double-layer structure,
and improved charge transfer resistance for the PEI-750k Ag NPs cathode electrode at
OCV. At a low overpotential (η < 100 mV), PEI-750k Ag NPs and Pt/C were demonstrated
to be kinetically analogous given their calculated Tafel slopes of 56 and 49 mV/dec,
respectively. Additionally, the Pt/C generated a current density of 10 mA cm–2 at –0.27 V
compared to –0.70 V (vs RHE - iR corrected) for PEI-750k Ag NPs, a potential difference
of only 430 mV. PEI Ag NPs electrodes with varying polymer molecular weights were then
studied in CO2-saturated bicarbonate electrolyte for the electrochemical reduction of CO2.
Independent of Mw, PEI-800, PEI-25k, and PEI-750k produced a current density of 10 mA
cm–2 at moderate overpotentials (η ≈ 600 mV) and produced H2 and CO as the sole
products after constant potential electrolysis. Finally, CO2 sorption-desorption studies
reveal that PEI Ag NPs can circumvent CO2 solubility limitations in aqueous systems by
exposing more CO2 to the interface of the electrode.
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Introduction
Given increasing energy demands and our dependance on fossil fuels, major
interests have been devoted towards innovative technologies for CO2 capture and
utilization.237,238 From both an environmental and economic lens, ECO2RR remains as
the industrially preferential method for remediation of excess carbon dioxide into its valueadded products.84,239,240 Specifically, the economic value of 2-to-1 H2 : CO syngas ratio
is substantial due to its crucial necessity as a chemical feedstock in Fischer-Tropsch and
methanol synthesis.241 Au and Ag have been demonstrated to be selective catalysts
towards the generation of syngas from CO2.
242–244 Given the higher costs for Au, however,
Ag-based materials remain the state-of-the-art electrocatalysts for reduction of CO2.
Thus, research efforts have mostly focused on synthesizing and investigating novel Ag
nanostructures with enhanced catalytic activity for this important product.122
NPs offer distinct advantages compared to other heterogeneous electrocatalysts
due to their unique physical and chemical properties. Their electronic properties can be
precisely tuned by manipulating size, shape, and composition. A high surface-to-volume
ratio provide more active sites and the larger presence of edges, corners, and defects
allow enhanced reactivity and selectivity.146,245,246 Additionally, Ag NPs have been
engineered on different supports and matrix materials to enhance stability and
durability.179,180,182 Nonetheless, Ag NPs present several limitations for their application in
an industrial scale. Firstly, many reports of these electrocatalytic systems often require
complex synthetic methodologies to achieve optimal performance. The use of anchoring
agents for immobilization onto carbon supports, the modification of ligands at the catalyst
surface, and the application of precise reaction conditions are challenging processes to
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scale up.143,148,150 Secondly, Ag NPs can aggregate over prolonged electrolytic times
resulting in catalyst degradation and thus, reduced activity and selectivity. Proper catalyst
supports must be taken into consideration to address this issue.247 Thirdly, because of the
low solubility of CO2 in aqueous electrolytes, high Ag catalyst loadings are sometimes
needed to achieve acceptable performances; a common challenge in ECO2RR
processes.248 Coupling carbon capture and ECO2RR can improve energy efficiency and
decrease overall production costs.249
Herein, we present the simple, green synthesis of PEI Ag NPs and their first studies
as cathode electrodes for the HER and ECO2RR for syngas generation.
Polyethyleneimine (PEI) has been shown to facilitate the reduction and stabilization of Ag
NPs due to its high density of amino functionalities. The cationic polymer provides a
positive surface charge, which prevents aggregation and enhances colloidal stability.250
The tunable degree of branching in PEI allows for tunability of Ag NP size and dispersity,
which has been optimized for both catalytic and antimicrobial applications.251–253
Additionally, the CO2 absorption capabilities of PEI-functionalized materials are well
documented in the literature.254–256 We believe PEI Ag NPs electrodes have the potential
to synergistically couple CO2 absorption (capture) and its electrochemical reduction
(utilization) into valuable products. Even though the electrocatalytic properties of PEI Ag
NP composites have been shown for different transformations, their application in
ECO2RR, to the best of our knowledge, has not yet been explored.257–260
Experimental Section
Synthesis of PEI encapsulated Ag NPs.– In a 20 mL glass vial, 0.84 g of branched PEI
solution (average Mn ~60,000 by GPC, average Mw ~750,000 by LS, 50 wt. % in H2O;
86
Aldrich) were dissolved in 3 mL DI H2O (18.2 MΩ cm).261 Then 42 mg of Vulcan carbon
(XC-72R; FuelCell Store) were added into the mix and stirred to ensure complete
dissolution. Meanwhile, 0.42 g of AgNO3 (≥ 99%; Oakwood Chemical) were dissolved in
1 mL DI H2O. The AgNO3 solution was slowly added to the PEI-C mixture and allowed to
complex by stirring for an additional 15 min. During this complexation process, 0.99 g Lascorbic acid (Aldrich) were dissolved in 5 mL DI H2O. One mL of the ascorbic acid
solution was added into the PEI–Ag–C complex for its chemical reduction into Ag NPs
and allowed to rest overnight. The PEI-750k Ag NPs were transferred to a centrifuge tube
and washed with DI H2O. The washing process was repeated a total of three times and
finally dried in oven at 100°C for 12 h. The isolated material was a black powder.
Analogues synthetic conditions were used for two additional branched PEI materials
(average Mn ~10,000 by GPC, average Mw ~25,000 by LS; average Mn ~600 by GPC,
average Mw ~800 by LS; Aldrich) but using 0.42 g instead.
Characterization of PEI Ag NPs.– The prepared catalysts were analyzed by UV-Vis
(Perkin Elmer Lambda 950), thermogravimetric analysis (Shimadzu TGA-50), X-ray
diffraction (XRD; Rigaku Ultima IV powder/thin film diffractometer), scanning electron
microscopy (SEM, FEI Nova NanoSEM 450), and transmission electron microscopy
(TEM) and high-resolution TEM (HRTEM; FEI Talos F200C).
Preparation of working electrode.– Catalyst inks were prepared by dissolving 30 mg of
PEI Ag NPs powder in 800/200 μL isopropanol/DI H2O solution. Then, 200 μL of
LIQUion™ 1105 binder solution (5% wt. Nafion; Ion Power) were added into the mixture
followed by sonication for three cycles of 480 s each. The catalyst ink was then
dropcasted onto the respective carbon substrate (glassy carbon electrode, IKA; Toray
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carbon paper 060 or Sigracet 36 BB carbon paper, FuelCell Store) and dried in oven at
60°C for 12 h. The active are of for each electrode was 4 cm2 and the catalyst loading
were 37.6, 34.4, and 32.3 mg for PEI-750K, PEI-25K, and PEI-800, respectively. The
exact methodology was followed to prepare the 5% Pt/C (FuelCell Store) electrodes.
Electrochemical measurements.–. Electrochemical testing was performed by using a
VersaSTAT3 Potentiostat Galvanostat (AMETEK Scientific Instruments) in a twocompartment, customized H-cell. Platinum wire and Ag/AgCl (saturated KCl) were used
as the counter electrode and the reference electrode, respectively. A Nafion 117 PEM
separated the catholyte and the anolyte of 125 mL each. The electrolyte solution of 0.5 M
KHCO3 was purged with either high purity N2 or CO2 gas for 45 min at 40 cc/min.: pH
values of 8.4 and of 7.4 were measured after saturation, respectively. Constant potential
electrolysis was performed at six different potentials for 15 min. Gaseous products were
collected after electrolysis at each applied potential and were then analyzed via gas
chromatography (GC, 1177 injector; Varian) equipped with a thermal conductivity detector
(TCD; at 100°C). The measured potentials were reported versus the reversible hydrogen
electrode (RHE) following the equation:
E (vs RHE) = E(vs Ag/AgCl) + 0.197 V + 0.0591 V × pH
CO2 adsorption-desorption studies. Samples previously dried in oven at 100°C were
loaded (15-21 mg) in a platinum crucible, placed in the instrument’s balance, and treated
with an N2 flow at 110°C for 30 min. The furnace temperature was then decreased to 25°C
followed by a gas flow switch of 100% CO2 for 3 h. After the elapsed time, a N2 gas flow
switch at 85°C followed for 30 min before the second and third cycle measurements. The
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change in sample mass was recorded over time to determine absorption capacity at 25,
55, and 85°C.
Results and Discussion
Physical Characterization.– UV-Vis data demonstrates the formation of silver
nanoparticles on the polyethyleneimine matrix. The blue absorption band represents that
of the PEI-Ag+ complex, in which the amino groups in PEI behave as an electron-pair
donor to chelate silver.262 After addition of ascorbic acid, a blue shift with maxima of 409
nm is observed. This intense absorption peak at around 400 nm indicates the surface
plasmon resonance of PEI Ag NPs.252 This also shows that Ag NPs are dispersed
uniformly in PEI. It has been demonstrated that as nanoparticle size increases, the
plasmon absorption band red shifts.250 Finally, the inset image in the spectra validates the
dispersity of this silver nanocomposite. Vulcan carbon was added into the PEI–Ag+
mixture before chemical reduction to Ag NPs help with isolation of the catalyst powder
and to enhance its conductive properties. XRD patterns of silver nanoparticles on different
PEI molecular weights are demonstrated in Fig. 12. The diffraction peaks at 38.2, 44,4,
Fig. 12 UV-Vis absorption spectra of PEI-750k-Ag+ complex (blue curve) and PEI Ag NPs (black
curve) after addition of the reducing agent; X-ray diffraction patterns of PEI Ag NPs with varying
PEI Mw of 800 (red), 25,000 (yellow), and 750,000 (green).
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64.5, 77.5, and 81.5 degrees were assigned to reflections from the (1 1 1), (2 0 0), (2 2
0), and (3 1 1) planes of metallic silver, respectively.251 Crystallite size was calculated
using the Debye-Scherrer equation for the (1 1 1) peak given its sharp intensity. The
crystallite size of 16.3, 14.6, and 12.3 nm corresponded to PEI-750k, PEI-25k, and PEI800. Thus, a relationship was determined in which decreasing the molecular weight of
PEI led to the formation of smaller crystallites.
The surface morphology of PEI-750k Ag NPs were observed by scanning electron
microscopy, as seen in Fig. 13. At low magnification, the porosity of the nanocomposite
can be appreciated. Differentiating between the Vulcan carbon particles (20-50 nm) to Ag
NPs (12-16 nm by XRD) through SEM imaging proved difficult. However, regions with
more brightness indicates areas of more electron density, which are likely due to higher
concentrations of silver at that location. EDS is necessary to differentiate between these
two components within the nanocomposite. At higher magnifications, the nanoporous
morphology can still be appreciated: highly porous materials are often desired given their
increased surface area. The microstructure of the PEI-750k Ag NPs was then studied
through TEM. As demonstrated in the first micrograph in Fig. 14, the high-contrast
spheres correspond to Ag NPs due to their higher electron density compared to the
polymer-carbon matrix. Additionally, silver NPs are uniformly dispersed across the porous
network. The high-resolution TEM (HR-TEM) image of a silver nanoparticle
representative shows lattice fringes with d-spacing of 2.31 Å. This value was attributable
to the (1 1 1) crystallographic plane.263
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Hydrogen Evolution Reaction.– After physical characterization, we tested the
electrocatalytic properties of PEI-750k Ag NPs towards HER in an H-cell. PEI-750k Ag
NPs were drop cast onto Toray carbon paper (TCP) and compared to the state-of-the-art
Pt/C in KHCO3. The anode chamber constituted of a Pt wire for the OER, and the cathode
chamber contained both the Ag/AgCl reference electrode and PEI Ag NPs working
electrode. For a more precise comparison of electrocatalytic activity, we conducted EIS
analysis to correct for the voltage loss (i.e., iR drop) caused by the electrolyte solution
between the reference and working electrode. Results from both galvanostatic and
potentiostatic EIS (GEIS and PEIS, respectively) modes were validated by KramersKroning (K-K) relations tests in the applied frequency range of 10 kHz to 1 Hz. In K-K
tests, residuals represent the difference between the measured impedance data and the
impedance predicted by the K-K relations. A lower sum of residuals (|Δ|max) generally
indicates reliable data, accurate fitting model, and that the system is stable and linear.
Fig. S21 demonstrates the Nyquist plots under galvanostatic mode at the open circuit
voltage (OCV) for (a) PEI-750k Ag NPs and (b) Pt/C at different sinusoidal current
perturbations and their corresponding K-K relations. For PEI-750k Ag NPs, the smallest
Fig. 13 SEM micrographs PEI-750k Ag NPs at magnifications of a) 65,000x and b) 150,000x with
accelerating voltage of 10 kV.
a
)
b
)
91
value of |Δ|max was demonstrated at perturbation amplitudes of 10, 25, and 50 mA. For 25
and 50 mA, the real and imaginary parts contributed near equal to discrepancies from the
expected value. In the case of Pt/C, the perturbation of 10 mA showed the highest |Δ|max,
which indicated poor compliance with the theoretical model. The perturbations at 1 mA
and 50 mA are considered acceptable for further EIS analysis. Figure S22 demonstrates
the Nyquist plots under potentiostatic mode at OCV for PEI-750k Ag NPs and Pt/C and
their corresponding K-K relations. For PEIS, perturbation amplitudes of 1, 10, 25, and 50
mV were applied to the system. The smallest |Δ|max values were determined at 10 mV for
PEI-750k Ag NPs and at 25 mV for Pt/C, as demonstrated in Fig. S23.
Given the relatively similar |Δ|max for the electrocatalysts at perturbation amplitudes
of 50 mA and 50 mV, the GEIS and PEIS plots were then modeled to the simplified
Randles circuit, and the results were tabulated for comparison, as shown in Fig. 15 and
Table 4. The Rs values for PEI-750k Ag NPs and 5% Pt/C catalysts in N2-saturated 0.5 M
KHCO3 electrolyte show distinct differences, highlighting the ionic conductivity variations
between the two materials. Overall, PEI-750k Ag NPs exhibited smaller Rs values in both
GEIS and PEIS modes compared to Pt/C, indicating better ionic conductivity.
Fig. 14 TEM image of the PEI-750k Ag NPs material and HR-TEM of an Ag NP representative
showing the calculated d-spacing.
92
Nonetheless, despite intrinsically different, the Rs values for both electrocatalysts
consistently fall within the 4-6 Ω range. More notable are the differences in charge transfer
resistance (Rct) between the two electrode materials. In GEIS, PEI-750k Ag NPs
demonstrated a significantly smaller Rct compared to that of Pt/C, indicating faster
reaction kinetics. This kinetic improvement can be attributed to the protonation of primary
amine functional groups in PEI-750k Ag NPs at the electrode interface, which facilitates
electron transfer.257 The net positive charge of PEI Ag NPs has been confirmed through
zeta potential measurements, suggesting enhanced interactions with negatively charged
species in the electrolyte.250,262 Even though opposing Rct results are demonstrated under
PEIS mode, the high percent error suggests model unreliability. In GEIS mode, PEI-750k
Ag NPs demonstrates significant non-ideal capacitive behavior compared to Pt/C based
on CPE-P values (0.371 and 0.642, respectively). This can be attributed to the high
degree of the porosity of the electrocatalyst as observed in TEM and SEM images.
Additionally, the higher CPE-T suggests a larger double-layer capacitance.264,265 The Rs
value under GEIS mode demonstrated lower error percentages for PEI-750k Ag NPs and
Pt/C compared to those outputted by the model under PEIS mode. Therefore, the values
of 5.468 and 5.991 were used for iR correction of PEI-750k Ag NPs and Pt/C, respectively.
The iR corrected LSVs under N2–saturated 0.5 M KHCO3 are observed in Fig. 16a. As
depicted, the electrochemical activity of Pt/C intensifies at a voltage of zero (vs RHE) and
quickly reaches 10 mA cm–2
(E = –0.27 V). In the case for the PEI-750k Ag NPs electrode,
the current enhancement begins at around –0.5 V and rapidly arrives at 10 mA cm–2 at (E
= –0.70 V). Thus, when compared to the state-of-the-art Pt/C under the exact same
conditions, PEI-750K Ag NPs demonstrates a potential difference of 430 mV to reach an
93
acceptable current density. Even though EIS modelling of PEI-750k Ag NPs calculated a
much smaller charge transfer resistance, it also displayed more non-ideal capacitive
behavior compared to Pt/C. This could explain why the electrocatalytic activity of Pt/C
surpasses that of PEI-750K Ag NPs at higher overpotentials.
Fig. 15 Nyquist plots for 5% Pt/C (blue) and PEI-750k Ag NPs (green) electrocatalysts in N2-
saturated 0.5 M KHCO3 electrolyte. The data was recorded over a frequency range from 1 kHz to
1 Hz at OCV. The left panel presents the galvanostatic mode with a perturbation amplitude of 50
mA, while the right panel shows the potentiostatic mode with a perturbation amplitude of 50 mV.
The experimental data are fitted to a Randles circuit model (inset), as shown by the black curves.
Table 4 EIS results derived from fitting the experimental data to the Randles circuit model for Pt/C
and PEI-750k Ag NPs catalysts in N2-saturated 0.5 M KHCO3 electrolyte. The parameters include
the solution resistance (Rs), the constant phase element-T (CPE-T), the constant phase elementP (CPE-P), and the charge transfer resistance (Rct). Results represent mean ± standard deviation
for both GEIS and PEIS mode.
Mode: Cat Rs (Ω) CPE-T (F cm- 2·sφ-1) CPE-P Rct (Ω)
GEIS: PEI-750k Ag NPs 5.468 ± 0.028 0.025 ± 0.000398 0.371 ± 0.004 9.607 ± 0.107
GEIS: 5% Pt/C 5.991 ± 0.049 0.003 ± 0.00007334 0.642 ± 0.005 37.080 ± 0.454
PEIS: 5% Pt/C 5.002 ± 0.043 0.079 ± 0.0004561 0.227 ± 0.003 3.944E+08 ± 6.982E+13
PEIS: PEI-750K Ag NPs 4.528 ± 0.031 0.023 ± 0.00003744 0.260 ± 0.001 1.255E+10 ± 1.196E+15
94
Figure 16 displays the derived Tafel slope (TS) at three overpotential HER regions
for PEI-750k Ag NPs and Pt/C in N2-saturated 0.5 M KHCO3. TS of 49 and 55 dec–1 were
observed under regions of low overpotential (η < 100) mV for Pt/C and PEI-750k Ag NPs,
respectively. These TS parallel those for the H3O+
reduction mechanism in acidic media
for platinum catalysts.266 When increasing the overpotential higher than 100 mV, a shift in
TS to 146 and 169 mV dec–1 was observed for Pt/C and PEI-750k Ag NPS, respectively.
This behavior demonstrates that the TS is potential dependent.267 At even higher
overpotentials (η > 200 mV), the TS drastically increase to 438 for Pt/C and 500
mV dec–1
for PEI-750k Ag NPs. The similarities in TS and exchange current density as
demonstrated in Table S13, validate that PEI-750k Ag NPs is kinetically analogous to the
state-of-the-art Pt/C. Interestingly, a two-step reduction mechanism was observed for the
Pt/C electrode under this near-neutral, aqueous system at higher overpotentials, as
observed in Fig. 16. At approximately η = 500 mV, Pt/C undergoes a mechanistic change
with enhanced kinetics compared to those of the PEI-750k Ag NPs electrode. The more
sluggish kinetics of PEI-750k Ag NPs leads to a larger current density generation by Pt/C.
Katsounaros et al. attributed the two-step mechanistic insight to the formation of a local
pH gradient. The current-potential relationship near-neutral pH is strictly determined by
the diffusion limitation of H+ and OH-
, as reproduced by the Nernst-Planck equation. They
claimed that all the reduction current contributions originate from hydronium ion (H3O+
)
reduction.268 Strmcnik et al. agree that the initial reduction events at low overpotentials
are due to diffusion-limited hydronium reduction. But with the increase of cathodic
overpotentials, rapid consumption of H3O+ at the liquid-solid interface shifts HER into a
diffusion-controlled process (i.e., limited by mass transport of H3O+ ions) causing the
95
reduction current to reach a plateau.269 The second step occurs exclusively at higher
overpotentials in which the primary reactants shift from H3O+
ions to H2O molecules,
leading to a steady increase in the reduction current.270 However, potassium
bicarbonate’s function is not exclusively pH regulation, it may also be considered a
possible proton donor via its own reductive pathway.271 Unfortunately, detailed
understanding of the HER mechanism in neutral or near-neutral electrolytes is still in its
early stages because these conditions were deemed unfavorable for electrochemical
water splitting.272,273 More intricate experiments and microkinetic modeling are needed to
understand proton sources for the mechanism of H2 generation.273,274 Finally, 15 min
electrolysis experiments using the PEI-750k Ag NPs electrode were performed at six
different potentials to confirm hydrogen as the sole product in 0.5 M KHCO3 electrolyte.
The chromatograms in Fig. S24 demonstrate three gas signals at retention times of 2.21,
2.51, and 5.42 min, which constitute to H2, N2, and CO2, respectively. As expected, even
at the lowest applied potential of E = –0.31 V, PEI-750k Ag NPs exclusively generated
H2 gas.
Electrochemical CO2 Reduction Reaction.– After HER studies, we tested PEI Ag NPs
for ECO2RR by saturating the 0.5 M KHCO3 electrolyte with high purity CO2. A glassy
carbon electrode (GCE) was first tested as the substrate support for PEI Ag NPs and an
initial reduction potential window of 0 V to –1.20 V (vs RHE) was screened. The LSVs for
pristine GCE and PEI-750k Ag NPs supported on GCE from 0 to –0.80 V in CO2-saturated
0.5 M KHCO3 are shown in Fig. S28. At a scan rate of 10 mV s
–1
, current densities of 1.60
and 4.22 mA cm–2 were generated by pristine GCE and PEI-750k Ag NPs supported on
GCE, respectively. This difference in current density was attributed to the electrochemical
96
activity of the PEI Ag NPs. Given this insight, constant potential electrolysis was
performed for 15 min at six different potentials. Figure S31 shows the chromatograms for
pristine GCE and PEI-750k Ag NPs after electrolysis for all measured potentials. The
peaks at 2.19, 2.53, and 5.33 min correspond to H2, N2, and CO2, respectively. At –0.87
V, a new peak with retention time of 2.70 min can be observed for the PEI-750k Ag NPs,
which corresponds to the CO product. The CO signal intensified at more negative
potential biases. As observed, pristine GCE generated CO at potentials more negative
than –1.12 V. This finding must be taken into consideration when using a GCE for studying
Fig. 16 a) LSVs after iR correction of PEI-750k Ag NPs (green) and Pt/C (blue) in N2-saturated
0.5 M KHCO3 at a scan rate of 10 mV s
–1
, Tafel plots of c) PEI-750k Ag NPs and d) Pt/C at three
overpotential regions, and d) Tafel analysis at higher overpotentials demonstrating the more
sluggish kinetics (inset) of PEI-750k Ag NPs compared to Pt/C.
a
)
b
c d
97
ECO2RR, specifically when calculating faradic efficiency and partial current density of CO.
Additionally, pristine GCE also generated H2 at the lowest reduction potential of –0.37 V.
The complications to effectively commercialize an electrolytic device which
employs a GCE are significant. Therefore, two gas diffusion electrodes (GDEs) were
studied as cathode electrodes: microporous layered carbon paper (MPL) and TCP. The
MPL carbon paper provides a PTFE (5% wet-proofed) coating for increased water
repulsion and allows for improved adhesion of the catalyst. The LSVs from 0 to –0.8 V of
pristine MPL and PEI-750k Ag NPs on MPL in CO2-saturated 0.5 M KHCO3 are shown in
Fig. S29. Current densities of 3.06 and 6.99 mA cm–2 were produced by pristine MPL and
the PEI-750k Ag NPs electrode, respectively. As observed in the chromatograms in Fig.
S31, pristine MPL did not participate in ECO2RR. It was only at the highest reduction
potential bias where a relatively negligible CO signal emerged. The chromatograms for
the PEI-750k Ag NPs on MPL carbon paper, however, effectively demonstrated a CO
signal at 2.70 min. The new peak emerged after 15 min electrolysis at –0.62 V and
became pronounced at more reductive potentials. Toray carbon paper 060 was then
studied as the carbon support for PEI-750k Ag NPs. The LSV comparisons between
pristine TCP and PEI-750k Ag NPs in CO2-saturated 0.5 M KHCO3 at various scan rates
are shown in Fig. S30. At –0.80 V, current densities of 1.61 and 7.80 mA cm–2 were
demonstrated by pristine TCP and PEI-750k Ag NPs on TCP electrode. Pristine TCP did
not yield the CO product, irrespective of the applied potential bias, seen in Fig. S33.
However, like the GCE and MPL supports, PEI-750k Ag NPs deposited on TCP also
generated a CO signal at 2.70 min after electrolysis at –0.62 V. Interestingly, at reductive
potentials larger than ca. –1.00 V, the electrocatalytic activity of pristine TCP was rapidly
98
enhanced and the current density separation to the PEI-750k Ag NPs electrode became
minimized. This increase in current density of pristine TCP electrode at higher potentials
was attributed to improved H2 generation. Nonetheless, GC data indicated that PEI-750k
Ag NPs can electrochemically reduce CO2 to CO.
Given the larger current density for PEI-750 Ag NPs on TCP electrode, its activity
was compared to polycrystalline Ag foil. Fig. 17a demonstrates the LSVs of PEI-750k Ag
NPs and its polycrystalline analogue. At a potential of –0.63 V, the silver nanoparticles
polymer composite generated 10 mA cm–2 while the Ag foil only produced 1.60 mA cm–2
,
a
)
b
)
c
)
d
)
Fig. 17 a) LSV curves of pristine TCP (dashed), polycrystalline Ag foil (blue) and PEI-750k Ag
NPs (green) in CO2-saturated 0.5 M KHCO3 at a scan rate of 50 mV s
–1
, b) LSVs of PEI-750k Ag
NPs at different scan rates, c) plot of the square root of scan rate vs. peak cathodic current density
and d) plot of log scan rate vs log peak cathodic current density for PEI-750k Ag NPs at three
different potential regions.
99
a 6.3-fold enhancement. This elucidates the high electrocatalytic activity of PEI-750k Ag
NPs. Additionally, the generated current density demonstrates a relationship with the
applied scan rate (ν; Fig. 17b) To gain insight into the mechanism occurring at the
electrode-electrolyte interface, three different potential regions were explored: high (E =
–0.6 V), medium (E = –0.4 V), and low potentials (E = –0.1 V). The plot of square root
of ν vs. peak current density (jp) follows a linear relationship for all three regions, which is
consistent with the Randles-Sevcik equation, shown in Fig. 17c. This plot confirms a
diffusion-controlled mechanism in the three regions. However, the plot of log v vs log ip as
seen in Fig. 17d, demonstrates fluctuation from exclusive diffusion control, in which the
ideal slope is α = 0.5. At the low potential region, the large slope of α = 0.72 (R2 = 0.999)
suggests that both diffusion and adsorption (mixed-controlled) contribute to the
mechanism. When the potential increases to –0.4 V, however, the mechanistic
contributions are predominantly diffusion-controlled given α = 0.50 (R2 = 0.986).
Interestingly, at the high potential region of –0.6 V, the small slope of α = 0.16 (R2 = 0.995)
suggests that the mechanism shifts to more complex phenomena.
We then studied PEI Ag NPs with different polymer molecular weights to
understand structure-activity relationships. The difference in Mw stems from variations in
the degree of branching, which has been reported to have different effects on the physical
and chemical properties in PEI.275,276 Fig. 18 shows the LSVs of Ag NPs in PEI of varying
molecular weights (750k, green; 25k, yellow; and 800, red) in CO2-saturated 0.5 M
KHCO3 at different scan rates. The electrochemical activity of the electrode was
independent of polymeric Mw, which reveals that the measured current density can be
primarily attributed to Ag NPs. The chromatograms in Fig. SI35 demonstrate product
100
analysis after constant potential electrolysis at three potentials. As observed, the three
electrocatalysts generate H2 and CO at –0.62 V with a total current density of
approximately 6 mA cm–2
. This enhancement in current density compared to pristine TCP
(~0.1 mA cm–2
), demonstrates that PEI Ag NPs electrodes can effectively generate
syngas.
CO2 adsorption-desorption studies.– One of the major challenges in electrochemical
CO2 reduction is the low solubility of carbon dioxide in aqueous electrolytes (~32 mM),
which often leads to mass transport limitations and therefore, poor energy efficiency.
Multiple strategies have been explored to bypass the solubility problem including the use
of GDEs and incorporating more complex electrolytes, such as ionic liquids. Recently,
research efforts have focused on coupling CO2 capture and conversion in a single
electrochemical cell to improve carbon utilization and save production energy.277 Given
that PEI-based materials have been demonstrated to be effective CO2 sorbents due to
their high uptake/release capacity and excellent thermal stability, we sought to study PEI
Ag NPs as a possible CO2 capture and utilization pathway. Figure S36-S38 present the
sorption-desorption isotherms for PEI Ag NPs at different temperatures in 100% carbon
b
)
a
)
Fig. 18 LSVs of Ag NPs in PEI with varying molecular weights (750k, green; 25k, yellow; 800,
red) in CO2–saturated 0.5 M KHCO3 at a) 10 mV s
–1 and b) 50 mV s
–1
.
101
dioxide. At room temperature, PEI-800, PEI-25k, and PEI-750k had a mass increase of
0.40, 0.17, and 0.1%, respectively. The mass change at elevated temperatures of 55 and
85°C decreased accordingly for the three electrocatalysts, indicating a degradation in CO2
uptake capacity. Nonetheless, these results demonstrate the thermal stability of PEI Ag
NPs, and that decreasing Mw of the polymer increases the CO2 sorption capacity of the
material.275 The CO2 uptake of the material can present more species available for
reduction at the electrode-electrolyte interface, further promoting ECO2RR.
Conclusion
This report demonstrated the simple and green synthesis of PEI encapsulated
silver nanoparticles and their application as cathode electrodes for syngas generation via
HER and ECO2RR studies. PEI-750k Ag NPs exhibited electrocatalytic activity
comparable to the state-of-the-art Pt/C. At OCV, K-K relations validated galvanostatic EIS
data demonstrated minimized charge transfer resistance for PEI-750k Ag NPs.
Additionally, at the two electrodes were shown to be kinetically analogous given their
respective Tafel slopes at overpotentials less than 250 mV. At higher overpotentials,
however, the electrocatalytic activity of Pt/C surpassed that of PEI-750k Ag NPs, which
was attributed to its enhanced behavior as an ideal capacitor (CPE-P) as calculated
through EIS. Nonetheless, a potential difference of only 430 mV was needed for PEI-750k
Ag NPs to achieve a current density of 10 mA/cm2
in N2-saturated 0.5 M KHCO3. In CO2-
saturated 0.5 M KHCO3, PEI-750k Ag NPs showed a six-fold current density
enhancement compared to polycrystalline Ag foil. Lower PEI Mw of 25k and 800 also
exhibited similar current density generation, reaching 10 mA cm–2 at overpotentials of
approximately 600 mV. Constant potential electrolysis at different potentials
102
demonstrated H2 and CO as the sole products in CO2-saturated electrolyte. Finally, CO2
adsorption-desorption measurements indicate the electrocatalysts’ potential to
circumvent CO2 solubility limitations by increasing CO2 near the electrode surface.
Future Directions
Figures of Merit.– A principal objective of this research is to precisely determine product
formation after constant potential electrolysis experiments. Firstly, a calibration curve of
both H2 and CO should be generated via GC to ensure precise quantification of gaseous
products generated after each applied potential. The mol of H2 and CO products are
critical to calculate important figures of merit including product selectivity, partial current
density, and faradic and energy efficiency. To generate syngas with a 2:1 ratio of H2 to
CO, the electrocatalyst should exhibit a selectivity of 67% and 33%, respectively.
Secondly, even though silver-based electrocatalysts are selective towards the CO
formation pathway, aliquots of the electrolyte should be analyzed via nuclear magnetic
resonance (NMR) to ensure that liquid byproducts (i.e., formate/formic acid) are not
formed by PEI Ag NPs.
Optimization of Electrolytic Device.– The H-cell system needs further optimization by
integrating continuous bubbling of CO2 gas into the electrolyte: constant CO2 bubbling is
critical for achieving high faradic efficiency and performance. Overall, this system should
be integrated into a zero-gap MEA-based electrolyzer. The zero-gap configurations
continue gathering significant research interest due to their ability to achieve high current
densities, enhance mass and charge transfer, and improve catalyst utilization. Zero-gap
devices are a promising approach for scalable and cost-effective CO2 conversion
technologies.
103
Catalyst Optimization.– Given their increased surface-to-area ratios, smaller
nanoparticles have been reported to enhance reactivity in numerous electrochemical
reactions. Therefore, the synthetic methodology of PEI Ag NPs should be further studied
to ideally promote the formation of smaller Ag NPs. It should be noted that PEI can act as
a reducing agent at elevated temperatures. Additionally, linear PEIs can also be explored
as a supporting matrix for Ag NPs and their electrochemical activity for syngas generation
compared to those of respective branched analogues. The isolation and electrochemical
studies of other PEI metal NPs, such as copper, remain an open area worth exploring.
104
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132
Appendix
Supplementary Information I
Synthesis of Ni–TiO2 Nanocomposites as Enzyme–less, Amperometric
Sensors for the Electrooxidation of Glucose
133
Ti K O K Ni K C K
Fig. S1 a) SEM image and EDS elemental mapping of the 20% Ni–TiO2/XC72R nanocomposite.
a
134
Fig. S1 b) SEM image and EDS elemental mapping of the 40% Ni–TiO2/XC72R nanocomposite.
Ti K O K Ni K C K
135
Fig. S1 c) SEM image and EDS elemental mapping of the 80% Ni/XC72R nanocomposite.
Ti K O K Ni K C K
136
Fig. S2 a) EDS spectrum and quantification of the 20% Ni–TiO2/XC72R nanocomposite.
137
Fig. S2 b) EDS spectrum and quantification of the 40% Ni–TiO2/XC72R nanocomposite.
138
Fig. S2 c) EDS spectrum and quantification of the 60% Ni–TiO2/XC72R nanocomposite.
139
Fig. S2 d) EDS spectrum and quantification of 80% Ni/XC72R nanocomposite.
140
Fig. S3 ICP-OES measurements for 20% Ni–TiO2/XC72R, 40% Ni–TiO2/XC72R, 60% Ni–
TiO2/XC72R, and 80% Ni/XC72R at a) 221.647 nm and b) 231.604 nm.
a
b
141
Fig. S4 XRD patterns for the four Ni–TiO2 nanocomposites: 20% Ni–TiO2/XC72R (green), 40%
Ni–TiO2/XC72R (blue), 60% Ni–TiO2/XC72R (red), and 80% Ni/XC72R (gray). Double dagger (‡)
signals correspond to TiO2 anatase and asterisk (*) signals correspond to the Ni nanoparticles.
142
Fig. S5 XPS survey spectra four Ni–TiO2 nanocomposites: 20% Ni–TiO2/XC72R (green), 40%
Ni–TiO2/XC72R (blue), 60% Ni–TiO2/XC72R (red), and 80% Ni/XC72R (gray).
143 Table S1 Ni wt. % quantification and comparison using (yellow) EDS, (green) ICP–OES and
(orange) XPS for the four synthesized nanocomposites.
Catalyst C O Ti Ni
Ni
221.647
nm
Ni
231.604
nm
Ni
Ni20-(TiO2)60 18.22 16.58 41.17 24.02 16.89 19.06 11.93
Ni40-(TiO2)40 18.96 10.64 25.09 45.30 39.03 41.11 16.58
Ni60-(TiO2)20 18.40 6.74 16.38 58.48 58.99 60.86 21.13
Ni80 19.20 0.10 0.10 80.60 77.63 79.34 34.63
Table S1 Ni wt. % quantification and comparison using (yellow) EDS, (green) ICP–OES and
(orange) XPS for the four synthesized nanocomposites.
Catalyst C O Ti Ni
Ni
221.647
nm
Ni
231.604
nm
Ni
Ni20-(TiO2)60 18.22 16.58 41.17 24.02 16.89 19.06 11.93
Ni40-(TiO2)40 18.96 10.64 25.09 45.30 39.03 41.11 16.58
Ni60-(TiO2)20 18.40 6.74 16.38 58.48 58.99 60.86 21.13
Ni80 19.20 0.10 0.10 80.60 77.63 79.34 34.63
144
Fig. S6 a) Activation of the 20% Ni–TiO2/XC72R electrode in 0.1 M NaOH at a scan rate of 50
mV s
–1
from 0 to 0.6 V.
145
Fig. S6 b) Activation of 40% Ni–TiO2/XC72R electrode in 0.1 M NaOH at a scan rate of 50 mV
s
–1
from 0 to 0.7 V.
146
Fig. S6 c) Activation of 60% Ni–TiO2/XC72R electrode in 0.1 M NaOH at a scan rate of 50 mV
s
–1
from 0 to 0.8 V.
147
Table S2 Summary of the 60% Ni–TiO2/XC72R nanocomposite activation in 0.1 M NaOH for
different potential windows (vs MMO) at a scan rate of 50 mV s
–1
E window (mV) cycles
j pa
(mA/cm2
)
600 90 0.69
700 45 8.80
800 30 12.65
900 8 13.83
148
Fig. S7 a) Activation after 8 cycles of 20% Ni–TiO2/XC72R electrode in 0.1 M NaOH at a scan
rate of 50 mV s
–1
from 0 to 0.9 V.
149
Fig. S7 b) Activation after 8 cycles of 40% Ni–TiO2/XC72R electrode in 0.1 M NaOH at a scan
rate of 50 mV s
–1
from 0 to 0.9 V.
150
Fig. S7 c) Activation after 8 cycles of 80% Ni/XC72R electrode in 0.1 M NaOH at a scan rate of
mV s
–1
from 0 to 0.9 V.
151
Fig. S8 Potentiostatic measurements for the Ni–TiO2/XC72R nanocomposites in 0.1 M NaOH at
0.7 V or 350 s.
152
Fig. S9 Amperometric studies for 60% Ni–TiO2/XC72R electrode at 0.7 V in 0.1 M NaOH (black
line) and 0.1 M NaOH + 0.1 M NaCl (green line) and its calibration curve with error bars (n = 3).
153
Supplementary Information II
NiSn Alloys as Electrocatalysts for the Urea Oxidation Reaction and
Integration into Direct Urea Fuel Cells
154
Fig. S10 a) XRF spectra for the Ni80/C powder material
0 1 2 4 6 8 10 20 30 40 50 60 70 80 90 100 200 300 KCps Rh KA1/Compton Rh KA1 Rh KB1/Compton Rh KB1 Rh LA1 Rh LB1 Ni KA1 Ni KB1 Ni LA1 Ni LB1 Ni KA1/Order 2
0.5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 30 40 50 52 54 56
KeV
Table S3 Elemental compositions for the Ni80/C material based on its XRF spectra
Formula Z Concentration Line 1 Net int. Stat. error LLD
N i 28 98.60% Ni KA1-HR-Tr 301.4 0.19% 122.8 PPM
P 15 0.24% P KA1-HR-Tr 0.06254 13.10% -
S 16 0.13% S KA1-HR-Tr 0.07002 13.60% 43.2 PPM
Ru 44 357 PPM Ru KA1-HR-Tr 0.01434 27.30% -
155
Fig. S10 b) XRF spectra for the Sn80/C powder material
Table S4 Elemental compositions for the Sn80/C material based on its XRF spectra
Formula Z Concentration Line 1 Net int. Stat. error LLD
Sn 50 98.30% Sn KA1-HR-Tr 76.71 0.38% 248.4 PPM
P 15 0.37% P KA1-HR-Tr 0.05869 13.50% -
S 16 0.33% S KA1-HR-Tr 0.1066 11.10% 48.1 PPM
0 1 2 3 4 5 6 7 8 9 10 20 30 40 50 60 70 80 KCps Sn KA1 Sn KB1 Sn LA1 Sn LB1 Rh KA1/Compton Rh KA1 Rh KB1/Compton Rh KB1 Rh LA1 Rh LB1
0.5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 30 40 50 51 52 53 54 55 56
KeV
156
Fig. S10 c) XRF spectra for the Ni60Sn20/C powder material
0 1 2 3 4 5 6 8 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150 160 KCps Sn KA1 Sn KB1 Sn LA1 Sn LB1Rh KA1/Compton Rh KA1 Rh KB1/Compton Rh KB1 Rh LA1 Rh LB1 Ni KA1 Ni KB1 Ni LA1 Ni LB1 Ni KA1/Order 2 Zn KA1 Zn KB1 Zn LB1 Cl KA1 Cl KB1
0.5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 30 40 50 51 52 53 54 55 56
KeV
Table S5 Elemental compositions for the Ni60Sn20/C material based on its XRF spectra
Formula Z Concentration Line 1 Net int. Stat. error LLD
N i 28 83.50% Ni KA1-HR-Tr 146.5 0.27% 38.4 PPM
Sn 50 12.90% Sn KA1-HR-Tr 19.2 0.77% 149.6 PPM
Cl 17 1.40% Cl KA1-HR-Tr 0.3519 5.75% 101.5 PPM
Zn 30 1.07% Zn KA1-HR-Tr 1.057 3.40% 62.0 PPM
P 15 0.35% P KA1-HR-Tr 0.06608 12.70% -
Cu 29 0.25% Cu KA1-HR-Tr 0.5436 5.12% 30.3 PPM
M g 12 0.23% Mg KA1-HR-Tr 0.01035 32.10% -
S 16 0.19% S KA1-HR-Tr 0.07414 13.40% 48.0 PPM
Ca 20 873 PPM Ca KA1-HR-Tr 0.0773 13.10% 71.8 PPM
Si 14 80.2 PPM Si KA1-HR-Tr 0.01792 24.40% -
157
Fig. S10 d) XRF spectra for the Ni40Sn40/C powder material
0 1 2 3 4 5 6 7 8 10 20 30 40 50 60 70 80 90 100 110 120 130 KCps Sn KA1 Sn KB1 Sn LA1 Cl KA1 Sn LB1 Cl KB1 Rh KA1/Compton Rh KA1 Rh KB1/Compton Rh KB1 Rh LA1 Rh LB1 Ni KA1 Ni KB1 Ni LB1 Ni KA1/Order 2 Sn KA1 Sn KB1 Sn LA1 Sn LB1 Ni KA1 Ni KB1 Ni LB1 Ni KA1/Order 2
0.5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 30 40 50 52 54 56
KeV
Formula Z Concentration Line 1 Net int. Stat. error LLD
N i 28 67.90% Ni KA1-HR-Tr 122.8 0.30% 6.4 PPM
Sn 50 28.70% Sn KA1-HR-Tr 25.67 0.66% 32.1 PPM
Cl 17 2.36% Cl KA1-HR-Tr 0.709 3.97% 15.0 PPM
P 15 0.27% P KA1-HR-Tr 0.06159 13.20% -
Zn 30 0.21% Zn KA1-HR-Tr 0.2762 7.63% 8.1 PPM
Cu 29 0.14% Cu KA1-HR-Tr 0.3179 7.13% 5.0 PPM
S 16 0.12% S KA1-HR-Tr 0.0555 16.60% 8.0 PPM
K 19 0.11% K KA1-HR-Tr 0.04236 15.90% -
Table S6 Elemental compositions for the Ni40Sn40/C material based on its XRF spectra
158
Fig. S10 e) XRF spectra for the Ni20Sn60/C powder material
0 1 2 3 4 5 6 7 8 9 10 20 30 40 50 60 70 80 90 100 KCps Rh KA1/Compton Rh KA1 Rh KB1/Compton Rh KB1 Rh LA1 Rh LB1 Sn KA1 Sn KB1 Sn LA1 Sn LB1 Ni KA1 Ni KB1 Ni LB1
0.5 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 30 40 50 51 52 53 54 55 56 KeV
Table S7 Elemental compositions for the Ni20Sn60/C material based on its XRF spectra
Formula Z Concentration Line 1 Net int. Stat. error LLD
Sn 50 84.10% Sn KA1-HR-Tr 93.37 0.34% 52.0 PPM
N i 28 14.50% Ni KA1-HR-Tr 27.04 0.63% 8.4 PPM
Cl 17 1.05% Cl KA1-HR-Tr 0.6232 4.24% 12.9 PPM
P 15 0.11% P KA1-HR-Tr 0.0607 13.30% -
S 16 782 PPM S KA1-HR-Tr 0.08026 13.20% 6.6 PPM
Fe 26 487 PPM Fe KA1-HR-Tr 0.05402 20.00% 9.5 PPM
159
Table S8 Ni and Sn wt. % quantification through XRD and XRF for the NiSn alloys and
monometallic analogues.
Catalyst Ni Sn Ni Sn
Ni80 100 - 98.6 -
Sn80 - 100 - 98.3
Ni60Sn20 91.3 8.7 83.5 12.9
Ni40Sn40 79.8 20.2 67.9 28.7
Ni20Sn60 60.8 39.2 14.5 84.1
XRD XRF
160
Ni20Sn60/C
Ni40Sn40/C
Ni80/C Sn80/C
a b
c d
Fig. S11 CVs at a scan rate of 50 mV s
–1
for a) Ni20Sn60/C, b) Ni40Sn40/C, c) Ni80/C, and d) Sn80/C
in 1 mol L–1 KOH demonstrating the Ni(OH)2/NiOOH redox couple.
161
Fig. S12 CVs of the Ni20Sn60/C, Ni40Sn40/C, Ni80/C, and Sn80/C electrodes showing the last
activation cycle (solid) in 1 M KOH and the oxidation of 0.33 mol L–1 urea (dashed) after 25 cycles
at a scan rate of 50 mV s
–1
Ni20Sn60/C Ni40Sn40/C
Ni80/C
Sn80/C
162
Table S9 Summary of electrochemical performance for Ni-based catalysts during activation,
including maximum current density (jmax), peak potential (E1/2), cycle number at jmax, and diffusion
coefficient (D) in alkaline media.
Cat
jmax
(mA · cm–2)
E1/2 (V) cycle #
D
(cm2
· s
–1)
N i20Sn60 2.18 0.52 15 1.46 × 10–6
N i40Sn40 14.37 0.55 15 8.68 × 10–5
N i60Sn20 22.91 0.53 7 1.21 × 10–4
N i80 10.37 0.50 5 5.73 × 10–5
163
Table S10 Tabulated ECSA results for Ni-based electrocatalysts in 1 M KOH post-activation. The
specific current density (jsp) at 0.65 V was calculated by normalizing to the mass of the catalyst
(m = 0.05 mg) instead of geometric surface area of RDE.
Cat q (µC)
ECSA
(cm2
)
ECSA
(cm2
· mgCAT
–1)
j sp @ 0.65 V
(mA · mgN i
–1 )
N i20Sn60 27.57 0.07 1.31 479
N i40Sn40 129.44 0.31 6.16 389
N i60Sn20 209.29 0.50 9.97 341
N i80 110.89 0.26 5.28 215
164
Table S11 Tabulated Tafel equation values for each electrocatalyst, the calculated exchange
current density (j0), and the correlation for the line of best fit when normalized to the geometric
surface area of the RDE.
Catalyst b (mV/dec) a j0 (mA/cm2
) R
2
N i20Sn60 67.3 0.985 9.64 0.992
N i40Sn40 70.9 0.976 9.48 0.995
N i60Sn20 77.8 0.967 9.27 0.995
N i80 90.0 0.951 8.92 0.998
165
Table S12 Tabulated Tafel equation values for each electrocatalyst, the calculated exchange
current density, and the correlation for the line of best fit when normalized to their respective
ECSA.
Catalyst b (mV/dec) a (V) j0 (mA/cm2
ECSA) R
2
N i20Sn60 69.5 1.070 11.74 0.994
N i40Sn40 71.8 1.103 12.71 0.996
N i60Sn20 81.3 1.539 34.47 0.996
N i80 94.7 1.450 28.2 0.997
166
Fig. S13 EIS for N20Sn60/C (blue), N40Sn40/C (red), N60Sn20/C (green), Ni80/C (black),and Sn80/C
(orange) electrodes with 0.33 mol L–1 urea in 1 mol L–1KOH at 0.5 V. The data was recorded over
a frequency range from 10 kHz to 100 mHz with perturbation amplitude of 25 mA.
167
Fig. S14 Polarization and power density curves for the DUFC at a) increasing anolyte and catholyte
flow rates under r.t. conditions and b) at different temperatures with the optimized flow rate of 580
mL min–1
a b
168
Fig. S15 Chronoamperometric response showing the current density versus time at a constant
applied potential of 0.2 V, illustrating the decay in current over time of the DUFC at r.t.
169
Fig. S16 Polarization and power density curves of the DUFC using mTPN1-TMA anion exchange
membrane at varying temperatures.
170
Supplementary Information III
Polyethyleneimine Encapsulated Silver Nanoparticles as Cathode
Electrodes for Syngas Generation
171
Table S13 Tafel slope (TS) analysis at three different overpotential regions: η < 100 mV (low),
100 mV< η < 250 mV (med), and η > 250 mV (med) for Pt/C and PEI-750k Ag NPs.
𝜂 Cat TS (mV dec-1) j0 (mA/cm2
) a (mV)
Low Pt/C 49 18.74 63
PEI Ag NPs 56 18.45 71
Med Pt/C 146 7.08 124
PEI Ag NPs 169 7.24 145
High Pt/C 438 0.359 195
PEI Ag NPs 500 0.309 255
172
c d
a b
Fig. S17 Nyquist plots under galvanostatic mode at OCV from 10 kHz to 1 Hz for a) PEI-750k Ag
NPs and b) Pt/C in N2-saturated 0.5 M KHCO3 at sinusoidal current perturbations of 1 mA (yellow),
10 mA (red), 25 mA (black), and 50 mA (blue). K-K relations for c) PEI-750k Ag NPs and d) Pt/C
comparing the real and imaginary components of impedance and their frequency dependance.
173
a b
c d
Fig. S18 Nyquist plots under potentiostatic mode at OCV from 10 kHz to 1 Hz for a) PEI-750k Ag
NPs and b) Pt/C in N2-saturated 0.5 M KHCO3 at sinusoidal current perturbations of 1 mV (yellow),
10 mV (red), 25 mV (black), and 50 mV (blue). K-K relations for c) PEI-750k Ag NPs and d) Pt/C
comparing the real and imaginary components of impedance and their frequency dependance.
174
c
a b
d
Fig. S19 Maximum deviation |Δ|max in the real and imaginary components of impedance as a
percentage for different perturbation amplitudes. Panels (a) and (b) show GEIS |Δ|max for current
perturbations (in mA) for PEI-750k Ag NPs and Pt/C, respectively. Panels (c) and (d) show PEIS
|Δ|max for voltage perturbations (in mV) for PEI-750k Ag NPs and Pt/C, respectively. Blue bars
represent the real part of the impedance, while orange bars represent the imaginary part. These
plots indicate the perturbation amplitudes at which the impedance data show minimal deviation,
thereby validating the data's consistency and suitability for further EIS analysis.
175
H2 N2 CO2
Fig. S20 Chromatograms after constant potential electrolysis of 15 min at six different potentials
(vs RHE) for PEI-750k Ag NPs in N2-saturated 0.5 M KHCO3. Hydrogen, nitrogen, and carbon
dioxide signals with corresponding retention times of 2.21, 2.51, and 5.42 min.
176
10 mV/s 50 mV/s
20 mV/s
25 mV/s
75 mV/s
100 mV/s
Fig. S21 LSVs of pristine glassy carbon electrode (GCE; dashed) and PEI-750k Ag NPs on GCE
(green) at a potential window from 0 to –0.8 V (vs RHE) in CO2-saturated 0.5 mol L–1 KHCO3 at
varying scan rates.
177
10 mV/s 50 mV/s
20 mV/s
25
mV/s
75 mV/s
100
mV/s
Fig. S22 LSVs of pristine microporous layer carbon paper (MPL; dashed) and PEI-750k Ag NPs
on MPL (green) at a potential window from 0 to –0.8 V (vs RHE) in CO2-saturated 0.5 mol L–1
KHCO3 at varying scan rates.
178
10 mV/s
20 mV/s
50 mV/s
75 mV/s
25
mV/s
100
mV/s
Fig. S23 LSVs of pristine Toray carbon paper (TCP; dashed) and PEI-750k Ag NPs on TCP
(green) at a potential window from 0 to –0.8 V (vs RHE) in CO2-saturated 0.5 mol L–1 KHCO3 at
varying scan rates.
179
Fig. S24 LSVs of pristine glassy carbon electrode (GCE; dashed) and PEI-750k Ag NPs on GCE
(green) at a potential window from 0 to –1.2 V (vs RHE) in CO2-saturated 0.5 mol L–1 KHCO3 at
varying scan rates.
10 mV/s
20 mV/s
50 mV/s
75 mV/s
25
mV/s
100
mV/s
180
10 mV/s 50 mV/s
20 mV/s
25 mV/s
75 mV/s
100 mV/s
Fig. S25 LSVs of pristine microporous layer carbon paper (MPL; dashed) and PEI-750k Ag NPs
on MPL (green) at a potential window from 0 to –1.2 V (vs RHE) in CO2-saturated 0.5 mol L–1
KHCO3 at different scan rates.
181
10 mV/s 50 mV/s
20 mV/s
25 mV/s
75 mV/s
100 mV/s
Fig. S26 LSVs of pristine Toray carbon paper (TCP; dashed) and PEI-750k Ag NPs on TCP
(green) at a potential window from 0 to –1.2 V (vs RHE) in CO2-saturated 0.5 mol L–1 KHCO3 at
different scan rates.
182
a
b
H2 N2 CO2
CO
Fig. S27 Chromatograms after constant potential electrolysis of 15 min at six different
potentials (vs RHE) for a) pristine MPL and b) PEI-750k Ag NPs in CO2-saturated 0.5
mol L–1 KHCO3. Hydrogen, nitrogen, carbon monoxide, and carbon dioxide signals with
corresponding retention times of 2.19, 2.53, 2.70 and 5.33 min.
183
a
b
H2 N2 CO2
CO
Fig. S28 Chromatograms after constant potential electrolysis of 15 min at six different
potentials (vs RHE) for a) pristine MPL and b) PEI-750k Ag NPs in CO2-saturated 0.5
mol L–1 KHCO3. Hydrogen, nitrogen, carbon monoxide, and carbon dioxide signals with
corresponding retention times of 2.19, 2.53, 2.70 and 5.33 min.
184
b
H2 N2
CO2
CO
a
Fig. S29 Chromatograms after constant potential electrolysis of 15 min at six different
potentials (vs RHE) for (a) pristine TCP and (b) PEI-750k Ag NPs in CO2-saturated 0.5
mol L–1 KHCO3. Hydrogen, nitrogen, carbon monoxide, and carbon dioxide signals with
corresponding retention times of 2.19, 2.53, 2.70 and 5.33 min.
185
Fig. S30 Chromatograms after constant potential electrolysis of 15 min at three
different potentials (vs RHE) for PEI-750k (green), PEI-25k (yellow), and PEI-800 (red)
Ag NPs in CO2-saturated 0.5 mol L–1 KHCO3. Hydrogen, nitrogen, carbon monoxide,
and carbon dioxide signals with corresponding retention times of 2.19, 2.53, 2.70 and
5.33 min.
186
Fig. S31 TGA isotherm of PEI-750k Ag NPs at three cycles with respective temperatures of 25,
55, and 85°C.
187
Fig. S32 TGA isotherm of PEI-25k Ag NPs at three cycles with respective temperatures of 25,
55, and 85°C.
188
Fig. S33 TGA isotherm of PEI-800 Ag NPs at three cycles with respective temperatures of 25,
55, and 85°C.
Abstract (if available)
Abstract
The interchangeable conversion between chemical energy and electrical energy is crucial for achieving a sustainable infrastructure. Fuel cells and electrolytic devices will serve as essential technologies for these energy conversions, enabling both efficient generation and storage of renewable power. While significant progress has been made, further innovation at the materials level, ranging from the micro to the macro scale, remains necessary. Specifically, catalytic materials used in anode and cathode electrodes continue to be a focus of intensive research. Ideally, these electrocatalysts should be non-precious metal-based (i.e., cheap and abundant), highly active, and stable under operational conditions. This dissertation presents practical synthetic methodologies for the development of innovative nanomaterials designed for electrochemical devices, which include glucose biosensors, AEM/PEM fuel cells, and CO2 electrolyzers. Different compositions of a targeted catalyst were synthesized, characterized using advanced analytical techniques (e.g., microscopy, spectroscopy, diffraction) and studied for their tailored redox transformations. The material’s structure-activity relationships were analyzed via a variety of methods and techniques, such as electrochemical kinetics (RDE and Tafel analysis), polarization (CSV/LSV), constant bias (CA/CP), and impedance spectroscopy (EIS), all aimed at enhancing reactivity. The optimized electrocatalyst was then integrated into a device, engineered to maximize energy output and operational stability. These findings demonstrate the feasibility of employing cost-effective, high-performance nanomaterials in market-ready, next-generation electrochemical technologies.
Linked assets
University of Southern California Dissertations and Theses
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Asset Metadata
Creator
de los Rios, Juan Pablo
(author)
Core Title
Design of nanomaterials for electrochemical applications in fuel cells and beyond
School
College of Letters, Arts and Sciences
Degree
Doctor of Philosophy
Degree Program
Chemistry
Degree Conferral Date
2024-12
Publication Date
01/07/2025
Defense Date
11/01/2024
Publisher
Los Angeles, California
(original),
University of Southern California
(original),
University of Southern California. Libraries
(digital)
Tag
biosensors,electrochemistry,electrolyzers,fuel cells,Materials,OAI-PMH Harvest,renewable energy
Format
theses
(aat)
Language
English
Contributor
Electronically uploaded by the author
(provenance)
Advisor
Prakash, Surya (
committee chair
), Picazo, Elias (
committee member
), Ronney, Paul D. (
committee member
)
Creator Email
delosrio@usc.edu,jp@hydrosonicsenergy.com
Unique identifier
UC11399F7S5
Identifier
etd-delosRiosJ-13717.pdf (filename)
Legacy Identifier
etd-delosRiosJ-13717
Document Type
Dissertation
Format
theses (aat)
Rights
de los Rios, Juan Pablo
Internet Media Type
application/pdf
Type
texts
Source
20250109-usctheses-batch-1231
(batch),
University of Southern California
(contributing entity),
University of Southern California Dissertations and Theses
(collection)
Access Conditions
The author retains rights to his/her dissertation, thesis or other graduate work according to U.S. copyright law. Electronic access is being provided by the USC Libraries in agreement with the author, as the original true and official version of the work, but does not grant the reader permission to use the work if the desired use is covered by copyright. It is the author, as rights holder, who must provide use permission if such use is covered by copyright.
Repository Name
University of Southern California Digital Library
Repository Location
USC Digital Library, University of Southern California, University Park Campus MC 2810, 3434 South Grand Avenue, 2nd Floor, Los Angeles, California 90089-2810, USA
Repository Email
cisadmin@lib.usc.edu
Tags
biosensors
electrochemistry
electrolyzers
fuel cells
renewable energy