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Understanding the mechanism of oxygen reduction and oxygen evolution on transition metal oxide electrocatalysts and applications in iron-air rechargeable battery

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Understanding the mechanism of oxygen reduction and oxygen evolution on transition metal oxide electrocatalysts and applications in iron-air rechargeable battery

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Content i

UNDERSTANDING THE MECHANISM OF
OXYGEN REDUCTION AND OXYGEN
EVOLUTION ON TRANSITION METAL OXIDE
ELECTROCATALYSTS AND APPLICATIONS IN
IRON-AIR RECHARGEABLE BATTERY

by

Phong Trinh

A Dissertation Presented to the
FACULTY OF THE USC GRADUATE SCHOOL
UNIVERSITY OF SOUTHERN CALIFORNIA
In Partial Fulfillment of the
Requirement for the Degree
DOCTOR OF PHILOSOPHY
(CHEMISTRY)

May 2016

ii

TABLE OF CONTENTS
ACKNOWLEDGMENTS ..................................................................................................................... iv
LIST OF TABLES ................................................................................................................................. v
LIST OF FIGURES .............................................................................................................................. vi
LIST OF ABBREVIATION AND SYMBOLS ......................................................................................... xii
SUMMARY .................................................................................................................................... xiii
Chapter 1. Introduction ................................................................................................................... 1
1.1 Metal-Air Batteries ................................................................................................................ 2
1.2 Water Electrolysis ................................................................................................................. 4
1.3 Alkaline Fuel Cell ................................................................................................................. 5
1.4 Oxygen Reduction Reaction.................................................................................................. 7
1.5 Oxygen Evolution Reaction ................................................................................................ 10
1.6 Problems and Challenges .................................................................................................... 11
Chapter 2. Experimental Techniques and Methodology for Analysis of Results ......................... 14
2.1 Catalyst Synthesis ............................................................................................................... 14
2.2 Physical Characterization .................................................................................................... 14
2.2.1 X-Ray Diffraction ......................................................................................................... 14
2.2.3 Surface Area ................................................................................................................. 15
2.2.4 X-ray Photoelectron Spectroscopy ............................................................................... 15
2.2.5 X-ray Absorption Near Edge Spectroscopy ................................................................. 16
2.3 Electrochemical Characterization ....................................................................................... 16
2.3.1 Catalyst Ink Preparation ............................................................................................... 16
2.3.2 Electrode Preparation ................................................................................................... 17
2.3.2.1 Preparation of Electrodes for the Oxygen Reduction Reaction and Oxygen
Evolution Reaction Studies .................................................................................... 17
2.3.3 Electrochemical Experimental Set up ........................................................................... 20
iii

2.3.4 Electrochemical Method of Analysis ........................................................................... 22
2.3.4.1 Rotating Disk Electrode ......................................................................................... 22
2.3.4.2 Rotating Ring Disk Electrode (RRDE) .................................................................. 28
2.3.4.3 Measurement of Oxygen Reduction Rates ............................................................ 30
2.3.4.4 Measurement of Oxygen Evolution Rates ............................................................. 31
2.3.4.5 Rate of Hydrogen Peroxide Reduction and Oxidation .......................................... 31
2.3.4.6 Direct Measurement of Rate of Hydrogen Peroxide Decomposition by the
Manometer Method ................................................................................................ 34
2.3.5 Iron-Air Cell Studies .................................................................................................... 35
2.3.5.1 ORR Electrode Preparation .................................................................................... 35
2.3.5.2 OER Electrode Preparation .................................................................................... 36
2.3.5.3 Iron-Air Cell Set up ............................................................................................... 37
Chapter 3. Understanding the Role of Carbon and Transition Metal Oxide in the Oxygen
Reduction Reaction ..................................................................................................... 39
Chapter 4. Oxygen Reduction and Hydrogen Peroxide Decomposition Activity of
Transition Metal Oxide Catalysts ............................................................................... 62
Chapter 5. Oxygen Evolution Activity of Perovskite and Spinel Transition Metal Oxide
Catalyst ....................................................................................................................... 77
5.1 Oxygen Evolution Activity of Perovskite Transition Metal Oxide on Rotating Ring
Disk Electrode .................................................................................................................... 78
5.2 Oxygen Evolution Activity of Spinel Transition Metal Oxide on Nickel Foam
Electrode ............................................................................................................................ .93
Chapter 6. Studies on Iron-Air Rechargeable Battery ................................................................ 104
Chapter 7. Conclusions ............................................................................................................... 109
References .................................................................................................................................. 113
Publications and Presentations .................................................................................................... 130

iv

ACKNOWLEDGMENTS
I would like to thank, first and foremost my advisor Professor Sri R Narayan for allowing me to
work on one of the most challenging problems in energy storage and renewable energy
technology. Working on these challenging topics under his guidance, I have learned a lot from
him. I appreciate his continuous support, patience and the investment of his time to guide me in
my academic life.
I also would like to thank Professor G. K. Surya Prakash, Dr. Robert Aniszfeld for their constant
support and encouragement.
I thank Dr. S. Malkhandi and Dr. Aswin K.Manohar for helping me during my time at University
of Southern California.
Finally, I would like to thank all the members of the Narayan group for their friendship, support
and kindness.

v

LIST OF TABLES
Table 1. Composition of catalyst mixtures consisting of LCCO and Acetylene Black (AB)
that were used in our studies .......................................................................................... 40
Table 2. Measured values of rate constant for decomposition of hydrogen peroxide on
LCCO at 25
o
C in 1M potassium hydroxide determined from the slope of I
disk
/I
ring

vs. 
-1/2
........................................................................................................................... 60
Table 3. Various manganese-substituted calcium-doped lanthanum cobalt oxides (LCCM) ...... 64
Table 4. Number of d-electrons for M
z+
and occupancy of antibonding orbital of M
z+
—OH
bond................................................................................................................................ 92
Table 5. Heat treatment temperature and catalysts weight of iron-doped nickel cobalt oxide
catalysts ......................................................................................................................... 95
Table 6. OER and ORR testing of nickel cobalt oxide electrodes ............................................. 107
Table 7. OER testing of lanthanum nickelate electrodes ........................................................... 108

vi

LIST OF FIGURES
Figure 1. Rotating-Ring Disk Electrode with 20 L droplet of the catalyst ink. ......................... 17
Figure 2. Rotating disk electrode with a thin film of catalyst after drying in air at 85
0
C. .......... 18
Figure 3. Catalyst-coated on nickel foam electrode (black coating is the oxide). ....................... 19
Figure 4. 3 cm x 5 cm Toray® paper with the coating of catalyst. .............................................. 20
Figure 5. Polyfluoroethylene cell setup for ORR and OER measurements. ................................ 21
Figure 6. Polyfluoroethylene cell setup for hydrogen peroxide measurements. .......................... 21
Figure 7. Electrochemical cell setup for the OER measurement of catalyst coating on
nickel foam electrode. ............................................................................................... 22
Figure 8. Schematic of a rotating disk electrode .......................................................................... 23
Figure 9. Relationship between diffusion layer thickness  and rotation rate ............................. 25
Figure 10. A expected plot of 1/I vs 1/ω
1/2
at a potential -200 mV for various rotation rate
of 400, 900, 1600, and 2500 rpm .............................................................................. 26
Figure 11. ORR polarization curve measured by rotating disk electrode at various rotation
rate ............................................................................................................................. 26
Figure 12. A typical Tafel plot ..................................................................................................... 27
Figure 13. Rotating ring disk electrode with glassy carbon disk and platinum ring electrode .... 28
Figure 14. Schematic of flow of oxygen to the RRDE when the electrode is rotated ................. 29
Figure 15. Hydrogen peroxide reduction and oxidation polarization experiments in argon
saturated 1M potassium hydroxide containing 50 mM hydrogen peroxide. ............. 32
Figure 16. Direct hydrogen peroxide decomposition experiment setup ...................................... 35
Figure 17. ORR electrode of NiCo
2
O
4
-vulcanXC-72 .................................................................. 36
Figure 18. OER electrode with Fe doped NiCo
2
O
4
coated on nickel foam. ................................ 37
Figure 19. (a) Configuration of iron-air cell (b) Iron-Air cell test under test .............................. 38
Figure 20. Separate electrode configuration for air electrode ...................................................... 38
Figure 21. Polarization curves for various catalyst ink compositions in oxygen-saturated 1
M potassium hydroxide at various rotation rates of 400-2500 rpm (a) LCCO
vii

without any added carbon, (b) acetylene black without any LCCO (c) mix of
LCCO (91%) and acetylene black (9%), (d) comparison of the mass activity of
the catalysts in (a), (b) and (c) at -0.25 V rotation rate of 400 r ................................ 41
Figure 22. Oxygen reduction activity in oxygen-saturated 1 M potassium hydroxide at
-0.10 V for 83% LCCO and 17% acetylene black, 83% LCCO and 17%
graphene, 83% LCCO and 17% carbon nanotubes, 83% LCCO and 17% gold
nanoparticles ............................................................................................................. 43
Figure 23. (a) and (b) Polarization curves for oxygen reduction in 1M potassium hydroxide
saturated with oxygen for various mass ratios of LCCO and acetylene black;
labels correspond to sample numbers in Table 1. Kinetic current at -0.1 V at
1600 rpm: (c) at various mass ratios of LCCO and acetylene black as in Table
1, (d) varying amounts of carbon and fixed amount of LCCO (8mg) and (e)
varying amounts of LCCO with a fixed amount of carbon (8mg). The dashed
lines in (d) and (e) are to aid the visualization of the trends. .................................... 45
Figure 24. (a) Polarization curves for oxygen reduction in 1M potassium hydroxide at
1600 rpm for various catalysts consisting of LCCO, lanthanum calcium cobalt
manganese oxide (La
0.6
Ca
0.4
Co
0.5
Mn
0.5
O
3-x
, LCCM), and nickel cobalt oxide
(NC). Curves correspond to I: LCCO, Ia: LCCO+C, II: LCCM, IIa:
LCCM+C; III: NC, IIIa: NC+C. (b) Oxygen reduction activity at -0.1 V with
various catalysts with and without carbon. Carbon content in the composite
catalyst was 16.7% (w/w) ......................................................................................... 47
Figure 25. Polarization experiments in oxygen-saturated 1M potassium hydroxide on a
rotating ring-disk electrode (a) LCCO, (b) Acetylene black, and (c) 83.3%
LCCO + 16.7% AB. The ring electrode was a platinum electrode held at +0.5
V vs. MMO. Each block is arranged vertically as disk current, ring current and
number of electrons transferred …... ........................................................................... 49
Figure 26. (a) change of limiting oxidation current at 0.2 V for hydroperoxide with time
in the presence of 10 mg of LCCO in 0.1M hydroperoxide in 1M potassium
hydroxide, (b) results of cathodic polarization (potential scanned at 5 mV s
-1
)
of LCCO-coated disk electrode in argon saturated 0.1 M hydroperoxide in 1M
potassium hydroxide. ................................................................................................. 51
viii

Figure 27. (a) Percentage of current used in hydroperoxide generation on various catalyst
compositions of Table 1 coated on the disk electrode, (b) The calculated
number of electrons transferred in the reaction determined from the amount of
hydroperoxide detected at the ring electrode. The percentage of hydroperoxide
and the number of electrons transferred are determined at -0.15 V. ......................... 52
Figure 28. Comparison of experimental and calculated values of hydroperoxide production
on LCCO-acetylene black composites of composition in Table 1 ........................... .54
Figure 29. Reaction scheme used in the analysis of the behavior of carbon –transition
metal oxide composite catalyst. ................................................................................. 56
Figure 30. Plot of I
disk
/I
ring
as a function of (1/rotation frequency)
1/2
(a) at various
potentials for composite catalyst (Carbon 8 mg–LCCO 8 mg); (b) for various
composite catalyst compositions with 8 mg of carbon and increasing amounts
of LCCO (1 mg, 2 mg, 4 mg and 8 mg) at electrode potentials -0.175 V. ................ 59
Figure 31. Kinetic current plotted vs. the potential of the disk electrode coated with
LCCO-carbon mixtures as indicated, (a) for compositions rich in LCCO (b)
compositions rich in acetylene black. ........................................................................ 61
Figure 32. ORR activity for various perovskite and spinel oxides system .................................. 63
Figure 33. (a) Polarization curves for acetylene black and LCCM (x=1) + AB in oxygen-
saturated 1M potassium hydroxide at rotation rate 2500 rpm. (b)
Hydroperoxide current detected at the ring for acetylene black and LCCM
(x=1) +AB. ................................................................................................................ 65
Figure 34. (a) Polarization curves for oxygen reduction in 1M potassium hydroxide at
2500 rpm for various compositions of the transition metal oxide (b)
Hydroperoxide current at the ring electrode for various compositions (c) ORR
activity vs. percentage of hydroperoxide observed at the ring electrode (d)
ORR activity vs. percentage of hydroperoxide for various value of manganese
fraction. ...................................................................................................................... 66
Figure 35. (a) Polarization curves of hydrogen peroxide reduction and oxidation kinetic
current for various oxides in 1M potassium hydroxide solution containing 50
mM hydrogen peroxide (b) Open-circuit potential (mixed potential) measured
and calculated for various compositions as indicated by ―x‖ values ........................ 70
ix

Figure 36. (a) Exchange current density of peroxide oxidation vs peroxide reduction for
various manganese fractions (b) Exchange current density of peroxide
oxidation vs ORR activity at -100 mV. ..................................................................... 72
Figure 37. (a) Peroxide oxidation rate vs ORR activity at -100 mV vs MMO (b) I
corr
vs.
ORR activity at -100 mV vs MMO. .......................................................................... 73
Figure 38. Exchange current density of hydrogen peroxide reduction and ORR activity at
-100 mV vs MMO for various manganese fractions. ................................................ 74
Figure 39. Tafel slope of hydrogen peroxide oxidation and reduction for various
manganese fractions .................................................................................................. 75
Figure 40. The decomposition rate I
corr
from peroxide polarization experiment vs
decomposition rate from peroxide direct decomposition study. ............................... 76
Figure 41. Oxygen evolution activity of various perovskite and spinel transition metal
oxides ......................................................................................................................... 77
Figure 42. (a) X-ray diffraction pattern for La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
for the various
compositions indicated (b) Magnified region 2  = 30
o
-35
o
, (c) Powder
diffraction file data (PDF#00-041-0496) for La
0.5
Ca
0.5
Mn
0.5
Co
0.5
O
3
. (d)
Crystallite size for various values of manganese fraction calculated from 2  =
33
o
(220). ................................................................................................................... 80
Figure 43. Scanning electron micrographs of La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
; (a) x = 0 (b) x = 0.1;
(c) x = 0.5 (d) x = 0.7 (e) x = 0.9 (f) x = 1.0 .......................................................... 81
Figure 44. X-ray photoelectron spectroscopy of (a) Mn-2p (b) Co-2p and (c) O-1s for the
various compositions indicated by the values of atomic ratio x = Mn/Co. XPS
spectra were corrected using carbon spectra as a standard........................................ 82
Figure 45. (a) Variation of concentration of Mn
2+
, Mn
3+
, Mn
4+
species with the manganese
fraction determined by deconvolution of Mn-2p
3/2
XPS peaks from data
presented in Figure 3. (b) Average oxidation state of manganese as a function
of manganese fraction, x. (c) Variation of concentration of Co
3+
and Co
4+

species with the manganese fraction as determined by deconvolution of Co-
2p
3/2
peaks in Figure 44. (d) Calculated average oxidation state of cobalt as a
function of manganese fraction, x ............................................................................ 83
x

Figure 46. (a) Manganese XANES for manganese fraction x = 0.1 to 1 in
La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
and LaMnO
3
(Marked as LMO). Inset has been include
for clarity. (b) Cobalt XANES for manganese fraction x = 0 to 0.9 in
La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
. ............................................................................................... 84
Figure 47. Potentiostatic polarization curves for oxygen evolution on the various catalysts
(indicated by x values) in 1 M potassium hydroxide at 25
o
C. E
MMO
vs. RHE =
+0.90 V. Oxide amount on the electrode was about 80 microgram for all the
samples. ..................................................................................................................... 87
Figure 48. Effect of manganese fraction on (a) specific activity at 550, 615 and 650 mV
vs. MMO reference, E
MMO
vs. RHE = +0.90 V (b) Tafel slope in the potential
region 550 to 650 mV .......................................................................................... ……88
Figure 49. Specific activity normalize for BET surface area at 650 mV vs. MMO
reference, E
MMO
vs. RHE = +0.90 V ......................................................................... 88
Figure 50. Molecular Orbital ordering schematic for M
z+
-OH bonds for cobalt and
manganese. ................................................................................................................ 91
Figure 51. Oxygen evolution activity and average number of electrons occupying in the
antibonding orbitals of M
z+
-OH at 650 mV vs. MMO. ............................................. 93
Figure 52. OER electrocatalyst activity of various metal oxide (adaped from
78
) ....................... 93
Figure 53. (a) X-ray diffraction pattern for iron-doped nickel cobalt oxide prepared at
various temperature values (b), (c) Magnification of the region in the box in
(a), (d) Powder diffraction file data PDF#00-020-0781 for NiCo
2
O
4
. ...................... 96
Figure 54. SEM picture of Fe doped NiCo
2
O
4
coated on nickel foam (first row), magnified
picture of metal oxide (second row). ......................................................................... 97
Figure 55. X-ray photoelectron spectroscopy of Ni 2p, Co 2p, Fe 2p, and O 1s for various
electrode at different temperatures of heat treatment. ............................................... 98
Figure 56. Example of deconvolution of XPS data of Ni 2p
3/2
and Co 2p
3/2
and variation of
concentration of Ni
2+
, Co
3+
, and Co
3+
as a function of the temperature of heat
treatment step. ........................................................................................................... 99
Figure 57. Variation of concentration of Fe
3+
,O
2
-
, and OH
-
as determined by deconvolution
of the XPS of Fe 2p
3/2
and O 1s at various temperatures. ........................................ 99
xi

Figure 58. (a) Steady-state polarization curves for various electrodes (b) Current density at
520 mV vs MMO for various electrodes (c) Electrode potentials at 10 mA/cm
2

(d) Tafel slope for various electrodes in the potential range from 500 mV to
650 mV. ................................................................................................................... 100
Figure 59. Schematic reaction mechanisms of oxygen evolution reaction on iron-doped
nickel cobalt oxide electrode. .................................................................................. 102
Figure 60. (a) Electrode potential vs time at 10 mA/ cm
2
current density for 65 hours.
(b) Micro volt per hour plot of various electrodes .................................................. 102
Figure 61. Polarization curves of commercial MnO
2
catalyst and NiCo
2
O
4
-VulcanXC-72
catalyst for ORR in iron-air cell containing 30% potassium hydroxide
electrolyte. ............................................................................................................... 105
Figure 62. (a) Polarization curve of 5% Fe doped NiCo
2
O
4
OER and NiCo
2
O
4
-Vulcan
XC72 ORR electrode (b) Long term testing of ORR and OER electrodes at
250 mA .................................................................................................................... 106
Figure 63. (a) Comparison of polarization curves of different OER electrodes (b) Long
term testing of OER electrodes at 250 mA .............................................................. 108

xii

LIST OF ABBREVIATION AND SYMBOLS
NHE: Normal Hydrogen Electrode
MMO: Mercury-Mercuric Oxide Electrode
RDE: Rotating Ring Disk Electrode
CV: Cyclic Voltammetry
ORR: Oxygen Reduction Reaction
OER: Oxygen Evolution Reaction
XPS: X-ray Photoelectron Spectroscopy
XANES: X-ray Absorption Near-Edge Spectroscopy
SEM: Scanning Electron Microscopy
LCCO: Lanthanum-calcium cobalt oxide
LCMO: Lanthanum-cobalt manganese oxide
LCCM: Lanthanum-calcium-cobalt manganese oxide
NC: Nickel cobalt oxide
AB: Acetylene Black
ω: Rotation rate-radian/second
rpm: rotation per minute

xiii

SUMMARY
In this dissertation I have focused on understanding the mechanism of the
electrochemical reduction of oxygen (ORR) and the oxygen evolution reaction (OER) on
transition metal oxides electrocatalysts in alkaline media, and the application of these
electrocatalysts in a rechargeable iron-air battery.
In Chapter 1, I provide an introduction to electrochemical energy storage systems, the
processes of oxygen reduction and oxygen evolution, the motivation and objectives of the
research, the technical challenges and the unique approach adopted in this work. In Chapter 2, I
describe the experimental methodologies for the preparation and characterization of various
catalytic materials, electrodes and cells.
Chapter 3 contains the description of the results of experiments on the role of carbon and
transition metal oxide in the catalysis of the oxygen reduction reaction. I found that conductive
carbon materials when added to transition metal oxides such as calcium-doped lanthanum cobalt
oxide, nickel cobalt oxide and calcium-doped lanthanum manganese cobalt oxide increase the
electrocatalytic activity of the oxide for oxygen reduction by a factor of five to ten. I have
studied rotating ring-disk electrodes coated with: (a) various mass ratios of carbon and transition
metal oxide, (b) different type of carbon additives and (c) different types of transition metal
oxides. Our experiments and analysis establish that in such a composite catalyst, carbon is the
primary electro-catalyst for the two electron electro-reduction of oxygen to hydroperoxide while
the transition metal oxide decomposes the hydroperoxide to generate additional oxygen that
enhances the observed current resulting in an apparent four-electron process.
xiv

Following up on the findings on the role of the transition metal oxide described in
Chapter 3, I present results and discussion on the relationship between oxygen reduction activity
and hydrogen peroxide decomposition activity of perovskite transition metal oxide in Chapter 4.
I found that the oxygen reduction activity of lanthanum-doped calcium manganese cobalt oxide
increased when the manganese fraction in the oxide increased. We also studied the mechanism of
decomposition of hydrogen peroxide on calcium-doped lanthanum cobalt manganese oxide and
established that this decomposition process occurs by an electrochemical pathway. We also
verified that the higher hydrogen peroxide decomposition rate on the oxide, the better was the
ORR activity of the carbon-transition metal oxide composite.
In chapter 5, I describe new insights for predicting and tuning the activity of transition
metal oxides for designing efficient and inexpensive electrocatalysts for the oxygen evolution
reaction. I have conducted a systematic investigation of nano-phase calcium-doped lanthanum
manganese cobalt oxide, an example of a mixed metal oxide that can be tuned for its
electrocatalytic activity by varying the composition of the transition metals. We have found that
the value of Tafel slopes are governed by the oxidation states and the bond energy of the surface
intermediates (such as Mn-OH and Co-OH bonds) while the catalytic activity increased with the
average d-electron-occupancy of the σ
*
orbital of the metal-OH bond. I have investigated the
electrocatalytic activity for oxygen evolution on iron-doped nickel cobalt oxide electrodes
prepared and treated at various temperatures. I found that the electrode prepared at 200
o
C is less
crystalline and has higher activity for oxygen evolution than the electrode prepared at 400
o
C.
In chapter 6, I studied the performance of iron-air rechargeable cells based on the most
catalytically active transition metal oxide-based oxygen reduction electrodes and oxygen
evolution electrodes resulting from my research. We have demonstrated that our NiCo
2
O
4
-
xv

Vulcan XC72 composite electrode exhibited considerably better activity for oxygen reduction
than the commercial MnO
2
electrode. I also tested the performance of iron-doped NiCo
2
O
4
-
based oxygen evolution electrode and found that the overpotentials were sufficiently low to
operate the cells at 10 mA/cm
2
. In these experiments, the oxygen reduction electrode and the
oxygen evolution electrode were operated as separated electrodes. Thus, I could demonstrate that
no noticeable change in overpotential or catalytic activity was observed for at least 100 hours in
the case of the oxygen reduction electrode and 200 hours in the case of the oxygen evolution
electrode.
Chapter 7 summarizes the understanding and conclusions arising from my research and the
implications of the findings for future development of iron-air rechargeable batteries.

1

Chapter 1
Introduction
Solar and wind power are well known as renewable energy resources that have the
potential to meet future demands for sustainable energy systems. However, the energy supply
from solar and wind generation is intermittent. Since the electricity must be continually available
for industrial and residential grid applications, fluctuations in electricity generation can be
detrimental to the integration of these renewable energy sources into the electricity grid.
Additionally, most of the suitable locations for solar and wind power collection are in desert
regions and sometimes far away from industrial and residential areas. To make solar and wind-
based electrical generation economical and practical, energy production must be coupled with a
reliable energy storage system. Thus, energy storage systems play a vital role in meeting the
demand for a sustainable energy solution in the future. Energy storage can provide more
flexibility and balance to the electricity grid by providing backup power during power outages
and by buffering the intermittency of generation from renewable energy sources. Apart from
these benefits to the electricity grid, there are at least 1.4 billion people globally that have no
access to electricity at all. Their lives can be transformed with the help of energy storage
combined with solar and wind systems. Rechargeable metal-air batteries, water electrolyzers and
alkaline fuel cells are promising technologies that can address the challenge of storing
inexpensively large amounts electrical energy generated form solar and wind resource; their high
energy density and environmental friendliness are particularly attractive for large-scale
deployment.
1-3
However, one the principal challenges with these systems is their efficiency and
cost. The following sections describe the processes in metal-air batteries, water electrolyzer and
2

alkaline fuel cells to present the relevance of my research on the topic of electrocatalysts for
oxygen reduction and oxygen evolution.
1.1 Metal-Air Batteries
Metal-air rechargeable batteries have received much attention for electrical energy
storage over the past several decades because of their significantly higher theoretical energy
density compared to other rechargeable batteries. In metal-air batteries, the electricity is
generated through the redox reaction between a metal at the negative electrode and oxygen at the
positive electrode.
Positive electrode : O
2
+ 2H
2
O + 4e
−
→ 4OH
−

(1)
Negative electrode : M M
n+
+ne
−
(2)
Metal-air batteries have an open cell structure as the cathode active material, oxygen,
must be supplied continuously from an atmospheric air source. This arrangement of drawing
oxygen from the surrounding air removes the need for storage of the positive active material
within the battery. This feature is one of the reasons that metal-air batteries have such
remarkably high theoretical energy density; both the weight and volume are significantly
reduced. Many metals have been considered as candidate negative electrode materials for metal-
air batteries, including zinc, magnesium, lithium, aluminum, and iron. Among this, the ―iron-air‖
rechargeable battery is particularly attractive because of the global abundance of iron, the
robustness of the iron electrode to charge/discharge cycling, and the eco-friendliness of the
battery materials.
4
Further, the practical specific energy of the iron-air battery being as high as
100-150 Wh/kg is quite attractive for stationary and mobile applications. Along with iron-air
battery, zinc-air battery has been used for a long time due to its high energy density and
3

abundance of raw material. The zinc-air primary battery has been commercialized as a hearing
aid battery. Even though zinc-air battery has theoretical specific energy of 1350 Wh/kg
5
the
zinc-air rechargeable battery still does not meet the requirement for large-scale energy storage
due to the poor cycle life caused by non-uniform zinc dissolution and deposition, and low
efficiency of the bifunctional air electrode.
6

Although metal-air batteries have higher theoretical energy density compared to other
rechargeable battery systems, they have many challenges from a scientific and technological
standpoint. The major issues are the low utilization efficiency of the negative electrode and the
slow kinetics at the positive electrode which lead to low practical energy density.
7-8
Designing a
low-cost, durable material for the bi-functional air electrode to reduce the overpotential losses
during oxygen reduction and oxygen evolution reaction is necessary for improving the state of
the art of metal-air battery.
7

Our group at USC has recently reported significant improvements to the charging
efficiency and rate capability of the iron electrode.
9-10
However, to achieve an iron-air battery
with long cycle life and high performance, it is also essential to have an efficient and robust bi-
functional air electrode. The charge and discharge reactions in an iron-air cell are as follows
Discharge reactions:
(+) Electrode: O
2
+ 2H
2
O + 4e
-
→ 4OH
¯
(3) E
o
= 0.401 V
(-) Electrode: Fe + 2OH
¯
→ Fe(OH)
2
+ 2e
-
(4) E
o
= -0.877 V
Charge reactions:
(+) Electrode: 4OH
¯
→ O
2
+ H
2
O + 4e
-
(5) E
o
= 0.401 V
4

(-) electrode: Fe(OH)
2
+ 2e
-
→ Fe + 2OH
¯
(6) E
o
= -0.877 V
During discharge, at the negative electrode, elemental iron is oxidized to iron (II)
hydroxide and electrons are released into the external circuit. At the same time, oxygen diffuses
into the positive electrode, and accepts the electrons from the negative electrode to produce
hydroxide ions. When the cell is charged the process is reversed with iron deposition at the
negative electrode and oxygen evolution at the positive electrode. Thus, the oxygen reduction
and oxygen evolution reactions form the basis of the operation of air-based batteries. For the
iron-air battery to operate efficiently we not only need a robust iron electrode but also a
bifunctional air electrode that can support both oxygen reduction and oxygen evolution. Further,
the material for this bifunctional electrode should be chemically stable at the positive electrode
potentials needed for oxygen reduction and oxygen evolution. Carbon combined with silver or
manganese dioxide has been used as an electrocatalyst for oxygen reduction in alkaline fuel cells
for many years.
11
Similarly, other carbon-metal oxide composites have shown good
electrocatalytic activity for oxygen reduction.
12
However, these catalyst composites are unstable
under oxygen evolution potentials due to the rapid loss of carbon by oxidation. Therefore,
carbon-based electrodes that are suitable for carrying out oxygen reduction are not durable as
oxygen evolution electrodes. To ensure durability it would be appropriate to separate the
functions of oxygen reduction and oxygen evolution on two electrodes, in which only the oxygen
reduction process occurs on the carbon-based catalysts.
1.2 Water Electrolysis
Water electrolysis is a promising method to produce hydrogen. However, today, most of
the hydrogen production is from hydrocarbon reforming and only about 4% of the world‘s
5

hydrogen production are from water electrolysis.
13
In water electrolysis, electricity is used to
split the water molecule into hydrogen and oxygen molecules. The reason that this process has
not become practical yet on a larger scale is due to the high cost of electricity and the poor
efficiency and high cost of electrocatalysts for the oxygen evolution reaction.
14
The half
reactions that occur at the positive and negative electrodes during water electrolysis in alkaline
medium are,
(+) Negative electrode: 2H
2
O

+ 2e
-
→ H
2
+ 2OH
¯
(7) E
o
= -0.84 V
(-) Positive electrode: 2OH
¯
→ 1/2O
2
+ H
2
O + 2e
-
(8) E
o
= +0.401 V
The overall reaction of water electrolysis process is H
2
O → H
2
+ 1/2O
2
(9)
Currently, alkaline electrolyzers are commercially available. The capital cost is estimated
to range from $1000 to $ 5000 /kW depending on the production capacity.
15
Using higher
temperatures of operation in alkaline water electrolyzers is one way to lower the losses due to the
slow kinetics of the oxygen evolution electrochemical reactions.
15
However, the efficient
operation of water electrolysis systems will benefit greatly from better electrocatalysts for the
oxygen evolution reaction.
1.3 Alkaline Fuel Cell
Along with metal-air batteries, the alkaline fuel cell is also interesting for large scale
energy storage because of its simplicity and low cost. Such fuel cells typically use hydrogen and
oxygen to produce water and electricity. The alkaline fuel cell can be combined with water
electrolysis to store energy in a regenerative mode.
16
The alkaline condition allows the kinetics
of the oxygen reduction to be facile, and a variety of inexpensive non-platinum metal catalysts
6

are stable in alkaline media.
17
The alkaline fuel cell also allows for a wider choice of fuels, such
as hydrazine, ammonia, and borohydride in addition to hydrogen.
18
In an alkaline fuel cell, the
oxygen is reduced to hydroxide ion at the positive electrode and then the hydroxide is
transported through an anion exchange membrane or an immobilized alkaline electrolyte such as
potassium hydroxide to the negative electrode where the hydroxide ions combine with hydrogen
to produce water. The reactions in an alkaline fuel cell are described below
(+) Positive electrode: 1/2O
2
+ H
2
O + 2e- → 2OH
¯
(10)
(-) Negative electrode: H
2
+ 2OH
¯
→ 2H
2
O + 2e
-
(11)
Overall reaction: H
2
+ 1/2O
2
→ H
2
O (12)
The alkaline fuel cell was developed and by Dr. T.Bacon between the 1930s and 1950s.
In 1952 Bacon completed construction and evaluated performance of a 5 kW hydrogen-oxygen
power plant.
19
In this fuel cell lithiated nickel oxide and nickel were used as electrocatalyst
materials for positive and negative electrodes, respectively. In 1962, NASA has used the alkaline
fuel cell to supply electric power and drinking water for the Apollo missions, and later continued
to use them in the Space Shuttle Orbiter.
20
These fuel cells used palladium and gold-based
electrocatalysts. The major disadvantages of the alkaline fuel cell are its high cost of
electrocatalysts and its limited lifetime due to the re-distribution of the alkaline electrolyte.
21

Recently, there has been renewed interest in the alkaline fuel cell due to the possibility of using
non-precious metal catalysts for facilitating the kinetics of oxygen reduction.
22-25
The efficiency
of the alkaline fuel cell is again determined to a large extent by the overpotential losses in the
oxygen reduction process. More efficient and low-cost electrocatalysts are needed to achieve
high efficiency of energy conversion.
25

7

1.4 Oxygen Reduction Reaction
As described in the previous sections, the oxygen reduction reaction (ORR) is the
principal process at the positive electrode during discharge of metal-air batteries and the alkaline
fuel cell. This reaction involves multiple electron transfer steps. In alkaline electrolyte ORR
occurs by two major pathways for the conversion of oxygen to hydroxide: the direct four-
electron reduction pathway and the series two-electron reduction pathway:
Direct Pathway:
O
2
+ 2H
2
O + 4e
-
→ 4OH
¯
(13)
Two-electron pathway:
O
2
+ H
2
O + 2e
-
→ HO
2
¯
+ OH
¯
(14)
The hydroperoxide (HO
2
¯
) can be further reduced electrochemically as,
HO
2
¯
+ H
2
O + 2e
-
→ 3OH
¯
(15)
or the hydroperoxide can undergo catalytic decomposition to produce hydroxide and oxygen,
HO
2
¯
→ OH
¯
+ 1/2O
2
(16)
The direct four-electron pathway does not result in any free hydroperoxide and is
therefore considered more efficient. In aqueous media, the oxygen reduction reaction is very
sluggish.
26-29
Thus, an electrocatalyst is needed to speed up the kinetics of oxygen reduction. The
oxygen reduction mechanism pathway is strongly dependent on the nature of the electrode
surface. Platinum favors the four-electron pathway, while gold favors the two-electron pathway
and produces peroxide.
30
Mukerjee et al.,
31
proposed, an electrocatalytic inner-sphere electron
8

transfer mechanism in which the oxygen is undergoes chemisorption on the surface of platinum
free oxide surface followed by four-electron transfer. In the case of electrocatalytic outer-sphere
electron transfer mechanism, electron transfer from the electrode surface occurs through an oxide
film layer to a solvated oxygen molecule.
Noble metals such as platinum and platinum-group metals are the most active
electrocatalysts for oxygen reduction due to their ability to support the direct 4-electron
process.
32-35
Platinum nanoparticles well-dispersed on a high surface area support like carbon
constitute the state of the art of ORR electrocatalysts.
36
Platinum-based alloys have been
investigated widely in acid media and these materials generally exhibit higher activity and
stability than platinum alone.
37-39
However, due to the high cost of platinum, the development of
noble-metal-free catalyst is viewed as the long-term solution.
40-43
Carbon doped with heteroatom
such as S, N, and P has been investigated as alternative electrocatalysts for ORR due to their low
cost.
44-50
Nitrogen-doped carbons are the most studied carbon-based catalysts for ORR.
Vertically- aligned nitrogen-doped carbon nanotubes are an example of non-noble metal ORR
catalyst that supports an apparent four-electron pathway in alkaline media.
45

Transition metal-based macrocycles have demonstrated good performance as
electrocatalysts for oxygen reduction.
43-44, 51-53
Recently, Mukerjee et al., reported an iron-based
non platinum group metal, FePhen@MOFAr-NH
3
that has similar ORR kinetics as Pt/C in
alkaline media.
54

Transition metal oxides of the perovskite, spinel and pyrochlore family are candidates as
electrocatalysts for oxygen reduction because of the redox activity of the transition metal sites
and the stability of the oxides in alkaline media.
55-59
Perovskite transition metal oxides with the
9

general formula ABO
3
have a crystal structure in which the A site is the rare-earth metal such as
lanthanum or calcium and B is the transition metal. The A site is often substituted partially with
calcium, barium or strontium. For example, partial substitution of lanthanum by calcium or
strontium as in the compositions, La
0.6
Ca
0.4
CoO
3
or La
0.8
Sr
0.2
CoO
3
, results in a mixture of Co
2+
,
Co
3+
, Co
4+
and generation of oxygen vacancies, accompanied by an increase in electrical
conductivity.
60-61
Also, when the transition metal in the B-site of the perovskite is substituted by
one or more of the first-row transition metals, a variety of d-electron configurations are presented
at the surface. The possibility of tuning the composition at the A and B site for electrocatalytic
activity has thus evoked considerable interest for the designing of electrocatalysts.
59, 62-64

Examples of compositions that have been reported to have high electrocatalytic activity for
oxygen reduction are Pr
0.6
Ca
0.4
MnO
3,
La
0.6
Ca
0.4
CoO
3
65
, Ca
0.9
La
0.1
MnO
3
66
, La
0.5
Sr
0.5
CoO
3,
La
0.99
Sr
0.01
NiO
3

67
, Ln
1-x
Ca
x
CoO
δ-x
(Ln=La,Er).
67
Hermann et al.,
68
demonstrated that
La
0.6
Ca
0.4
CoO
3
-activated carbon reduces oxygen by two electron-pathway and HO
2
-
is further
reduced or chemically decomposed by transition metal oxide. Carbonio et al.,
69
showed that
oxygen reduction activity of LiFe
x
Ni
1-x
O
3
on carbon-based electrode correlates very well with
the catalytic activity for hydroperoxide decomposition. Along with perovskite oxides, transition
metal spinel oxide with the formula AB
2
O
4
have also shown significant activity. In the spinel,
the B site is reported to play an important role in determining the activity of the catalyst.
70-71
.
Similar to the perovskite-type transition metal oxide, the spinel oxide, MnCo
2
O
4
also followed
the 2 electron pathway with the formation of hydroperoxide.
72-73
However, the fundamental
catalytic processes that rely on the transition metal in these oxide systems has not been
understood.
10

1.5 Oxygen Evolution Reaction
The oxygen evolution reaction (OER) is the process that occurs at the positive electrode
during the charging of metal-air batteries and in water electrolysis.
4OH
¯
→ O 2 + 2H 2O + 4e
-
(17)
Today, iridium oxide, bismuth ruthenate, and ruthenium oxide are quite active as
electrocatalysts for OER.
74-77
A recent benchmarking of electrocatalysts for oxygen evolution
reported in 1 M NaOH at 10 mA/cm
2
showed that iridium oxide exhibited an overpotential of
325 mV and is still best catalyst.
78
Several different mechanisms for OER have been proposed by
Bockris
79
Matsumoto
80
and Trasatti
76
and Sunde et al.
81
The most probable mechanism on the
surface of mixture IrO
2
-RuO
2
is as follows.
Step 1 is the adsorption of OH- on the active surface S
S + H
2
O → S-OH + H+ + e-
The adsorbed OH
-
will follow an ―oxide‖ path or ―electrochemical oxide‖ path to lead to the
formation of the oxygen-covered surface.
Step 2a: Oxide path
2S-OH → S-O + S + H2O
Step 2b: Electrochemical oxide path
S-OH → S-O + H+ +e-
In step 3, two unstable S-O surface species will allow for the formation of oxygen
11

S-O → O
2
+ 2S
Despite noble metals and noble metal oxides showing good electrocatalytic activity for
oxygen evolution reaction, the high cost of these precious metal-based materials preclude cost
effective deployment on a large scale.
74-77
Therefore, research has been focused on the transition
metal perovskites and spinel oxides. Nickel cobalt oxide, NiCo
2
O
4
with a spinel structure is one
of the most promising electrocatalyst materials for oxygen evolution in alkaline media. NiCo
2
O
4

also exhibited higher OER activity than Co
3
O
4
.
82-86
Recently, iron has been shown to have a
great effect on the OER activity in nickel-iron oxyhydroxide, and NiCo
2
O
4
.
87-90
Zhao et al.,
87

have developed a high efficiency OER electrode by electrodepositing Ni-Fe on to macroporous
nickel foam. This electrode shown an overpotential of 240 mV at 500 mA/cm2 in 10M KOH.
The understanding of the mechanism of oxygen evolution and the role of the transition metal are
still inadequately understood to permit the design of new electrocatalysts.
1.6 Problems and Challenges
As described above, the research on electrocatalysts for the cathodic oxygen reduction
reaction (ORR) and the oxygen evolution reaction (OER) are crucial for achieving the cost,
efficiency and durability of metal-air batteries, fuels cells and water electrolysis systems.
91
,
92

The slow kinetics of the ORR at the cathode limits the efficiency of metal-air batteries and fuel
cells.
26-29
The undesirable voltage losses resulting from the slow kinetics result in the reduction
in the round-trip energy efficiency of the battery. Approximately, 70% of voltage loss in metal-
air batteries, fuel cells and water electrolysis system originates from the oxygen electrode. An
example of this is the bifunctional air electrode in the iron-air battery that was reported by the
Swedish National Development Company.
93
Further, the process of oxygen evolution causes
degradation of the electrode structure and catalyst materials due to the oxidation of carbon
12

material in electrode structure. Therefore, increasing the energy efficiency and durability of the
air electrode continues to be a major focus area for improving the performance of metal-air
batteries, fuel cells, and water electrolyzers.
40, 94-95

For a rechargeable metal-air battery, the catalyst materials or electrodes must be ―bi-
functional‖ in that they must support both oxygen reduction and oxygen evolution. A variety of
inexpensive materials such as manganese dioxide, cobalt oxide, metal phthalocyanines and high
surface area carbon have been used in oxygen reduction in zinc-air batteries and alkaline fuel
cells.
47, 59, 61-62, 68, 96-108
However, many of these catalyst materials are not active or chemically
stable when used in the oxygen evolution mode. Therefore, development of electrocatalyst
materials for the oxygen reduction and oxygen evolution reactions that are cheap, highly
efficient, and stable is one of the prime challenges for fuel cells and metal-air batteries.
52

Transition metal oxide perovskite and spinel oxide are highly promising as electrocatalyst
for oxygen reduction and oxygen evolution reactions because of the following compelling
reasons: (a) low cost resulting from the use of abundantly-available materials (b) environment-
friendliness (c) ability to support both OER and ORR activity (d) the possibility of tuning the
activity by varying the transition metal composition, and (e) good stability in alkaline media.
Thus, in the present study, we have focused on fundamental studies to understand the mechanism
of oxygen reduction and oxygen evolution on transition metal oxides.
1.7 Objectives of the Research
The overall objective is to find a pathway to the design of transition metal oxides based
catalysts for efficient reduction of oxygen and oxygen evolution. The specific objectives are:
13

1. Understand the mechanism of oxygen reduction on transition metal oxides catalyst to
improve their catalytic activity.
2. Understanding the relationship between the decomposition rate of hydroperoxide and
the oxygen reduction activity of the metal oxide catalysts.
3. Understanding the factors affecting oxygen evolution on selected spinel and
perovskite oxides.
4. Developing methods of preparation of nanocrystalline oxides and coated electrodes.
5. Demonstrating electrodes based on the optimally performing catalysts and testing in
an iron-air cell for at least 100 hours.

14

Chapter 2
Experimental Techniques and Methodology for
Analysis of Results
2.1 Catalyst Synthesis
The transition metal oxides studied in this work were synthesized in-house by a modified
Pechini process.
3, 109
Stoichiometric amounts of metal nitrates were dissolved in deionized water
(18.2 MOhm cm
-1
) along with citric acid. This solution mixture was slowly evaporated at 80
o
C
to form a sol-gel. This sol-gel was left at room temperature for 18 hours and was then heated in a
vacuum oven at 90
o
C for 12 hours. The resulting sol-gel product then underwent heat treatment
in air in two sequential steps. First, the material was heated at 150 °C for 30 minutes upon which
a black, high-surface-area oxide was rapidly formed, accompanied by combustion. This black
powder was ground gently with a mortar and pestle for about 15 minutes, and then the powder
was heated at 700°C for 2 hours to yield the required final sample of metal oxide. The second
step of heat treatment was varied between 200 and 700
o
C to study the effect of preparation
temperature.
2.2 Physical Characterization
2.2.1 X-Ray Diffraction
The phase analysis of the samples was performed by X-ray Diffraction Analysis (Rigaku
Ultima IV, using Cu K
α
). The crystallite size was calculated from the full-width at half maximum
of the (220) reflections using the Scherrer formula (Eq. 18) after correcting for instrument
broadening.
15

Where B is full-width at half maximum (FWHM), L is the crystallite size, λ is the X-ray
wavelength, θ is the Bragg angle, and K is the Scherrer constant. The constant K depends on how
the breadth B is determined and the actual shape of the crystallite. The most common value of
constant K is 0.94.
110-111

2.2.2 Scanning Electron Microscopy
The morphology of the catalysts was viewed and recorded using a scanning electron
microscope (SEM, JEOL JSM 7001).
2.2.3 Surface Area
The surface area of the sample was determined by the Brunauer–Emmet–Teller (BET)
method using the instrument Quantachrome NOVA 2200e. Approximately 120 mg of metal
oxide powder was used for measuring the BET surface area. The sample was degassed at 250
o
C
for 6 hours before transfering to the analyzer chamber, and nitrogen was used as the adsorbate.
2.2.4 X-ray Photoelectron Spectroscopy
The oxidation states of the metals in the catalysts were studied using X-ray photoelectron
spectroscopy (XPS) with a magnesium X-ray source (1253.6 eV, SPECS XPS at NETL). The
oxide samples were heated in the vacuum oven at 150
o
C overnight and then coated on carbon
tape before loading into the chamber. The XPS data was analyzed by using CASA software and
all the XPS data was corrected based on the carbon peak at 284.6 eV.
16

2.2.5 X-ray Absorption Near Edge Spectroscopy
X-ray Absorption Near Edge Spectroscopy (XANES) studies were performed at Sector
20-bending magnet beam-line in the Advanced Photon Source with the assistance of
Balasubramanian‘s group at Argonne National Laboratory. Before the X-ray absorption
measurements, the oxide samples were ground into fine powders and an appropriate quantity was
mixed thoroughly with boron nitride and cold-pressed into pellets. Care was taken to minimize
distortions to the absorption spectra from thickness effects. The measurements were carried out
in the transmission mode using a Si (111) mono-chromator. A rhodium-coated harmonic
rejection mirror was used to minimize harmonic contamination. Reference manganese and cobalt
foils were used for energy calibration. The threshold energies (defined using the position of the
first inflection point) at the manganese and cobalt K-edges were taken as 6537.7 eV and 7708.8
eV, for the respective foils.
112
The relative uncertainty in energy between the various samples
was estimated to be   0.05 eV. Data reduction followed standard procedures using the Athena
program in the IFEFFIT suite of software.
113

2.3 Electrochemical Characterization
2.3.1 Catalyst Ink Preparation
The catalyst ink was prepared by mixing 8 mg of catalyst, 2 mg of acetylene black, and 2
ml of a mixture containing of 89.5 vol% water, 10 vol% isopropanol, and 0.5 vol% Nafion
solution (using 5% solution of Nafion® 1100 EW ionomer, Sigma Aldrich). The mixture was
subjected to ultrasonic agitation for 40 minutes to form a uniformly dispersed ink.
17

2.3.2 Electrode Preparation
2.3.2.1 Preparation of Electrodes for the Oxygen Reduction Reaction and Oxygen Evolution
Reaction Studies
Preparation of catalyst on Rotating Disk and Rotating Ring Disk Electrodes
A 20 µl droplet of the catalyst ink solution (prepared as in section 2.3.1) was coated on
the clean surface of a glassy carbon disk electrode. For ORR studies, a ring-disk electrode was
used in which the disk electrode was concentric to a platinum ring electrode. OER studies were
carried out a glassy carbon rotating disk electrode without a ring electrode. (Figures 1 and 2).
After coating, the electrode was dried in air for 15 minutes at 85
o
C and then cooled to room
temperature. In all the experiments, the loading of the oxide catalyst on the glassy carbon disk
electrode was approximately 80 µg. This choice of amount is consistent with the range of
catalyst loading on rotating disk electrodes (RDE) used by others.
114-116

Figure 1. Rotating-Ring Disk Electrode with 20 L droplet of the catalyst ink.
18

Figure 2. Rotating disk electrode with a thin film of catalyst after drying in air at 85
0
C.
Preparation of Spinel Iron-doped Nickel Cobalt Oxide on Nickel Foam Electrodes for Oxygen
Evolution Studies
A stock solution of the precursor for the coating of electrodes was prepared by mixing
100mL of 0.2M of nickel nitrate (Ni(NO
3
)
2
Sigma Aldrich) solution, 100mL of 0.4M cobalt
nitrate (Co(NO3)
2
Sigma Aldrich) solution, 25mL acetone and 25mL isopropanol. This solution
was used for the preparation of the nickel cobalt oxide. 10 mL of stock solution was mixed with
1 mL of 0.1M ferric nitrate (Fe(NO
3
)
3
Sigma Aldrich) to prepare the coating solution for the
iron-doped nickel cobalt oxide.
19

Figure 3. Catalyst-coated on nickel foam electrode (black coating is the oxide).
A rectangular piece of nickel foam of size 10 cm x 3 cm with a pore size 800 micron was
used to prepared the electrode. An area of 7 cm x 3 cm was then coated with the oxide catalyst
(Figure 3). The coating was made by the following procedure: First the nickel foam was placed
on a hot plate. The temperature of the hot plate was maintained at 250
o
C. The coating solution
was divided in two portions for coating in two steps. The first half of the coating solution was
added dropwise on the foam and dried. After adding a few drops of solution, the coating was
allowed to dry before application of the next layer of the coating solution. Approximately 1 mL
of solution was added each time. Spillover of the solution on to the hot plate must be avoided to
prevent the loss of the catalysts materials during the preparation process. After completing the
coating and drying operation with half the amount of starting ink, the foam was placed in a
furnace and heat treated for 30 minutes at a designated temperature in the range of 200 to 700
o
C.
The coating and drying operation on the same nickel foam with the rest of the ink was repeated
to complete the coating of catalyst. Subsequently, the foam was heat treated again in the furnace.
20

The coating process and the temperatures for heat treatment steps were kept the same for each
step.
2.3.2.2 Electrode Preparation for Hydrogen Peroxide Reduction and Oxidation Study
A mixture containing 50 mg of catalyst, 291 mg of 5% Nafion
®
solution, and 166 mg of
deionized water was sonicated for 20 minutes. This well-dispersed mixture was coated on a 3 x 5
cm Toray
®
paper electrode using a paint brush (Figure 4). The electrode was then heated in the
furnace at 125
o
C for 10 minutes.

Figure 4. 3 cm x 5 cm Toray® paper with the coating of catalyst.
2.3.3 Electrochemical Experimental Set up
Electrochemical testing was conducted in a three-electrode polyfluoroethylene cell
(Figure 5). A mercury/mercuric oxide (MMO) reference electrode (20% potassium hydroxide
solution, E
o
= +0.098 V) for all experiments. A platinum wire counter electrode was used for
ORR and OER testing and a glassy carbon rod was used as the counter electrode for the
hydrogen peroxide study (Figure 5 and 6). For the ORR and OER measurements the rotating
ring disk electrode and the glassy carbon rotating disk electrode were used as the working
21

electrodes, respectively. For the hydrogen peroxide studies, the Toray
®
paper coated with catalyst
was used as a working electrode (Figure 6). High-purity water (18.2 M ohm cm
-1
, 4ppb total
organic carbon) was used in all experiments. A 1 M solution of potassium hydroxide was used as
the electrolyte in all tests. The electrolyte was continuously purged with high purity oxygen or
argon (UHV grade 99.999%) as applicable to the experiment.

Figure 5. Polyfluoroethylene cell setup for ORR and OER measurements.

Figure 6. Polyfluoroethylene cell setup for hydrogen peroxide measurements.
22

For studying the catalyst-coated nickel foam electrode, electrochemical testing was
performed in a three-electrode polyethylene cell with a mercury/mercuric oxide (MMO)
reference electrode (20% potassium hydroxide solution, E
o
= +0.098 V) and nickel mesh as the
counter electrode (Figure 7). High-purity water (18.2 M Ohm cm
-1
, 4ppb total organic carbon)
was used in all the experiments. 6 M potassium hydroxide solution was used as the electrolyte.

Figure 7. Electrochemical cell setup for the OER measurement of catalyst coating on nickel
foam electrode.
2.3.4 Electrochemical Method of Analysis
2.3.4.1 Rotating Disk Electrode
The Rotating Disk Electrode (RDE) consists of a circular disk electrode embedded in an
insulator rod connected to a rotating shaft. The rotation rate of the shaft is controlled by an
electric motor with a speed controller (Figure 8). As the electrode rotates a laminar flow of
solution is set up perpendicular to the electrode (along the axis of rotation) accompanied by
removal of liquid radially across the electrode. (The flow pattern set up at the RDE can be seen
23

at this video link that I have recorded and posted. https://youtu.be/WfRbCOXANrk) This type of
steady state laminar flow across the surface of the RDE will allow us to develop a boundary
layer of constant thickness (Figure 9).
117

Figure 8. Schematic of a rotating disk electrode
In the boundary layer, diffusion is the only means of transport. Therefore the mass-
transport-limited current at a rotating disk electrode is determined by the thickness of the
diffusion layer, the diffusion coefficient and the concentration of the electro-active species by the
relationship I = nFA D (C
o
*
-C
o
)/δ where n is the number of electron transferred in the reaction, F
is a Faraday constant, A is the area of the electrode, D is the oxygen diffusion coefficient, C
o
*
is
the concentration of oxygen in the bulk, C
o
is the concentration of oxygen at the surface of the
electrode, and  is the diffusion layer thickness. The diffusion layer thickness,  is determined by
24

the rotation rate, the diffusion coefficient, and the viscosity by according to δ= 1.61D
1/3

-1/2

1/6
,
where D is the diffusion coefficient of reactant in the electrolyte,  is the rotation rate of the
electrode (in radian/s), and  is the kinematic viscosity of the electrolyte. Therefore, the mass-
transport limited current at a rotating disk electrode is obtained by combining these foregoing
equations and setting the concentration at the surface to zero. We thus obtain the Levich
equation,

Where n is the number of electrons involved in the reaction, F is Faraday constant, A is the area
of the electrode.
When the rate of reaction at the surface of the electrode is limited by mass transport and slow
kinetics, it is possible to determine the kinetic parameters for charge-transfer and mass transport
by using the Koutecky-Levich equation to analyze the data from the rotating disk measurement.

Here I is the measured current, I
k
is charge transfer current, and I
lim
is limited mass-transport
current or diffusion-limited current. By substituting Eq. 19 into the Koutecky-Levich equation
we have the relationship between current and rotation rate in terms of both charge-transfer and
mass transport current (Eq. 21)

25

Figure 9. Relationship between diffusion layer thickness  and rotation rate
At a constant bulk concentration and at a given electrode potential, the observed current
will depend on the rotation rate as per Eq.21. When the rotation rate increases the concentration
at the interface will also increase, hence also the observed current. Therefore as ω tends to
infinity the observed current may be considered to be just the ―kinetic current‖. By plotting 1/I
vs 1/ ω
1/2
we can determine from the intercept the kinetic current and from the slope of the line
the number of electrons transferred, n (Figure 10)
26

Figure 10. A expected plot of 1/I vs 1/ω
1/2
at a potential -200 mV for various rotation rate of
400, 900, 1600, and 2500 rpm

Figure 11. ORR polarization curve measured by rotating disk electrode at various rotation rate
A typical ORR polarization curves at various potential generated by rotating disk
electrode are presented in Figure 11. The region from 0 mV to -300 mV is the region of mixed
kinetic control and mass-transport control but the plateau region form -300 mV to -450 mV is
27

just the mass-transport-limited current. Thus, kinetic currents can be calculated from the
measured current-potential curves by correcting for mass transport effects. The kinetic current is
given by:

Where I
obs
is the measured current and I
lim
is the mass-transport limited current.
From the kinetic currents I
k
that were obtained after the mass-transport correction
(equation 22), the Tafel parameters were also determined by plotting the overpotential η versus
the log I
k
(kinetic current) by using the Tafel equation (Eq. 23). The Tafel slope can be obtained
from the slope of this line and the exchange current density I
o
can be calculated from the
intercept of this line with the X axis at η = 0 (Figure 12).
(23)

Figure 12. A typical Tafel plot
28

2.3.4.2 Rotating Ring Disk Electrode (RRDE)
Use of a rotating ring disk electrode is a well-known method to study the intermediate
species of a chemical reaction. In this work the RRDE was used to study hydrogen peroxide
formation during the oxygen reduction reaction process (Figure 13).
118

Figure 13. Rotating ring disk electrode with glassy carbon disk and platinum ring electrode
As described earlier, when the disk electrode is rotated the solution will be drawn upward
then sent radially and laminar flow is achieved. (Figure 14) Thus, oxygen from the bulk will
follow the laminar flow and arrive at the surface of the disk electrode. When the electrode is
negatively polarized, on certain surfaces hydroperoxide will be produced. The hydroperoxide
will then be convectively transported to the platinum ring electrode and can be detected by an
oxidation current when the ring electrode is polarized appropriately. This method is used widely
for detecting peroxide intermediates in oxygen reduction.
29

Figure 14. Schematic of flow of oxygen to the RRDE when the electrode is rotated
So, during the rotating ring disk experiments, the disk electrode was scanned to potentials
negative to the reversible potential for oxygen reduction at a scan rate of 2 mV/s.
Simultaneously, the ring electrode was held constant at + 0.5 V vs. MMO for oxidizing any
hydroperoxide intermediate. Such an experiment is conducted using a bipotentiostat that has the
ability to control the potential of two electrodes simultaneously using a single reference electrode
and counter electrode. In the RRDE measurement only a fraction of concentration of the
intermediate species are detected at the ring electrode due to the radius and size of the ring-disk
assembly, resulting in a parameter called collection efficiency N,
(24)
The collection efficiency was measured by using the reversible half reaction
ferrocyanide/ferricyanide in which the reduction of potassium ferricyanide at the disk will
30

generate potassium ferrocyanide that is detected at the ring by an oxidation current. Details of
this procedure are available in the literature reference
119
.
The electron-transfer number (n =2 to 4) and the percentage of hydroperoxide were determined
by the following equation.
120

(25)

2.3.4.3 Measurement of Oxygen Reduction Rates
The electrolyte was saturated with oxygen at 1 atm for at least 50 minutes prior to the
oxygen reduction study. The oxygen reduction activity was studied by linear-sweep voltammetry
at scan rate of 2 mV/s between 0.1 V and -0.5 V vs. MMO while the working electrode was
rotated at 400, 900, 1600, and 2500 rotations per minute (rpm). Before recording the polarization
curves, the working electrode was cycled between 0.1V and -0.5 V at 20 mV/s for 20 cycles to
obtain an invariant voltammogram. All measurements were performed with a bi-potentiostat
(Autolab PGS30). A rotating ring disk electrode (Pine Instruments Inc.) consisting of a 5 mm
diameter glassy carbon disk and a 7 mm diameter platinum ring was used for studying the
hydrogen peroxide intermediates produced during oxygen reduction. The ring electrode was
held at +0.5 V vs. MMO for oxidizing the peroxide arriving from the disk electrode. The
collection efficiency of ring-disk combination was determined to be 0.20 by polarization
experiment that used a potassium ferricyanide solution purged with argon.
31

2.3.4.4 Measurement of Oxygen Evolution Rates
The oxygen evolution reaction (OER) study used a setup similar to that for the ORR
study but a glassy carbon rotating disk electrode (RDE) was used as the working electrode
instead. The electrolyte was continuously purged with high-purity argon (Airgas Ultra-high
purity grade 99.999%) to remove dissolved air from the electrolyte and to maintain a carbon
dioxide-free environment. The oxygen evolution activity was investigated by steady-state
potentiostatic polarization measurements in the electrode potential range of 400 mV to 650 mV
vs. MMO by holding at each of the potential values for at least 300 seconds before recording the
steady state value of current.
For the catalyst-coated on nickel foam electrode, steady-state potentiostatic experiments
were carried out at 400 mV, 450 mV, 475 mV, 500 mV, 525 mV, 550 mV, 575 mV, 600 mV,
615 mV, 630 mV, and 650 mV. The duration of measurement varied between 300 s and 900 s
depending on stability of the currents. The ohmic resistance contribution from the electrolyte was
determined using electrochemical impedance spectroscopy from the high-frequency intercept of
the Nyquist plot. Each of the potentiostatic measurements were corrected for the ohmic potential
drop using the resistance values determined from impedance spectroscopy. The durability tests
were conducted as a chronopotentiometric measurement for 65 hours at a constant current
density of 10 mA/cm
2
.
2.3.4.5 Rate of Hydrogen Peroxide Reduction and Oxidation
The electrolyte was first saturated with argon at 1atm for 50 minutes before the hydrogen
peroxide study was started. The 3 x 5 cm Toray® paper with the appropriate catalyst coating
was used as the working electrode. Typically, 1.283 mL of 9.79 M hydrogen peroxide (hydrogen
peroxide 30% Suprapur®, EMD Millipore) was added to 250 mL of 1M potassium hydroxide
32

solution to make a 50 mM hydrogen peroxide solution. The solution was continuously stirred
during the measurements to maintain a steady-state mass transport. Prior to the polarization
measurements with hydrogen peroxide, the open-circuit potential of the working electrode was
measured. Based on the open-circuit potential for each catalyst, the steady-state polarization
measurements were planned such that hydrogen peroxide reduction and hydrogen peroxide
oxidation could be investigated.

Figure 15. Hydrogen peroxide reduction and oxidation polarization experiments in argon
saturated 1M potassium hydroxide containing 50 mM hydrogen peroxide.
33

The decomposition of hydrogen peroxide mediated by redox couples can be viewed as
two conjugate reactions occurring on the electrode surface. These two reactions will result in a
mixed potential as governed by the kinetics of the two processes (Figure 15).
Hydrogen peroxide reduction reaction
In base: HO
2
-
+ H
2
O + 2M
(n-1)+
→ 2M
n+
+ 3OH
-
(27)
Hydrogen peroxide oxidation reaction
In base: OH
-
+ HO
2
-
+ 2M
n+
→ O
2
+ H
2
O + 2M
(n-1)+
(28)

The kinetics of these two reactions (peroxide oxidation to oxygen and the peroxide
reduction to hydroxide) are determined by the facility which they are mediated on the surface of
the metal oxide. Under open circuit conditions when decomposition is occurring, there is no net
flow of current external to the electrode. Therefore, the oxidation and reduction currents are
equal and the electrode exhibited a mixed potential E
mixed
. The polarization curves of
hydroperoxide reduction and oxidation were obtained by polarizing cathodically and anodically
to the value of E
mixed
. The Tafel lines for hydrogen peroxide reduction and oxidation were
generated by plotting the potential vs. Log I using Eqs. 29 and 30.

The subscripts 1 and 2 refer to the oxidation and reduction processes, respectively.
From Eqs. 29 and 30 the value of decomposition rate I
corr
and the calculated open circuit
potential was determined by the following relationship
34

The exchange current density for the hydrogen peroxide oxidation and reduction was obtained by
the Eqs. 33 and 34, respectively.

Where , , b
1,
and b
2
are standard reduction potential of hydrogen peroxide
oxidation reaction, standard reduction potential of hydrogen peroxide reduction, Tafel slope of
peroxide oxidation and Tafel slope of peroxide reduction, respectively.
2.3.4.6 Direct Measurement of Rate of Hydrogen Peroxide Decomposition by the Manometer
Method
The hydrogen peroxide decomposition reaction was carried out in a round bottom flask
which was connected to a manometer tube to measure the volume of oxygen evolved that
displaced water in the tube (Figure 16). A 150 mL of 1M potassium hydroxide containing
approximately 2mg of transition metal oxide was saturated with oxygen prior to the experiment.
The catalyst was placed in a vial with water and then placed in ultra-sonicator for 5 minutes
before the suspension was added into the electrolyte solution. An appropriate amount of 30%
hydrogen peroxide was added into the electrolyte to make the 50 mM hydrogen peroxide
solution. The volume of oxygen produced by decomposition of peroxide was determined by the
35

amount of water displaced in the manometer. The solution was stirred throughout the
decomposition experiment.

Figure 16. Direct hydrogen peroxide decomposition experiment setup
2.3.5 Iron-Air Cell Studies
2.3.5.1 ORR Electrode Preparation
ORR electrode was prepared by coating catalyst ink on a 5.6 cm x 5.6 cm Toray® paper
using a paint brush (Figure 17). The electrode was constructed of two layers. The first layer
contained 30 mg of catalyst, 60 mg of PTFE (balls MP1300), 84 mg of 5% Nafion solution, 630
mg of deionized water, and 160 mg of isopropanol. The second layer contained 60 mg of
catalyst, 280 mg of 5% Nafion solution, 730 mg of deionized water, and 220 mg of isopropanol.
36

The solution mixture for each layer was subjected to 15 minutes of ultrasonic agitation before the
coating processes. After coating the electrode was held in a furnace at 125
o
C for 10 minutes.

Figure 17. ORR electrode of NiCo
2
O
4
-vulcanXC-72
2.3.5.2 OER Electrode Preparation
The method for preparation of the OER for the iron-air cell studies is similar to the
procedure described in section 2.3.2.1 with the desired stoichiometric amounts of metal nitrates
and chosen heat treatment temperature. The OER electrodes in the iron-air cell studies used a
5cm x 5cm nickel foam as the substrate for the oxide coating (Figure 18).
37

Figure 18. OER electrode with Fe doped NiCo
2
O
4
coated on nickel foam.
2.3.5.3 Iron-Air Cell Set up
The configuration of test cell for the iron-air battery experiments is shown in Figure 19. The cell
was made of polypropylene. The negative (iron) electrode was assembled in the middle
compartment of the cell. The air electrodes for oxygen evolution were separated from that used
for oxygen reduction (Figure 20). The two OER electrodes stayed on both sides of the iron
electrode (Figure 20). The two ORR electrodes were held in the two stainless steel compartments
on both sides of the cell and these two stainless steel compartments had ports through which
oxygen was supplied to the electrodes. A mercury/mercuric oxide (MMO) (20 w/v% potassium
hydroxide solution, E
o
= +0.098 V) was used as a reference electrode and 30 w/v% potassium
hydroxide was as the electrolyte. The charge/discharge studies were carried out using a 16-
channel MACCOR 4200 battery cycler.
38

Figure 19. (a) Configuration of iron-air cell (b) Iron-Air cell test under test

Figure 20. Separate electrode configuration for air electrode
39

Chapter 3
Understanding the Role of Carbon and Transition
Metal Oxide in the Oxygen Reduction Reaction
In most of the previous studies on the use of transition metal oxides as catalysts for the
oxygen reduction reaction, these oxides have been mixed with a highly-conducting carbon such
as acetylene black or graphite. Typically, a mixture consisting of approximately 80% (w/w)
oxide and 20% (w/w) of high surface area carbon is used.
61, 104, 121
The carbon additive is
described as necessary to enhance the electrical conductivity of the poorly-conducting transition
metal oxide catalyst.
105-107
Further, these studies on ―carbon-containing composite catalysts‖
ascribe the observed variations in catalytic activity for the electro-reduction of oxygen entirely to
the properties of the transition metal oxides, and the electrocatalytic activity of carbon is often
completely ignored. Our results show that the role of carbon as an electrocatalyst is primary to
the function of such composite catalysts.
59, 62
Therefore, in the present study, we focus on
understanding the catalytic activity of the transition metal oxide when used with or without
carbon additives.
We have observed that the oxygen reduction activity of even conducting oxides such as
calcium-doped lanthanum cobalt oxide (LCCO) is substantially enhanced by adding carbon. In
the case of LCCO, since the oxide is already a good conductor of electrons, the substantial
enhancements in oxygen reduction activity caused by the addition of carbon becomes intriguing.
Also, we find that the oxygen evolution activity of these oxide catalysts was not improved by the
presence of carbon additives. It appears from the foregoing observations that the role of carbon
in oxygen reduction goes beyond enhancing the electrical conductivity of the composite layer.
40

Recent reports from other research groups also suggest a synergistic interaction between carbon
and the transition metal oxide in catalyzing the oxygen reduction reaction.
68, 108, 115-116, 122-123

Such a finding has significant implications for the interpretation of oxygen reduction reaction
studies on transition metal oxides with and without carbon. Therefore, the objective of our study
was to determine the specific roles of the transition metal oxide and carbon in the composite
catalysts. To this end we have focused largely on the perovskite oxide of the molecular formula,
La
0.6
Ca
0.4
CoO
3-x
(LCCO) combined with a conductive carbon, acetylene black (AB).
Table 1. Composition of catalyst mixtures consisting of LCCO and Acetylene Black (AB) that
were used in our studies

LCCO was mixed with acetylene black in various proportions shown in Table 1 to
prepare the composite catalysts. The catalyst inks were prepared using the procedure described in
Chapter 2. In three other separate preparations, 8 mg of LCCO was combined with either (a) 8
mg gold nanoparticles (nanopowder, <100 nm particle size, 99.9% trace metal basis, Sigma
Aldrich), (b) 8 mg of carbon nanotubes (multi-walled > 98% carbon -
1 8.0 100.0 0.0 0.0
2 8.0 90.9 0.8 9.1
3 8.0 83.3 1.6 16.7
4 8.0 66.7 4.0 33.3
5 8.0 50.0 8.0 50.0
6 4.0 33.3 8.0 66.7
7 2.0 20.0 8.0 80.0
8 1.0 11.1 8.0 88.9
9 0.5 5.9 8.0 94.1
10 0.0 0.0 8.0 100.0
Composition
#
Mass%
of LCCO
Mass %
of AB
Mass of
LCCO, mg
Mass of
AB, mg
41

20 µm, Sigma Aldrich), or (c) 8 mg of graphene (8 nm flakes Grade AO-2, Graphene
Supermarket) to produce inks and catalyst layers using the same procedure in Chapter 2.

Figure 21. Polarization curves for various catalyst ink compositions in oxygen-saturated 1 M
potassium hydroxide at various rotation rates of 400-2500 rpm (a) LCCO without any added
carbon, (b) acetylene black without any LCCO (c) mix of LCCO (91%) and acetylene black
(9%), (d) comparison of the mass activity of the catalysts in (a), (b) and (c) at -0.25 V rotation
rate of 400 rpm.
(d)
(b)
AB
(c)
LCCO+AB
(a)
LCCO
42

The results of the effect of carbon on the oxygen reduction activity of LCCO are
presented in Figure 21. The kinetics of oxygen reduction on LCCO in the absence of any added
carbon appeared to be limited largely by the charge-transfer process even at -0.40 V, and the
familiar mass-transport limited plateau current was not observed even at significantly negative
electrode potentials (Figure 21a). Acetylene black exhibited much faster charge-transfer kinetics
compared to LCCO and a mass-transport limited current was observed even at -0.25 V (Figure
21b). At similar values of electrode potential, the current observed with acetylene black was
about one order of magnitude greater than that with LCCO. However, with the mixture of 9%
acetylene black and 91% LCCO (Figure 21c), the kinetic current (for example, in the potential
range of -0.1 to -0.2 V) and the mass-transport limited currents were about two times greater than
with acetylene black catalyst and twenty-fold more than with the ―LCCO-only‖ catalyst. We also
note that the ―half-wave potentials‖ for carbon, LCCO and the composite LCCO-carbon catalysts
are different. These differences are discussed later here in terms of the Tafel analysis of the
kinetic currents.
The doubling of the mass-transport limited current with the combination of acetylene
black and LCCO suggested a doubling of the number electrons transferred in the overall oxygen
reduction reaction, because the geometric area of all the electrodes tested was the same. To
determine if the enhancement in catalyst performance observed with acetylene black addition to
LCCO was caused by just an increase in the electrical conductivity of the catalyst layer, we also
conducted experiments with various other conductive additives such as carbon nanotubes,
graphene and gold nanoparticles. The results of these tests (Figure 22) show that the addition of
other carbon-based additives, namely, carbon nanotubes and graphene also enhanced the
observed activity of LCCO similar to the addition of acetylene black. However, among the
43

additives tested, the gold nanoparticles despite their good electrical conductivity had the least
impact on the oxygen reduction activity of LCCO. Thus, the enhancement in activity appeared to
be quite specific to the carbon-based additives. Also the significant differences observed among
the various types of carbon additives suggest that carbon plays a direct role in the catalysis
beyond just increasing the electrical conductivity of the composite catalyst.

Figure 22. Oxygen reduction activity in oxygen-saturated 1 M potassium hydroxide at -0.10 V
for 83% LCCO and 17% acetylene black , 83% LCCO and 17% graphene, 83% LCCO and 17%
carbon nanotubes, 83% LCCO and 17% gold nanoparticles.
As part of our effort to develop bi-functional catalysts, we also investigated the activity
of the LCCO based catalysts toward oxygen evolution. We observed that the oxygen evolution
activity was not influenced by the addition of carbon. If the role of carbon during oxygen
reduction was merely to increase the electrical conductivity of the catalyst layer, then one should
44

have observed a similar enhancement in oxygen evolution activity with carbon addition. Since
the results of oxygen evolution studies did not support this expectation, it became clear that the
carbon additives in the LCCO-carbon composite has a distinct role in enhancing the oxygen
reduction activity beside the commonly perceived function of just increasing the electronic
conductivity of the composite catalyst.
In further studies on the role of carbon we found that the oxygen reduction activity of the
composite catalyst was dependent on the mass ratio of the acetylene black to LCCO (Figure 23).
For the all the compositions that included both acetylene black and LCCO, the activity of the
catalysts exceeded that of the values predicted by the weighted sum of the activity of the
individual constituents (Figure 23c). The higher than predicted values of activity for all the
compositions of the carbon-containing catalysts confirmed the synergistic catalytic role for
LCCO and acetylene black. In comparing the activity of composition 8 and 9, we found that the
even 6 to 11 weight% of LCCO enhanced the activity of the composite catalyst significantly. In
comparing compositions 2, 3, 4 and 5 (in which we keep the mass of LCCO constant) we found
that the activity of the composite catalysts scaled linearly with the amount of carbon suggesting a
first-order dependence on the amount of carbon for all the compositions (Figure 23d). The
current also increased rapidly with small increases in the amount of LCCO, but the rise in
activity was less significant at higher fraction of LCCO (Figure 23e). Thus, the dependence of
activity of the composite catalysts on the amount of LCCO changed from a first-order
dependence for a small amount of LCCO to a near zero-order dependence at high fractions of
LCCO. Such a change in order could be because at higher fractions of LCCO in the catalyst, not
all the mass of LCCO can make direct contact with the carbon material.
45

Figure 23. (a) and (b) Polarization curves for oxygen reduction in 1M potassium hydroxide
saturated with oxygen for various mass ratios of LCCO and acetylene black; labels correspond to
sample numbers in Table 1. Kinetic current at -0.1 V at 1600 rpm: (c) at various mass ratios of
LCCO and acetylene black as in Table 1, (d) varying amounts of carbon and fixed amount of
LCCO (8mg) and (e) varying amounts of LCCO with a fixed amount of carbon (8mg). The
dashed lines in (d) and (e) are to aid the visualization of the trends.
46

The first-order dependence on carbon suggested that carbon functioned in a primary
catalyst role whereby the larger the amount of carbon, the larger was the reduction current.
However, the enhancement in activity with LCCO leveling off at 11% of LCCO suggested that
LCCO was not the primary surface for the electrochemical reaction and that LCCO functioned as
a co-catalyst that enhanced the role of carbon.
Savinova et al.,
122
have recently reported that this co-catalyst role of the perovskite oxide
is due to the decomposition of hydrogen peroxide produced on the surface of carbon. For the first
time their studies drew direct attention to higher catalytic activity of mixtures of carbon with
lanthanum cobalt oxide and strontium-doped lanthanum manganese oxide compared to that of
the individual constituents. They concluded that the oxygen reduction reaction on carbon
containing oxide composite cathodes must be considered as coupled reactions where the
individual contributions cannot always be separated. They have recommended further validation
of this mechanism by rotating ring disk methods. We also recognize that in the study of cobalt
containing non-platinum catalysts reported by Atanassov et al.,
116, 123
suggest a dual-site
synergistic process. Here, the authors identify the cobalt oxide on the nanoparticle to be the site
for the destruction of peroxide.
Therefore, in the present study we have focused on understanding the mechanism of
interaction between the oxides and carbon in greater detail by extensive analysis of ring-disk
results.
In addition to LCCO, we have observed similar enhancements in oxygen reduction
activity when lanthanum calcium cobalt manganese oxide (La
0.6
Ca
0.4
Co
0.5
Mn
0.5
O
3-x
, LCCM),
and nickel cobalt oxide spinel are combined with carbon (Figure 24). We find that when nickel
47

cobalt oxide is combined with carbon it is at least an order of magnitude more active compared
to LCCO and LCCM. It is therefore clear that the chemical constitution of the transition metal
oxide plays a crucial role in the observed kinetics of oxygen reduction. Thus our results also
indicated that the degree of enhancement in activity varied with the type of transition metal oxide
(Figure 24) and the type of carbon additive (Figure 22). Therefore, while studying the oxygen
reduction activity of transition metal oxides mixed with carbon it is important not to ascribe the
observed activity differences to just the electrochemical processes on the transition metal oxide.

Figure 24. (a) Polarization curves for oxygen reduction in 1M potassium hydroxide at 1600 rpm
for various catalysts consisting of LCCO, lanthanum calcium cobalt manganese oxide
(La
0.6
Ca
0.4
Co
0.5
Mn
0.5
O
3-x
, LCCM), and nickel cobalt oxide (NC). Curves correspond to I:
LCCO, Ia: LCCO+C, II: LCCM, IIa: LCCM+C; III: NC, IIIa: NC+C. (b) Oxygen reduction
activity at -0.1 V with various catalysts with and without carbon. Carbon content in the
composite catalyst was 16.7% (w/w).
We describe below our efforts to confirm the mechanism by which the combination of
carbon and transition metal oxides gave rise to the enhanced oxygen reduction activity.
48

The oxygen reduction reaction in alkaline media is known to occur by three familiar
pathways
124
represented by the following chemical reactions:
(1) Direct four-electron pathway
O
2
+ 2H
2
O + 4e
−
→ 4OH
−
[35]
(2) ―Series‖ pathway with two consecutive electrochemical steps
O
2
+ H
2
O + 2e
−

→ HO
2
−
+ OH
−
[36]
HO
2
−
+ H
2
O + 2e
−
→ 3OH
−
[37]
(3) Electron transfer followed by decomposition of hydrogen peroxide
O
2
+ H
2
O + 2e
−
→ HO
2
−
+ OH
−
[38]
HO
2
−
→ OH
−
+ ½ O
2

[39]
The formation of hydroperoxide (Eqs. 36 and 38) can be monitored using the rotating-ring disk
electrode arrangement wherein the hydroperoxide produced on the rotating disk electrode is
convectively transported to the ring electrode and detected by an oxidation current.
125-126
The
ratio of the disk and ring currents can be used to diagnose the reaction pathway.
127
Therefore, in
this work, the disk electrode was coated with the composite catalyst of interest, and the
hydroperoxide formed at the disk electrode was oxidized at the platinum ring electrode. The disk
and ring currents were used to determine the percentage of the current involved in the production
of hydroperoxide and the number of electrons transferred per mole of oxygen (Figure 25).
49

Figure 25. Polarization experiments in oxygen-saturated 1M potassium hydroxide on a rotating
ring-disk electrode (a) LCCO, (b) Acetylene black, and (c) 83.3% LCCO + 16.7% AB. The ring
electrode was a platinum electrode held at +0.5 V vs. MMO. Each block is arranged vertically as
disk current, ring current and number of electrons transferred.
On LCCO, the formation of hydroperoxide accounted for approximately 5% of the
oxygen reduction current at the disk. The number of electrons transferred was calculated to be
3.9. Consequently, it was concluded that the oxygen reduction reaction on LCCO occurred by
either the direct four-electron pathway (Eq. 35) or by the ―series‖ pathway with a slow two-
electron transfer step followed by a rapid two-electron reduction of the hydroperoxide (Eqs. 36
50

and 37), or the slow electrochemical formation of hydroperoxide followed by rapid
decomposition to oxygen (Eqs. 38 and 39). For all these situations hydroperoxide should be
detected only in significant amounts at the ring electrode and the number of electrons transferred
will approximate to four. Further identification of the exact pathway could be carried out only
with the following additional experiments.
To determine which of the two ―four-electron‖ pathways were operative in the reduction
of oxygen on LCCO, we measured the currents for the direct reduction of 0.1M hydroperoxide
on the LCCO-coated disk electrode to estimate the contribution of the electrochemical reduction
in Eq. 36, and also measured the decomposition rate of hydroperoxide in contact with LCCO.
When LCCO powder was added to a 0.1M hydrogen peroxide solution a rapid decrease in the
hydrogen peroxide oxidation current at the glassy carbon disk electrode (Figure 26a) suggested a
decrease in the concentration of hydrogen peroxide and confirmed that LCCO rapidly
decomposed the hydroperoxide anion to produce oxygen. Thus, we were able to confirm that the
process of decomposition according to Eq. 39 readily occurred on LCCO. Further, when a disk
coated with 80 g of LCCO was polarized in 0.1 M solution of hydroperoxide (Figure 26b) the
reduction current observed was similar to the current observed with oxygen-saturated solutions
(Figure 25a). This result is consistent with the occurrence of two-electron electro-reduction of
oxygen to hydroperoxide (Eq. 38) followed by rapid decomposition of the hydroperoxide to
oxygen (Eq. 39). However, if there was a small current contribution on LCCO by the direct
four-electron reduction of oxygen (Eq. 35), it could not be ruled out.
51

Figure 26. (a) change of limiting oxidation current at 0.2 V for hydroperoxide with time in the
presence of 10 mg of LCCO in 0.1M hydroperoxide in 1M potassium hydroxide, (b) results of
cathodic polarization (potential scanned at 5 mV s
-1
) of LCCO-coated disk electrode in argon
saturated 0.1 M hydroperoxide in 1M potassium hydroxide.
On acetylene black, the rate of generation of hydroperoxide was at least an order of
magnitude higher than that on LCCO (Figure 25b). The hydroperoxide measurements on the ring
electrode showed that almost 95% of the reduction current was directed through the two-electron
pathway (Eq. 36). The number of electrons transferred in the reaction of 2.1 to 2.5 was consistent
with Eq. 36. This number of electrons being slightly higher than 2 suggested a minor amount of
decomposition of hydroperoxide on acetylene black or a small contribution from subsequent
reduction steps through electron transfer.
On the composite catalyst consisting of acetylene black and LCCO, unlike the acetylene
black electrode, the hydroperoxide generation was just 5% of the total reduction current at the
disk electrode. Thus, the number of electrons transferred was calculated to be 3.9 suggesting a
net four-electron process. This value of approximately 4 for the number of electrons was
consistent with the mass transport-limited disk currents being twice as much as that on acetylene
black electrode. Since we had independently confirmed the near quantitative generation of
52

hydroperoxide on acetylene black (Figure 25b) and the rapid decomposition of hydrogen
peroxide on LCCO (Figure 26a and 26b), it was reasonable to expect that in the composite
catalyst, the hydroperoxide formed at the carbon electrode would be rapidly decomposed by the
LCCO. In the steady state, the electro-reduction to hydroperoxide on carbon and the subsequent
decomposition to oxygen on LCCO would lead to a net observation of four electrons transferred
per oxygen molecule, readily seen by adding Eq. 38 and 39.

Figure 27. (a) Percentage of current used in hydroperoxide generation on various catalyst
compositions of Table 1 coated on the disk electrode, (b) The calculated number of electrons
transferred in the reaction determined from the amount of hydroperoxide detected at the ring
electrode. The percentage of hydroperoxide and the number of electrons transferred are
determined at -0.15 V.
The generation of hydroperoxide was studied for varying amounts of LCCO and
acetylene black (Figure 27). When the amount of LCCO was large as compared to the amount of
acetylene black (composition 2,3,4,5 and 6 in Table 1) almost all the hydroperoxide produced on
the surface of acetylene black is decomposed, and a very small amount of free hydroperoxide
53

was detected at the ring electrode. However, when the carbon-LCCO composite was made of a
very large fraction of carbon as in compositions 7, 8 and 9, the hydroperoxide was not
completely decomposed allowing significant amounts of hydroperoxide to be detected at the ring
electrode. Thus, the calculated value for the number of electrons transferred varied between 3
and 4 when the LCCO concentration was low.
Based on the results presented in Figure 25 and Figure 26b it was certain that LCCO
produced almost no hydroperoxide during oxygen reduction and the acetylene black catalyst
produced hydroperoxide almost quantitatively. We have estimated the expected production of the
amount of hydroperoxide on the composite catalyst using a weighted sum of the independent
contributions of LCCO and acetylene black to the production of hydroperoxide. We have
compared (Figure 28) the calculated results of hydroperoxide production with the experimental
results (Figure 27a)
If the LCCO and carbon acted independently the amount of hydroperoxide detected on
the LCCO-carbon composite would have to be considerably larger than that experimentally
observed (Figure 28). Therefore, the decreased amount of hydroperoxide detected at the ring
when both acetylene black and LCCO are present is simply the result of the decomposition by
LCCO of the hydroperoxide produced on acetylene black. Thus, LCCO acted synergistically
with acetylene black in catalyzing the electro-reduction of oxygen and the rapid decomposition
of the hydroperoxide produced on the acetylene black when sufficient LCCO was present, and
thus yielded the net four electrons transferred per oxygen molecule.
54

Figure 28. Comparison of experimental and calculated values of hydroperoxide production on
LCCO-acetylene black composites of composition in Table 1.
LCCO, by itself, was found to be a relatively poor catalyst for the oxygen reduction
reaction compared to the composite of LCCO and acetylene black. Acetylene black was not a
spectator that simply enhanced electrical conductivity, but in fact had a primary role in
determining the catalytic activity of the composite. Thus, LCCO may be termed a co-catalyst as
it enhanced the observed currents by decomposing the hydroperoxide species produced on
acetylene black. This role of LCCO is similar to that of silver and manganese dioxide
(commonly termed peroxide decomposers) used widely with carbon-based electrodes in alkaline
fuel cells and primary metal –air batteries. The findings presented above require that the analysis
of the electrocatalytic activity of composites of LCCO and carbon do not ignore the primary
function of LCCO as a peroxide decomposer.
55

We have determined that many other transition metal oxide perovskites and spinels
behave in a similar manner to LCCO. In the other examples of transition metal oxides the
enhanced activity of composite catalysts correlated well with the reduced amount of
hydroperoxide detected at the ring electrode, as in the case of LCCO. Consequently, the
differences in the observed electrocatalytic activity of various transition metal oxides when
combined with carbon, although apparently a four-electron transfer process, depends
significantly on the ability of the transition metal oxide to decompose the hydroperoxide
generated by the electro-reduction of oxygen on the carbon constituent of the composite catalyst
(Eqs. 38 and 39). It is quite likely that the ability of the transition metal oxide to decompose
hydroperoxide (Eq. 39) is related to the ability of the oxide surface to support a direct four-
electron transfer (Eq. 35). However, we must make a distinction between these two net ―four
electron‖ pathways while interpreting the observed activity of transition metal oxide-carbon
composites.
When the acetylene black in the composite catalyst was substituted by carbon nanotubes
or graphene, a similar enhancement in catalytic activity was observed (Figure 22). However,
substitution of acetylene black with gold nanoparticles did not produce such an enhancement.
While carbon-based materials are known for their ability to perform efficient two electron
reduction to hydroperoxide, gold is generally a poor catalyst for oxygen reduction.
128-129

Consequently, gold does not substitute the primary catalyst role of carbon in the production of
hydroperoxide necessary for producing the enhancement in activity in conjunction with LCCO.
To further verify that the LCCO-carbon composite catalyst operates through the pathway
involving steps of electrochemical generation of the hydroperoxide and the decomposition of
hydroperoxide to oxygen, we have analyzed the relationship between the disk and ring currents
56

using the formalism similar to that developed by Damjanovic and Bockris.
130-131
We have
performed the analysis of the oxygen reduction process that involves the following three main
steps (also depicted in Figure 29):

Figure 29. Reaction scheme used in the analysis of the behavior of carbon-transition metal oxide
composite catalyst.
Step 1: Oxygen is electrochemically reduced to hydroperoxide by a two-electron process.
O
2
+H
2
O +2e
−
 HO
2
−
(disk)

+ OH
−

Step 2: The hydroperoxide generated on the carbon surface is decomposed on the surface of
LCCO to produce half a mole of oxygen that will contribute to the electrochemical step of
reduction.
HO
2
−
(disk)
OH
−
+ ½ O
2

Not all the hydroperoxide may be decomposed as this will depend on the kinetics of the
decomposition reaction and the amount of LCCO present. We associate this reaction with a
heterogeneous rate constant k
p
.
O
2
( bulk) O
2
(disk)
+H
2
O +2e-
HO
2
-
(disk) + OH
-
HO
2
-
(ring)
HO
2
-
(boundary layer)
Decomposition
By Transition Metal Oxide
HO
2
-
(bulk)
I
1
I
2
I
3
I
ring
k
p
57

Step 3: The hydroperoxide that is not decomposed to oxygen diffuses to the bulk and is detected
on the ring electrode
HO
2
−
(disk)
 HO
2
−

(boundary layer)
 Partly ―collected‖ at the ring electrode.
The current observed at the disk electrode, I
disk
, results from the current for the reduction
of oxygen to hydroperoxide from Step1, I
1
, and the current generated from the oxygen produced
by decomposition of hydroperoxide from the Step 2, I
2
.

The rate of generation of hydroperoxide in the volume of the boundary layer, I
3
is the difference
between the rate at which hydroperoxide is produced at the disk and the rate at which it is
decomposed by the transition metal oxide. Consequently, the current in Step 3 is given by,

In the steady state, the rate of diffusion of hydroperoxide across the boundary layer of the disk
electrode is given by,

Where A
disk
is the area of the electrode where hydroperoxide is produced, D
HO2
−
is the diffusion
coefficient for hydroperoxide, and C
HO2
−
is the concentration of hydroperoxide at the surface of
the disk,  is the boundary layer thickness at the rotating electrode, n is the number of electrons
in the oxidation of hydroperoxide, and F is the Faraday constant.
The value of I
2
is determined by the rate of decomposition of hydroperoxide to oxygen
and can be expressed in terms of the rate constant for decomposition by,
58

(43)
Where A
i
is the area of the perovskite participating in the decomposition process and k
p
is the
heterogeneous rate constant for decomposition with the unit cm s
-1
. Thus A
i
will be determined
by the mass of catalyst and its dispersion on the carbon.
The current observed at the ring, I
ring
, is obtained from the mass transport of un-
decomposed hydroperoxide across the boundary layer into the bulk and also the collection
efficiency N at the ring electrode.
(44)
From equations 42 and 43 we obtain:

Therefore, by combining equations 40, 41, 44 and 45 we obtain,

For a rotating disk electrode,
(47)
Where ν

is the kinematic viscosity of the electrolyte, and 

is the rotation rate.
Therefore, from equations 46 and 47 we obtain
59

Therefore, a plot of I
disk
/I
ring
vs. 
-1/2
is expected to be linear. The slope of this line will
depend on the heterogeneous rate constant for the transition metal oxide and the effective
interfacial area involved in decomposition as determined by the amount of the decomposer
(transition metal oxide). Thus, the slope will also depend on the dispersion of the transition metal
oxide to the carbon catalyst, as determined by A
i.
However, the heterogeneous rate constant k
p

will be specific to the type of oxide material and its surface area.
We find that plot of I
disk
/I
ring
vs. 
-1/2
could be fitted with a line for various values of
electrode potential (Figure 30a), and the slope of this line at any particular potential increases in
proportion to the amount of LCCO for a fixed amount of carbon (Figure 30b) and is consistent
with Eq. 48.

Figure 30. Plot of I
disk
/I
ring
as a function of (1/rotation frequency)
1/2
(a) at various potentials for
composite catalyst (Carbon 8 mg–LCCO 8 mg); (b) for various composite catalyst compositions
with 8 mg of carbon and increasing amounts of LCCO (1 mg, 2 mg, 4 mg and 8 mg) at electrode
potentials -0.175 V.

(b)
(a)
60

Table 2. Measured values of rate constant for decomposition of hydrogen peroxide on LCCO at
25
o
C in 1M potassium hydroxide determined from the slope of I
disk
/I
ring
vs. 
-1/
.
2

Composition # Slope ( rpm
½
) Rate Constant, cm s
-1

5 1336 0.0009
6 1071 0.0015
7 915 0.0025
8 237 0.0013

Table 2 lists the value of the rate constant for decomposition of peroxide by LCCO as
calculated from the slope of the lines in Figure 30b and Eq. 48. The calculations used the
following in values for the constants: the diffusion coefficient of hydroperoxide D
HO2-
of 1.65 x
10
-5
cm
2
s
-1
, kinematic viscosity of the electrolyte, ν

= 0.0095 cm
2
s
-1
, the collection efficiency, N
= 0.2, A
disk
=0.1925

cm
2
. The interfacial contact area, A
i
, for each sample was calculated from the
mass of LCCO and previously measured values of electrochemically active surface area of 10
m
2
/g reported in our group earlier publication.
60
The rate constant values were in the range of
0.0009 to 0.0025 cm s
-1
. These values of rate constants are in close agreement with those for
platinum surfaces in alkaline media.
132

The Tafel slopes were determined from the polarization results for various catalyst
compositions (Figure 31)
When just LCCO was present in the catalyst layer, the Tafel slope was 94 mV/decade.
With just acetylene black in the catalyst layer, the Tafel slope was 60 mV/decade. These values
of Tafel slope are consistent with those reported in the literature.
115
For the composite catalysts
with a higher fraction of LCCO than acetylene black, the Tafel slope was in the range of 80-95
mV/decade consistent with that observed with LCCO. However, with compositions rich in
acetylene black, the values were closer to 60 mV/decade. The apparent exchange current density
61

calculated from the intercept of the Tafel line was one order of magnitude lower for pure LCCO
compared to acetylene black, suggesting that the direct reduction of oxygen to peroxide was
significantly slower on the LCCO compared to carbon.

Figure 31. Kinetic current plotted vs. the potential of the disk electrode coated with LCCO-
carbon mixtures as indicated, (a) for compositions rich in LCCO (b) compositions rich in
acetylene black.

62

Chapter 4
Oxygen Reduction and Hydrogen Peroxide
Decomposition Activity of Transition Metal Oxide
Catalysts
In the previous section we have been demonstrated that the composite catalysts prepared
by physical mixing of electrically conductive transition metal oxide with acetylene black exhibit
two to ten times higher electrocatalytic activity compared to catalysts consisting of just transition
metal oxide or carbon. We establish that there is a synergistic effect between carbon and
transition metal oxide. In such composite catalysts, carbon is the primary electro-catalyst for the
two electron electro-reduction of oxygen to hydroperoxide while the transition metal oxide
decomposes the hydroperoxide to generate additional oxygen that enhances the observed current
resulting in an apparent four–electron process. We expect this finding to help the design of
improved transition-metal-based catalysts for oxygen reduction. Our hypothesis is that when the
same kind of carbon is mixed with various transition metal oxides, the oxygen reduction reaction
activity will vary depending on how the hydroperoxide decomposition process occurs on the
transition metal oxide.
Many transition metal oxides have been studied by Bockris and others
121, 133-134
for
catalyzing oxygen reduction. In the current study we have focused on manganese and cobalt
based oxides. We have investigated on many oxides of perovskite and spinel structure (Figure
32). Among these oxides catalysts, the lanthanum calcium cobalt oxide (LCCO) perovskite has
shown promise as an electrocatalyst for oxygen reduction. However, cobalt is more expensive
than manganese and therefore if we replace cobalt by manganese we can reduce the cost of the
63

catalyst. Further, by varying the amount of cobalt and manganese we expect to also tune the
electrocatalyst activity. The list of perovskite oxide samples for this study prepared by the
method in Chapter 2 is shown in Table 3. The results of SEM, XRD, XPS, XANES will be
discussed later in Chapter 5.

Figure 32. ORR activity for various perovskite and spinel oxides system
In addition to LCCO, lanthanum calcium cobalt manganese oxide (LCCM) containing
various amounts of manganese and cobalt showed similar enhancement in oxygen reduction
activity when combined with acetylene black (Figure 33). When acetylene black alone was used
at the catalyst, the oxygen reduced to hydroperoxide and this hydroperoxide was detected at the
ring electrode (Figure 33b). No decomposition of the hydroperoxide occurred on acetylene black.
64

However, when LCCM was combined with acetylene black the mass-transport limited current at
the disk was twice that of the current with acetylene black. The hydroperoxide current at the ring
was also about 19% of that observed with just acetylene black. This observation confirmed that
most of the hydroperoxide generated at the carbon surface was decomposed by LCCM. These
results are in good agreement with the findings reported in Chapter 3.
Table 3. Various manganese-substituted calcium-doped lanthanum cobalt oxides (LCCM)

The reversible potential for the electro-reduction of oxygen to hydroperoxide in alkaline media
(Eqs. 36 or 38 Chapter 3) can be calculated using the Nernst equation,

Where E is reversible electrode potential, E
o
is standard reduction potential of eq. 38, P
O2
is the
partial pressure of oxygen, and a
HO2
-
, a
OH
-
, a
H2O
are the activities of HO
2
¯
, OH
¯
, H
2
O,
respectively. The activity of hydroperoxide refer to that present at the surface of the electrode.
X Formula La
0.6
Ca
0.4
Mn
x
Co
1-x
O
3

0 La
0.6
Ca
0.4
CoO
3

0.1 La
0.6
Ca
0.4
Mn
0.1
Co
0.9
O
3

0.3 La
0.6
Ca
0.4
Mn0.3Co
0.7
O
3

0.5 La
0.6
Ca
0.4
Mn
0.5
Co
0.5
O
3

0.7 La
0.6
Ca
0.4
Mn
0.7
Co
0.3
O
3

0.9 La
0.6
Ca
0.4
Mn
0.9
Co
0.1
O
3

1 La
0.6
Ca
0.4
MnO
3

65

When LCCM was combined with acetylene black we noted that the onset potential
(which is the potential where the oxygen reduction begins to occur during a potential sweep in
the negative direction) was + 30 mV vs MMO while with just acetylene black the onset potential
was at approximately –70 mV. This 100 mV increase in onset potential again confirmed that in
the case of transition metal oxide composite catalyst, the transition metal oxide decomposed the
hydroperoxide generated on the carbon surface. Depending on the effectiveness of the transition
metal oxide in decomposing the hydroperoxide, the concentration of hydroperoxide at the disk
electrode will be altered and the electrode potential will shift. Thus, we expected to observe
variation in the oxygen reduction activity for various compositions of the transition metal oxide.

Figure 33. (a) Polarization curves for acetylene black and LCCM (x=1) + AB in oxygen-
saturated 1M potassium hydroxide at rotation rate 2500 rpm. (b) Hydroperoxide current detected
at the ring for acetylene black and LCCM (x=1) +AB.
The results of studies of oxygen reduction on various transition metal oxides are
presented in Figure 34a. The disk current was not significantly different when the manganese
fraction (x value) was varied from 0 to 0.5. However, there was an increase in disk current when
66

the x value increased to 0.7, 0.9, and 1. The onset potential also shifted to more positive values
for the higher manganese fractions. In addition, the hydroperoxide current at the ring electrode
(Figure 34b) decreased as the manganese fraction increased.

Figure 34. (a) Polarization curves for oxygen reduction in 1M potassium hydroxide at 2500 rpm
for various compositions of the transition metal oxide (b) Hydroperoxide current at the ring
electrode for various compositions (c) ORR activity vs. percentage of hydroperoxide observed at
the ring electrode (d) ORR activity vs. percentage of hydroperoxide for various value of
manganese fraction.
The oxygen reduction activity at -100 mV increased and the percentage of hydroperoxide
decreased (Figure 34c, d) as the manganese fraction increased. These results confirmed that for
67

the transition metal oxide-carbon-composite catalyst, the oxygen reduction activity with any
given type of carbon will depend on the decomposition rate of hydroperoxide by the transition
metal oxide. This observation suggested that with an increased activity for hydroperoxide
decomposition, higher oxygen reduction activity can be realized. Therefore, a strong dependence
of oxygen reduction activity on the composition of oxides is to be expected.
We investigated the difference in the decomposition characteristics for hydrogen
peroxide on various transition metal oxides using polarization studies and direct chemical
analysis (hydrogen peroxide reduction and oxidation by polarization studies, and measurement of
the amount of oxygen generated during the direct decomposition of hydrogen peroxide by
transition metal oxides). All the electrodes were prepared according to method described in
Chapter 2.
According to the literature on the subject of decomposition of hydrogen peroxide on
transition metal oxides, the reaction is mediated by the redox states in the metal oxides.
135-137

Catalytic decomposition of hydrogen peroxide has been studied by many researchers. Weiss
studied the decomposition of hydrogen peroxide on metals such as platinum, gold, palladium,
silver and zinc and proposed an explanation of hydrogen peroxide decomposition on the metal
surface.
138
According to Weiss, there are two principal reactions that involved in the hydrogen
peroxide decomposition process where the hydrogen peroxide acts as an oxidizing by accepting
electron from the metal as well as reducing agent by donating electron to the metal.
The work by Roy
135
on the decomposition of hydrogen peroxide by some oxide catalysts
suggested that the potential of the oxide system with two oxidation states of the element plays an
important role in the catalytic activity of the oxide in the hydrogen peroxide decomposition
process.
68

Also, many authors have investigated the decomposition of hydrogen peroxide over
perovskite catalysts such as LaFe
x
Ni
1-x
O
3
, La
0.9
Sr
0.1
Ni
1-x
Cr
x
O
3
, La
1-x
Ca
x
MnO
3
, LaCoO
3
,
LaMnO
3
and La
0.8
Sr
0.2
MnO
3
.
69
,
139
,
140
,
141
However, it is still unclear if the catalytic
decomposition of hydrogen peroxide occurs by chemical or by an electrochemical pathway or if
both pathways can occur simultaneously. In general, these mediated processes involve the
simultaneous oxidation of hydrogen peroxide to oxygen and the reduction of hydrogen peroxide
to hydroxide or water. The metal-ion redox species mediate the shuttling of electrons between
the oxidation and reduction processes. The valence state of the metal ion thus remains unchanged
while hydrogen peroxide is decomposed. These reactions involved in the mediation process in
alkaline media is given by:
HO
2
-
+ H
2
O + 2M
(n-1)+
→ 2M
n+
+ 3OH
-
(50)
OH
-
+ HO
2
-
+ 2M
n+
→ O
2
+ H
2
O + 2M
(n-1)+
(51)
In the reduction of hydroperoxide the O-O bond is broken. Therefore, we can expect a
high activation barrier for this reaction. Such bond breakage most likely involves the formation
of M
(n-1)+
-OH bonds. To perform this reaction, the standard reduction potential of the metal-ion
couples should be negative to + 0.88 V.
The second conjugate reaction in the decomposition of hydroperoxide involves the
oxidation of peroxide to oxygen and to restore the oxidation state of the metal ion.
Note that the O-O bond is preserved in the oxidation of peroxide to oxygen. To perform
this oxidation process, the standard reduction potential of the metal ion couple should be more
positive to – 0.14 V in base. One would expect that the formation of the M-O-O···H intermediate
as important for reaching the transition state.
69

We aimed to address the underlying reasons for the differences in ORR activity to the
ability of the oxide to decompose peroxide. We noted that with an increase in manganese
fraction, the ORR activity increased and the amount of hydroperoxide generated decreased
(Figure 34d). We also observed that the oxide exhibited a characteristic open circuit potential in
the solution of 50 mM hydrogen peroxide dissolved in 1 M potassium hydroxide. This open-
circuit potential varied with the composition of the oxide. Considering the mediator role played
by the transition metal ion (M
n+
), we postulated that an electrochemical mechanism could be
operating, mediated by the metal ions. The conjugate reactions of oxidation and reduction of
peroxide that would govern the process are:
HO
2
-
+ H
2
O + 2e-  3OH
-
(52)
OH
-
+ HO
2
-
→ O
2
+ H
2
O+ 2e-

(53)

M
(n+)
+e-  M
(n-1)+
(54)

E
o
= +0.21 V for Co
3+
/Co
4+
, E
o
= +0.11 V for Mn
3+
/Mn
4+
.
M
(n-1)
 M
(n)+
+e- (55)
When the decomposition of peroxide occurs via an electrochemical mechanism the four
reactions listed above must occur simultaneously. The observed open-circuit potential of the
oxide in hydroperoxide solutions will then be determined by the kinetics of these four reactions.
The kinetics of the metal ion oxidation and reduction could be expected to be kinetically
reversible as it involves only exchange of electrons unlike the oxidation and reduction of
peroxide that involves adsorption and bond-breaking events. Therefore, the kinetics of the
peroxide reactions will determine the steady state ratio of M
n+
and M
(n-1)+
and hence the open
circuit potential. If the oxidation of peroxide is facile, then the mixed potential will move in the
negative direction and the concentration of metal ions will shift more towards the reduced state.
Thus, the observed electrode potential under open circuit conditions is determined by the kinetics
70

of the peroxide oxidation and reduction and also the potential of the metal ion redox couple.
Further when the electrode is polarized to carry out either oxidation or reduction of peroxide, the
ratio of the number of metal ions in the oxidized and reduced state will change according to the
applied potential as the kinetics of the metal ion oxidation and reduction is expected to be more
reversible than peroxide oxidation and reduction.

Figure 35. (a) Polarization curves of hydrogen peroxide reduction and oxidation kinetic current
for various oxides in 1M potassium hydroxide solution containing 50 mM hydrogen peroxide (b)
Open-circuit potential (mixed potential) measured and calculated for various compositions as
indicated by ―x‖ values.
The observed electrode potential under open circuit conditions is the range of +0.11 to
+0.21 V vs. NHE. This potential range corresponds to the 3+/4+ reactions for manganese and
cobalt at pH =14 (Pourbaix‘s diagram for manganese and cobalt in Pourbaix‘s Atlas of
Electrochemical Equilibria in Aqueous Solutions, pages 290 and 325). Such oxidation states are
consistent with the presence of 3+ oxidation state for the transition metal in the oxides of the
formula La
0.6
Ca
0.4
Co
x
Mn
1-x
O
3
. The variation of the observed open-circuit potential with
71

composition of the oxide is thus influenced by the nature of the metal ions and also by the
kinetics of peroxide oxidation and reduction.
To verify that the mixed potential observed under open circuit conditions was dependent
on the kinetics of peroxide oxidation and reduction for any particular oxide composition we
measured the steady-state current during polarization positive and negative to the mixed
potential. The resulting current–potential data for oxidation and reduction of peroxide was fitted
to Tafel lines. The point of intersection of the Tafel lines corresponds to the potential at which
the oxidation and reduction current have the same values. If the electrochemical mechanism was
operating, this value of potential at the intersection of the Tafel polarization lines must be the
same as the open-circuit mixed potential. Comparison of the potential observed from the
intersection of the Tafel lines (Figure 35a) with that of the observed open circuit potential shows
that these two values were very close for the all the oxide compositions (Figure 35b), lending
direct support to the electrochemical mechanism for the decomposition of peroxide on these
oxides.
If the decomposition of peroxide proceeded via an electrochemical mechanism, then the
two conjugate reactions which are hydrogen peroxide reduction and hydrogen peroxide oxidation
(Eqs. 52 and 53) would both occur on the surface of the metal oxide producing both oxygen and
hydroxide ions. The rate of this decomposition reaction on the oxide can be determined from the
currents for the conjugate processes of oxidation and reduction of peroxide occurring at the
mixed potential. This decomposition rate would correspond to the current determined from the
point of intersection of the two Tafel lines. We have determined the value of current for the
oxides with various cobalt and manganese fractions and called it I
corr
, as it bears similarity to the
72

current associated with two conjugate reactions occurring under open circuit conditions in a
galvanic corrosion process.
142

Figure 36. (a) Exchange current density of peroxide oxidation vs peroxide reduction for various
manganese fractions (b) Exchange current density of peroxide oxidation vs ORR activity at -100
mV.
We have described here in Chapter 3 that the decomposition of hydroperoxide on the
transition metal oxide surface enhances the of oxygen reduction activity of carbon-transition
metal oxide composites. We had shown that oxygen is reduced to hydroperoxide on the carbon
surface and the decomposition of hydrogen peroxide on the transition metal oxide leads to
further production of oxygen. This extra oxygen generated by the decomposition of hydrogen
peroxide enhances the oxygen reduction current on the surface of carbon. Thus, the efficacy of
the transition metal oxide to decompose or oxidize hydroperoxide to oxygen at any chosen
potential would predict the oxygen reduction activity of the carbon-transition metal oxide
composite. We can determine the rate of oxygen generation from hydroperoxide at various
electrode potentials on the transition metal oxides from the steady-state polarization data (Figure
35a). We have determined the exchange current density I
o
for peroxide oxidation for various
73

values of manganese fractions in the oxide (Figure 36a). We find that the oxygen generation rate
went through a minimum value and then rapidly increased as the manganese fraction was
increased. The minimum occurred at composition of x=0.3. The open-circuit potential for this
oxide was also the most positive value at this composition. These results are consistent with the
understanding that lower the tendency for the oxidation of hydroperoxide to oxygen, the more
positive would be the open-circuit mixed potential.

Figure 37. (a) Peroxide oxidation rate vs ORR activity at -100 mV vs MMO (b) I
corr
vs. ORR
activity at -100 mV vs MMO.
We have compared the exchange current density of hydroperoxide oxidation on various
oxides with the oxygen reduction activity observed at the potential of -100 mV on the various
carbon-transition metal oxide composites (Figure 36b). We find that the higher rate of oxidation
of hydroperoxide correlated directly with the improved oxygen reduction activity in the
composite catalyst (Figure 36b). Similarly, an increase in the rate of hydroperoxide oxidation at -
100 mV or the current at the mixed potential (I
corr
) also resulted in increase of oxygen reduction
activity (Figures 37a, b). We also found that the exchange current density for the reduction of
hydroperoxide did not correlate with the oxygen reduction activity (Figure 38). This finding is
74

understandable in that the peroxide reduction does not produce oxygen needed for improving the
oxygen reduction activity of carbon. Also, the exchange current density values for reduction
were about four orders of magnitude smaller than that for the oxidation, suggesting that the
reduction of hydroperoxide is quite hindered under the conditions of the oxygen reduction
experiment. Since reduction of hydroperoxide involves the cleaving of the O-O bond we can
expect the activation barrier for the reduction process to be high. Therefore, the oxidation of
hydroperoxide to oxygen is preferred over the reduction to hydroxide. Consequently, a
significant overpotential is needed to achieve the decomposition of peroxide. These findings
verify that role of the surface of the metal oxide in the decomposition of hydroperoxide and
enhancing the oxygen reduction activity of the carbon-transition metal oxide composite catalysts.
The results presented here also support an electrochemical mechanism for the decomposition of
peroxide on the metal oxide.

Figure 38. Exchange current density of hydrogen peroxide reduction and ORR activity at -100
mV vs MMO for various manganese fractions.
75

We determined the Tafel slopes for the oxidation and reduction of hydroperoxide on the
metal oxide surfaces of different compositions. We found no specific trend in the Tafel slopes
with composition. The values of Tafel slopes varied from 120 mV/decade to 150 mV/decade for
both the oxidation and reduction reactions (Figure 39).

Figure 39. Tafel slope of hydrogen peroxide oxidation and reduction for various manganese
fractions.
While we have found evidence for the decomposition of peroxide by the electrochemical
mechanism, we do not know if other mechanisms such as the free-radical mechanism operate on
the same surface. One method of proving this would be measure the decomposition rate of
hydroperoxide directly by determining the rate of oxygen evolution rate and then comparing it
with the rate estimated from the I
corr
value. We have described the direct method of measuring
76

decomposition rate in Chapter 2. We could not find the same trend between amount of oxygen
evolved and the decomposition of peroxide as determined by the I
corr
values (Figure 40). In fact
an opposite trend was observed. We also recognized that the two experiments may not be
comparable because the utilization of transition metal oxide could be different, as the
electrochemical tests were carried out on catalyst coated on Toray® paper while in the direct
measurement the peroxide decomposition rate the catalyst was in powder form and floating
around in the solution. Therefore, we need to re-design this experiment to account for these
differences before we can determine if other mechanisms besides the electrochemical mechanism
operate during the decomposition of peroxide on a transition metal oxide surface.

Figure 40. The decomposition rate I
corr
from peroxide polarization experiment vs decomposition
rate from peroxide direct decomposition study.

77

Chapter 5
Oxygen Evolution Activity of Perovskite and Spinel
Transition Metal Oxide Catalyst
We have studied oxygen evolution activity on many perovskite and spinel transition
metal oxides (Figure 41). Among these oxides, lanthanum calcium cobalt manganese perovskite
oxide and nickel cobalt spinel oxide exhibited high OER activity. Therefore, this chapter will
focus on the OER activity for these oxides.

Figure 41. Oxygen evolution activity of various perovskite and spinel transition metal oxides.
78

5.1 Oxygen Evolution Activity of Perovskite Transition Metal Oxide
on Rotating Ring Disk Electrode
Transition metal oxides of the perovskite and spinel family present an opportunity for
designing new low-cost catalysts because of the variety of compositions that are conceivable.
61,
79-80, 121, 133-134, 143-154
The perovskite oxides have the general formula ABO
3
, where A is a rare
earth metal ion and B is a transition metal ion. Typically, the transition metal ion at the B site is
the catalytically active center for oxygen evolution.
147
Since most of the transition metals form
perovskite oxides, the compositions with different metals in the B-site are numerous. The variety
that is possible in the composition of oxides, presents a significant opportunity for tuning and
enhancing the electrocatalytic activity.
79, 134, 147, 155
In the present study, we provide new insights
into the factors controlling the electrocatalytic activity of such oxides.
A popular approach to alter the electrocatalytic properties of a perovskite oxide, is to
partially substitute the A and B atoms with different elements A' and B', to achieve compositions
of the general formula, A
y
A'
1-y
B
x
B'
1-x
O
3
. This type of metal substitution has been adopted by
solid-state physicists to tune the electrical and magnetic properties of such oxides.
156-168
When A
and A‘ have different valences as in lanthanum(III) and calcium(II), an increase in the
oxidation state of the B atom or an increase in the number of oxygen vacancies occurs to
maintain overall charge neutrality in the lattice. As a result of these changes, the electrical,
magnetic and electrocatalytic properties are usually modified. For example, lanthanum cobalt
oxide (LaCoO
3
), lanthanum manganese oxide (LaMnO
3
) and lanthanum nickelate (LaNiO
3
) have
been widely investigated, and the catalytic activity of these oxide catalysts is dependent on the
type and oxidation state of the transition metal.
59, 62-64
For example, partial substitution of
lanthanum by calcium or strontium as in the compositions, La
0.6
Ca
0.4
CoO
3
or La
0.8
Sr
0.2
CoO
3
,
79

results in a mixture of Co
2+
, Co
3+
, Co
4+
and generation of oxygen vacancies, accompanied by an
increase in electrical conductivity.
60-61
Also, when the transition metal in the B-site of the
perovskite is substituted by one or more of the first-row transition metals, a variety of d-electron
configurations are presented at the surface. The possibility of tuning the composition at the A
and B site for electrocatalytic activity has thus evoked considerable interest.
59, 62-64

By studying the oxygen evolution activity of substituted transition metal oxides we can
gain further insights into the role of the B-site in modifying electrocatalytic activity. For this
purpose, we have focused on calcium-doped lanthanum cobalt oxide perovskite (of the formula
La
0.6
Ca
0.4
CoO
3
)

with various amounts of manganese substituting for cobalt (as in the formula,
La
0.6
Ca
0.4
Mn
x
Co
1-x
O
3
). The inspiration for the choice of this perovskite arose from the attributes
of manganese, namely: a) a low-cost and globally-abundant material b) environment-friendly,
and c) ability to achieve a range of oxidation states of +2 to +7. Therefore, we have investigated
the effect of the systematic replacement of cobalt with equivalent amounts of manganese on the
electrocatalytic activity of the perovskite oxide towards oxygen evolution. Such a study was
aimed at providing new insights into how the transition metal site may be modified to alter
activity and also reduce the cost of the electrocatalyst. The transition metal oxide compositions
were outline in Table 3 in Chapter 4.
X-ray diffraction analysis confirmed that all the composition in Table 3 were in a pure
perovskite phase consistent with the powder diffraction file data (PDF#00-041-0496) for
La
0.5
Ca
0.5
Mn
0.5
Co
0.5
O
3
(Figure 42). The XRD peaks shifted to lower 2θ values with increasing
manganese content (Figure. 42b), consistent with an increase of the cell dimensions that is to be
expected from the higher ionic radii of manganese ions.
162
The peak width at half-maximum was
independent of the cobalt and manganese fraction, and crystallite size varied between 11 and 22
80

nm. Thus, the relatively low temperature of processing allowed us to successfully produce nano-
crystalline oxide materials. The nano-particulates were thin flakes about 100-200 nm thick
(Figure 43). This morphology was similar for all the ratios of manganese to cobalt outline in
Table 3.

Figure 42. (a) X-ray diffraction pattern for La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
for the various compositions
indicated (b) Magnified region 2  = 30
o
-35
o
, (c) Powder diffraction file data (PDF#00-041-
0496) for La
0.5
Ca
0.5
Mn
0.5
Co
0.5
O
3
. (d) Crystallite size for various values of manganese fraction
calculated from 2  = 33
o
(220).

81

Figure 43. Scanning electron micrographs of La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
; (a) x = 0 (b) x = 0.1; (c) x =
0.5 (d) x = 0.7 (e) x = 0.9 (f) x = 1.0
Oxidation state of manganese and cobalt in the various oxides
The binding energy values for the manganese-2p and cobalt-2p levels obtained by X-ray
photoelectron spectroscopy (Figure 44) confirmed that both of these transition metals were in the
oxidized state. In the case of oxygen 1s, two distinct binding energy values were found, a small
peak that corresponded to elemental oxygen possibly from absorbed oxygen species, and a more
prominent peak at lower binding energy that indicated the presence of anionic oxygen
corresponding to an oxide.
The binding energy of the oxidized states of manganese 2p
3/2
spanned from 641 to 645
eV and the peaks were asymmetric suggesting that many oxidation states were present on the
surface. Similarly with cobalt, the binding energy associated with the 2p
3/2
peak spanned from
82

779 to 784 eV. To obtain further insight into the distribution of the oxidation states, we de-
convoluted the Mn 2p
3/2
and Co 2p
3/2
peaks for each catalyst and estimated the contributions of
Mn
2+
, Mn
3+
, Mn
4+
, Co
3+
and Co
4+
(Figure 45).

Figure 44. X-ray photoelectron spectroscopy of (a) Mn-2p (b) Co-2p and (c) O-1s for the
various compositions indicated by the values of atomic ratio x = Mn/Co. XPS spectra were
corrected using carbon spectra as a standard.
The binding energy values of the individual oxidation states used for the de-convolution
were Mn
2+
: 641.4 eV, Mn
3+
: 642.5 eV, Mn
4+
: 644.1 eV, Co
3+
: 779.6 eV, and Co
4+
: 781.5 eV.
These assignments are based on the values of the oxides of cobalt and manganese reported in
various compounds.
169
The sample without manganese shows nearly 80 % of Co
3+
, whereas
introduction of manganese lowers the surface concentration of Co
3+
to nearly 55 % and
83

subsequently increased to about 70% as the manganese fraction approaches x=0.9. A large
fraction of the surface manganese ions was distributed between oxidation states of +3 and +4. A
significant amount of Mn
4+
was present even in the samples with low manganese content, while
the Mn
2+
content approached 50% at x=1.0. The distribution of manganese among the lower
oxidation states (as for example 2+) suggested the presence of oxygen vacancies.

Figure 45. (a) Variation of concentration of Mn
2+
, Mn
3+
, Mn
4+
species with the manganese
fraction determined by deconvolution of Mn-2p
3/2
XPS peaks from data presented in Figure 3.
(b) Average oxidation state of manganese as a function of manganese fraction, x. (c) Variation of
concentration of Co
3+
and Co
4+
species with the manganese fraction as determined by
deconvolution of Co-2p
3/2
peaks in Figure 44. (d) Calculated average oxidation state of cobalt as
a function of manganese fraction, x.

84

Figure 46. (a) Manganese XANES for manganese fraction x = 0.1 to 1 in La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3

and LaMnO
3
(Marked as LMO). Inset has been include for clarity. (b) Cobalt XANES for
manganese fraction x = 0 to 0.9 in La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
.

The XANES data corresponding to manganese is presented in Figure 46a. In the pre-edge
region some residual oscillations from the lanthanum L1-edge are present. These oscillations
marginally distort the manganese XANES for the x = 0.1 sample and contributes to a slowly-
varying and small background. For higher manganese fractions the distortion of the XANES
spectra is greatly reduced. With respect to the pure LaMnO
3
, the main edge position of the x=1.0
sample is shifted to larger values by ~ 0.5 eV. This observation is consistent with the increased
average oxidation state of manganese resulting from the replacement of some lanthanum with
calcium.
For samples containing both manganese and cobalt, the charge compensation for calcium
replacing lanthanum can be accommodated by both the transition metals. From the manganese
K-edge XANES it is clear that when cobalt replaces manganese (while maintaining the La/Ca
ratio) the average oxidation state of manganese increases. Specifically, as the manganese content
is reduced from x=1.0 to x=0.3, the manganese edge position shifts progressively to higher
85

energies. The edge position for the x=0.1 and x=0.3 sample is coincident. With respect to x=1.0
sample the edge position of x=0.1 sample is shifted to larger values by ~ 0.5 eV. These
observations indicate the enhanced presence of Mn
4+
even as the manganese content decreases.
This observation is consistent with the XPS results that also indicated the presence of Mn
4+
in the
samples x=0.1 to 0.8.
Earlier studies have explored the relationship of the edge position with the Mn oxidation
state and oxygen content of samples in other manganites.
170-172
A shift of ~ 3.3 eV / valence unit
has been reported in one study; however, a much smaller shift has also been seen in another
study. These differences are attributed to differences in oxygen stoichiometry of the various
sample sets. Finally, the pre-edge peaks are broad and largely similar with only minor changes in
position and shape. This XANES observation reveals the absence of any significant amount of
Mn
2+
in the bulk of the samples. This observation is consistent with the XPS results that indicate
minimal concentration of Mn
2+
in almost all the samples except for x=0.9 and x=1.
Vashook et al.,
162
have reported from XANES studies of La
0.6
Ca
0.4
Co
1-x
Mn
x
O
3
that the
oxidation state of manganese is dependent on fraction x. These studies indicate that changes in
oxidation state of manganese can be induced by cobalt substitution. Our XANES studies indicate
that the relative amount of Mn
4+
increases with as the ratio of cobalt to manganese increases.
(Figure 46). Also, the trend of oxidation state of manganese with x (Mn fraction) is opposite to
that of XPS. Analysis of the XANES data shows that the average bulk oxidation state increases
slightly with increasing manganese content, whereas XPS indicates a slight decrease in surface
oxidation states. This inconsistency can arise from XPS being a surface-sensitive measurement,
whereas XANES is primarily sampling the bulk of the material. This difference between XANES
and XPS results suggested a higher concentration of oxygen vacancies on the surface relative to
86

the bulk leading to the lower oxidations state of manganese on the surface. Consequently,
examining XPS data in conjunction with the XANES data provided insight into the difference
between the bulk and surface compositions.
The results of the cobalt XANES (Figure 46b) suggested that the main edge position for
samples with high cobalt content (x=0-0.3) was largely coincident but as the cobalt content
further diminished (x = 0.5-0.9) the edge position shifted systematically to lower energies. In
prior studies, Sikora et al
171
reported that addition of manganese to LaCoO
3
induced a reduction
of Co
3+
to Co
2+
, with a shift in the edge position by ~3.2 eV/unit valence. In samples where
cobalt could potentially exist as 4+, such as in La
1-x
Sr
x
CoO
3
, very little change in edge position
was seen.
173
It has been argued that the higher formal valence of cobalt did not lead to a
significant edge shift, in this system, presumably due to the higher covalency of the Co-O bond.
Thus the Co K-edge position might not have much direct sensitivity to the presence of Co
4+
. A
comparison of the Co XANES of x=0 and x=0.9 samples indicates a lowering of edge position
by ~ 1.5 eV, which can be ascribed to the presence of Co
2+
in samples as the amount of
manganese increases. This increased presence of Co
2+
can be interpreted from the evolution of
white line (first main peak), which is also similar to that seen in the data of Sikora et al.
171

In contrast to the results from the XANES studies, the results from XPS indicated a
nearly constant average oxidation state close to +3.4 for cobalt in all the samples. The de-
convoluted peaks of cobalt indicated "U-shaped" and ―inverted-U" shaped trend for Co
3+
and
Co
4+
variation with fraction x, respectively. These differences suggested that the surface and
bulk compositions of the perovskites could vary significantly. Thus, care must be taken in
attributing the observed trends in catalytic activity simply to the bulk properties. Also, the
surface properties could undergo change during operation of the electrode in copious amounts of
87

oxygen and water. The changes in surface properties due to the presence of water and oxygen
can be studied using in situ techniques such as vibrational spectroscopy.
Steady-state polarization studies were conducted in the electrode potential range of 400
mV to 650 mV by holding the electrode at each value of potential for 300 seconds. The values of
electrode potential were corrected for the ohmic drop. The oxidation current was not significant
below 450 mV (Figure 47). The polarization plots showed a distinct Tafel region in the potential
range of 550-650 mV, spanning over almost two decades in current.

Figure 47. Potentiostatic polarization curves for oxygen evolution on the various catalysts
(indicated by x values) in 1 M potassium hydroxide at 25
o
C. E
MMO
vs. RHE = +0.90 V.
Oxide amount on the electrode was about 80 microgram for all the samples.
The specific activity measurements (current/mass of catalyst) at 550 mV, 615 mV and
650 mV indicated that a manganese fraction beyond 0.3 led to a significant reduction in activity
(Figure 48a). The determination of the electrochemically active surface area of oxide catalysts is
not readily possible for these catalysts. However, we did measure the BET surface area of the
materials and normalized the activity for this BET area. We notice that the trends in activity did
not change (Figure 49). Also, the value of the Tafel slope in the potential region of 550-625 mV
88

increased from about 55 mV/decade (or approximately, 2.3RT/F) to progressively larger
values, and reached about 120 mV/decade ( or approximately 4.6 RT/F) at a manganese fraction
of 1.0 (Figure 48b), where, R is the universal gas constant, T is the absolute temperature and F is
the Faraday constant.

Figure 48. Effect of manganese fraction on (a) specific activity at 550, 615 and 650 mV vs.
MMO reference, E
MMO
vs. RHE = +0.90 V (b) Tafel slope in the potential region 550 to 650
mV
0.0 0.2 0.4 0.6 0.8 1.0
0.00
0.03
0.06
0.09
0.12
OER Current at 650 mV / mA cm
-2
Mn fraction x

Figure 49. Specific activity normalize for BET surface area at 650 mV vs. MMO reference,
E
MMO
vs. RHE = +0.90 V

89

The significant change of Tafel slope with increasing manganese fraction suggests a
substantial change in the mechanism of oxygen evolution upon replacing cobalt by manganese.
As indicated by Bockris and Otagawa a couple of decades ago
79, 134, 147
, the most plausible
mechanistic pathway on perovskite oxides was the pathway consisting of elementary steps of
adsorption of hydroxide ions at the B-site (Eq. 56) followed by a rate-determining
electrochemical desorption step to generate adsorbed oxygen (Eq. 57). The next step of re-
combination of the oxygen atoms to desorb as molecular oxygen was considered to be relatively
fast compared to step 2 (Eq. 58).
Adsorption of OH
-
: M
z+
+ OH
─
⇌ M
z+
─ OH + e
─
(56)
Electrochemical desorption(rate-determining) : M
z+
─ OH + OH
─
 M
z+
─ O + H
2
O + e
─

(57)
Desorption : 2 M
z+
─ O ⇌ M
z+
+ O
2
(58)
In the above chemical equations, M
z+
is the transition metal ion with valence state z+ at
the catalyst surface. In our case, M
z+
is a mixture of cobalt and manganese, and z+ is the
oxidation state of the metal ions on the surface. An alternate rate-determining electrochemical
desorption step was also proposed by Bockris and Otagawa that involved formation of peroxo
species (Eq.59) followed by release of oxygen (Eq. 60).
M
z+
─ OH + OH
─
 M
z+
—H
2
O
2
+ e
─
(59)
M
z+
—H
2
O
2 ad
+ M
z+
—HO
2
-

 H
2
O + OH
-
+O
2
(60)
Bockris and Otagawa
79, 134, 147
reported that the Tafel slope at an overpotential of 0.3 V
(or 0.6 V vs. MMO) for the un-doped cobalt perovskite (LaCoO
3
) was 2.3 RT/F, while that for
the un-doped manganese perovskite (LaMnO
3
) was 4.6 RT/F. In our measurements of calcium-
doped perovskites (La
0.6
Ca
0.4
CoO
3
and La
0.6
Ca
0.4
MnO
3
) we found that the Tafel slope values
followed the same trend as reported by Bockris and Otagawa for the un-doped perovskites. As
90

we substitute the cobalt progressively with manganese, the Tafel slope increased from 2.3 RT/F
to 4.6 RT/F (Figure 48b).
On the basic of analysis provided by Bockris and Otagawa
79, 134, 147
, we suggest that the
difference in Tafel slope values arises from the difference in coverage of the surface by the OH
ads

species on the two oxides. When the heat of adsorption for the formation of M
z+
─ OH is low,
the surface coverage of oxygen atoms will be low (with  values ranging from 0.2 to 0.8) with
adsorbate-adsorbate interactions limiting the equilibrium coverage. Under these conditions the
application of the Temkin isotherm is appropriate and this leads to a Tafel slope value of 2.3
RT/F. When the M
z+
─ OH bond is very strong, the coverage is extensive and the application of
the Langmuir isotherm leads to a Tafel slope of 4.6 RT/F.
79
Thus, the experimentally-determined
value of Tafel slope correlated with the strength of the M
z+
─ OH interaction. The average bond
energy for a Co
3+
—OH and Mn
3+
—OH are 544 kJ/mole and 627 kJ/mole, respectively.
147
This
difference in bond strength of about 83 kJ/mole is large enough to result in a difference in
adsorption energy and surface coverage of the OH species and could explain the observed
change in Tafel slopes with the composition of the oxide.
More generally, for transition metal ions with a larger number of d-electrons, the
occupancy of the anti-bonding orbitals in the M
z+
—OH bond increases and hence the bond is
expected to be much weaker. Consequently, the higher oxidation states of any particular
transition metal will have fewer d-electrons and this will lead to stronger M
z+
—OH and thereby a
higher value of Tafel slope. The spin state of the metal ion will also have to be considered in
such bond strength estimates.
174
Thus, the bond strength estimated from the electron occupancy
of the molecular orbital scaffold for the M
3+
—OH (Figure 50), will be lower for cobalt compared
to manganese. In the case of mixed perovskites with both cobalt and manganese, we can expect a
91

smooth increase in the values of Tafel slopes as we move from surface compositions rich in Co
3+

to those richer in Mn
3+
and Mn
4+
.

Figure 50. Molecular Orbital ordering schematic for M
z+
-OH bonds for cobalt and manganese.
92

For the calcium-doped mixed metal perovskites studied here, the XPS results indicated
that the surface consisted of various oxidation states of manganese and cobalt and possibly
oxygen vacancies (Figure 45). As per our XPS results, we have a distribution of surface
oxidation states for the various catalysts. Thus, knowing the distribution of surface oxidation
states we have calculated the average value of electron occupancy of the anti-bonding orbitals for
each of the catalysts studied. This average value of electron occupancy for each of the oxide
compositions of Table 4 is the weighted average of the oxidation state values derived from XPS
data after deconvolution (Figure 45).
Table 4. Number of d-electrons for M
z+
and occupancy of antibonding orbital of M
z+
—OH bond
Metal ion d electron
configuration
Electron-occupancy of anti-bonding orbital of M
z+
—
OH bond
Mn
2+
d
5
3
Mn
3+
d
4
2
Mn
4+
d
3
1
Co
3+
(h.s.) d
6
4
Co
3+
(l.s.) d
6
3
Co
4+
d
5
3

For Co
3+
since the high spin state t
2g
4
e
g
2
is not feasible, we use the intermediate spin
case (t
2g
5
e
g
1
) but denote it as high-spin (h.s.).
175
We expect that with increasing electron
occupancy in the anti-bonding orbital the M
z+
-OH bond weakens, and OER becomes more facile.
Consistent with this expectation, we find a direct correlation between the observed activity and
the calculated average value for electron occupancy (Figure 51). The molecular orbital diagrams
used for the calculation of the average value of d-electron occupancy are provided as supporting
material.

93

Figure 51. Oxygen evolution activity and average number of electrons occupying in the
antibonding orbitals of M
z+
-OH at 650 mV vs. MMO.
5.2 Oxygen Evolution Activity of Spinel Transition Metal Oxide on
Nickel Foam Electrode

Figure 52. OER electrocatalyst activity of various metal oxide (adaped from
78
)
Jaramillo et al.,
78
has offered a benchmarking of OER electrocatalysts of various metal
oxides in 1M sodium hydroxide at 10 mA/ cm
2
(Figure 52). Since noble metals based catalysts,
are not suitable for large-scale application like grid-scale energy storage the research on low-
94

cost OER catalysts based on transition metal oxides is an important area of research. Among the
various non-noble metal based catalysts, nickel and cobalt-based mixed oxides, especially the
spinel, nickel cobalt oxide NiCo
2
O
4
, has shown significantly activity for OER.
176
In addition,
recently iron has shown a benificial effect in enhancement the oxygen evolution reaction
activity.
88-89

The transition metal oxides were prepared by a sol-gel method, also known as the Pechini
process, described in the previous section. All the catalysts were synthesized in powder form
with particle size ranging from sub-microns to few nanometers depending on the synthesis
condition. This kind of synthesis of powders is a time and energy intensive process, and these
powder catalysts must be subsequently used to form the oxygen evolution electrode using
various other methods. This type of catalyst and electrode fabrication will be expensive from a
manufacturing standpoint.
In this work, we have developed a single-step, low-cost, synthetic approach for making
catalytic electrodes. The proposed method is scalable without major modifications. The catalysts
are synthesized right on the surface of the electrode bypassing separate steps of synthesis of the
powder catalysts and application of a coating on the electrode. The new synthetic approach
consists of coating of catalyst from a precursor solution containing the relevant metal ions on to a
nickel foam electrode. This process is followed by two steps of heat treatment to form the oxide-
coated electrode. The first step of the heat treatment helps to dry out coating solution on the
nickel foam. The second step of the heat treatment involves decomposition of the precursor in air
to form the final catalyst layer (details described in chapter 2). In the present study four
electrodes were prepared in this method, in which the preparation processes was maintained and
the nickel cobalt oxide contained 10 mole% of iron. However, the temperature during the second
95

step of heat treatment was varied for each electrode to be at 200
o
C, 250
o
C, 300
o
C and 400
o
C
(Table 5). All of four electrodes were characterized by X-ray diffraction, scanning electron
microscopy, X-ray photoelectron spectroscopy, and electrochemical tests.
Table 5. Heat treatment temperature and catalysts weight of iron-doped nickel cobalt oxide
catalysts.
Temperature of Preparation Measured Catalysts Weight
200
o
C 270 mg
250
o
C 197 mg
300
o
C 174 mg
400
o
C 163 mg

The X-Ray diffractogram of the catalysts shows two typical peaks of nickel metal arising
from the nickel foam at 2θ = 45
o
and 53
o
(Figure 53a) for all catalysts prepared at indicated
temperature. According to the PDF file of NiCo
2
O
4
(Figure 53d), the peaks of interest were to be
at 2θ = 31.14
o
, 36.69
o
, 59.09
o
, and 64.98
o
. However, we did not see any typical pattern peaks for
NiCo
2
O
4
in the XRD (Figure 53a), this may due to the amount of catalyst on the nickel foam
being very small. Therefore, we magnified these regions and we found some peaks that
corresponded to NiCo
2
O
4
that increased in magnitude with increase in temperature of preparation
(Figure 53b,c). These peaks became more crystalline as the heat treatment temperature was
increased from 200
o
C to 400
o
C but the deposits were still largely amorphous. These observations
can be explained by the dehydration of nickel hydroxide and cobalt hydroxide during heat
treatment (Eq. 61). This result can also be confirmed by the progressive weight loss of the
electrode with increasing heat treatment temperature (Table 5).
Ni(OH)
2
+ 2Co(OH)
2
→ NiCo
2
O
4
+ xH
2
O (61)
96

Figure 53. (a) X-ray diffraction pattern for iron-doped nickel cobalt oxide prepared at various
temperature values (b), (c) Magnification of the region in the box in (a), (d) Powder diffraction
file data PDF#00-020-0781 for NiCo
2
O
4
.
The surface morphology for all of four catalytic electrodes prepared at 200, 300, 350, and
400
o
C are very similar. The structure of nickel foam and the magnified pictures of the oxide-
coated on nickel foam are presented in Figure 54.
97

Figure 54. SEM picture of Fe doped NiCo
2
O
4
coated on nickel foam (first row), magnified
picture of metal oxide (second row).
The X-ray photoelectron spectroscopy of nickel, cobalt, iron and oxygen did not show
any systematic trend with heat treatment temperatures (Figure 55). The binding energy of nickel
2p
3/2
and cobalt 2p
3/2
spanned from 650 to 668 eV and 775 to 787 eV, respectively. However, the
XPS peaks were asymmetric. This asymmetry suggested a mixture of oxidation states. Therefore,
we performed a deconvolution of the peaks and confirmed that Ni
2+
, Co
3+
, and Co
4+
were present
on the surface of all the electrodes (Figure 56).
The binding energy of the iron 2p
3/2
and oxygen 1s peaks spanned from 702 to 728 eV
and from 526 to 535 eV, respectively (Figure 57). Upon deconvolution these peaks we confirmed
the presence of Fe
3+
on all the electrode surfaces. In the case of oxygen 1s, we found that there
was oxygen from oxides and oxygen from hydroxide species (Figure 57).
98

Figure 55. X-ray photoelectron spectroscopy of Ni 2p, Co 2p, Fe 2p, and O 1s for various
electrode at different temperatures of heat treatment.
99

Figure 56. Example of deconvolution of XPS data of Ni 2p
3/2
and Co 2p
3/2
and variation of
concentration of Ni
2+
, Co
3+
, and Co
3+
as a function of the temperature of heat treatment step.

Figure 57. Variation of concentration of Fe
3+
,O
2
-
, and OH
-
as determined by deconvolution of
the XPS of Fe 2p
3/2
and O 1s at various temperatures.
100

Figure 58. (a) Steady-state polarization curves for various electrodes (b) Current density at 520
mV vs MMO for various electrodes (c) Electrode potentials at 10 mA/cm
2
(d) Tafel slope for
various electrodes in the potential range from 500 mV to 650 mV.
Steady state polarization studies of all four electrodes were conducted in the electrode
potential range of 400 mV to 650 mV vs MMO (details of the experimental setup and
electrochemical measurements are described in Chapter 2). The steady-state polarization curves
indicated increasing catalytic activity for OER activity as the heat treatment temperature for the
electrodes was reduced (Figure 58a). At the electrode potential of 520 mV vs. MMO the
electrode prepared at 200
o
C had a current density 5 times that of the electrode prepared at 400
o
C
(Figure 58b). The enhancement of OER activity was more apparent at current density of
10mA/cm
2
(Figure 58c). The electrode potential decreased steadily as the heat treatment
temperature was decreased. There was about 80 mV difference in electrode potential between the
101

electrodes prepared at 200
o
C and 400
o
C. The best performing electrode was the one prepared at
200
o
C that had an overpotential of 225 mV at current density of 10 mA/cm
2
. However, the Tafel
slope in the electrode potential range 500 mV to 650 mV increased from 32 mV to 47 mV as the
heat treatment temperature of the electrode decreased from 400
o
C to 200
o
C (Figure 58d).
Recently, there are many reports on the high electrocatalytic activity of amorphous materials for
oxygen evolution reaction.
87, 177-178
Based on the Tafel slope and the OER activity observed for
the four electrodes prepared at various temperatures, we postulate that there are two plausible
mechanisms for oxygen evolution on these oxide-coated surfaces (Figure 59). The first step is
the adsorption of OH
-
on the active surface site (S) which is very fast because of a significant
heat of adsorption. When the adsorbed OH
-
is formed on the surface (S), it can be followed by
either step 2a or step 2b, and for these the Tafel slope will be 2.3*2RT/3F and 2.3*RT/2F,
respectively. The absorbed oxygen S-O species undergo association and desorption to form
oxygen
Step 2a will be the rate determining step, when the surface group does not have high
surface mobility. However, when the surface groups have high surface mobility the mechanism
involving Step 2b as the rate determining step will be favored. Amorphous materials are likely to
have lower surface mobility for the hydroxide compared to the crystalline materials because of
the lower level of ordering on the surface. We also observed that amorphous materials have
higher OER activity than the crystalline material. This result may be due to the lower oxygen
coordination on the surface of the amorphous materials compared to the crystalline phases,
allowing an abundance of hydroxyl groups to be formed on the surface of amorphous materials.
Thus, both the number of hydroxyl groups formed on the surface by adsorption and their surface
mobility appear to be different as the heat treatment temperature is varied.
102

Figure 59. Schematic reaction mechanisms of oxygen evolution reaction on iron-doped nickel
cobalt oxide electrode.

Figure 60. (a) Electrode potential vs time at 10 mA/ cm
2
current density for 65 hours. (b)Micro
volt per hour plot of various electrodes
103

The results of long-term chronopotentiometric studies of all four electrodes for 65 hours
are shown in Figure 60a. All of the electrodes were relatively stable. However, the electrode heat
treated at 200
o
C was not only stable but also showed some improvement in activity over time
(Figure 60b). These results confirmed that we have successfully synthesized catalytic electrodes
that are robust to oxygen evolution reaction.

104

Chapter 6
Studies on Iron-Air Rechargeable Battery
As mentioned in Chapter 1 the iron-air rechargeable battery has many advantages such as
high theoretical energy density, low material cost and environmental friendliness, properties that
are much needed for large scale energy storage. However, the round-trip energy efficiency
during charge and discharge cycle is still low. The low energy efficiency arises from poor
charging efficiency at the iron electrode and voltage losses at the air electrode.
179
Further the
durability of the air electrodes is not adequate for grid-scale applications. In the last five years,
Prof. Narayan‘s group has demonstrated a low cost and robust iron electrode with over 1200
cycles.
180
However, the challenge of reducing the voltage losses and increasing the durability at
the air electrode must now be resolved to realize a low cost, high efficiency, and durability
bifunctional air electrode.
We reported in Chapters 4 and 5 that carbon was the primary catalyst that generated
hydroperoxide and the transition metal oxide decomposed the hydroperoxide during oxygen
reduction process. However, carbon is oxidized during oxygen evolution process. Therefore,
using the same electrode for oxygen reduction and oxygen evolution would result in rapid
degradation of both ORR and OER activity of the electrode. Thus, to ensure durability for large-
scale energy storage, we have chosen to design an iron-air cell with separated ORR and OER
electrode even though this may reduce the energy density of the cell.
All electrode preparation steps and iron-air cell configuration are described in Chapter 2.
105

We have tested the behavior of ORR and OER electrodes based on nickel cobalt oxide.
We have used NiCo
2
O
4
_Vulcan XC72 on Toray® paper as the ORR electrode and the iron-
doped NiCo
2
O
4
on nickel foam as the OER electrode and combined these with the iron electrode.
NiCo
2
O
4
-Vulcan XC-72 catalyst for ORR studies was synthesized by using the same
procedure in Chapter 2. However, the Vulcan XC-72 carbon black was added into the mixture of
nitrate salt solution during the synthesis and not physical mixed with the transition metal oxide.
The addition of the carbon black to the synthesis mixture ensured that a coating of transition
metal oxide was formed on the carbon surface. Such an intimate contact between the oxide and
the carbon was expected will facilitate more rapid rate of decomposition of hydroperoxide
compared to a catalyst where the oxide and the carbon are just physically mixed. The OER
electrodes were prepared on the nickel foam electrode as described in Chapter 4.

Figure 61. Polarization curves of commercial MnO
2
catalyst and NiCo
2
O
4
-VulcanXC-72 catalyst
for ORR in iron-air cell containing 30% potassium hydroxide electrolyte.
106

The ORR activity of electrodes made from nickel cobalt oxide was coated on Vulcan XC-
72 carbon black was compared with commercial carbon-manganese dioxide electrodes (used in
the manufacture of zinc-air batteries). We found that the nickel cobalt oxide-Vulcan XC-72
electrodes showed a significantly higher ORR activity as compared to commercial MnO
2
electrode (Figure 61). In fact the NiCo
2
O
4
-Vulcan-XC-72 was 150 mV more positive to that of
the commercial MnO
2
electrode at a current density of 10 mA/cm
2
. Further the polarization
curves indicated that the electrodes were not limited by the mass transport of oxygen to the
surface of the catalyst.

Figure 62. (a) Polarization curve of 5% Fe doped NiCo
2
O
4
OER and NiCo
2
O
4
-Vulcan XC72
ORR electrode (b) Long term testing of ORR and OER electrodes at 250 mA
Similarly, the oxygen evolution reaction was studied on the nickel foam coated with iron-
doped nickel cobalt oxide. These polarization curves were obtained by switching from oxygen
reduction during discharge to oxygen evolution during charge using a diode arrangement.
Despite the improvements in catalytic activity achieved with the ORR and OER
electrodes, the overpotential losses at 10 mA/cm
2
were quite significant. At 10 mA/cm
2
there was
a 724 mV difference in the electrode potentials during ORR and OER (Figure 62b). Ignoring the
107

minor losses due to the resistance of the electrolyte and overpotential losses at the iron electrode
the difference of 724 mV in the potential of the oxygen electrode during charge and discharge,
yields a round-trip energy efficiency of 50% for the iron-air cell. Consequently, further
improvement in catalytic activity is needed to achieve a round trip efficiency of 75-80%.
Both of the OER and ORR electrodes showed a relatively stable electrode potential
during 100 hours stability test at 10 mA/cm
2
(Figure 62b) suggesting that the oxides do not
undergo any irreversible transformation during cycling. Such stability was achievable only
because we were able to use separated electrodes for ORR and OER. The test results are shown
in Table 6.
Table 6. OER and ORR testing of nickel cobalt oxide electrodes

We have also studied different OER electrodes by coating the perovskite oxide
lanthanum calcium nickel oxide on a sintered nickel substrate. These electrodes were also
prepared in the same way as the iron-doped nickel cobalt oxide electrode. The electrode GE 39
had double the catalyst loading of electrode GE 34. The polarization curves, stability testing and
results are shown in Figure 63 and Table 7. At 10 mA/cm
2
the electrode GE 39 had lowest OER
overpotential compared to all the electrodes tested.
108

The electrodes were tested for 200 hours at a constant current of 10 mA/cm
2
and no
noticeable increase in overpotential was observed. Therefore, it is likely that the catalysts used in
this study can provide the durability required of an iron-air cell. Another important feature of the
electrodes is that they are configured to operate separately allowing the oxygen evolution to
occur only on the foam or sintered nickel structures. Thus, the oxidation of carbon and
subsequent wetting of the electrode is avoided.

Figure 63. (a) Comparison of polarization curves of different OER electrodes (b) Long term
testing of OER electrodes at 250 mA
Table 7. OER testing of lanthanum nickelate electrodes

109

Chapter 7
Conclusions
This dissertation work is aimed at understanding the principles underlying operation of
transition-metal-oxide-based electrocatalysts for carrying out oxygen reduction and oxygen
evolution efficiently in aqueous alkaline media. The focus has been on perovskite and spinel
oxides containing transition metals such as cobalt, nickel, manganese and iron. The
understanding is important for the development of a long-life and efficient iron-air rechargeable
battery for grid-scale energy storage applications. Such batteries are essential for the integration
of the solar photovoltaic and wind energy into electricity grid.
The results and analysis of the electro-reduction of oxygen in aqueous alkaline media
establish that the composite catalysts prepared by mixing electrically-conductive lanthanum
calcium cobalt oxide (LCCO) with acetylene black exhibit as high as ten times more
electrocatalytic activity compared to catalysts consisting of just LCCO or carbon. We establish
that the apparent synergistic effect arises from the acetylene black serving the role of the primary
electrocatalyst for oxygen reduction and LCCO acting as a co-catalyst. Specifically, the surface
of acetylene black and other conductive carbon additives supports the electro-generation of
hydroperoxide anion while the transition metal oxide rapidly decomposes the hydroperoxide
anion to oxygen to enhance the total current. We conclude that the role of carbon in the
composite catalyst is substantially different from the commonly perceived additive that increases
the electrical conductivity of the catalyst layer. We have verified the well-defined roles of the
transition metal oxide and carbon by the analysis of the results of rotating-ring disk experiments
and determined the heterogeneous rate constants for the decomposition of hydrperoxide on the
110

surface of LCCO. We find that LCCO is about as active as platinum for decomposing
hydroperoxide.
We also see the ORR activity is strongly dependent on the transition metal oxide
composition. This effect has been tested with manganese and cobalt containing perovskites
(La
0.6
Ca
0.4
Mn
x
Co
1-x
O
3
) where the ratio of cobalt to manganese can be varied. The ORR activity
increased with the value of the manganese fraction. We found a strong dependence of ORR
activity on the ability of transition metal oxide to decompose hydroperoxide to oxygen. By
studying the hydrogen peroxide reduction and hydrogen peroxide oxidation on the above
perovskite family of oxides, we proved that the mixed electrode potential of various transition
metal oxides in hydrogen peroxide solution was governed by the two conjugate reactions of
hydroperoxide oxidation to oxygen and hydroperoxide reduction to hydroxide mediated by
oxidation and reduction of the transition metal ion. With polarization studies I was able to
demonstrate for the first time that the mixed electrode potential under open circuit conditions
could be predicted from the kinetics of the conjugate reactions. These findings confirm that an
electrochemical mechanism operates in the decomposition of hydroperoxide. The electrode
potential observed in oxygen reduction and peroxide decomposition are therefore dependent on
the properties of the transition metal ion in the oxides. Further investigation on direct hydrogen
peroxide decomposition must be performed to confirm if the electrochemical decomposition of
hydrogen peroxide is the only pathway, as other pathways that operate via formation of free
radicals are also known in other contexts.
Electrochemical studies also demonstrated that nanocrystalline oxides of the general
formula La
0.6
Ca
0.4
Mn
x
Co
1-x
O
3
of the perovskite structure are active electrocatalysts for the
oxygen evolution reaction. We have shown that the multiple type of B sites in ABO
3
type
111

perovskite oxides can be tuned systematically for OER activity, extending the understanding
from previous studies. XPS and XANES studies indicated that the oxidation state of cobalt in the
cobalt-rich compositions was largely in the 3+ state, while a more reduced cobalt surface was
seen in the manganese-rich composition. The manganese ions in all the samples were distributed
between oxidation states of +2, +3 and +4 on the surface as per XPS, while the bulk of the
samples contained mainly manganese in the 4+ state even at low manganese content. While
manganese is relatively inexpensive, substitution of cobalt by manganese beyond x=0.3 led to an
order of magnitude decrease in electrochemical activity and increase of Tafel slope. Analysis of
the kinetics based on surface coverage by M
z+
-OH species suggested that the relative strength of
the cobalt-oxygen and manganese-oxygen bonds determined the surface activity; higher bond
strengths of the manganese (III) —OH and manganese (IV) —OH most likely led to extensive
coverage of the surface by hydroxide in the manganese rich samples, reducing activity, and
increasing the Tafel slope value of 2.3RT/F to 4.6RT/F. The average electron occupancy of the
σ* anti-bonding orbital of the M
z+
—OH bond correlated with the increase in bond energy and
reduction of activity for OER confirming that the higher oxidation state of the manganese site is
not desirable for increased activity. This insight on the role of the surface oxidation states of the
transition metal ions, bond energies of the M
z+
—OH bonds, and the population of the σ* anti-
bonding orbitals of the M
z+
—OH can be used to predict and tune the activity of the catalysts for
OER.
Noting the beneficial effect of iron on catalytic activity of nickel oxides for oxygen
evolution we studied oxygen evolution activity on iron-doped nickel cobalt oxide catalyst. We
have developed a single step, low cost, synthesis approach for an oxygen evolution electrode that
uses a nickel foam or a sintered nickel plaque as a substrate. The oxide coating on the substrates
112

was prepared by thermal decomposition of the precursor on the substrate. Such electrodes when
used for oxygen evolution had an overpotential of just 225 mV at current density of 10 mA/cm
2
.
These values of overpotential are among the lowest values reported for non-precious metal or
metal oxide catalysts in alkaline media. We have found that the OER overpotential is strongly
dependent on the temperature of heat treatment during preparation. The OER activity is five time
higher for the electrode prepared at 200
o
C as compared to the electrode prepared at 400
o
C. We
attributed higher activity to the amorphous nature of the oxides prepared at the lower
temperature. The amorphous nature of the oxides leads to a larger number of active sites.
We have designed an iron-air cell to test the performance of the ORR and OER catalysts.
Our NiCo
2
O
4
-Vulcan-XC72 exhibited a 150 mV better in potential as compared to commercial
MnO
2
electrode. We demonstrated at 10 mA/cm
2
that the La
0.6
Ca
0.4
NiO
3
OER electrode and
NiCo
2
O
4
-Vulcan-XC72 ORR electrode operated at an overpotential of 100 mV and 300 mV for
oxygen evolution and oxygen reduction reaction, respectively. Continuous operation at 10
mA/cm
2
in the iron-air cell did not show any sign of degradation for 200 hours for the OER
electrodes and for 100 hours for the ORR electrodes. These studies also prove the benefit of
operating the air electrodes as separate electrodes instead of a single bi-functional electrode.

113

References
1. Armand, M.; Tarascon, J. M., Building Better Batteries. Nature 2008, 451, 652-657.
2. Bruce, P. G.; Hardwick, L. J.; Abraham, K. M., Lithium-Air and Lithium-Sulfur Batteries.
MRS Bulletin 2011, 36, 506-512.
3. Neburchilov, V.; Wang, H.; Martin, J. J.; Qu, W., A Review on Air Cathodes for Zinc–Air
Fuel Cells. Journal of Power Sources 2010, 195, 1271-1291.
4. Narayanan, S. R.; Prakash, G. K. S.; Manohar, A.; Yang, B.; Malkhandi, S.; Kindler, A.,
Materials Challenges and Technical Approaches for Realizing Inexpensive and Robust
Iron–Air Batteries for Large-Scale Energy Storage. Solid State Ionics 2012, 216, 105-109.
5. Goldstein, J.; Brown, I.; Koretz, B., New Developments in the Electric Fuel Ltd. Zinc/Air
System. Journal of Power Sources 1999, 80, 171-179.
6. Li, Y.; Dai, H., Recent Advances in Zinc–Air Batteries. Chemical Society Reviews 2014, 43,
5257-5275.
7. Cheng, F.; Chen, J., Metal–Air Batteries: From Oxygen Reduction Electrochemistry to
Cathode Catalysts. Chemical Society Reviews 2012, 41, 2172-2192.
8. Rahman, M. A.; Wang, X.; Wen, C., High Energy Density Metal-Air Batteries: A Review.
Journal of The Electrochemical Society 2013, 160, A1759-A1771.
9. Manohar, A. K.; Malkhandi, S.; Yang, B.; Yang, C.; Prakash, G. S.; Narayanan, S., A High-
Performance Rechargeable Iron Electrode for Large-Scale Battery-Based Energy Storage.
Journal of The Electrochemical Society 2012, 159, A1209-A1214.
10. Manohar, A. K.; Yang, C.; Malkhandi, S.; Yang, B.; Prakash, G. S.; Narayanan, S.,
Understanding the Factors Affecting the Formation of Carbonyl Iron Electrodes in
Rechargeable Alkaline Iron Batteries. Journal of The Electrochemical Society 2012, 159,
A2148-A2155.
11. Kordesch, K. V., 25 Years of Fuel Cell Development/1951-1976. Journal of the
Electrochemical Society 1978, 125, 77.
12. Falcon, H.; Carbonio, R., Study of the Heterogeneous Decomposition of Hydrogen
Peroxide: Its Application to the Development of Catalysts for Carbon-Based Oxygen
Cathodes. Journal of Electroanalytical Chemistry 1992, 339, 69-83.
13. Dunn, S., Hydrogen Futures: Toward a Sustainable Energy System. International journal of
hydrogen energy 2002, 27, 235-264.
114

14. Stojić, D. L.; Marčeta, M. P.; Sovilj, S. P.; Miljanić, Š. S., Hydrogen Generation from Water
Electrolysis—Possibilities of Energy Saving. Journal of Power Sources 2003, 118, 315-319.
15. Ursua, A.; Gandia, L. M.; Sanchis, P., Hydrogen Production from Water Electrolysis:
Current Status and Future Trends. Proceedings of the IEEE 2012, 100, 410-426.
16. Zeng, K.; Zhang, D., Recent Progress in Alkaline Water Electrolysis for Hydrogen
Production and Applications. Progress in Energy and Combustion Science 2010, 36, 307-
326.
17. Slade, R. C.; Kizewski, J. P.; Poynton, S. D.; Zeng, R.; Varcoe, J. R., Alkaline Membrane
Fuel Cells. In Fuel Cells, Springer: 2013; pp 9-29.
18. Winter, M.; Brodd, R. J., What Are Batteries, Fuel Cells, and Supercapacitors? Chemical
reviews 2004, 104, 4245-4270.
19. Blomen, L. J.; Mugerwa, M. N., Fuel Cell Systems; Springer Science & Business Media,
2013.
20. Kordesch, K., Alkaline Fuel Cells Applications, Innovative Energy Technology. Austria:
Institute of High Voltage Engineering, U Graz 1999.
21. Farooque, M.; Maru, H. C., Fuel Cells-the Clean and Efficient Power Generators.
Proceedings of the IEEE 2001, 89, 1819-1829.
22. Savy, M., Oxygen Reduction in Alkaline Solution on Semiconducting Cobalt Oxide
Electrodes. Electrochimica Acta 1968, 13, 1359-1376.
23. Gasteiger, H. A.; Kocha, S. S.; Sompalli, B.; Wagner, F. T., Activity Benchmarks and
Requirements for Pt, Pt-Alloy, and Non-Pt Oxygen Reduction Catalysts for PEMFCs.
Applied Catalysis B: Environmental 2005, 56, 9-35.
24. Asazawa, K.; Yamada, K.; Tanaka, H.; Oka, A.; Taniguchi, M.; Kobayashi, T., A Platinum ‐
Free Zero ‐Carbon ‐Emission Easy Fuelling Direct Hydrazine Fuel Cell for Vehicles.
Angewandte Chemie 2007, 119, 8170-8173.
25. Nagai, T.; Yamazaki, S. I.; Fujiwara, N.; Asahi, M.; Siroma, Z.; Ioroi, T., Non-Precious
Metal Catalyst That Combines Perovskite-Type Oxide and Iron Phthalocyanine for Use as a
Cathode Catalyst in an Alkaline Fuel Cell. Journal of The Electrochemical Society 2016,
163, F347-F352.
26. Yang, R.; Leisch, J.; Strasser, P.; Toney, M. F., Structure of Dealloyed PtCu
3
Thin Films
and Catalytic Activity for Oxygen Reduction. Chemistry of Materials 2010, 22, 4712-4720.
115

27. Cheng, F.; Chen, J., Metal-Air Batteries: From Oxygen Reduction Electrochemistry to
Cathode Catalysts. Chemical Society Reviews 2012, 41, 2172-2192.
28. Yang, R.; Bian, W.; Strasser, P.; Toney, M. F., Dealloyed PdCu
3
Thin Film Electrocatalysts
for Oxygen Reduction Reaction. Journal of Power Sources 2013, 222, 169-176.
29. Gasteiger, H. A.; Kocha, S. S.; Sompalli, B.; Wagner, F. T., Activity Benchmarks and
Requirements for Pt, Pt-Alloy, and Non-Pt Oxygen Reduction Catalysts for PEMFCs.
Applied Catalysis B: Environmental 2005, 56, 9-35.
30. Lu, Y. C.; Gasteiger, H. A.; Shao-Horn, Y., Catalytic Activity Trends of Oxygen Reduction
Reaction for Nonaqueous Li-Air Batteries. Journal of the American Chemical Society 2011,
133, 19048-19051.
31. Ramaswamy, N.; Mukerjee, S., Fundamental Mechanistic Understanding of Electrocatalysis
of Oxygen Reduction on Pt and Non-Pt Surfaces: Acid Versus Alkaline Media. Advances in
Physical Chemistry 2012, 2012.
32. Vogel, W., The Reduction of Oxygen on Rhodium and Nickel in Concentrated Potassium
Hydroxide. Electrochimica Acta 1968, 13, 1821-1826.
33. Yang, Y. F.; Zhou, Y. H.; Cha, C. S., Electrochemical Reduction of Oxygen on Small
Palladium Particles Supported on Carbon in Alkaline Solution. Electrochimica acta 1995,
40, 2579-2586.
34. Pattabiraman, R., Electrochemical Investigations on Carbon Supported Palladium Catalysts.
Applied Catalysis A: General 1997, 153, 9-20.
35. Sepa, D.; Vojnovic, M.; Vracar, L. M.; Damjanovic, A., Different Views Regarding the
Kinetics and Mechanisms of Oxygen Reduction at Pt and Pd Electrodes. Electrochimica
Acta 1987, 32, 129-134.
36. Yamamoto, K.; Imaoka, T.; Chun, W. J.; Enoki, O.; Katoh, H.; Takenaga, M.; Sonoi, A.,
Size-Specific Catalytic Activity of Platinum Clusters Enhances Oxygen Reduction
Reactions. Nature chemistry 2009, 1, 397-402.
37. Lim, B.; Jiang, M.; Camargo, P. H.; Cho, E. C.; Tao, J.; Lu, X.; Zhu, Y.; Xia, Y., Pd-Pt
Bimetallic Nanodendrites with High Activity for Oxygen Reduction. science 2009, 324,
1302-1305.
116

38. Ramesh, K.; Shukla, A., Carbon-Based Electrodes Carrying Platinum-Group Bimetal
Catalysts for Oxygen Reduction in Fuel Cells with Acidic or Alkaline Electrolytes. Journal
of power sources 1987, 19, 279-285.
39. Chatenet, M.; Aurousseau, M.; Durand, R.; Andolfatto, F., Silver-Platinum Bimetallic
Catalysts for Oxygen Cathodes in Chlor-Alkali Electrolysis: Comparison with Pure
Platinum. Journal of the Electrochemical Society 2003, 150, D47-D55.
40. Wang, B., Recent Development of Non-Platinum Catalysts for Oxygen Reduction Reaction.
Journal of Power Sources 2005, 152, 1-15.
41. Zhang, L.; Zhang, J.; Wilkinson, D. P.; Wang, H., Progress in Preparation of Non-Noble
Electrocatalysts for PEM Fuel Cell Reactions. Journal of Power Sources 2006, 156, 171-
182.
42. Jaouen, F.; Proietti, E.; Lefèvre, M.; Chenitz, R.; Dodelet, J. P.; Wu, G.; Chung, H. T.;
Johnston, C. M.; Zelenay, P., Recent Advances in Non-Precious Metal Catalysis for
Oxygen-Reduction Reaction in Polymer Electrolyte Fuel Cells. Energy & Environmental
Science 2011, 4, 114-130.
43. Chen, Z.; Higgins, D.; Yu, A.; Zhang, L.; Zhang, J., A Review on Non-Precious Metal
Electrocatalysts for PEM Fuel Cells. Energy & Environmental Science 2011, 4, 3167-3192.
44. Lefèvre, M.; Proietti, E.; Jaouen, F.; Dodelet, J. P., Iron-Based Catalysts with Improved
Oxygen Reduction Activity in Polymer Electrolyte Fuel Cells. science 2009, 324, 71-74.
45. Gong, K.; Du, F.; Xia, Z.; Durstock, M.; Dai, L., Nitrogen-Doped Carbon Nanotube Arrays
with High Electrocatalytic Activity for Oxygen Reduction. science 2009, 323, 760-764.
46. Liu, R.; Wu, D.; Feng, X.; Müllen, K., Nitrogen ‐Doped Ordered Mesoporous Graphitic
Arrays with High Electrocatalytic Activity for Oxygen Reduction. Angewandte Chemie
2010, 122, 2619-2623.
47. Wu, G.; More, K. L.; Johnston, C. M.; Zelenay, P., High-Performance Electrocatalysts for
Oxygen Reduction Derived from Polyaniline, Iron, and Cobalt. Science 2011, 332, 443-447.
48. Wu, J.; Yang, Z.; Li, X.; Sun, Q.; Jin, C.; Strasser, P.; Yang, R., Phosphorus-Doped Porous
Carbons as Efficient Electrocatalysts for Oxygen Reduction. Journal of Materials Chemistry
A 2013, 1, 9889-9896.
117

49. Choi, C. H.; Park, S. H.; Woo, S. I., Heteroatom Doped Carbons Prepared by the Pyrolysis
of Bio-Derived Amino Acids as Highly Active Catalysts for Oxygen Electro-Reduction
Reactions. Green Chemistry 2011, 13, 406-412.
50. Bao, X.; von Deak, D.; Biddinger, E. J.; Ozkan, U. S.; Hadad, C. M., A Computational
Exploration of the Oxygen Reduction Reaction over a Carbon Catalyst Containing a
Phosphinate Functional Group. Chemical Communications 2010, 46, 8621-8623.
51. Elzing, A.; Van der Putten, A.; Visscher, W.; Barendrecht, E., The Cathodic Reduction of
Oxygen at Cobalt Phthalocyanine: Influence of Electrode Preparation on Electrocatalysis.
Journal of electroanalytical chemistry and interfacial electrochemistry 1986, 200, 313-322.
52. Morozan, A.; Jousselme, B.; Palacin, S., Low-Platinum and Platinum-Free Catalysts for the
Oxygen Reduction Reaction at Fuel Cell Cathodes. Energy & Environmental Science 2011,
4, 1238-1254.
53. Dodelet, J. P.; Chenitz, R.; Yang, L.; Lefèvre, M., A New Catalytic Site for the
Electroreduction of Oxygen? ChemCatChem 2014, 1866-1867.
54. Strickland, K.; Miner, E.; Jia, Q.; Tylus, U.; Ramaswamy, N.; Liang, W.; Sougrati, M. T.;
Jaouen, F.; Mukerjee, S., Highly Active Oxygen Reduction Non-Platinum Group Metal
Electrocatalyst without Direct Metal-Nitrogen Coordination. Nature communications 2015,
6.
55. Jadhav, H. S.; Kalubarme, R. S.; Jadhav, A. H.; Seo, J. G., Iron-Nickel Spinel Oxide as an
Electrocatalyst for Non-Aqueous Rechargeable Lithium-Oxygen Batteries. Journal of Alloys
and Compounds 2016.
56. Yan, W.; Bian, W.; Jin, C.; Tian, J. H.; Yang, R., An Efficient Bi-Functional Electrocatalyst
Based on Strongly Coupled CoFe
2
O
4
/Carbon Nanotubes Hybrid for Oxygen Reduction and
Oxygen Evolution. Electrochimica Acta 2015, 177, 65-72.
57. De Koninck, M.; Poirier, S. C.; Marsan, B., Cu
x
Co
3−x
O
4
Used as Bifunctional
Electrocatalyst Physicochemical Properties and Electrochemical Characterization for the
Oxygen Evolution Reaction. Journal of The Electrochemical Society 2006, 153, A2103-
A2110.
58. Nikolova, V.; Iliev, P.; Petrov, K.; Vitanov, T.; Zhecheva, E.; Stoyanova, R.; Valov, I.;
Stoychev, D., Electrocatalysts for Bifunctional Oxygen/Air Electrodes. Journal of Power
Sources 2008, 185, 727-733.
118

59. Suntivich, J.; Gasteiger, H. A.; Yabuuchi, N.; Nakanishi, H.; Goodenough, J. B.; Shao-
Horn, Y., Design Principles for Oxygen-Reduction Activity on Perovskite Oxide Catalysts
for Fuel Cells and Metal–Air Batteries. Nature chemistry 2011, 3, 546-550.
60. Malkhandi, S.; Yang, B.; Manohar, A.; Manivannan, A.; Prakash, G. S.; Narayanan, S.,
Electrocatalytic Properties of Nanocrystalline Calcium-Doped Lanthanum Cobalt Oxide for
Bifunctional Oxygen Electrodes. The Journal of Physical Chemistry Letters 2012, 3, 967-
972.
61. Shimizu, Y.; Uemura, K.; Matsuda, H.; Miura, N.; Yamazoe, N., Bi ‐Functional Oxygen
Electrode Using Large Surface Area La
1−x
Ca
x
CoO
3
for Rechargeable Metal ‐Air Battery.
Journal of The Electrochemical Society 1990, 137, 3430-3433.
62. Suntivich, J.; May, K. J.; Gasteiger, H. A.; Goodenough, J. B.; Shao-Horn, Y., A Perovskite
Oxide Optimized for Oxygen Evolution Catalysis from Molecular Orbital Principles.
Science 2011, 334, 1383-1385.
63. Man, I. C.; Su, H. Y.; Calle ‐Vallejo, F.; Hansen, H. A.; Martínez, J. I.; Inoglu, N. G.;
Kitchin, J.; Jaramillo, T. F.; Nørskov, J. K.; Rossmeisl, J., Universality in Oxygen Evolution
Electrocatalysis on Oxide Surfaces. ChemCatChem 2011, 3, 1159-1165.
64. Rossmeisl, J.; Qu, Z. W.; Zhu, H.; Kroes, G. J.; Nørskov, J. K., Electrolysis of Water on
Oxide Surfaces. Journal of Electroanalytical Chemistry 2007, 607, 83-89.
65. Hyodo, T.; Hayashi, M.; Mitsutake, S.; Miura, N.; Yamazoe, N., Praseodymium–Calcium
Manganites (Pr
1−x
Ca
x
MnO
3
) as Electrode Catalyst for Oxygen Reduction in Alkaline
Solution. Journal of applied electrochemistry 1997, 27, 745-745.
66. Kiros, Y.; Myrén, C.; Schwartz, S.; Sampathrajan, A.; Ramanathan, M., Electrode R&D,
Stack Design and Performance of Biomass-Based Alkaline Fuel Cell Module. International
journal of hydrogen energy 1999, 24, 549-564.
67. Manoharan, R.; Shukla, A., Oxides Supported Carbon Air-Electrodes for Alkaline Solution
Power Devices. Electrochimica Acta 1985, 30, 205-209.
68. Hermann, V.; Dutriat, D.; Müller, S.; Comninellis, C., Mechanistic Studies of Oxygen
Reduction at La
0.6
Ca
0.4
CoO
3
-Activated Carbon Electrodes in a Channel Flow Cell.
Electrochimica Acta 2000, 46, 365-372.
119

69. Falcón, H.; Carbonio, R. E., Study of the Heterogeneous Decomposition of Hydrogen
Peroxide: Its Application to the Development of Catalysts for Carbon-Based Oxygen
Cathodes. Journal of Electroanalytical Chemistry 1992, 339, 69-83.
70. Ortiz, J.; Gautier, J., Oxygen Reduction on Copper Chromium Manganites. Effect of Oxide
Composition on the Reaction Mechanism in Alkaline Solution. Journal of Electroanalytical
Chemistry 1995, 391, 111-118.
71. Ríos, E.; Reyes, H.; Ortiz, J.; Gautier, J., Double Channel Electrode Flow Cell Application
to the Study of HO
2
−
Production on Mn
x
Co
3−x
O
4
(0≤ x ≤ 1) Spinel Films. Electrochimica
Acta 2005, 50, 2705-2711.
72. Restovic, A.; Rıos, E.; Barbato, S.; Ortiz, J.; Gautier, J., Oxygen Reduction in Alkaline
Medium at Thin Mn
X
Co
3−X
O
4
(0≤ X≤ 1) Spinel Films Prepared by Spray Pyrolysis. Effect
of Oxide Cation Composition on the Reaction Kinetics. Journal of Electroanalytical
Chemistry 2002, 522, 141-151.
73. Singh, R.; Pandey, J.; Singh, N.; Lal, B.; Chartier, P.; Koenig, J. F., Sol-Gel Derived Spinel
M
x
Co
3−x
O
4
(M= Ni, Cu; 0≤ x≤ 1) Films and Oxygen Evolution. Electrochimica Acta 2000,
45, 1911-1919.
74. Ardizzone, S.; Trasatti, S., Interfacial Properties of Oxides with Technological Impact in
Electrochemistry. Advances in colloid and interface science 1996, 64, 173-251.
75. Trasatti, S., Electrocatalysis by Oxides—Attempt at a Unifying Approach. Journal of
Electroanalytical Chemistry and Interfacial Electrochemistry 1980, 111, 125-131.
76. Trasatti, S., Electrocatalysis in the Anodic Evolution of Oxygen and Chlorine.
Electrochimica Acta 1984, 29, 1503-1512.
77. Lee, Y.; Suntivich, J.; May, K. J.; Perry, E. E.; Shao-Horn, Y., Synthesis and Activities of
Rutile IrO
2
and RuO
2
Nanoparticles for Oxygen Evolution in Acid and Alkaline Solutions.
The journal of physical chemistry letters 2012, 3, 399-404.
78. McCrory, C. C.; Jung, S.; Peters, J. C.; Jaramillo, T. F., Benchmarking Heterogeneous
Electrocatalysts for the Oxygen Evolution Reaction. Journal of the American Chemical
Society 2013, 135, 16977-16987.
79. Bockris, J. O.; Otagawa, T., Mechanism of Oxygen Evolution on Perovskites. The Journal
of Physical Chemistry 1983, 87, 2960-2971.
120

80. Matsumoto, Y.; Sato, E., Electrocatalytic Properties of Transition Metal Oxides for Oxygen
Evolution Reaction. Materials chemistry and physics 1986, 14, 397-426.
81. Tunold, R.; Marshall, A. T.; Rasten, E.; Tsypkin, M.; Owe, L. E.; Sunde, S., Materials for
Electrocatalysis of Oxygen Evolution Process in PEM Water Electrolysis Cells. ECS
Transactions 2010, 25, 103-117.
82. Haenen, J.; Visscher, W.; Barendrecht, E., Characterization of NiCo
2
O
4
Electrodes for O
2

Evolution: Part I. Electrochemical Characterization of Freshly Prepared NiCo
2
O
4

Electrodes. Journal of electroanalytical chemistry and interfacial electrochemistry 1986,
208, 273-296.
83. Haenen, J. G. D.; Visscher, W.; Barendrecht, E., Oxygen Evolution on NiCo
2
O
4
Electrodes.
Journal of applied electrochemistry 1985, 15, 29-38.
84. Lee, D. U.; Kim, B. J.; Chen, Z., One-Pot Synthesis of a Mesoporous NiCo
2
O
4
Nanoplatelet
and Graphene Hybrid and Its Oxygen Reduction and Evolution Activities as an Efficient Bi-
Functional Electrocatalyst. Journal of Materials Chemistry A 2013, 1, 4754-4762.
85. Castro, E.; Gervasi, C., Electrodeposited Ni–Co-Oxide Electrodes: Characterization and
Kinetics of the Oxygen Evolution Reaction. International journal of hydrogen energy 2000,
25, 1163-1170.
86. Prabu, M.; Ketpang, K.; Shanmugam, S., Hierarchical Nanostructured NiCo
2
O
4
as an
Efficient Bifunctional Non-Precious Metal Catalyst for Rechargeable Zinc–Air Batteries.
Nanoscale 2014, 6, 3173-3181.
87. Lu, X.; Zhao, C., Electrodeposition of Hierarchically Structured Three-Dimensional Nickel–
Iron Electrodes for Efficient Oxygen Evolution at High Current Densities. Nature
communications 2015, 6.
88. Trotochaud, L.; Young, S. L.; Ranney, J. K.; Boettcher, S. W., Nickel–Iron Oxyhydroxide
Oxygen-Evolution Electrocatalysts: The Role of Intentional and Incidental Iron
Incorporation. Journal of the American Chemical Society 2014, 136, 6744-6753.
89. Klaus, S.; Cai, Y.; Louie, M. W.; Trotochaud, L.; Bell, A. T., Effects of Fe Electrolyte
Impurities on Ni(OH)
2
/NiOOH Structure and Oxygen Evolution Activity. The Journal of
Physical Chemistry C 2015, 119, 7243-7254.
90. Corrigan, D. A., The Catalysis of the Oxygen Evolution Reaction by Iron Impurities in Thin
Film Nickel Oxide Electrodes. Journal of the Electrochemical Society 1987, 134, 377-384.
121

91. Debe, M. K., Electrocatalyst Approaches and Challenges for Automotive Fuel Cells. Nature
2012, 486, 43-51.
92. Gasteiger, H. A.; Marković, N. M., Just a Dream—or Future Reality? Science 2009, 324,
48-49.
93. Öjefors, L.; Carlsson, L., An Iron—Air Vehicle Battery. Journal of Power Sources 1978, 2,
287-296.
94. Spendelow, J. S.; Wieckowski, A., Electrocatalysis of Oxygen Reduction and Small Alcohol
Oxidation in Alkaline Media. Physical Chemistry Chemical Physics 2007, 9, 2654-2675.
95. Feng, Y.; Alonso ‐Vante, N., Nonprecious Metal Catalysts for the Molecular Oxygen ‐
Reduction Reaction. physica status solidi (b) 2008, 245, 1792-1806.
96. Yeager, E., Electrocatalysts for O
2
Reduction. Electrochimica Acta 1984, 29, 1527-1537.
97. Yeager, E., Dioxygen Electrocatalysis: Mechanisms in Relation to Catalyst Structure.
Journal of Molecular Catalysis 1986, 38, 5-25.
98. Shukla, A.; Kannan, A.; Hegde, M.; Gopalakrishnan, J., Effect of Counter Cations on
Electrocatalytic Activity of Oxide Pyrochlores Towards Oxygen Reduction/Evolution in
Alkaline Medium: An Electrochemical and Spectroscopic Study. Journal of power sources
1991, 35, 163-173.
99. Kannan, A.; Shukla, A.; Sathyanarayana, S., Oxide-Based Bifunctional Oxygen Electrode
for Rechargeable Metal/Air Batteries. Journal of Power Sources 1989, 25, 141-150.
100. Prakash, J.; Tryk, D.; Yeager, E., Electrocatalysis for Oxygen Electrodes in Fuel Cells and
Water Electrolyzers for Space Applications. Journal of power sources 1990, 29, 413-422.
101. Li, B.; Amiruddin, S.; Prakash, J., 212th ECS Meeting, Oct. 7–12, 2007, Washington DC.
ECS Trans 2007, 11, 235-240.
102. Li, B.; Prakash, J., Oxygen Reduction Reaction on Carbon Supported Palladium–Nickel
Alloys in Alkaline Media. Electrochemistry Communications 2009, 11, 1162-1165.
103. Bockris, J. M.; Minevski, Z., Electrocatalysis: A Futuristic View. International journal of
hydrogen energy 1992, 17, 423-444.
104. Wang, X.; Sebastian, P.; Smit, M. A.; Yang, H.; Gamboa, S., Studies on the Oxygen
Reduction Catalyst for Zinc–Air Battery Electrode. Journal of power sources 2003, 124,
278-284.
122

105. Müller, S.; Striebel, K.; Haas, O., La
0.6
Ca
0.4
CoO
3
: A Stable and Powerful Catalyst for
Bifunctional Air Electrodes. Electrochimica Acta 1994, 39, 1661-1668.
106. Weidenkaff, A.; Ebbinghaus, S. G.; Lippert, T., Ln
1-x
A
x
CoO
3
(Ln=Er, La; A=Ca,
Sr)/Carbon Nanotube Composite Materials Applied for Rechargeable Zn/Air Batteries.
Chemistry of materials 2002, 14, 1797-1805.
107. Lee, C. K.; Striebel, K. A.; McLarnon, F. R.; Cairns, E. J., Thermal Treatment of
La
0.6
Ca
0.4
CoO
3
Perovskites for Bifunctional Air Electrodes. Journal of The Electrochemical
Society 1997, 144, 3801-3806.
108. Sunarso, J.; Torriero, A. A.; Zhou, W.; Howlett, P. C.; Forsyth, M., Oxygen Reduction
Reaction Activity of La-Based Perovskite Oxides in Alkaline Medium: A Thin-Film
Rotating Ring-Disk Electrode Study. The Journal of Physical Chemistry C 2012, 116, 5827-
5834.
109. Zhang, H. M.; Teraoka, Y.; Yamazoe, N., Preparation of Perovskite-Type Oxides with
Large Surface Area by Citrate Process. Chemistry Letters 1987, 16, 665-668.
110. Patterson, A., The Scherrer Formula for X-Ray Particle Size Determination. Physical review
1939, 56, 978.
111. Langford, J. I.; Wilson, A., Scherrer after Sixty Years: A Survey and Some New Results in
the Determination of Crystallite Size. Journal of Applied Crystallography 1978, 11, 102-
113.
112. Kraft, S.; Stümpel, J.; Becker, P.; Kuetgens, U., High Resolution X ‐Ray Absorption
Spectroscopy with Absolute Energy Calibration for the Determination of Absorption Edge
Energies. Review of Scientific Instruments 1996, 67, 681-687.
113. Ravel, B.; Newville, M., Athena, Artemis, Hephaestus: Data Analysis for X-Ray Absorption
Spectroscopy Using IFEEEIT. Journal of synchrotron radiation 2005, 12, 537-541.
114. Malkhandi, S.; Trinh, P.; Manohar, A. K.; Jayachandrababu, K. C.; Kindler, A.; Prakash, G.
K. S.; Narayanan, S. R., Electrocatalytic Activity of Transition Metal Oxide-Carbon
Composites for Oxygen Reduction in Alkaline Batteries and Fuel Cells. Journal of the
Electrochemical Society 2013, 160, F943-F952.
115. Li, X.; Qu, W.; Zhang, J.; Wang, H., Electrocatalytic Activities of La
0.6
Ca
0.4
CoO
3
and
La
0.6
Ca
0.4
CoO
3
-Carbon Composites toward the Oxygen Reduction Reaction in Concentrated
Alkaline Electrolytes. Journal of The Electrochemical Society 2011, 158, A597-A604.
123

116. Olson, T. S.; Pylypenko, S.; Fulghum, J. E.; Atanassov, P., Bifunctional Oxygen Reduction
Reaction Mechanism on Non-Platinum Catalysts Derived from Pyrolyzed Porphyrins.
Journal of the Electrochemical Society 2010, 157, B54-B63.
117. Levich, B., The Theory of Concentration Polarisation. Discussions of the Faraday Society
1947, 1, 37-49.
118. Filinovsky, V. Y.; Pleskov, Y. V., Rotating Disk and Ring-Disk Electrodes. Electrodics:
Experimental Techniques”, eds. E. Yeager, JO Bockris, BE Conway, S. Sarangapani, New
York: Plenum 1984, 293-352.
119. Instrumentation, P. R. Voltammetry.Net.
http://wiki.voltammetry.net/pine/rotating_electrode/theory (accessed 16 Jan).
120. Jiao, Y.; Zheng, Y.; Jaroniec, M.; Qiao, S. Z., Origin of the Electrocatalytic Oxygen
Reduction Activity of Graphene-Based Catalysts: A Roadmap to Achieve the Best
Performance. Journal of the American Chemical Society 2014, 136, 4394-4403.
121. Shimizu, Y.; Matsuda, H.; Miura, N.; Yamazoe, N., Bi-Functional Oxygen Electrode Using
Large Surface Area Perovskite-Type Oxide Catalyst for Rechargeable Metal-Air Batteries.
Chemistry Letters 1992, 1033-1036.
122. Poux, T.; Napolskiy, F.; Dintzer, T.; Kéranguéven, G.; Istomin, S. Y.; Tsirlina, G.; Antipov,
E.; Savinova, E., Dual Role of Carbon in the Catalytic Layers of Perovskite/Carbon
Composites for the Electrocatalytic Oxygen Reduction Reaction. Catalysis Today 2012,
189, 83-92.
123. Olson, T. S.; Pylypenko, S.; Atanassov, P.; Asazawa, K.; Yamada, K.; Tanaka, H., Anion-
Exchange Membrane Fuel Cells: Dual-Site Mechanism of Oxygen Reduction Reaction in
Alkaline Media on Cobalt− Polypyrrole Electrocatalysts. The Journal of Physical Chemistry
C 2010, 114, 5049-5059.
124. Ludwig, J., Bifunctional Oxygen/Air Electrodes. Journal of Power Sources 2006, 155, 23.
125. Genshaw, M.; Damjanovic, A.; Bockris, J. M., Hydrogen Peroxide Formation in Oxygen
Reduction at Gold Electrodes: I. Acid Solution. Journal of Electroanalytical Chemistry and
Interfacial Electrochemistry 1967, 15, 163-172.
126. Paulus, U.; Schmidt, T.; Gasteiger, H.; Behm, R., Oxygen Reduction on a High-Surface
Area Pt/Vulcan Carbon Catalyst: A Thin-Film Rotating Ring-Disk Electrode Study. Journal
of Electroanalytical Chemistry 2001, 495, 134-145.
124

127. Bard, A. J.; Faulkner, L. R., Electrochemical Methods: Fundamentals and Applications;
Wiley New York, 1980; Vol. 2.
128. Sarapuu, A.; Nurmik, M.; Mändar, H.; Rosental, A.; Laaksonen, T.; Kontturi, K.; Schiffrin,
D. J.; Tammeveski, K., Electrochemical Reduction of Oxygen on Nanostructured Gold
Electrodes. Journal of Electroanalytical Chemistry 2008, 612, 78-86.
129. Strbac, S.; ić, R., The Influence of OH
−
Chemisorption on the Catalytic Properties of
Gold Single Crystal Surfaces for Oxygen Reduction in Alkaline Solutions. Journal of
Electroanalytical Chemistry 1996, 403, 169-181.
130. Damjanovic, A.; Genshaw, M.; Bockris, J. M., Distinction between Intermediates Produced
in Main and Side Electrodic Reactions. The Journal of Chemical Physics 1966, 45, 4057-
4059.
131. Wroblowa, H. S.; Razumney, G., Electroreduction of Oxygen: A New Mechanistic
Criterion. Journal of Electroanalytical Chemistry and Interfacial Electrochemistry 1976, 69,
195-201.
132. Venkatachalapathy, R.; Davila, G. P.; Prakash, J., Catalytic Decomposition of Hydrogen
Peroxide in Alkaline Solutions. Electrochemistry communications 1999, 1, 614-617.
133. Carbonio, R.; Fierro, C.; Tryk, D.; Scherson, D.; Yeager, E., Perovskite-Type Oxides:
Oxygen Electrocatalysis and Bulk Structure. Journal of Power Sources 1988, 22, 387-398.
134. Bockris, J. M.; Otagawa, T.; Young, V., Solid State Surface Studies of the Electrocatalysis
of Oxygen Evolution on Perovskites. Journal of Electroanalytical Chemistry and Interfacial
Electrochemistry 1983, 150, 633-643.
135. Roy, C. B., Catalytic Decomposition of Hydrogen Peroxide on Some Oxide Catalysts.
Journal of Catalysis 1968, 12, 129-133.
136. Goldstein, J.; Tseung, A., The Kinetics of Hydrogen Peroxide Decomposition Catalyzed by
Cobalt-Iron Oxides. Journal of Catalysis 1974, 32, 452-465.
137. Falcon, H.; Carbonio, R.; Fierro, J., Correlation of Oxidation States in LaFe
x
Ni
1-x
O
3+ẟ

Oxides with Catalytic Activity for H
2
O
2
Decomposition. Journal of Catalysis 2001, 203,
264-272.
138. Weiss, J., The Catalytic Decomposition of Hydrogen Peroxide on Different Metals.
Transactions of the Faraday Society 1935, 31, 1547-1557.
125

139. Ariafard, A.; Aghabozorg, H. R.; Salehirad, F., Hydrogen Peroxide Decomposition over
La
0.9
Sr
0.1
Ni
1-
xCr
x
O
3
Perovskites. Catalysis Communications 2003, 4, 561-566.
140. Soleymani, M.; Moheb, A.; Babakhani, D., Hydrogen Peroxide Decomposition over
Nanosized La
1–x
Ca
x
MnO
3
(0 ≤ x ≤ 0.6) Perovskite Oxides. Chemical Engineering &
Technology 2011, 34, 49-55.
141. Poux, T.; Bonnefont, A.; Ryabova, A.; Kerangueven, G.; Tsirlina, G. A.; Savinova, E. R.,
Electrocatalysis of Hydrogen Peroxide Reactions on Perovskite Oxides: Experiment Versus
Kinetic Modeling. Physical Chemistry Chemical Physics 2014.
142. Perez, N., Electrochemistry and Corrosion Science; Springer Science & Business Media,
2004.
143. Hardin, W. G.; Slanac, D. A.; Wang, X.; Dai, S.; Johnston, K. P.; Stevenson, K. J., Highly
Active, Nonprecious Metal Perovskite Electrocatalysts for Bifunctional Metal–Air Battery
Electrodes. The Journal of Physical Chemistry Letters 2013, 4, 1254-1259.
144. Otagawa, T.; Bockris, J. M., Lanthanum Nickelate as Electrocatalyst: Oxygen Evolution.
Journal of the Electrochemical Society 1982, 129, 2391-2392.
145. Tseung, A.; Jasem, S., Oxygen Evolution on Semiconducting Oxides. Electrochimica Acta
1977, 22, 31-34.
146. Tseung, A.; Rasiyah, P. In Factors Governing the Choice of Semiconducting Oxides for the
Evolution of Oxygen in Alkaline Media, JOURNAL OF THE ELECTROCHEMICAL
SOCIETY, ELECTROCHEMICAL SOC INC 10 SOUTH MAIN STREET,
PENNINGTON, NJ 08534: 1981; pp C131-C131.
147. Bockris, J. O. M.; Otagawa, T., The Electrocatalysis of Oxygen Evolution on Perovskites.
Journal of the Electrochemical Society 1984, 131, 290-302.
148. Jiang, Z.; Tryk, D.; Yeager, E. In Oxygen Generation and Reduction on Nickel-Based
Intermetallic Compounds, JOURNAL OF THE ELECTROCHEMICAL SOCIETY,
ELECTROCHEMICAL SOC INC 10 SOUTH MAIN STREET, PENNINGTON, NJ
08534: 1982; pp C116-C116.
149. Tryk, D.; Aldred, W.; Yeager, E. In Electrocatalysts for Bifunctional Oxygen Electrodes for
Rechargeable Metal-Air Cells, JOURNAL OF THE ELECTROCHEMICAL SOCIETY,
ELECTROCHEMICAL SOC INC 10 SOUTH MAIN STREET, PENNINGTON, NJ
08534: 1985; pp C336-C336.
126

150. Carbonio, R.; Aldred, W.; Tryk, D.; Yeager, E. In Oxygen Electrocatalysis on Perovskites,
JOURNAL OF THE ELECTROCHEMICAL SOCIETY, ELECTROCHEMICAL SOC INC
10 SOUTH MAIN STREET, PENNINGTON, NJ 08534: 1986; pp C300-C300.
151. Singh, R. N.; Hamdani, M.; Koenig, J. F.; Poillerat, G.; Gautier, J.; Chartier, P., Thin Films
of Co
3
O
4
and NiCo
2
O
4
Obtained by the Method of Chemical Spray Pyrolysis for
Electrocatalysis III. The Electrocatalysis of Oxygen Evolution. Journal of applied
electrochemistry 1990, 20, 442-446.
152. Hardin, W. G.; Mefford, J. T.; Slanac, D. A.; Patel, B. B.; Wang, X.; Dai, S.; Zhao, X.;
Ruoff, R. S.; Johnston, K. P.; Stevenson, K. J., Tuning the Electrocatalytic Activity of
Perovskites through Active Site Variation and Support Interactions. Chemistry of Materials
2014, 26, 3368-3376.
153. Risch, M.; Grimaud, A.; May, K. J.; Stoerzinger, K. A.; Chen, T. J.; Mansour, A. N.; Shao-
Horn, Y., Structural Changes of Cobalt-Based Perovskites Upon Water Oxidation
Investigated by EXAFS. The Journal of Physical Chemistry C 2013, 117, 8628-8635.
154. May, K. J.; Carlton, C. E.; Stoerzinger, K. A.; Risch, M.; Suntivich, J.; Lee, Y. L.; Grimaud,
A.; Shao-Horn, Y., Influence of Oxygen Evolution During Water Oxidation on the Surface
of Perovskite Oxide Catalysts. The Journal of Physical Chemistry Letters 2012, 3, 3264-
3270.
155. Matsumoto, Y.; Manabe, H.; Sato, E., Oxygen Evolution on La
1−x
Sr
x
CoO
3
Electrodes in
Alkaline Solutions. Journal of The Electrochemical Society 1980, 127, 811-814.
156. Ali, K. S.; Saravanan, R.; Pashchenko, A.; Pashchenko, V., Local Distortion in Co-Doped
LSMO from Entropy-Maximized Charge Density Distribution. Journal of Alloys and
Compounds 2010, 501, 307-312.
157. Pishahang, M.; Bakken, E.; Stølen, S.; Larring, Y.; Thomas, C. I., Oxygen Non-
Stoichiometry and Redox Thermodynamics of LaMn
1−x
Co
x
O
3−ẟ
. Solid State Ionics 2013,
231, 49-57.
158. Pecchi, G.; Campos, C.; Jiliberto, M. G.; Moreno, Y.; Peña, O., Doping of Lanthanum
Cobaltite by Mn: Thermal, Magnetic, and Catalytic Effect. Journal of Materials Science
2008, 43, 5282-5290.
127

159. Pawlak, D. A.; Ito, M.; Dobrzycki, L.; Wozniak, K.; Oku, M.; Shimamura, K.; Fukuda, T.,
Structure and Spectroscopic Properties of (AA′)(BB′)O
3
Mixed-Perovskite Crystals. Journal
of materials research 2005, 20, 3329-3337.
160. Rao, C., The World of Perovskite Oxides: From Dielectrics to Superconductors. Physica C:
Superconductivity 1988, 153, 1762-1768.
161. Rao, C., Perovskite Oxides and High-Temperature Superconductivity. Ferroelectrics 1990,
102, 297-308.
162. Vashook, V.; Franke, D.; Zosel, J.; Vasylechko, L.; Schmidt, M.; Guth, U., Electrical
Conductivity and Oxygen Nonstoichiometry in the Double B Mixed
La
0.6
Ca
0.4
Mn
1−x
Co
X
O
3−ẟ
Perovskite System. Journal of Alloys and Compounds 2009, 487,
577-584.
163. Campos, C.; Pecchi, G.; Moreno, Y.; Moure, C.; Gil, V.; Barahona, P.; Pena, O., Mn-
Substituted Perovskites RECo
x
Mn
1-x
O
3
: A Comparison between Magnetic Properties of
LaCo
x
Mn
1-x
O
3
and GdCo
x
Mn
1-x
O
3
. Boletín de la Sociedad Española de Cerámica y Vidrio
2008, 47, 207-212.
164. Pecchi, G.; Campos, C.; Peña, O.; Cadus, L. E., Structural, Magnetic and Catalytic
Properties of Perovskite-Type Mixed Oxides LaMn
1−y
CoyO
3
(y= 0.0, 0.1, 0.3, 0.5, 0.7, 0.9,
1.0). Journal of Molecular Catalysis A: Chemical 2008, 282, 158-166.
165. Ghosh, K.; Ogale, S.; Ramesh, R.; Greene, R.; Venkatesan, T.; Gapchup, K.; Bathe, R.;
Patil, S., Transition-Element Doping Effects in La
0.7
Ca
0.3
MnO
3
. Physical Review B 1999,
59, 533.
166. Kolat, V.; Gencer, H.; Gunes, M.; Atalay, S., Effect of B-Doping on the Structural,
Magnetotransport and Magnetocaloric Properties of La
0.67
Ca
0.33
MnO
3
Compounds.
Materials Science and Engineering: B 2007, 140, 212-217.
167. Phuc, N.; Bau, L.; Khiem, N.; Son, L.; Nam, D., Magnetic and Transport Properties of
La
0.7
Sr
0.3
Co
1−y
Mn
y
O
3
—No Double Exchange between Mn and Co. Physica B: Condensed
Matter 2003, 327, 177-182.
168. Jonker, G., Magnetic and Semiconducting Properties of Perovskites Containing Manganese
and Cobalt. Journal of Applied Physics 1966, 37, 1424-1430.
128

169. Chastain, J.; King Jr, R., Handbook of X-Ray Photoelectron Spectroscopy: A Reference
Book of Standard Spectra for Identification and Interpretation of XPS Data. Physical
Electronics, Eden Prairie 1995.
170. Bridges, F.; Booth, C.; Anderson, M.; Kwei, G.; Neumeier, J.; Snyder, J.; Mitchell, J.;
Gardner, J.; Brosha, E., Mn K-Edge Xanes Studies of La
1−x
A
x
MnO
3
Systems (A=Ca, Ba,
Pb). Physical Review B 2001, 63, 214405.
171. Sikora, M.; Kapusta, C.; Knížek, K.; Jirák, Z.; Autret, C.; Borowiec, M.; Oates, C.;
Procházka, V.; Rybicki, D.; Zajac, D., X-Ray Absorption near-Edge Spectroscopy Study of
Mn and Co Valence States in LaMn
1−x
Co
x
O
3
(x= 0–1). Physical Review B 2006, 73,
094426.
172. Croft, M.; Sills, D.; Greenblatt, M.; Lee, C.; Cheong, S.-W.; Ramanujachary, K.; Tran, D.,
Systematic Mn D-Configuration Change in the La
1− x
Ca
x
MnO
3
System: A Mn K-Edge XAS
Study. Physical Review B 1997, 55, 8726.
173. Jiang, Y.; Bridges, F.; Sundaram, N.; Belanger, D.; Anderson, I.; Mitchell, J.; Zheng, H.,
Study of the Local Distortions of the Perovskite System La
1−x
Sr
x
CoO
3
(0≤x≤ 0.35) Using
the Extended X-Ray Absorption Fine Structure Technique. Physical Review B 2009, 80,
144423.
174. Wolfram, T.; Ellialtioğlu, Ş., Concepts of Surface States and Chemisorption on D-Band
Perovskites. In Theory of Chemisorption, Springer: 1980; pp 149-181.
175. Qian, D.; Hinuma, Y.; Chen, H.; Du, L. S.; Carroll, K. J.; Ceder, G.; Grey, C. P.; Meng, Y.
S., Electronic Spin Transition in Nanosize Stoichiometric Lithium Cobalt Oxide. Journal of
the American Chemical Society 2012, 134, 6096-6099.
176. Rasiyah, P.; Tseung, A., A Mechanistic Study of Oxygen Evolution on NiCo
2
O
4
II.
Electrochemical Kinetics. Journal of the Electrochemical Society 1983, 130, 2384-2386.
177. Lian, K.; Thorpe, S.; Kirk, D., The Electrocatalytic Activity of Amorphous and Crystalline
Ni Co Alloys on the Oxygen Evolution Reaction. Electrochimica Acta 1992, 37, 169-175.
178. Tsuji, E.; Imanishi, A.; Fukui, K. I.; Nakato, Y., Electrocatalytic Activity of Amorphous
RuO
2
Electrode for Oxygen Evolution in an Aqueous Solution. Electrochimica Acta 2011,
56, 2009-2016.
179. Linden, D.; Reddy, T. B., Handbook of Batteries. 3rd. McGraw-Hill: 2002.
129

180. Manohar, A. K.; Yang, C.; Narayanan, S., The Role of Sulfide Additives in Achieving Long
Cycle Life Rechargeable Iron Electrodes in Alkaline Batteries. Journal of The
Electrochemical Society 2015, 162, A1864-A1872.

130

Publications and Presentations
Publications
● S. Malkhandi, P. Trinh, Aswin K. Manohar, A. Manivannan, M. Balasubramanian, G. K. Surya
Prakash, S.R. Narayanan, ―Design Insights for Tuning the Electrocatalytic Activity of Perovskite
Oxides for the Oxygen Evolution Reaction‖. J. Phys. Chem. C. 2015.
● Maria Abreu-Sepulveda, Phong Trinh, S. Malkhandi, S. R. Narayanan, Jacob Jorne, David J.
Quesnel, James A. Postonr Jr, A. Manivannan, ― Investigation of Oxygen Evolution Reaction of
LaRuO
3
, La
3.5
Ru
4
O
13
, La
2
RuO
5
‖. Electrochimica Acta, 2015, 180, 401-408.
● S. Malkhandi, P. Trinh, Aswin K. Manohar, K. C. Jayachandrababu, A. Kindler, G. K. Surya
Prakash and S. R .Narayanan, ―Electrocatalytic Activity of Transition Metal Oxide-Carbon
Composites for Oxygen Reduction in Alkaline Batteries and Fuel Cells”. J. Electrochem. Soc.
2013, 160, 943-952.
Presentations
Conference Presentations
Primary Presenter
● P. Trinh, S. Malkhandi, A. K. Manohar, A. Manivannan and S. R. Narayanan ―Iron-Doped
Nickel Cobalt Oxide Based High Performance Oxygen Evolution Reaction Electrode in Alkaline
Media‖ at ECS in Phoenix, AZ (October 2015)
● P. Trinh, S. Malkhandi, N. Moreno, A. K. Manohar, G. K. Surya Prakash and S. R. Narayanan
―Substitution of Co with Mn and its influence on OER activity and stability of La
0.6
Ca
0.4
CoO
3

perovskite in alkaline medium‖ at ECS-PRiME in Honolulu Hawaii (October 2012)
131

Co-author
● Maria Abreu-Sepulveda, Phong Trinh, S. Malkhandi, David J. Quesnel, S. R. Narayanan, and
A. Manivannan, ‗Comparison of Pyrochlore and Perovskite Electrodes Toward Oxygen the
Evolution in Alkaline Media‖ at ECS in Chicago (May 2015)
● S. Malkhandi, P. Trinh, A. K. Manohar, G. S. Prakash, and S.R. Narayanan, ―The Effect of
Dispersion of Metal Oxides on Carbon on the Electrocatalytic Activity for Oxygen Reduction
Reaction in Alkaline Media‖ at ECS in Orlando (May 2014)
● S. Malkhandi, P. Trinh, A. K. Manohar, G. S. Prakash, and S.R. Narayanan, ―Investigation
of Various Calcium-Based Transition Metal Oxides Compounds for the Oxygen Evolution
Reaction in Alkaline Media‖ at ECS in Orlando (May 2014)
● S. Malkhandi, P. Trinh, A. K. Manohar, K. C. Jayachandrababu,A. Kindler, G. S. Prakash,
and S.R. Narayanan, ―Understanding the Performance of Transition Metal Oxide-Carbon
Composite Catalysts in Air Electrodes for Metal-Air Batteries and Alkaline Fuel Cells‖ at ECS
in San Francisco (October 2013)
● S. Malkhandi, P. Trinh, A. K. Manohar, G. K. Surya Prakash and S. R. Narayanan ―Effect of
Substitution of Cobalt by Manganese on the Properties of Calcium Doped Lanthanum Cobalt
Oxide for Oxygen Reduction Reaction in Alkaline Medium‖ at ECS-PRiME in Honolulu Hawaii
(October 2012)

132

Poster Presentations
● P. Trinh, S. Malkhandi, A. K. Manohar, A. Manivannan, G. K. S. Prakash, and S. R. Narayan
―Understanding the Oxygen Reduction Activity of Composite Catalysts Based on Carbon and
Calcium-Doped Lanthanum Cobalt Manganese Oxide‖ at ECS in Chicago (May 2015)
● S. Malkhandi, P. Trinh, Aswin K. Manohar, K. C. Jayachandrababu, A. Kindler, G. K. Surya
Prakash and S. R. Narayanan ―Oxygen Reduction of Transition Metal Oxide-Carbon Composite
Catalysts in Air Electrodes for Metal-Air Batteries and Alkaline Fuel Cells‖, Engineering
Conferences International at Newport Beach (June 2013)

Abstract (if available)

Abstract In this dissertation I have focused on understanding the mechanism of the electrochemical reduction of oxygen (ORR) and the oxygen evolution reaction (OER) on transition metal oxides electrocatalysts in alkaline media, and the application of these electrocatalysts in a rechargeable iron-air battery. ❧ In Chapter 1, I provide an introduction to electrochemical energy storage systems, the processes of oxygen reduction and oxygen evolution, the motivation and objectives of the research, the technical challenges and the unique approach adopted in this work. In Chapter 2, I describe the experimental methodologies for the preparation and characterization of various catalytic materials, electrodes and cells. ❧ Chapter 3 contains the description of the results of experiments on the role of carbon and transition metal oxide in the catalysis of the oxygen reduction reaction. I found that conductive carbon materials when added to transition metal oxides such as calcium-doped lanthanum cobalt oxide, nickel cobalt oxide and calcium-doped lanthanum manganese cobalt oxide increase the electrocatalytic activity of the oxide for oxygen reduction by a factor of five to ten. I have studied rotating ring-disk electrodes coated with: (a) various mass ratios of carbon and transition metal oxide, (b) different type of carbon additives and (c) different types of transition metal oxides. Our experiments and analysis establish that in such a composite catalyst, carbon is the primary electro-catalyst for the two electron electro-reduction of oxygen to hydroperoxide while the transition metal oxide decomposes the hydroperoxide to generate additional oxygen that enhances the observed current resulting in an apparent four-electron process. ❧ Following up on the findings on the role of the transition metal oxide described in Chapter 3, I present results and discussion on the relationship between oxygen reduction activity and hydrogen peroxide decomposition activity of perovskite transition metal oxide in Chapter 4. I found that the oxygen reduction activity of lanthanum-doped calcium manganese cobalt oxide increased when the manganese fraction in the oxide increased. We also studied the mechanism of decomposition of hydrogen peroxide on calcium-doped lanthanum cobalt manganese oxide and established that this decomposition process occurs by an electrochemical pathway. We also verified that the higher hydrogen peroxide decomposition rate on the oxide, the better was the ORR activity of the carbon-transition metal oxide composite. ❧ In chapter 5, I describe new insights for predicting and tuning the activity of transition metal oxides for designing efficient and inexpensive electrocatalysts for the oxygen evolution reaction. I have conducted a systematic investigation of nano-phase calcium-doped lanthanum manganese cobalt oxide, an example of a mixed metal oxide that can be tuned for its electrocatalytic activity by varying the composition of the transition metals. We have found that the value of Tafel slopes are governed by the oxidation states and the bond energy of the surface intermediates (such as Mn-OH and Co-OH bonds) while the catalytic activity increased with the average d-electron-occupancy of the σ* orbital of the metal-OH bond. I have investigated the electrocatalytic activity for oxygen evolution on iron-doped nickel cobalt oxide electrodes prepared and treated at various temperatures. I found that the electrode prepared at 200 ℃ is less crystalline and has higher activity for oxygen evolution than the electrode prepared at 400 ℃. ❧ In chapter 6, I studied the performance of iron-air rechargeable cells based on the most catalytically active transition metal oxide-based oxygen reduction electrodes and oxygen evolution electrodes resulting from my research. We have demonstrated that our NiCo₂O₄-Vulcan XC72 composite electrode exhibited considerably better activity for oxygen reduction than the commercial MnO₂ electrode. I also tested the performance of iron-doped NiCo₂O₄-based oxygen evolution electrode and found that the overpotentials were sufficiently low to operate the cells at 10 mA/cm². In these experiments, the oxygen reduction electrode and the oxygen evolution electrode were operated as separated electrodes. Thus, I could demonstrate that no noticeable change in overpotential or catalytic activity was observed for at least 100 hours in the case of the oxygen reduction electrode and 200 hours in the case of the oxygen evolution electrode. ❧ Chapter 7 summarizes the understanding and conclusions arising from my research and the implications of the findings for future development of iron-air rechargeable batteries.

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Asset Metadata

Creator Trinh, Phong (author)

Core Title Understanding the mechanism of oxygen reduction and oxygen evolution on transition metal oxide electrocatalysts and applications in iron-air rechargeable battery

School College of Letters, Arts and Sciences

Degree Doctor of Philosophy

Degree Program Chemistry

Publication Date 04/21/2016

Defense Date 03/08/2016

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

Tag OAI-PMH Harvest,oxygen evolution,oxygen reduction,transition metal oxide

Format application/pdf (imt)

Language English

Contributor Electronically uploaded by the author (provenance)

Advisor Prakash, G. K. Surya (committee chair), Narayan, Sri. R. (committee member), Shing, Katherine (committee member)

Creator Email thanhphongus72@yahoo.com

Permanent Link (DOI) https://doi.org/10.25549/usctheses-c40-237530

Unique identifier UC11276879

Identifier etd-TrinhPhong-4333.pdf (filename),usctheses-c40-237530 (legacy record id)

Legacy Identifier etd-TrinhPhong-4333-0.pdf

Dmrecord 237530

Document Type Dissertation

Format application/pdf (imt)

Rights Trinh, Phong

Type texts

Source 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 a...

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